thermal and catalytic kinetics of charcoal oxidation · i thermal and catalytic kinetics of...
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Thermal and Catalytic kinetics of Charcoal Oxidation
Abdullah Saqib
Thesis to obtain the Master of Science Degree in
Chemical Engineering
Supervisor: Prof. Francisco Lemos
Examination Committee
Chairperson: Prof. Sebastiao Alves
Supervisor: Prof. Francisco Lemos
Member of the committee: Prof. Joao Bordado
June 2016
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“In the name of Allah who is the most beneficent and merciful”
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To my deceased special child
Zill-e-Noor
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Acknowledgements
Firstly, I would like to express my heartiest gratitude and thanks to my supervisor Prof.Francisco
Lemos, whose invaluable guidance and encouraging attitude enable me to finish the said tasks in
limited time. He gave his input starting from experimental set up to excel graphics till kinetic modeling
and review.
I also like to thank Prof. Maria Amelia for the provision of useful data and links along with special
thanks extended to Prof. Luis Sousa Lobo and Dr. Sonia Carabineiro for the activated charcoal
samples with the impregnated catalysts.
I would also like to thank Dr. Khurram; my supervisor at home university along with Dr. Javed Akhtar
and Dr. Rizwan Haider. They always gave me motivation to become the candidate for scholarship.
I would like to pay my special thanks to Miss Ana Barbosa at IST international office; coordinator for
Experts-Sustain program. I remember her warm welcome and all extended favors from creation of my
bank account, dealing with reimbursements, extension of visa, referring to a doctor till timely
arrangement of residence at university hostel.
I also remember the time to time help gained through my colleagues specially Mr. Everton from Brazil
and Tejas Tankaria from India.
I feel blessed to have a nice roommate at hostel, Ramices Igor from Florianapolis, Brazil. I remember
his portugese lessons and sharing of wonderful experiences.
During my six months stay in Lisbon I would like to thank Mr.Tariq who made living and lodging
arrangements upon my arrival. I enjoyed three months stay at his residence. I also remember the
BBQ evenings and jokes by Mr. Faisal. All these efforts made me feel good.
I can never forget the continuous support and encouragement of my family in Pakistan.
Finally, I would like to thank EUROPEAN COMMISSION and Experts Sustain team of Erasmus
Mundus at Goettingen University, Germany who granted me the scholarship and provided me an
opportunity to gain an international experience.
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Abstract
In order to meet the large increase in energy demand, nations are on a way for the exploration of
natural resources and the development of new technologies. Utilization of renewables is also being
considered seriously along with the efficient and environment friendly utilization of major natural
resources, like the coal in its various forms and biomass.
Now a days co-firing and co-gasification of biomass with coal has gained much attention because of
the need to reduce CO2, SOx and NOx emissions. Both biomass and coal have an inherent content of
mineral matter in varying degree and origin; one relevant area of research has become the study of
the catalytic effects of the mineral matter in char combustion and gasification. Catalytic effects of
alkali and alkaline earth metals in char conversion processes have been observed. However, the
catalytic effects of transition metals like Vanadium, Copper etc. have not gathered much attention,
although the catalytic effects of Vanadium on various oxidation processes have been recognized.
Hence, the present study is focused on the observation of the catalytic effects of Vanadium and
Copper as well as their mixtures, on the charcoal oxidation.
Thermogravimetric (TG), Differential Thermogravimetric (DTG) and Differential Scanning Calorimetry
(DSC) techniques have been used to provide insight into the catalytic effects induced by Vanadium,
Copper and their mixtures on Charcoal oxidation. Combustion and gasification experiments were
carried out using simultaneous Thermal Analysis and Differential Scanning Calorimetry utilizing pure
air as an oxidant. Isothermal kinetic data were gathered for a temperature range of 400 – 800 degree
Celsius. DSC data was also acquired simultaneously and has been used for the estimation of reaction
products utilizing enthalpy of formation of the product gases. Charcoal samples impregnated with
1%V, 1%Cu and a mixture of the two with the corresponding amounts were used to observe the
kinetic effects associated with the catalysts.
The data revealed that the charcoal impregnated with 1% vanadium induced an exothermic oxidation
reaction for temperatures above around 400°C. In case of copper and the mixture of catalysts, this
threshold temperature was around 450°C. This contrasts with the non-impregnated charcoal that had
a threshold temperature for the combustion and gasification of around 500°C. These results show
lowering of the temperatures for combustion and gasification reactions with impregnation of vanadium
and copper. The DTG results showed that the rate of reaction increases progressively with
temperature for all samples. DSC results were used for the estimation of reaction products and
indicate that, for a given supply of air, carbon monoxide formation occurs preferentially both at low
temperatures and high temperatures for reaction kinetic reasons and diffusion limitations respectively;
however, carbon monoxide formation was relatively low for middle temperature range. A kinetic model
has also been suggested which incorporated the catalytic effects and was based on Langmiur-
Hinshelwood type kinetics.
Keywords: Charcoal, combustion, gasification, catalyst, isothermal kinetics
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Resumo
A fim de atender ao grande aumento na procura de energia , todas as nações estão num caminho de
exploração dos recursos naturais e do desenvolvimento de novas tecnologias. A utilização de
energias renováveis também está a ser sendo seriamente considerada , juntamente com a utilização
eficiente e ambientalmente adequada de grandes recursos naturais, como sejam, o carvão nas suas
várias formas, e a biomassa.
Hoje em dia a co-incineração e co-gaseificação de biomassa com carvão tem ganho muita atenção
devido à necessidade de reduzir as emissoões de CO2, SOx e NOx. Tanto a biomassa como o
carvão têm um conteúdo inerente de matéria mineral em vários graus e de origem; uma área de
investigação que tem ganho relevância é o estudo dos efeitos catalíticos da matéria mineral em
combustão e gaseificação de carvão. Foram observados efeitos catalíticos de metais alcalinos e
alcalino-terrosos em processos de conversão de carvão animal. No entanto- os efeitos catalíticos de
metais de transição como o vanádio não têm recolhido muita atenção, apesar dos -efeitos catalíticos
de vanádio em vários processos de oxidação serem conhecidos. Assim,o presente estudo é focado
na observar dos efeitos catalíticos de vanádio e de cobre, bem como de uma mistura dos dois, sobre
a oxidação do carvão vegetal .
As técnicas de Termogravimetria (TG) ,Termogravimetria Diferencial (DTG) e Calorimetria Diferencial
de Varrimento (DSC) foram usadas para desenvolver um entendimento dos efeitos catalíticos
induzidos por vanádio , cobre e suas misturas sobre a oxidação do carvão vegetal. Os ensaios de
combustão e gaseificação foram realizados usando análise térmica e calorimetria diferencial de
varrimento simultânea utilizando ar puro como um oxidante- Os dados cinéticos isotérmicos foram
obtidos para uma gama de temperaturas de 400-800 graus Celsius . Os dados de DSC, que também
foram adquiridos simultanmente, forami usados para a estimativa dos produtos de reacção ,
utilizando a entalpia de formação dos produtos gasosos. Amostras de carvão impregnados com 1%
de V,1% Cu e uma mistura dos dois com os montantes correspondentes foram usadas para observar
os efeitos cinéticos associados a utilização dos catalisadores.
Os dados revelaram que, no carvão impregnado com 1% de vanádio as temperaturas acima de
cerca de 400°C. No caso do cobre e as suas misturas estade oxidação exotérmicas ocorrem
temperatura limiar foi de cerca de 450°C. Isto contrasta com o carvão vegetal não-impregnado tinha
uma temperatura de limiar para a combustão e gaseificação de cerca de 500°C. Estes resultados
mostram uma diminuição das temperaturas para as reacções de combustão e gaseificação com a
impregnação de vanádio e de cobre. Os resultados de DTG mostraram que a taxa de reacção
aumenta com a progressivamente temperatura para todas as amostras. Os resultados de DSC foram
utilizadas para a estimativa dos produtos de reacção e indicam que, para uma dada alimentação de
ar, a formação de monóxido de carbono ocorre preferncialmente a baixas temperaturas e a altas
temperaturas, por motivos cinéticos ou de limitações difusionais, respectivamente; no entanto, a
formação de monóxido de carbono foi relativamente baixa para o temperaturas intermédias. Foi
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também desenvolvido um modelo cinético, que incorporou os efeitos catalíticos e baseou numa
cinética tipo Langmiur-Hinshelwood.
Palavras-chave: carvão vegetal, combustão, gaseificação, catalisador, cinética isotérmicos
VIII
Table of Contents
Acknowledgements…………………………………………………................................................... IV
Abstract…....................................................................................................................................... V
Resumo…........................................................................................................................................ VI
Table List…..................................................................................................................................... X
Figure List…................................................................................................................................... XI
Abbreviations List…………………………………………………………………………………………. XIV
1. Introduction……………………………………………………….…………...………………….. 1
2. Literature Study…………..….………………………………………….…...…………………... 5
2.1. Biomass and charcoal……………………………………….……………………………….. 5
2.1.1. Biomass Conversion Processes………………….............................................. 5
2.1.2. Catalytic Effect of Mineral Matter……………………………………………….…. 6
2.1.3. Effect of AAEM……………………………….......................................………..… 6
2.1.4. Char Conversion Rate………………………………………………………………. 7
2.1.5. Intrinsic Reaction Rate……………..……………………………………………….. 7
2.2. Chemical Kinetics…………………....................................……………………………….. 8
2.2.1. Definition……………………………………………………………………………… 8
2.2.2. History………………………………………………………………………………… 8
2.2.3. Reaction Mechanism and Rate……………………………………………………. 9
2.2.4. Factors Affecting Reaction Rate………………………………………………….. 9
2.3. Solid state kinetics…………………………………………………………......................... 12
2.3.1 Rate Law……………………………………………………………………………. 12
2.3.2 Gas-Solid Reactant Systems……………….................................................... 13
2.3.3 Intrinsic Kinetics of Gas-Solid Systems…………............................................ 15
2.3.4 Models and Mechanisms in Solid State Kinetics……………………………….. 16
2.3.5 Data Collection and Interpretation………………………………………………... 18
2.3.6 Controversies in Solid State Kinetics……………………………………..……… 20
2.4. Catalysis………………………………………………………………………….……………. 20
2.4.1. Heterogeneous Catalysis………………………………………………….……….. 21
2.4.2. Applied Catalysis……………………………………………................................ 22
2.4.3. Heterogeneous Catalyst……………………………………................................ 23
2.5. Thermal Analysis……………………………………………………………………………… 25
2.5.1. TGA………………………………………………………………………….………... 25
2.5.2. DTA/DSC…………………………………………………….......................... ……. 26
2.5.3. Thermo-analytical Methods Vs Kinetics…………………………………………... 28
IX
3. Research Methodology……………………………………….……………………………........ 30
4. Results and Discussions…………………………………………..…………………………… 40
4.1. Charcoal (Raw)………………………………………………………………………………... 40
4.2. Charcoal (1% V Impregnated)………………………………...…………………………….. 45
4.3. Charcoal (1% Cu Impregnated)……………………………….…………………………...... 57
4.4. Charcoal (1%V+1%Cu Impregnated)………………………...…………………………….. 64
5. Conclusions…………………………….……………………………………………………........ 73
6. References………………………..……………………………..……………………………....... 75
X
Table List
2.1 Examples of Catalytic Oxidation Processes…….......................................................................... 23
2.2 Thermal Analysis Techniques…………………………………………….......................................... 25
3.1 Proximate Analysis........................................................................................................................ 30
3.2 Procedure adopted for kinetics study………………………………………………...………………... 31
3.3 Shomate Equation Constants………………………………………………………………………....... 34
3.4 Heat of formation…….................................................................................................................... 34
4.1 Gas mix composition at low temperature…………………………………………………………....... 41
4.2 Gas mix composition at high temperature……………………………………………………............. 42
4.3 Estimated Kinetic Parameters………………………………………………………………………….. 44
4.4 Gas mix composition at low temperature…………………………………………………………....... 52
4.5 Gas mix composition at high temperature…………………………………………………………….. 53
4.6 Estimated Kinetic Parameters………………………………………………………………………...... 56
4.7 Gas mix composition at low temperature…………………………………………………………….... 59
4.8 Gas mix composition at high temperature…………………………………………………………...... 60
4.9 Estimated Kinetic Parameters…………………………………………………………………….......... 63
4.10 Gas mix composition at low temperature…………………………………………………….…........ 66
4.11 Gas mix composition at high temperature……………………………………….………………...... 68
4.12 Estimated Kinetic Parameters……………………………………………………………………........ 70
5.1 Estimated Kinetic Parameters……………………………………………………………………….... 72
XI
Figure List
1.1 Processes in Gasification……………………………………………………………………………..... 3
2.1 Role of Biomass in electricity generation…………………………………………………………....... 7
2.2 Reaction Rate variations………………………………………………………………………………... 10
2.3 Generic potential energy diagram showing the effect of a catalyst in a hypothetical endothermic
chemical reaction................................................................................................................................. 12
2.4 The Arrhenius plot for different temperature regime during heterogeneous chemical reactions... 15
2.5 Mechanisms for Heterogeneous Catalysis……………………………………………………………...23
2.6 Block Diagram of Thermo-balance…………………………………………………………………….. 27
2.7 Typical DSC Curve……………………………………………………………………………………..... 28
2.8 An idealized DSC curve showing the shapes associated with particular phase transitions.......... 29
3.1 Apparatus………………………………………………………………………………………………..... 31
3.2 Enthalpy of formation for CO and CO2……………………………………………………………….... 33
3.3 Langmiur-Hinshelwood Model....................................................................................................... 38
4.1 Fractional mass loss at various temperatures………………………………………………………... 40
4.2 Conversion at various temperatures………………………………………………………………….... 41
4.3 Time Derivative of char mass fractions at low temperature……………………………………........ 41
4.4 Time Derivative of char mass fractions at high temperature……………………………………....... 42
4.5 Specific HF at low temperatures……………………………………………………………………...... 42
4.6 Specific HF at higher temperatures………………………………………………………………..…... 43
4.7 Percentage Carbon monoxide formation (Charcoal)………………………………………….……... 43
4.8 Model Fit at 550 oC…………………………………………………………………………….…..…... 43
4.9 Model Fit at Higher Temperatures……………………………………………………………..……..... 43
4.10 Raw Experimental data at 400 0C for different sample masses…………………………..……..... 44
4.11 Fractional mass loss at various temperatures…………………………………………………......... 46
XII
4.12 Conversion at various temperatures………………………………………………………………..... 46
4.13 Conversions at low temperatures................................................................................................ 47
4.14 Conversion at high temperatures…………………………………………………………………...... 48
4.15 Time Derivative of char mass fractions at low temperatures…………………………………..….. 48
4.16 Time derivative of char mass fractions at high temperatures………………………………..…..... 49
4.17 Specific HF and Fractional mass loss as function of time…………………………………..……... 49
4.18 Specific HF and Fractional mass loss as function of time……………………………………......... 50
4.19 Specific HF and Fractional mass loss as function of time………………………………………..... 50
4.20 Effect of mass on specific HF at various temperatures…………………………………………...... 51
4.21 Specific HF at low temperatures…………………………………………………………………........ 51
4.22 Specific HF at high temperatures……………………………………………………………………... 52
4.23 Percentage Carbon monoxide formation (1% V)………………………………………………........ 53
4.24 Model Fit at 400 0C………………………………………………………………………………........ 54
4.25 Model Fit at 450 oC………………………………………………………………………………........ 54
4.26 Model Fit at 500 oC………………………………………………………………………………........ 55
4.27 Model Fit at Higher Temperatures…………………………………………………………………..... 55
4.28 Fractional mass loss at various temperatures……………………………………………………..... 57
4.29 Conversion at various temperatures………………………………………………………………..... 57
4.30 Time derivative of char mass fractions at low temperatures…………………………………......... 58
4.31 Time derivative of char mass fractions at high temperatures…………………………………….... 58
4.32 Specific HF at low temperatures…………………………………………………………………........ 59
4.33 Specific HF at high temperatures……………………………………………………………………... 60
4.34 Percentage Carbon monoxide formation (1% Cu)………………………………………………...... 61
XIII
4.35 Model Fit at 450 oC………………………………………………………………………………........ 62
4.36 Model Fit at 500 and 550 oC……………………………………………………………………….... 62
4.37 Model Fit at Higher Temperatures…………………………………………………………………..... 63
4.38 Fractional mass loss at various temperatures…………………………………………………......... 64
4.39 Conversion at various temperatures………………………………………………………………..... 64
4.40 Time derivative of char mass fractions at low temperatures……………………………………..... 65
4.41 Time derivative of char mass fractions at high temperatures…………………………………….... 65
4.42 Specific HF at low temperatures…………………………………………………………………........ 66
4.43 Specific HF at high temperatures……………………………………………………………………... 67
4.44 Percentage Carbon monoxide formation (1%V+1%Cu)………………………………………........ 68
4.45 Model Fit at low Temperatures……………………………………………………………………...... 69
4.46 Model Fit at Higher Temperatures…………………………………………………………………..... 69
5.1 Percentage Carbon monoxide formation…………………………………………………………….... 70
XIV
Abbreviations List
TGA; Thermogravimetric Analysis
DTA; Differential Thermal Analyzer
DSC; Differential Scanning Calorimetry
TG; Thermogravimetric
DTG; Differential Thermogravimetric
UNFCC; United Nations framework convention on climate change
AAEM; Alkali and alkaline earth minerals
ri; rate of species i
dt; change in time
V; Volume
Ea; Activation energy
A; Pre exponential factor
K; rate constant
R; universal gas constant
T; Absolute Temperature
X; Conversion
T; time
mo; initial mass
mt; mass at time t
ma; mass of ash
mT; mass at certain temperature
HF; heat flow
∆H; Enthalpy change
n; order of reaction
XV
1
1 Introduction
The world´s energy demand has increased steadily since the industrial revolution, in the nineteenth
century, till today. Nations have explored many of the available natural resources to meet the day by
day increase in the need for energy in all the various forms. Major natural resources used include
different sources, like coal in its various forms, natural gas, petroleum, hydro-power, biomass etc. As
a matter of fact, most of the natural resources are limited and due to the tremendous volume of
consumption, these reserves are now being consumed at a much faster pace as compared to the
onset of industrial revolution in the nineteenth century and, although new discoveries are constantly
being made and improvements in technology are allowing the exploitation of previously unreachable
resources, most of these energy sources are non-renewable and will eventually be exhausted. The
research is still on the way for the exploration of new resources and significant achievements have
also been made regarding the use of solar, tidal, geothermal, shale gas, nuclear energy etc. along
with the development of various alternative fuels.
Among the natural fossil fuel resources; it has also been predicted that in comparison to the
utilization of petroleum and natural gas reserves; coal will still have reserves that will allow it to be
one of the major sources in the energy mix for the next 100 to 150 years. However, most scenarios
and in particular the ones that are aimed at more stringent control on CO2 emissions, imply a
decrease in coal consumption since coal utilization in traditional ways poses serious environmental
concerns regarding emissions levels, in particular of SOx, NOx and CO2, but also of trace elements in
the form of As, Hg etc. Hence, despite of all efforts being made to address the energy crisis relying
on the development and exploration of new technologies; it has also been understood that an
efficient and environment friendly utilization of the existing fuels and energy resources is also very
important for the sustainable energy use.
Regarding coal utilization, a lot of efforts have been made for the efficient and environment friendly
use of coal which includes:
1. Clean Coal Technologies
2. Coal Liquefaction
3. Coal Gasification
4. Power Generation using IGCC i.e. Integrated Gasification Combined Cycle.
One of the applications of Clean Coal Technology to reduce the greenhouse gas emissions, i.e. CO2
is the co-firing of biomass based fuels in coal fired power plants [1-3], as biomass is the key
renewable energy source and its combustion is CO2 neutral. Many of the biomass fuels used today
come in the form of wood products, dried vegetation, crop residues, and aquatic plants which,
unfortunately, have an energy density much lower than coal. Biomass was the traditional energy
source before the industrial revolution and, although it was replaced by fossil fuels after the
2
industrialization, it has become again one of the most commonly used renewable sources of energy in
the last two decades, second only to hydropower in the generation of electricity.[4]
The use of biomass is likely to become particularly relevant for the production of transportation fuels
since these are all hydrocarbon based and, in order to replace fossil carbon, resorting to biogenic
carbon will be needed.
The production of hydrocarbon fuels from biomass can be made by a variety of methods, in particular
by thermochemical conversions, like pyrolysis and gasification. The use of pyrolysis produces a bio-oil
that requires extensive further refining while gasification leads to the formation of syngas that can be
further converted into hydrocarbons by the Fischer-Tropsch process.
All of these processes can be thermally driven or promoted with catalysts so that the equipment can
be operated at lower temperatures.
In the light of above, the present study is focused on the kinetics of biomass based char i.e. charcoal
thermochemical transformation, both thermal and promoted by metal cation catalyst. The study has
been made using an activated charcoal produced by the de-volatilization or pyrolysis of biomass
material. Understanding of charcoal combustion and gasification kinetics can contribute in the overall
optimization of the design parameters of biomass based co fired or gasification units.
The catalytic effects of transition metals, like Vanadium, in oxidation processes are well known for
different reactions, e.g. the conversion of SO2 into SO3 in sulfuric acid preparation. Hence, catalytic
activity of Vanadium and Copper on the thermochemical conversion of the activated charcoal was
made a subject of the current kinetic study.
Catalyst impregnation was done through incipient wetting technique by Dr. Dr. Sónia Carabineiro.
The kinetics was studied by a combined TG-DSC analysis under isothermal conditions to observe the
kinetics for temperatures ranging from 400 to 800 degree Celsius.
The TGA system was used for data collection, and data interpretation was done through conventional
i.e. TG and DTG kinetics along with utilization of heat flow data obtained by DSC to analyze the
thermicity of the reactions taking place and for the estimation of reaction products. The TG/DTG and
DSC results for catalyst impregnated charcoal have also been compared with the non-impregnated
charcoal. Finally, an effort has been made to propose a suitable model governing the charcoal
conversion at range of temperatures that was investigated.
The experimental results obtained suggest that the reaction of charcoal with oxygen in air is possible
at temperatures as low as 400 °C with the use of 1% Vanadium catalyst impregnation. Without
catalyst, oxidation of char in air only occurred at temperatures in the vicinity of 500 °C. From the
analysis of the heat flow signal it was possible to observe that at low temperature the formation of
carbon monoxide was more significant.
3
For future work, the utilization of other gasifying agents like CO2, H2, pure O2 or steam are envisaged.
Also, the comparison with other transition metal catalysts and their mixtures can be tested.
1.2 Motivation
The main motivation for this study is to further the understanding of the thermochemical conversions
that occur on charcoal both in inert and oxidative conditions, thus including pyrolysis, gasification and
also combustion. All of these reactions are relevant regardless of the end process that we will
consider since, for example inside a gasifier, different regions will spontaneously occur where all of
these types of reactions will occur.
Fig. 1.1 Processes in Gasification
Char oxidation is a complex heterogeneous process which often governs the overall rate of
combustion and gasification. Oxidation rates are partially governed by surface properties of the char
and reactions catalyzed by minerals within the char matrix. Biomass chars have inherently some alkali
and alkaline-earth minerals and their catalytic effects have also been observed during pyrolysis [5].
However, the effect of transition metals on char reactivity during combustion had not gathered much
attention. Hence, the present effort has been made to look for catalytic effects of transition metal i.e.
Vanadium, Copper and their mix on the char reactivity during air gasification at various temperatures.
4
1.3 Aim
The overall aim of this study is to investigate the possible catalytic activity of Vanadium, Copper and
their mixtures for char gasification in air at various temperatures.
1.4 Objectives
The specific objectives of this work include:
To investigate how char conversion is affected by the presence of added catalyst in the form
of 1%V, 1%Cu and their mixtures.
To observe the effect of added catalysts on the DTG curve of the samples.
To analyze the reactions those occur on the samples from DSC data.
To develop the suitable kinetic model to fit the experimental data.
To compare the results with raw charcoal.
1.5 Scope
In order to achieve the above mentioned major objectives, the following was undertaken:
Raw charcoal was analyzed and proximate analysis obtained through TGA.
Char samples impregnated with desired catalyst and different concentrations were obtained
through incipient wetness technique. (Supplied by Dr. Sonia Carabineiro and Prof. Luis Sousa
Lobo).
The gasification and combustion reactivity of the chars were evaluated using simultaneous
Thermogravimetric Analyzer and Differential Scanning Calorimetry (TG/DSC).
Isothermal experiments were performed in the presence of fixed amount of air, utilizing
samples mass less than 10mg.
Char conversion and overall reaction rate with the associated parameters were determined
using the experimental results.
Appropriate kinetic model was evaluated by including rate controlling resistances.
The model was tested against the experimental results and kinetic parameters estimated.
1.6 Thesis Structure
This work has been divided into six chapters. The first chapter is about the introduction of the topic
background and gives the aim, objectives and scope of research.The second chapter deals with the
literature survey and includes the relevant knowledge related to the research methodology of the
current work.
The third chapter discusses the methods utilized for the research purpose. It gives description of the
samples, apparatus, testing procedure etc. The fourth chapter deals in detail with the experimental
results. The approach adopted for discussing the results is as follows: first TG and DTG data have
5
been discussed followed by DSC and kinetic modeling. The same approach has been applied for all
the samples tested.
Conclusions of the study have been discussed in the fifth chapter and the sixth chapter gives the
references to the literature.
6
2 Literature Study
2.1 Biomass and Charcoal
Biomass
The term "biomass" refers to organic matter that has stored energy through the process of
photosynthesis. It exists in many forms as plants may be transferred through the food chain to the
body of animals and their wastes, all of which can be converted for everyday human use through
processes such as combustion, which releases the carbon stored in the form of carbon dioxide.
Also, Biomass refers to any organic materials that is derived from plants or animals. A generally
accepted definition is difficult to find. However, the one used by the United Nations Framework
Convention on Climate Change (UNFCCC, 2005) is:
A non-fossilized and biodegradable organic material originating from plants, animals and micro-
organisms. This shall also include products, by-products, residues and waste from agriculture,
forestry and related industries as well as the non-fossilized and biodegradable organic fractions of
industrial and municipal wastes.
Charcoal a Biochar
Charcoal is a light, black residue, consisting mainly of carbon and any remaining ash, obtained by
removing water and other volatile constituents from animal and vegetation substances. It is usually
produced by slow pyrolysis, the heating of wood or other substances in the absence of oxygen.
2.1.1 Biomass Conversion Processes
Biomass already has a significant contribution for the production of electricity in many countries,
mainly by combustion, as it can be seen in figure 2.1.
However, biomass can also be converted into different gaseous, liquid or solid fuels by a variety of
processes, whic include:
1. ThermoChemical Processes
Gasification
Pyrolysis
Torrefaction
2. Chemical Processes
Direct Liquefaction
Hydrothermal Carbonization
7
3. Biochemical Processes
Hydrolysis
Fermentation
In thermal conversions heat is the dominant driving force that is used to convert the biomass into
another chemical form. The basic alternatives are separated mainly by the extent to which the
chemical reactions involved are allowed to proceed and by the atmosphere in which they are
conducted. Combined heat and Power CHP and co-firing are also the applications of thermal
processes.
Figure 2.1 Role of Biomass in Electricity Generation (Top Five Countries)
2.1.2 Catalytic Effect of Mineral Matter in Biomass
The oxidation of carbonaceous material can be catalysed by naturally occurring inorganic materials
present in the fuels[5]. In coal, inorganic materials reside as minerals, whereas in biomass they are
generally present as salts or organically bound cations. Alkali metals, potassium, and sodium are
active catalysts in reactions with oxygen-containing species. Dispersed alkali metals in biomass
contribute to the high catalytic activity of inorganic materials in biomass. Inorganic matter also affects
pyrolysis, giving char varying morphological characteristics. Potassium and sodium catalyze the
polymerization of volatile matter, increasing the char yield; at the same time they produce solid
materials that deposit on the char pores, blocking them. During subsequent oxidation of the char, the
alkali metal catalyzes this process. Polymerization of volatile matter dominates over the pore-blocking
effect. A high pyrolysis temperature may result in thermal annealing or loss of active sites and thereby
loss of char reactivity [5].
8
2.1.3 Effect of Alkali and Alkaline-Earth Metallic Species (AAEM)
Extensive research has been carried out on the effect of alkali and other metals on the decomposition
behavior of biomass by many researchers of which Shafizadeh, Hsishengeng, Raaveendran have
made significant contributions [7]. All suggested that the inorganic species present in biomass
become one of the influencing factors which determine the behavior of the biomass under thermal
degradation which in turn affects the quality and conversion during pyrolysis, combustion and
gasification. The inorganic constituents of biomass include more than nineteen metals including
alkali metals [9]; some of which act as catalysts that can influence the rate of degradation and yield of
char in pyrolysis. Many of the inorganic components are retained in the char and these can catalyze
the combustion or gasification of the solid residue. It has been shown that the main components
which affect pyrolysis degradation are sodium, potassium, magnesium and silicates [10-11].
Furthermore, in combustion; sulfur, sodium, chlorine and potassium in particular, influence the ash
chemistry and hence also dictate corrosion, slagging and fouling characteristics.
Shafizadeh et al [8] realized that the conversion of organic matter to gas and char varies as the
inorganic content varies, and that higher inorganic contents promote secondary reactions; breaking
down higher molecular compounds to smaller ones [11] and developed the ‘’waterloo model’’ in which
cellulose has two major alternative routes for degradation, depending upon the amount of alkali
metals present. If high levels of alkali metals are present, the degradation mechanism favors
fragmentation i.e. ring scission producing lower molecular weight compounds such as hydroxyl
acetaldehyde, while lower alkali metal contents promotes a de-polymerization mechanism resulting in
higher molecular weight compounds such as beta-d-fructose. However, optimum yield temperatures
vary as the alkali metals in the biomass vary.
2.1.4 Char Conversion Rate
Heterogeneous rate of char conversion depends upon the following factors:
1. Surface Area
2. Surface Accessibility
3. Carbon Active Sites
4. Catalytic Active Sites
5. Gaseous Reactant Concentration
Consequently, the reactivity depends on three chief characteristics of the sample:
1. Chemical Structure
2. Inorganic Constituents
3. Porosity
9
2.1.5 Intrinsic Reaction Rate
Char gasification takes place on the surface of solid char particles, which is generally taken to be the
outer surface area. However, char particles are highly porous, and the surface areas of the inner pore
walls are several orders of magnitude higher than the external surface area. For example, the actual
surface area (BET) Brunauer–Emmett–Teller of an internal pore of a 1-mm-diameter beechwood char
is 660 cm2 while its outer surface area is only 3.14 cm2. Thus, if there is no physical restriction, the
reacting gas can potentially enter the pores and react on their walls, resulting in a high overall char
conversion rate. For this reason, two char particles with the same external surface area (size) may
have widely different reaction rates because of their different internal structure. From a scientific
standpoint, it is wise to express the surface reaction rate on the basis of the actual surface on which
the reaction takes place rather than the external surface area. The rate based on the actual pore wall
surface area is the intrinsic reaction rate; the rate based on the external surface area of the char is the
apparent reaction rate. The former is difficult to measure, so sometimes it is taken as the reactive
surface area determined indirectly from the reaction rate instead of the total pore surface area
measured by the physical adsorption of nitrogen BET area [12].
2.2 Chemical Kinetics
2.2.1 Definition
Chemical kinetics, also known as reaction kinetics, is the study of rates of chemical processes.
Chemical kinetics includes investigations of how different experimental conditions can influence the
speed of a chemical reaction and yield information about the reaction's mechanism and transition
rates, as well as the construction of mathematical models that can describe the characteristics of a
chemical reaction.
2.2.2 History
The general theory of the dependence of reaction rates on concentrations; or the fundamental
equation of chemical kinetics was presented by Cato Guldberg and Peter Waage as a kinetic mass
law in 1864 [13-15].
Van't Hoff studied chemical dynamics and published in 1884 his famous "Etudes de dynamique
chimique"[16]. In 1901 he was awarded by the first Nobel Prize in Chemistry "in recognition of the
extraordinary services he has rendered by the discovery of the laws of chemical dynamics and
osmotic pressure in solutions [17]. After van't Hoff, chemical kinetics deals with the experimental
determination of reaction rates from which rate laws and rate constants are derived. Relatively simple
rate laws exist for zero order reactions (for which reaction rates are independent of concentration),
first order reactions, and second order reactions, and can be derived for others. Elementary reactions
follow the law of mass action, but the rate law of stepwise reactions has to be derived by combining
the rate laws of the various elementary steps, and can become rather complex. In consecutive
10
reactions, the rate-determining step often determines the kinetics. The activation energy for a reaction
is experimentally determined through the Arrhenius equation and the Eyring equation. The main
factors that influence the reaction rate include: the physical state of the reactants, the concentrations
of the reactants, the temperature at which the reaction occurs, and whether or not any catalysts are
present in the reaction.
2.2.3 Reaction Mechanism and Rate
Mechanism:
Molecules or atoms of reactants must collide with each other.
The molecules must have sufficient energy (activation energy) to initiate the reaction.
In some cases orientation of molecules during the collision must also be considered.
Reaction Rate:
The specific rate of consumption or production of any reaction species i, ri, is the rate of change of the
number of molecules of species i with time per unit volume of reaction medium [39]:
𝒓𝒊 = 𝟏
𝑽 𝒅𝒏𝒊
𝒅𝒕 2.1
The rate is negative when i represent a reactant and positive when i represent a product.
Fig.2.2 Reaction Rate variation
11
2.2.3 Factors Affecting Reaction Rate
Nature of the reactants
Depending upon what substances are reacting, the reaction rate varies. Acid/base reactions, the
formation of salts, and ion exchange are fast reactions. When covalent bond formation takes place
between the molecules and when large molecules are formed, the reactions often tend to be very
slow. Nature and strength of bonds in reactant molecules greatly influence the rate of its
transformation into products.
Physical state
The physical state (solid, liquid, or gas) of a reactant is also an important factor of the rate of change.
When reactants are in the same phase, as in aqueous solution, thermal motion brings them into
contact. However, when they are in different phases, the reaction is limited to the interface between
the reactants. Reaction can occur only at their area of contact; in the case of a liquid and a gas, at the
surface of the fluid. Vigorous shaking and stirring may be needed to bring the reaction to completion.
This means that the more finely divided a solid or liquid reactant the greater its surface area per unit
volume and the more contact it with the other reactant, thus the faster the reaction. Further,
homogenous reactions take place faster than heterogeneous reactions.
Concentration
The reactions are due to collisions of reactant species. The frequency with which the molecules or
ions collide depends upon their concentrations. The more crowded the molecules are, the more likely
they are to collide and react with one another. Thus, an increase in the concentrations of the
reactants will usually result in the corresponding increase in the reaction rate, while a decrease in the
concentrations will usually have a reverse effect. For example, combustion that occurs in air (21%
oxygen) will occur more rapidly in pure oxygen.
For elementary reactions, the law of mass action states that the rate is proportional to the
concentrations of the reactants raised to the power of their respective molecularity. Thus for an
elementary irreversible reaction such as the rate equation is:
𝒓 = 𝒌𝑪𝒂𝒑𝑪𝒃𝒒 2.2
The exponents p and q correspond to the order of the reaction in relation to each of the reactants,
and these coincide with the stoichiometric coefficients when the stoichiometric equation truly
represents the mechanism of reaction, i.e., when the reactions are elementary. Thus, for elementary
reactions order and molecularity are the same.
12
Temperature
Temperature usually has a major effect on the rate of a chemical reaction. Molecules at a higher
temperature have more thermal energy. Although collision frequency is greater at higher
temperatures, this alone contributes only a very small proportion to the increase in rate of reaction.
Much more important is the fact that the proportion of reactant molecules with sufficient energy to
react (energy greater than activation energy: E > Ea) is significantly higher and is explained in detail
by the Maxwell–Boltzmann distribution of molecular energies.
The Arrhenius equation relates the specific rate constant to the absolute temperature.
𝒌 = 𝑨 𝒆−𝑬𝒂
𝑹𝑻 2.3
where Ea is called the activation energy and A is the preexponential factor. As seen from equation
2.3, the rate of the reaction can increase very sharply (exponentially) as a function of temperature,
depending on the magnitude of the activation energy E. This equation works well for elementary
reactions, and it also works reasonably well for global reactions over a relatively narrow range of
temperatures in the absence of mass-transfer limitations. The Arrhenius form represents an energy
barrier on the reaction pathway between reactants and products that has to be overcome by the
reactant molecules.
Catalysts
Fig. 2.3 Generic potential energy diagram showing the effect of a catalyst in a hypothetical
endothermic chemical reaction.
13
A catalyst is a substance that accelerates the rate of a chemical reaction but remains mostly
unchanged afterwards. The catalyst increases the rate of the reaction by providing a different reaction
mechanism to occur with lower activation energy. In autocatalysis a reaction product is itself a catalyst
for that reaction leading to positive feedback. Proteins that act as catalysts in biochemical reactions
are called enzymes.
Pressure
Increasing the pressure in a gaseous reaction will increase the number of collisions between
reactants, increasing the rate of reaction. This is because the activity of a gas is directly proportional
to the partial pressure of the gas. This is similar to the effect of increasing the concentration of a
solution.
2.3 Solid State Kinetics
Most solid-state kinetic principles were derived from those for homogenous phases in the past
century.
Solid-state kinetic reactions can be mechanistically classified as:
Nucleation
Geometrical contraction/expansion
Diffusion
2.3.1 Rate Law
Using conversion fraction, rate expression for a first order process can be expressed as:
𝒅𝑿
𝒅 𝒕 = 𝒌 (𝟏 − 𝑿) 2.3
Also, by taking integral and re-arranging:
−𝒍𝒏 ( 𝟏 − 𝑿 ) = 𝒌 𝒕 2.4
Unlike rate laws in homogenous kinetics which usually depend on reaction order (i.e. first, second,
etc.), a rate law for an elementary solid state reaction could depend on factors such as rate of nuclei
formation, interface advance, diffusion, and/or geometrical shape of solid particles.These factors lead
to several decomposition models that do not have a similar counterpart in homogenous kinetics.
Kinetic equations can be generally expressed as:
𝒅𝑿
𝒅 𝒕 = 𝒌 𝒇(𝑿) 2.5
14
By taking integral:
𝒈(𝑿) = 𝒌 𝒕 2.6
Where, f(X) and g(X) are the differential and integral kinetic-model dependent functions respectively,
which can describe the changes in the physical or chemical properties of the sample during
gasification or combustion; t is the time. Assuming that the partial pressure of gasifying agent remains
constant during the process, the reaction rate constant can be expressed using the Arrhenius
equation, as follows:
𝒌 = 𝑨 𝒆−𝑬𝒂
𝑹𝑻 2.7
Where, A is the pre-exponential or frequency factor, Ea is activation energy and T is absolute
temperature and R is the gas constant.
Finally rate equation is:
𝒅𝑿
𝒅𝒕 = 𝑨 𝒆
−𝑬𝒂
𝑹𝑻 𝒇(𝑿) 2.8
Also,
𝒈(𝑿) = 𝑨 𝒆−𝑬𝒂
𝑹𝑻 𝒕 2.9
2.3.2 Gas-Solid Reactant Systems
Multiphase systems correspond to a vast field in which substances react in different phases and may
involve two or three (or more) phases. Gas-Solid reactant system is a multiphase system which
includes two phases i.e. gas and solid. Such systems can be further classified as:
Type 1:
The solid is reacted with gas to form another solid or solids.
Examples:
2ZnS(s) + 3O2(g) → 2 ZnO(s) + 2SO2(g) 2.10
Fe3O4(s) + 4H2(g) → 3Fe(s) + 4H2O(g) 2.11
CaC2(s) + N2(g) → CaCN2(s) + C(s) 2.12
2CaO(s) + 2SO2(g) → 2CaSO4(s) 2.13
15
Type 2:
In this type the products are all gaseous, and the solid shrinks and may eventually disappear.
C(s) + O2(g) → CO2(g) 2.14
C(s) + H2O(g) → CO(g) + H2(g) 2.15
The heterogeneous char-gas chemical reactions takes place in different temperature zones or
regimes which determines; which resistance is rate-controlling. This depends on the particle size,
reactor type, reaction temperature and the reactants.
Figure 2.4 The Arrhenius plot for different temperature regime during heterogeneous chemical
reactions [19]
Figure 2.4 illustrates the Arrhenius plot for different temperature regimes or zones during
heterogeneous oxidation reactions [19]. During char-gas reaction, when the intrinsic reactivity is rate
controlling, the reactant gases diffuse through the porous system to the internal surface of the char
particle; the overall particle size remains constant for a certain period of time, while the density of the
particles is decreased. But if the reaction rate is very fast at high temperature, the gaseous reactants
are consumed rapidly as if they approache the inner surface of the particle [20]. The chemical
structure of the char promotes intrinsic reactivity by providing dislocations, crystalline edges and
heterocyclic centers. The inorganic constituents of the char create further dislocations and promote
16
catalytic activity. The char pore structure or pore network controls the rate of diffusion and the
concentration of reactant gases by fixing the total accessible surface area [22].
Regime I
For reactions at low temperature, the rate of reaction is controlled by the chemical reactivity of the
char. The chemical reaction rate is relatively slow compared to the diffusion rate of the reactant gases
to the internal surface of the particle. The reaction gases diffuse freely into the interior of the porous
char and react uniformly. The particle is converted internally and the particle size might change or
remain constant but the density of the particle decreases. Under this condition, the activation energy
obtained is the true activation energy, and the order of this reaction is also true, since chemical
reaction is the reaction rate determining step. Intrinsic reaction rates at this condition are defined as
the chemical reaction rates, when there is an absence of pore and film diffusion[19-21]. This regime
can be predominant; in fine particles where the diffusion resistance is negligible (very small) and
when the temperature is low with slow kinetic rate [22]
Regime II
At an intermediate temperature the intrinsic rate of reaction and the consumption of the gaseous
reactant is higher than the internal diffusion rate of the reactant gas. The reaction gas does not
penetrate through the pores to the interior of the reacting solid particle, which limits the rate of
reaction. The reaction gas is then consumed in the reaction zone on the surface of the particle,
leaving an unreacted core. The char particle burns internally and externally with decreasing particle
size and particle density. Since internal diffusion limitations occur, the observed activation energy is
about one half of the true activation energy value, while the apparent reaction order is similar to the
true reaction order [21-22].
Regime III
At very high temperatures, the reaction gas does not diffuse through the particle surface. Reaction
occurs at the surface of the char particles only, due to the fast intrinsic reaction. Rate of reaction
depends mostly on the gaseous diffusion through the boundary layer to the particle surface. The
particle diameter decreases as reaction proceeds and the particle density remains more or less
constant with no effect on chemical reactivity or porosity. The activation energy obtained in this zone
will be very small, corresponding to the apparent activation energy for the diffusion process [23].
2.3.3 Intrinsic kinetics of gas-solid reactions
The mechanisms, and hence theoretically derived rate laws, for noncatalytic heterogeneous reactions
involving solids are even less well understood than those for surface catalyzed reactions. This arises
because the solid surface changes as the reaction proceeds, unlike catalytic surfaces which usually
reach a steady-state behavior.
17
Gasification reactions of solids: The reactions of solids with gas-phase reactants to
form gaseous products are generally described in the same manner as are surface catalyzed
reactions. The reaction of carbon with water vapor is an example:
C(s) + H2O(g) → H2(g) + CO(g) 2.16
This reaction is important in such processes as the decoking of catalysts, the manufacture of
activated carbon for adsorption, and the gasification of carbonaceous materials for production of
hydrogen or fuel gas [38].
A two-step mechanism and resulting rate law can be developed as follows. Reactive carbon sites, C*
(total number NC.), are assumed to exist on the surface of the solid. These can be oxidized reversibly
by water vapor:
C∗ + H2O k− 1 ←
k1 → C∗O + H2 2.17
where C*O is an oxidized carbon site. The oxidized site can then “decompose” to produce CO(g):
C*O 𝑘2 → CO(g) 2.18
In addition to CO(g) formation, step (2) exposes a variable number, n, of previously inactive carbon
atoms, C, thus producing C* to continue the reaction. The average value of n is close to unity, so that
NC* varies slowly as reaction proceeds.
If elementary rate laws are assumed for each step, and if NC* is essentially constant over a short time,
a (pseudo-) steady-state rate law can be developed:
𝑟 = 𝑁𝐶∗
𝑁𝐶
𝑘1𝑘2𝑐𝐻2𝑂
𝑘1𝑐𝐻2𝑂+ 𝑘−1𝑐𝐻2+ 𝑘2 2.19
The above equation is similar to Langmuir-Hinshelwood kinetics. The rate is expressed on the basis
of the instantaneous number of solid carbon atoms, Nc. The rate r (measured at one gas
composition) typically goes through a maximum as the carbon is converted. This is the result of a
maximum in the intrinsic activity (related to the fraction of reactive carbon atoms, NC*/NC) because of
both a change in NC* and a decrease in NC.
Since both NC* and NC change as the reaction proceeds, r can be expressed as a function of
fractional conversion of carbon (XC) or of time (t).
18
2.3.4 Models and Mechanisms in Solid State kinetics
A model is a theoretical, mathematical description of what occurs experimentally. In solid-state
reactions, a model can describe a particular reaction type and translate that mathematically into a rate
equation. Many models have been proposed in solid-state kinetics and these models have been
developed based on certain mechanistic assumptions. Other models are more empirically based and
their mathematics facilitates data analysis with less mechanistic meaning. Therefore, different rate
expressions are produced from these models.
Commonly Employed Models
Three models are commonly implemented to interpret the experimental result which are as follows:
1. VM (Volume Model)
2. GM (Grain Model)
3. RPM (Random Pore Model)
These models have a theoretical basis and involve fewer parameters and give different formulations
of the term f (X).
Volume Model
The VM assumes that a homogeneous reaction occurs throughout the char bed and that it results in a
linear decrease in the reaction surface area with conversion [24]. The overall reaction rate is given by:
𝒅𝑿
𝒅𝒕= 𝒌𝑽𝑴 (𝟏 − 𝑿) 2.20
Grain Model
The GM considers that the gasifying agents react on the surface of the non-porous grains or in pore
surfaces within the solid [25]. According to different assumptions, the reaction rates in the regime of
chemical control can be expressed as:
𝒅𝑿
𝒅𝒕= 𝒌𝑮𝑴 (𝟏 − 𝑿)𝟐/𝟑 2.21
Random Pore Model
The RPM considers the overlapping of pore surfaces, which results in the reduction of surface area
available for the reaction [26]. The general rate equation for this model is:
𝒅𝑿
𝒅𝒕= 𝒌𝑹𝑷𝑴 (𝟏 − 𝑿)√𝟏 −𝝍 𝒍𝒏 (𝟏 − 𝑿) 2.22
This model can predict a maximum for the reactivity during the reaction, as it considers the competing
effects of pore growth during the initial stages of gasification, and the destruction of the pores due to
19
the coalescence of neighboring pores during the reaction. The RPM contains two parameters, ψ ,
which is related to the initial pore structure ofthe char sample (X = 0) and the reaction rate constant, k.
𝝍 = 𝟒 𝝅 𝑳𝒐 (𝟏−𝓔𝒐)
𝑺𝒐𝟐 2.23
where S0, L0, and ℰ0 represent the pore surface area, pore length, and solid porosity, respectively.
2.3.5 Data collection and Interpretation
Experimentally, solid-state kinetics can be studied either isothermally or nonisothermally. Many
mathematical methods have been developed to interpret experimental data for both heating protcols.
These methods generally fall into one of two categories:
Model-fitting
Model-free.
Model Fitting Methods:
For these methods, different models are fit to the data and the model giving the best statistical fit is
chosen as the model of choice from which the activation energy (Ea) and frequency factor (A) are
calculated. Model fitting approach is often use for pyrolysis reaction description. As a result a one set
of kinetic parameters is estimated for entire range of temperatures and extension of the reaction.
These are called apparent kinetic parameters.
Isothermal model-fitting method:
This method is identical to that in homogenous kinetics. It involves two fits: The first, determines the
rate constant (k) of the model that best fits the data for single temperature experiments while the
second determines specific kinetic parameters such as the activation energy (Ea) and frequency
factor (A) using the Arrhenius equation using data from different temperatures.
Nonisothermal model-fitting method:
There are many model fitting methods that extract the complete set of kinetic parameters known as
the kinetic triplet (A, Ea and model) from nonisothermal data. These methods were used extensively
earlier in solid-state kinetic analysis and they continue to be developed. These methods have been
critically evaluated and it’s been shown that the sole use of these methods is not recommended
because:
They assume a constant kinetic triplet (A, Ea and model).
They involve fitting three parameters (A, Ea and model) which are determined from a single
run (for example a single heating rate) which is not always sufficient to determine reaction
kinetics.
20
Examples:
Direct Differential Method, Coats Red fern Method
Model-free/isoconversional methods
• Model-free methods calculate the reaction activation energy (Ea) without model assumptions, which
is usually done by grouping terms such as the frequency factor (A) and model into the intercept of a
linear equation and using the slope of that equation to calculate the activation energy (Ea).The
frequency factor (A) can be calculated from the intercept of the linear equation but requires modelistic
assumptions for such a determination. Therefore, model-free methods usually report only activation
energies.
• Isoconversional methods are model-free methods that evaluate kinetic parameters, namely the
activation energy (Ea) at progressive conversion values (X). These methods require several kinetic
curves to perform the analysis and have therefore been called, “multi-curve” methods.
• Calculations from several curves at different heating rates are performed on the same value of
conversion (X), thus, the name isoconversional. As a result, these methods calculate the
activation energy for each conversion point (Ea, X), resulting in an isoconversional plot (Ea vs. X).
•The terms, “model-free” and “isoconversional” are sometimes used interchangeably; however, not all
model-free methods are isoconversional.
Isoconversional approaches can be used to analyze both isothermal and nonisothermal data, as
described below:
Isothermal
Standard
Friedman
Non Isothermal
Kissinger
Ozawa, Flyn-Wall (Linear Doyle Approximation; Less accurate)
Vyazovkin (Non Linear Senum yang Approximation; More accurate)
Historically, model-fitting methods were widely used because of their ability to directly determine the
kinetic triplet (i.e., frequency factor [A], activation energy [Ea] and model). However, these methods
suffer from several problems among which is their inability to uniquely determine the reaction model.
This has led to the decline of these methods in favor of isoconversional (model-free) methods that
evaluate kinetics without model assumptions. However, isoconversional methods do not compute a
21
frequency factor nor determine a reaction model which are needed for a complete and accurate
kinetic analysis. A new approach has to be proposed that combines the power of isoconversional
methods with model-fitting methods.
2.3.6 Controversies in Solid State kinetics.
Solid-state kinetics was developed from reaction kinetics in homogenous systems (i.e. gases and
liquids). The Arrhenius equation relates the rate constant of a simple one-step reaction to the
temperature through the activation energy (Ea) and preexponential factor (A).
It has been generally assumed that activation energy (Ea) and frequency factor (A) remain constant,
however, it’s been shown in solid-state reactions these kinetic parameters may vary with the reaction
progress (). This variation can be detected by isoconversional methods. While this variation appears
to be in conflict with basic chemical kinetic principles, in reality, it may not be.
2.4 Catalysis
Catalysis is the increase in the rate of a chemical reaction due to the participation of an additional
substance called a catalyst. With a catalyst, reactions occur faster and require less activation energy.
Because catalysts are not expected to be consumed in the catalyzed reaction, they can continue to
catalyze the reaction of further quantities of reactant. Often only tiny amounts are required.
2.4.1 Heterogenous Catalysis
In chemistry, heterogeneous catalysis refers to the form of catalysis where the phase of the catalyst
differs from that of the reactants. Phase here refers not only to solid, liquid or gas, but also immiscible
liquids, e.g. oil and water. The great majority of practical heterogeneous catalysts are solids and the
great majority of reactants are gases or liquids [27]. Heterogeneous catalysis is of paramount
importance in many areas of the chemical and energy industries.
Adsorption
Adsorption is commonly an essential first step in heterogeneous catalysis. Adsorption is when a
molecule in the gas phase or in solution binds to atoms on the solid or liquid surface. The molecule
that is binding is called the adsorbate, and the surface to which it binds is the adsorbent. The process
of the adsorbate binding to the adsorbent is called adsorption. The reverse of this process (the
adsorbate splitting from adsorbent) is called desorption.
22
Types of adsorption
Two types of adsorption are recognized in heterogeneous catalysis, although many processes fall into
an ambiguous range between the two extremes. In the first type, physisorption; induces only small
changes to the electronic structure of the adsorbate. Typical energies for physisorption are from 2 to
10kcal/mol. The second type is chemisorption, in which the adsorbate is strongly perturbed, often with
bond-breaking and the formation of chemical bonds between the adsorbate and the adsorbent.
Energies for typical chemisorption range from 15 to 100 kcal/mol.
For physisorption, the adsorbate is attracted to the surface atoms by van der wall’s forces. A
mathematical model for physisorption was developed by London to predict the energies of basic
physisorption of non-polar molecules. The analysis of physisorption for polar or ionic species is more
complex.
Chemisorption results in the sharing of electrons between the adsorbate and the adsorbent.
Chemisorption is traditionally described by the Lennard-Jones potential, which considers various
cases, two of which are.
Molecular adsorption, where the adsorbate remains intact. An example is alkene binding by
platinum.
In dissociation adsorption, one or more bonds break concomitantly with adsorption. In this
case the barrier to dissociation affects the rate of adsorption. An example of this the binding
of H2, where the H-H bond is broken upon adsorption [27] by hydrogen spillover.
Surface Reactions
With catalyst supports, the reaction that occurs often occurs on the surface of either the catalyst or
the support. In terms of surface reactions there are three mechanisms.
Langmuir-Hinshelwood mechanism. The two molecules A and B both adsorb to the surface.
While adsorbed to the surface, the A and B "meet," bond, and then the new molecule A-B
desorbs.
Rideal-Eley mechanism. One of the two molecules, A, adsorbs to the surface. The second
molecule, B, meets A on the surface, having never adsorbed to the surface, and they react
and bind. Then the newly formed A-B desorbs.
Precursor mechanism. One of the two molecules, A, is adsorbed on the surface. The second
molecule, B, collides with the surface, forming a mobile precursor state. The molecule B then
collides with A on the surface, they react, bind and the new molecule desorbs.
Any surface reaction can be described as following one of these mechanisms, or some combination
of these mechanisms. In addition, all of these above mechanisms can occur in reverse. In general,
the pathway for a reaction on a surface is as follows. First the reactants adsorb onto the surface.
23
Through a series of bonds being formed and being broken, adsorbed intermediates are produced and
destroyed. Then the final product(s) is produced and it desorbs from the solid. Most metal surface
reaction occurs by chain propagation [27].
Fig. 2.5 Mechanisms for Heterogeneous Catalysis
2.4.2 Applied Catalysis
Oxidation catalysis is conducted by both heterogeneous catalysis and homogeneous catalysis. In the
heterogeneous processes, gaseous substrate and oxygen/air are passed over solid catalysts. Typical
catalysts are platinum, redox-active oxides of iron, vanadium, and molybdenum. In many cases,
catalysts are modified with a host of additives or promoters that enhance rates or selectivity’s.
Catalytic Oxidation
Catalytic oxidations are processes that oxidize compounds using catalysts. Common applications
involve oxidation of organic compounds by the oxygen in air. Such processes are conducted on a
24
large scale for the remediation of pollutants, production of valuable chemicals, and the production of
energy.
Illustrative catalytic oxidation processes are presented in the table:
Table. 2.1 Examples of Catalytic Oxidation processes
Substrate Process
Catalyst
(homogeneous or
heterogeneous
Product
Application
sulfur dioxide
contact
process
vanadium pentoxide
(heterogeneous) sulfuric acid fertilizer production
ammonia
Ostwald
process
platinum
(heterogeneous) nitric acid basic chemicals, TNT
hydrogen
sulfide
Claus process
vanadium pentoxide
(heterogeneous) sulfur
remediation of byproduct
of
oil refinery
methane,
ammonia
Andrussow
process
platinum
(heterogeneous)
hydrogen
cyanide
basic chemicals, gold
mining extractant
ethylene epoxidation
mixed Ag oxides
(heterogeneous) ethylene oxide
basic chemicals,
surfactants
cyclohexane K-A process Co and Mn salts
(homogeneous)
cyclohexanol
cyclohexanone nylon precursor
ethylene
Wacker
process
Pd and Cu salts
(homogeneous) acetaldehyde basic chemicals
2.4.3 Heterogenous Catalysts
Heterogeneous catalysts act in a different phase than the reactants. Most heterogeneous catalysts
are solids that act on substrates in a liquid or gaseous reaction mixture. Diverse mechanisms for
reactions on surfaces are known, depending on how the adsorption takes place (Langmuir-
Hinshelwood, Eley-Rideal, and Mars-van Krevelen) [28]. The total surface area of solid has an
important effect on the reaction rate. The smaller the catalyst particle size, the larger the surface area
for a given mass of particles.
25
Active Sites
A heterogeneous catalyst has active sites, which are the atoms or crystal faces where the reaction
actually occurs. Depending on the mechanism, the active site may be either a planar exposed metal
surface, a crystal edge with imperfect metal valence or a complicated combination of the two. Thus,
not only most of the volume, but also most of the surface of a heterogeneous catalyst may be
catalytically inactive. Finding out the nature of the active site requires technically challenging
research. Thus, empirical research for finding out new metal combinations for catalysis continues.
For example, in the Haber process, finely divided iron serves as a catalyst for the synthesis of
ammonia from nitrogen and hydrogen. The reacting gases adsorb onto active sites on the iron
particles. Once physically adsorbed, the reagents undergo chemisorption that results in dissociation
into adsorbed atomic species, and new bonds between the resulting fragments form in part due to
their close proximity. In this way the particularly strong triple bond in nitrogen is broken, which would
be extremely uncommon in the gas phase due to its high activation energy. Thus, the activation
energy of the overall reaction is lowered, and the rate of reaction increases. Another place where a
heterogeneous catalyst is applied is in the oxidation of sulfur dioxide on Vanadium(V) oxide for the
production of sulfuric acid.
Catalyst Support
Heterogeneous catalysts are typically "supported," which means that the catalyst is dispersed on a
second material that enhances the effectiveness or minimizes their cost. Supports prevent or reduce
agglomeration and sintering of the small catalyst particles, exposing more surface area, thus catalysts
have a higher specific activity (per gram) on a support. Sometimes the support is merely a surface on
which the catalyst is spread to increase the surface area. More often, the support and the catalyst
interact, affecting the catalytic reaction. Supports are porous materials with a high surface area, most
commonly alumina, zeolites or various kinds of activated carbon. Specialized supports include silicon
dioxide, titanium dioxide, calcium carbonate, and barium sulfate.
Catalyst Preparation
Incipient wetness impregnation (IW or IWI), also called capillary impregnation or dry impregnation, is
a commonly used technique for the synthesis of supported heterogeneous catalysts. Typically, the
active metal precursor is dissolved in an aqueous or organic solution. Then the metal-containing
solution is added to a catalyst support containing the same pore volume as the volume of the solution
that was added. Capillary action draws the solution into the pores. Solution added in excess of the
support pore volume causes the solution transport to change from a capillary action process to a
diffusion process, which is much slower. The catalyst can then be dried and calcined to drive off the
volatile components within the solution, depositing the metal on the catalyst surface. The maximum
loading is limited by the solubility of the precursor in the solution. The concentration profile of the
26
impregnated compound depends on the mass transfer conditions within the pores during
impregnation and drying [29,30,31]
2.5 Thermal Analysis
There are several types of techniques that involve thermal analysis and some of them are presented
in table 2.2.
Table 2.2 Thermal Analysis Techniques
2.5.1 TGA
Thermogravimetric Analysis (TGA) measures the amount and rate of change in the weight of a
material as a function of temperature or time in a controlled atmosphere. Measurements are used
primarily to determinethe composition of materials and to predict their thermal stability. The technique
can characterize materials that exhibit weight loss or gain due to decomposition, oxidation, or
dehydration.
27
What is it for?
• Composition of Multicomponent Systems
• Thermal Stability of Materials
• Oxidative Stability of Materials
• Estimated Lifetime of a Product
• Decomposition Kinetics of Materials
•The Effect of Reactive or Corrosive Atmospheres on Materials
• Moisture and Volatiles Content of Materials
Fig. 2.6 Block Diagram of Thermobalance
2.5.2 DTA/DSC
Differential thermal analysis or DTA is the tecnique in which heat flow to the sample and reference
remains the same rather than the temperature. When the sample and reference are heated
identically, phase changes and other thermal processes cause a difference in temperature between
the sample and reference.
Differential scanning calorimetry or DSC is a thermoanalytical technique in which the difference in
the amount of heat required to increase the temperature of a sample and reference is measured as a
function of temperature. Both the sample and reference are maintained at nearly the same
temperature throughout the experiment.
28
DSC, in theory allows for the measurement of a change in enthalpy. DSC measures the energy
required to keep both the reference and the sample at the same temperature whereas DTA measures
the difference in temperature between the sample and the reference when they are both put under the
same heat.
Fig. 2.7 Typical DSC Curve
Detection of phase transitions in DSC
The basic principle underlying this technique is that when the sample undergoes a physical
transformation such as phase transitions, more or less heat will need to flow to it than the reference
to maintain both at the same temperature. Whether less or more heat must flow to the sample
depends on whether the process is exothermic or endothermic. For example, as a solid sample melts
to a liquid, it will require more heat flowing to the sample to increase its temperature at the same
rate as the reference. This is due to the absorption of heat by the sample as it undergoes the
endothermic phase transition from solid to liquid. Likewise, as the sample undergoes exothermic
pr°Cesses (such as crystallization) less heat is required to raise the sample temperature. By observing
29
the difference in heat flow between the sample and reference, differential scanning calorimeters are
able to measure the amount of heat absorbed or released during such transitions.
Fig. 2.8 An idealized DSC curve showing the shapes associated with particular phase
transitions.
2.5.3 Thermal analytical methods vs. kinetics
Solid-state kinetics can be studied with thermal analytical methods by measuring a sample property
as it is heated or held at a constant temperature. If a reaction involves weight loss, then weight is
followed and the kinetics are usually studied by thermogravimetry(TGA). Heat (evolved or consumed)
is another measurable property that is used for kinetic evaluation using differential scanning
calorimetry (DSC) or differentia lthermal analysis (DTA).
Weight loss or heat flow data are converted to a normalized form called conversion fraction (X).
The conversion fraction ranges from 0 and 1 and is a measure of reaction progress as a function of
time or temperature.For isothermal thermogravimetric analysis, the conversion fraction (X) at any time
is:
30
𝑿 = 𝒎𝒐−𝒎𝒕
𝒎𝒐−𝒎∞ 2.24
Where, mo is the initial sample weight, mt is the sample weight at time, t, and m is the final sample
weight i.e. ash content.
Non-isothermally, the conversion fraction (X) at any temperature is:
𝑿 = 𝒎𝒐−𝒎𝑻
𝒎𝒐−𝒎∞ 2.25
The analysis of this fractional conversion as a function of time or temperature will be the basis for
the linetic analysis.
31
3 Research Methodologies
3.1 Materials
1. Charcoal (Raw)
2. Charcoal (1% V Impregnated)
3. Charcoal (1% Cu Impregnated)
4. Charcoal (1% V + 1% Cu Impregnated)
Table 3.1 Proximate Analysis
Parameter Percentage
Moisture 10
Volatile Matter 3.2
Fixed Carbon 83
Ash 4.3
3.2 Sample Preparation
Incipient wetness technique was used to impregnate the desired catalyst concentration to the raw
charcoal sample.
3.3 Equipment
A TGA measures the mass change of a sample in relation to temperature as it is subjected to a
specified heating programme in a controlled atmosphere. The equipment consists of a sample pan
loaded on to a highly sensitive mass balance, a furnace and a thermocouple.
Perkin-Elmer series STA 600
The Perkin-Elmer TGA contains a standard furnace, which can be heated 1000° C. However, the
highest temperature that samples were heated to was to avoid burnout of the furnace. The heating
rates used ranged from 20 K/min to 100 K/min. A cooling unit was connected to the TGA in order to
cool the equipment between runs. Aluminium Oxide crucibles were used for loading the samples. A
picture of the equipment is presented in Figure 3.1. The TGA was linked to a computer loaded with
Pyris software, which was used to control the equipment and display the results of the experiments.
32
Figure 3.1: Apparatus
3.4 Procedure
Initially an empty alumina crucible was placed in the TGA holder. The crucible mass was set to zero.
The crucible was then removed from the TGA and the sample was placed inside it. Less than 10mg
samples was used for each run. The crucible and sample were first weighed on an analytical balance.
After that it was gently positioned into the apparatus, taking precaution not to place too much
pressure on the balance stem. Once the sample mass reading had stabilised, the weight percent was
set to 100 %. The details of the experimental runs were set up with the software as follows:
Program
Table 3.2 Procedure adopted for kinetics study
Step Description Stage
1 Flow of Pure Air (20ml/min) Initial
2 Ramp to Tisothermal with specified heat rate Heating
3 Hold at Tisothermal for specified time Isothermal
4 Cool to room temperature Cooling
33
3.5 Calculations
TG/DTG
Thermogravimetric analysis for char gasification takes the general form of mass loss over time for a
specified temperature profile. Fractional mass loss curves are developed utilizing the mass loss data
as follows:
𝐅𝐫𝐚𝐜𝐭𝐢𝐨𝐧𝐚𝐥 𝐦𝐚𝐬𝐬 𝐥𝐨𝐬𝐬 = 𝒎𝒕−𝒎𝒂
𝒎𝒐−𝒎𝒂 3.1
The isothermal segments were exported and normalized from 0% to 100% char conversion according
to equation (3.2).
𝑿 = 𝒎𝒐−𝒎𝒕
𝒎𝒐−𝒎𝒂 3.2
Where, mo denotes the sample mass at the start of gasification, mt the sample mass at time t and ma
the mass of ash remained after complete gasification.
The experimental and processed conversion results for charcoal impregnated with 1% Vanadium are
presented in the next section. Fractional char conversions have been plotted against time. At low
temperatures, char conversion under air atmosphere, with accompanying mass loss took place in two
different regimes i.e.
a. Exothermic mass loss
b. Endothermic mass loss
Both regions have been marked separately, where applicable and should correspond, respectively, to
oxidation (gasification) and pyrolysis. Separation of gasification and pyrolysis regions were based on
heat flow data.
Exothermic regions were considered as gasification in which CO, CO2 or a mixture of the two, was
expected since the gasifying agent is air and the sample being tested is charcoal.
DSC
Specific heat flow data was determined using the following relation:
𝐒𝐩𝐞𝐜𝐢𝐟𝐢𝐜 𝐇𝐅 = 𝐇𝐅𝒕
𝒎𝒐− 𝒎∞ 3.3
34
Where, HFt is the heat flow at time t.
Thermodynamic data has been used to calculate the percentage of carbon monoxide and carbon
dioxide in the product gas mix during the gasification stage. The standard enthalpy of formation for
char gasification products in the presence of air are as follows:
∆H for CO2 = - 393.2 KJ/mol at 298 oK
∆H for CO = - 110.2 KJ/mol at 298 oK
Fig. 3.2 Enthalpy of formation for CO and CO2
These values were corrected for the working temperature ranges i.e. 400 – 800 °C by using Shomate
equation which is explained below [32].
Ho = A*t + B*t2/2 + C*t3/3 +D*t4/4 – E/t + F Eq: 3.4
Where,
H° = standard enthalpy (kJ/mol)
t = temperature (K) / 1000.
Table 3.2 displays the constants for the shomate equation for [CO] while Table 3.3 presents the heat
of formation computed from these equations for CO and CO2 as a function of temperature.
35
Table 3.3 Shomate Equation Constants
Temperature (K) 298. - 1200. 1200. - 6000.
A 24.99735 58.16639
B 55.18696 2.720074
C -33.69137 -0.492289
D 7.948387 0.038844
E -0.136638 -6.447293
F -403.6075 -425.9186
G 228.2431 263.6125
H -393.5224 -393.5224
Reference [33-34] Chase, 1998 Chase, 1998
For a particular isothermal experiment; the total energy released with a given mass was calculated by
integrating the heat flow data in the exothermic region. Energy released per unit mass was then
calculated in KJ/g.
Table 3.4 Heat of Formation
Temperature (°C) Heat of Formation (Theoretical) KJ/g
CO CO2
400 -8.278 -31.151
450 -8.278 -31.151
500 -8.149 -30.883
550 -8.017 -30.603
600 -7.750 -30.007
650 -7.615 -29.692
700 -7.478 29.365
750 -7.340 -29.026
800 -7.200 -28.675
850 -7.06 -28.314
36
It was assumed that the gas mixture resulting from the gasification process was a mixture of CO and
CO2 and thus, the fraction of CO produced could be estimated by comparing the observed heat
released with the enthalpies of formation of the two components. CO contribution in the total energy
released was calculated using the enthalpy of formation and the experimental energy released (Y)
data as follows:
For ‘Y’ KJ/g of energy released:
𝐘 = 𝐗 . ∆ 𝐇 𝐂𝐎 + (𝟏 − 𝐗). ∆ 𝐇 𝐂𝐎𝟐 3.5
Applying the algebraic manipulations, we got fraction of carbon converted to CO as:
𝐗 = ∆𝐇 𝐂𝐎𝟐−𝐘
∆𝐇 𝐂𝐎𝟐− ∆𝐇 𝐂𝐎 3.6
3.6 Kinetic Modelling
Surface Catalysis: Intrinsic Kinetics
Surface catalysis is involved in a large majority of industrial catalytic reactions. The rate
laws developed are based on the following assumptions [38]:
(1) The surface of the catalyst contains a fixed number of sites.
(2) All the catalytic sites are identical.
(3) The reactivities of these sites depend only on temperature. They do not depend on
the nature or amounts of other materials present on the surface during the reaction.
These assumptions are the basis of the simplest rational explanation of surface catalytic kinetics and
models for it. The preeminent of these, formulated by Langmuir and Hinshelwood, makes the further
assumption that for an overall (gas-phase) reaction, for example, A(g) + . . . + product(s), the rate
determining step is a surface reaction involving adsorbed species, such as A ● s. Despite the fact that
reality is known to be more complex, the resulting rate expressions find wide use in the chemical
industry, because they exhibit many of the commonly observed features of surface-catalyzed
reactions.
Surface Reaction Steps:
Central to surface catalysis are reaction steps involving one, or more than one, surface bound
(adsorbed) intermediate species. In case of unimolecular surface reaction, we have:
A ● s → B ● s 3.7
37
Where, A ● s is a surface bound species involving A and site s. The rate of this reaction is given by:
(-rA) = k ᶿA 3.8
Where, ᶿA is the fraction of the surface covered by adsorbed species A.
Langmuir-Hinshelwood (LH) Kinetics
By combining surface-reaction rate laws with the Langmiur expressions for surface coverages,
Langmuir-Hinshelwood (LH) rate laws for surface-catalysed reactions are obtained, although we focus
on the intrinsic kinetics of the surface-catalysed reaction, the LH model should be set in the context of
a broader kinetics scheme to appreciate the significance for this.
A kinetic scheme for an overall reaction expressed as:
A(g) → B(g) 3.9
Where A is a gas-phase reactant and B a gas-phase product, is as follows:
A(g) 𝑘Ag → A (surface vicinity); mass transfer (fast) Step:1
A (surface vicinity) + s 𝑘𝑑A ←
𝑘𝑎A → A ● s; adsorption-desorption (fast) Step:2
A ● s 𝑘 → B (surface vicinity) + s; surface reaction (slow, rds) Step:3
B (surface vicintiy) 𝑘Bg → B(g); mass transfer (fast) Step:4
Here A(g) and B(g) denote reactant and product in the bulk gas at concentrations CA and CB,
respectively; kAg and kBg are mass-transfer coefficients, s is an adsorption site, and A ● s is a surface
reaction intermediate. In this scheme, it is assumed that B is not adsorbed. In focusing on step (3) as
the rate-determining step, we assume that kAg and kBg are relatively large, and step (2) represents
adsorption-desorption equilibrium.
Following Langmuir isotherm for competing species:
θA = KA 𝐶𝐴
1+ KACA+ KBCB 3.10
38
For the overall reaction A → B, if the rate determining step is the unimolecular surface reaction by
eq(3.7), then the rate of reaction is obtained by using eq(3.10) for θA in eq(3.8) to result in:
(−𝑟A) = k KA 𝐶A
1+ KACA+ KBCB 3.11
Fig. 3.3 Langmiur-Hinshelwood Model
Above explained L-H type kinetics has been made the basis in order to develop an appropriate kinetic
model for the explanation of the experimental mass loss data.
As discussed in section 2.3.1 for reactions at low temperatures, the rate of reaction is controlled by
the chemical reactivity of the char and hence chemical reaction rate is relatively slow compared to the
diffusion rate of the reactant gases to the internal surface of the particles. Under these conditions the
rate of chemical reaction can be expressed as:
𝑟𝑐 = 𝑘 𝑃𝑂2(𝑆) 3.12
Where, 𝑃𝑂2(𝑆) is the partial pressure of oxygen at the reactant surface when there is no diffusion
limitation. At low temperatures this partial pressure of oxygen will be the same as in the bulk gas
phase 𝑃𝑂2(𝑆)=𝑃𝑂2(𝑔).
39
It is noteworthy that all the experiments have been carried-out at a constant air flow, and so the bulk
oxygen concentration can be considered as being the same at all times and the influence of oxygen in
the kinetic data will only be relevant in the context of diffusion limitations of the oxidant.
Apart from the order of the reaction in relation to the oxygen, one all has to consider the apparent
order in relation to the carbonaceous material itself. In this respect, it is clear from the experimental
data that there is significan segment in the beginning of the reaction where there is a linear trend in
the mass loss, indicating that the rate of reaction is not directly proportional to the amount of carbon
material. This can be explained according to several mechanisms but there are two main
interpretaations that can be put forward. On one hand the reaction can occur mostly catalyzed, either
by the added catalyst of by the inorganic contaminants present beforehand and the reaction proceeds
by a CASA (contact active surface area) mechanism. In this case the reaction rate will depend only on
the amount of catalyst present until the carbon amount is relatively low. On the other hand, if the
combustion occurs in the microporous surface, as discussed above, the reaction will be mostly
dependent on the outer surface of the particle and not directly related to the inner surface and this will
also reduced the dependence on the amount of the carbon present in the sample.
In order to describe this type of relationship the dependence on the weight of carbon material was
introduced in the kinetic rate expression in the following form:
𝑟𝑐 = 𝑘𝑤
𝑤0+𝑤 𝑃𝑂2(𝑆) 3.12a
Let us now consider what will happen as the temperature increases. With the increase in temperature,
the chemical reaction rate and hence the consumption of the gaseous reactant will be higher than the
diffusion rate of the reactant gas. The reactant gas will not be able to penetrate through the pores to
the interior of the reacting solid particle. This diffusion related phenomenon will start limiting the rate
of reaction. This can happen only due to internal diffusuion limitations, thus reducing the apparent
activation energy as explained above, or even due to external diffusion limitations, where the process
will be fully controlled by the diffusion of the gaas from the bulk of the gas to the surface of the
material. To understand this effect another reaction path is considered which is as:
𝑟𝑑 = 𝑘𝑔(𝑃𝑂2(𝑔) − 𝑃𝑂2(𝑆)) 3.13
At very high chemical reaction rate, the 𝑃𝑂2(𝑆) = 0 at extreme diffusion limitation.
Furthermore, at quasi steady-state;
𝑟 = 𝑟𝐶 = 𝑟𝑑 3.14
By equating equations 3.12 and 3.13 and following algebraic manipulations; 𝑃𝑂2(𝑆) was evaluated as:
40
𝑃𝑂2(𝑆) = 𝑘𝑔𝑃𝑂2(𝑔)
𝑘+ 𝑘𝑔 3.15
Utilizing eq 3.15; equation 3.12 becomes:
𝑟 =𝑘𝑘𝑔𝑃𝑂2(𝑔)
𝑘+ 𝑘𝑔 3.16
Equation 3.16 yielded [k/] as 𝑘 𝑘𝑔
𝑘+ 𝑘𝑔 .
It can be observed that overall rate constant is governed by individual k and kg. To evaluate the k/, the
concept of resistances was used. By this, individual rates were added by converting them to the
reciprocal form as:
1
𝑘/=
1
𝑘+
1
𝑘𝑔 3.17
Curve Fitting
In order to fit the predicted model curve following the experimental data; kinetic parameters were first
tried as per theoretical bases to get the approaching fit. After getting close the near close fit was
obtained with the help of SOLVER function offered by Excel. To run the solver least square method
was used for input data.
41
4 Results and Discussion
4.1. Charcoal (non-impregnated)
4.1.1 TG/DTG
Fig. 4.1 Fractional mass loss at various temperature
Fig. 4.2 Conversion at various temperatures
In the above two figures fractional mass loss and conversion at various temperatures have been
plotted against time. It was observed that the carbon conversion started to occur at 500 °C with a very
slow pace. It increases with a faster pace at 550 °C till 600 °C. After that the mass loss occurred in a
42
more uniform way till 850 °C. These trends showed that the combustion of charcoal follows complete
combustion pattern from 600 °C. At low temperatures till 500 °C the combustion is partial and very
slow.
Fig. 4.3 Time Derivative of char mass fractions at low temperature
Fig. 4.4 Time Derivative of char mass fractions at high temperature
In the above two figures time derivative of char mass fractions have been plotted as a function of
time. It was observed that the reactivity of char increases with increasing temperature. The char
oxidation reaction begins at around 500 °C and increases rapidly with time till 650 °C. After reaching a
peak value, the reactivity starts to decrease due to the combustion of the less reactive portion of the
char. Char reactivity was more or less same for temperatures 700 to 850 °C.
43
4.1.2 DSC
Fig. 4.5 Specific HF at low temperatures
In the above figure, heat flow data has been plotted as a function of time for the three lowest
temperatures. It was observed that charcoal showed exothermic reaction at around 500 °C. However,
it shifted to endothermic in the midway indicating that the mass loss started to occur via pyrolysis
instead of combustion. The same pattern occurred also, although in a more limited way at 550 °C and
it disappeared at 600 °C. This trend of partial combustion and pyrolysis may be due to combination of
various factors e.g. low reaction temperature, low rate of reaction, inhibiting effect of mineral
impurities and added catalyst etc.
Table: 4.1 Gas mix composition at low temperatures
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
500 -8.017 -30.603 -8.94 96
550 -7.885 -30.311 -16.61 61
600 -7.75 -30.007 -19.88 45
Table 4.1 shows the conversion of carbon to carbon monoxide at low temperatures in the exothermic
portion of the transformation, as calculated by the equation 3.6 described above. The results suggests
that with the given air flow rate the CO formation was at maximum for 500 °C and it decreses with the
increase of temperature.
44
Fig. 4.6 Specific HF at higher temperatures
In figure 4.6 the specific heat flow data has been plotted against time for moderate to high
temperatures. It was observed from the figure that at temperature of 650 °C, the heat flow peak was at
the max i.e at 13mW. Further increase of temperature decreases the heat flow peak values and it was
at the lowest for temperature 850 °C. It was also observed that the isothermal segments were of same
duration for high temperatures i.e. 700 to 850 °C. It suggests that despite of increasing the
temperature there was no noticeable increase in the rate of reaction. This phenomenon suggests that
the reaction may be diffusion limited at high temperatures.
Table: 4.2 Gas mix composition at high temperatures
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
650 -7.615 -29.692 -22.86 31
700 -7.478 -29.365 -17.18 55
750 -7.340 -29.026 -15.23 63
800 -7.200 -28.675 -14.69 65
850 -7.06 -28.314 -9.63 88
Table 4.2 shows the selectivity of the conversion of carbon to carbon monoxide at high temperatures
as a function of temperature as measured by the heat-flow in the exothermic portion of the DSC
curves. It suggests that CO contribution in the gas mix decreases till 650 °C and it starts to increase
beyond this value attaining the maximum observed value at 850 °C. This pattern suggested that the
fixed amount of air used was enough for low temperautres but not sufficient for higher temperatures,
again reinforcing the idea that the reaction may be diffusion limited at higher temperatures.
45
Furthermore, the observed lowering of the peaks of specific heat flow also suggested an increase in
the formation of CO in the gas mix as the temperature increased which was validated through the
heat flow calculations discussed in previous chapter.
Fig.4.7 Percentage Carbon monoxide formation (Charcoal)
In figure 4.7, the percentage of CO formation has been plotted for the whole temperature range. It
indicates that, with a given air flow, the CO formation rate decreases uniformly in the beginning. After
reaching its lowest value at around 650 °C it started to increase uniformly again. IT seems likely that
this behaviour is linked to the fact that, at low temperature the reaction rate is low and the full
oxidation may not be attained; as the temperature increases the formation of CO2 becomes more
likely but, when the reaction begins to be diffusion limited the combustion is, again, incomplete,
increasing the amount of CO that is formed.
4.1.3 Kinetic Modeling
To see if the developed model was able to describe the behaviour of the experimental data, the model
was fitted to the experimental data. For non-impregnated charcoal exothermic reaction began at 550
oC. In fig. 4.8 the experimental and model predicted mass loss have been plotted against time.
46
Fig.4.8 Model Fit at 550 C
Fig. 4.8 shows that the proposed model dictates higher reactivity at low temperature. As the model
was developed utilizing assumption of uniform reaction and the experimental mass loss data was
comprised of partial combustion and pyrolysis segments which dictates global reactions. Perhaps, it
may became the reason for model deviation at low temperature.
Fig.4.9 Model Fit at Higher Temperatures
Figures 4.9 represent the model fitting at high temperatures. It is observed that the model fits very
well for temperatures above 600 °C. However, at low temperature the model predicts a higher rate of
reaction which was not in accordance with experimental data.
47
Table.4.3 Estimated Kinetic Parameters
Temperature
‘T’
Frequency
Factor
‘k’
Activation
Energy
‘EA’
Frequency
Factor
‘kg’
Activation
Energy
‘Ea(g)’
Reaction
order
‘n’
°C Kcal/mol Kcal/mol
600-850 0.068 1540 0.0000005 21000 1.1
48
4.2 Charcoal (1% V Impregnated)
4.2.1 TG/DTG
Fig. 4.10 Raw Experimental data at 400 0C for different sample masses.
In the above figure experimental results at 400 °C for different sample masses have been plotted
against time. It can be observed from the figure that the mass loss curves are almost overlapping till
the temperature rise of 110 °C. It shows that all three of different masses had more or less the same
amount of moisture content that initially came out during temperature ramp. Furtherore all curves are
almost completely overlapping with only negligible differences. However, at the isothermal
temperature of 400 °C the sample with large mass showed a relatively smaller mass loss at the mid of
conversion. Also, the difference between 5 and 9 mg sample is small and they overlap at the
completion of reaction.
49
Fig. 4.11 Fractional mass loss at various temperatures
Fig. 4.12 Conversion at various temperatures
In the above two figures fractional mass loss and corresponding conversion at various temperatures
have been plotted against time. It can be observed that the mass loss occurs even at 400 °C albeit
with a very slow pace. It increases with a faster pace at 450 °C till 500 °C. After that the mass loss
occurred in a more uniform way till 850 °C. These trends show that the combustion of charcoal
impregnated with 1% Vanadium follows a fast complete combustion pattern above temperatures 500
°C. At low temperatures till 400 °C the combustion is partial and seems to follow by a complex
mechanism, as it will be seen by the analysis of the heat-flow signal.
50
The plot of the carbon conversion against time for charcaol impregnated with 1% Vanadium during air
gasification at low temperature is presented below in Figure 4.13.
Fig. 4.13 Conversions at low temperatures
It can be observed from Fig. 4.13 that, as expected, the time required for complete carbon conversion
decreases with the increase of temperature. The sample tested at 400 °C took the longest time.
Furthermore at lower temperature, mass loss occurred at a fairly constant rate till 80 percent of
conversion after that remaining mass loss occurred with a decreasing rate. The decrease in mass
loss at fractional conversion greater than 80 percent may be attributed due to decrease in available
surface area, but may also be associated with a change in mechanism because, as indicated in figure
4.13, the thermicity changes from exo- to endo-thermal. With the progress in mass loss; the pores
may collapse and coalesce; thus limiting the reaction of oxygen with carbon. The phenomenon of
decrease in carbon conversion tendency was very limited at temperature 450 °C and absent at
temperature 500 °C.
Fig. 4.14 Conversion at high temperatures
51
The plot of the carbon conversion against time at high reaction temperatures for charcoal
impregnated with 1% Vanadium during air gasification is presented in Figure 4.14. It was observed
from the figure that conversion time decreases with increase in temperature and the least time was for
isothermal 800 °C.
Plot of DTG versus time for low temperatures has been presented in the figure 4.15. It was evident
from the figure that char impregnated with 1% vanadium showed reactivity at 400 °C.
Fig. 4.15 Time Derivative of char mass fractions at low temperatures
Fig. 4.16 Time derivative of char mass fractions at high temperatures
In the above two figures time derivative of char mass fractions have been plotted as a function of
time. It was observed that the reactivity of char increases with increasing temperature. The char
52
oxidation reaction begins at around 400 °C and increases rapidly with time. After reaching a peak
value, the reactivity starts to decrease due to the combustion of the less reactive portion of the char.
Char reactivity was more or less same for temperatures from 500 up to 700 °C. However, at higher
temperatures it showed sudden remarkable increase.
Furthermore, as already indicated above, the deviation from linearity in mass loss at higher
conversion corresponds to an endothermic process, it may be considered due to the known factors
having a detrimental effect on char gasification rate at higher conversions.
Over time, carbonaceous material remaining in the char is gradually annealed. Annealing reduces
char reactivity by increasing the ordering of the char structure, destroying carbon edges, and reducing
structural defects. As carbon is simultaneously depleted from the char, micropores coalesce into
meso and macropores, reducing char reactivity by effectively reducing the available surface area for
gasification.
Finally, deactivation of the inherent catalytic inorganic species may occur over time. As conversion
increases, each of these interrelated processes has an increasing effect on the gradually decelerating
gasification rate. [35-37]
4.2.2 DSC
Fig. 4.17 Specific HF and Fractional mass loss as function of time
53
Fig. 4.18 Specific HF and Fractional mass loss as function of time
In the above two and the following figure, specific heat flow and mass loss data have been plotted as
a function of time. The plot of 400 °C suggests that with the given amount of air, charcoal impregnated
with 1% Vanadium gives combustion gases. However, in the midway it turned to endothermic region
indicating that the mass loss started to occur via pyrolysis instead of combustion. This trend was very
limited at 450 °C and it disappeared altogether above 500 °C. This trend of partial combustion and
pyrolysis may be due to combination of various factors e.g. low reaction temperature, low rate of
reaction, inhibiting effect of mineral impurities and added catalyst etc.
Fig. 4.19 Specific HF and Fractional mass loss as function of time
54
Fig. 4.20 Effect of mass on specific HF at various temperatures
In the above figure specific HF has been plotted against time. The experiments at 750 and 800 °C
were performed with twice mass, in order to have a significant mass loss at isothermal section. It can
be seen that due to the increase in mass the isothermal segment took longer time.
Fig. 4.21 Specific HF at low temperatures
In the above figure, heat flow data has been plotted as a function of time for three low temperatures. It
was observed that charcoal impregnated with 1% Vanadium showed exothermic reaction at
temperature of 400 °C. However, it shifted to endothermic in the midway. Further, at little higher
temperatures it was observed that CO2 formation started to increase which resulted in higher
exothermic reaction confirmed from the shape of higher heat flow peaks and following table. At 500 °C
55
heat flow peak reaches to the maximum at 20mW, it was observed that instead of having stability in
that region, curve turned more quickly towards endothermic as the rate started to decrease. This
effect was not observed at 450 °C where the oxidation curve after reaching to the highest value of
heat flow, stayed there for some time and followed steady state reaction.
Table: 4.4 Gas mix composition at low temperatures
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
400 -8.278 -31.151 -8.380 99.55
450 -8.149 -30.883 -22.778 35.65
500 -8.017 -30.603 -23.587 31.06
The above table shows that the conversion of carbon to carbon monoxide occurs mostly at low
temperatures. It is observed that at 400 °C during the exothermic regime, total energy released
corresponded to the heat of formation of CO. However, with the increase of temperature it dropped
quickly to 31% at 500 °C.
Fig. 4.22 Specific HF at high temperatures
In the above figure, heat flow data have been plotted against time for moderate to high temperatures.
It was observed from the figure that at temperatures of 550 to 650 °C, the heat flow peaks are at the
max i.e at 20mW. Further increase of temperature decreases the heat flow peak values and it was at
the lowest for temperature 850 °C. Consequently, CO contribution in the gas mix decreases till 650 °C
and it started to increase again and it was at max at 850 °C. This pattern suggested that the fixed
amount of air used was enough for complete combustion at low temperautres but not sufficient for
56
higher temperatures. Furthermore, having low peaks of specific heat flow also suggested increase in
the formation of CO in the gas mix which was validated through heat flow calculations discussed in
previous chapter. It was also observed that with progressive increase in temperature the isothermal
segments decreases till 800 °C and it increased a little again at 850 °C.
Table 4.5 Gas mix composition at high temperatures
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
550 -7.885 -30.311 -25.05 23
600 -7.750 -30.007 -26.22 17
650 -7.615 -29.692 -27.92 8
700 -7.478 -29.365 -20.11 42
750 -7.340 -29.026 -16.98 55
800 -7.200 -28.675 -15.81 60
850 -7.06 -28.314 -8.981 91
The above table shows that the conversion of carbon to carbon monoxide occurs again at high
temperatures. It was observed that CO formation decreases till 650 °C and it started to rise again;
reaching to the maximum at 850 °C.
Fig. 4.23 Percentage Carbon monoxide formation (1% V)
In figure 4.23, the percentage of CO formation has been plotted against the temperature. It suggests
that with a given air flow rate the CO formation decreases rapidly in the beginning and then at slower
57
pace. After reaching its lowest value at around 650 °C it started to increase with a higher pace and
become maximum at 850 °C.
4.2.3 Kinetic Modelling
Fig. 4.24 Model Fit at 400 C
Fig. 4.25 Model Fit at 450 C
In figures 4.24 and 4.25 calculated model and experimental mass loss data have been plotted against
time. Figures show that at low temperature where combustion was followed by pyrolysis like non-
impregnated sample; model shows higher reactivity. However, model fits more or less with increase of
temperature as in figure 4.25.
58
Fig. 4.26 Model Fit at 500 C
In figure 4.26, the experimental and calculated model mass losses have been plotted against time. It
is observed that the model shows little less reactivity at 500 oC. It may be due to the transition from
partial oxidation to complete oxidation.
Fig. 4.27 Model Fit at Higher Temperatures
Figure 4.27 shows the model fitting at moderate to higher temperatures. It can be seen that the
proposed model gives best fits.
59
Table: 4.6 Esitmated Kinetic Parameters at high temparatures.
Temperature
‘T’
Frequency
Factor
‘k’
Activation
Energy
‘EA’
Frequency
Factor
‘kg’
Activation
Energy
‘EA’(g)
Reaction
order
‘n’
°C Kcal/mol Kcal/mol
550-800 0.08 1380 0.000006 20000 1.16
60
4.3 Charcoal (1% Cu Impregnated)
4.3.1 TG/DTG
Fig. 4.28 Fractional mass loss at various temperatures
Fig. 4.29 Conversion at various temperatures
In the above two figures fractional mass loss and conversion at various temperatures have been
plotted against time. It was observed that there is mass loss started at 450 °C with a very slow pace. It
increases with a faster pace at 500 °C. After that the mass loss occured in a more uniform way till 850
°C. These trends showed that the combustion of charcoal impregnated with 1% Copper follows
complete combustion pattern from 550 °C. At low temperatures till 450 °C the combustion is partial
and seems to follow a complex mechanism.
61
Fig.4.30 Time derivative of char mass fractions at low temperatures
Fig. 4.31 Time derivative of char mass fractions at high temperatures
In the above two figures time derivative of char mass fractions have been plotted as a function of
time. It was observed that the reactivity of char increases with increasing temperature. The char
oxidation reaction begins at around 450 °C and increases rapidly with time. After reaching a peak
value, the reactivity starts to decrease due to the combustion of the less reactive portion of the char.
Char reactivity was more or less same for temperatures 650 to 750 °C. However, at higher
temperatures it showed sudden remarkable increase.
62
4.3.2 DSC
Fig. 4.32 Specific HF at low temperatures
In the above figure, specific heat flow data has been plotted as a function of time for three low
temperatures. It was observed that charcoal impregnated with 1% Copper showed exothermic
reaction at temperature 450 °C. However, it shifted to endothermic in the midway. As for the previous
samples we can observe that CO is the most likely product at lower temperatures. Furthermore, at
little higher temperatures it was observed that CO2 formation started to increase which resulted in
higher exothermic reaction confirmed from the shape of higher heat flow peaks and following table. At
550 °C heat flow peak reaches to the maximum at 18mW.
Table. 4.7 Gas mix composition at low temperature
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
450 -8.149 -30.883 -14.21 73
500 -8.017 -30.603 -23.65 31
550 -7.885 -30.311 -23.91 28
The above table shows the conversion of carbon to carbon monoxide. It is seen that CO formation
was at maximum at lowest temperature of 450 °C and then it decreases with the increase of
temperature.
63
Fig. 4.33 Specific HF at high temperatures
In the above figure, heat flow data have been plotted against time for moderate to high temperatures.
It was observed from the figure that at temperatures of 600 to 650 °C, the heat flow peaks are at the
max i.e at 20mW. Further increase of temperature decreases the heat flow peak values and it was at
the lowest for temperature 850 °C. Consequently, CO contribution in the gas mix decreases till 650 °C
and it started to increase again and it was max at 850 °C. This pattern suggests as before that the
fixed amount of air used was enough for low temperatures but not sufficient for higher temperatures.
Further, having low peaks of specific heat flow also suggested increase in the formation of CO in the
gas mix which was validated through heat flow calculations discussed in previous chapter and
presented in the following table. It was also observed that with progressive increase in temperature
the isothermal segments decreases till 850 °C.
Table. 4.8 Gas mix composition at high temperature
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
600 -7.750 -30.007 --25.19 21
650 -7.615 -29.692 -22.51 32
700 -7.478 29.365 -19.79 43
750 -7.340 -29.026 -16.05 60
800 -7.200 -28.675 -11.17 82
The above table explains the conversion of carbon to carbon monoxide at high temperatures. It was
observed that CO formation decreases till 600 °C and it started to rise again; reaching to the
maximum at 850 °C. This trend of increase in CO with the increase of temperature suggests that
64
diffusion limitations started to develop after 600 °C and this phenomenon decreases the formation of
CO despite of increase in temperature.
Fig. 4.34 Percentage Carbon monoxide formation (1% Cu)
In the above figure, percentage of CO formation has been plotted against the temperature. It suggests
that with a given air flow rate the CO formation decreases rapidly in the beginning and then followed
by a constant segment it decrease with a slower pace. After reaching to a limiting value at around 600
°C it started to increase uniformly and become maximum at 850 °C.
4.3.3 Kinetic Modeling
Fig.4.35 Model Fit at 450 C
65
Fig.4.36 Model Fit at 500 and 550 C
Fig.4.37 Model Fit at Higher Temperatures
In figures 4.35 – 4.37, predicted results and experimental data have been plotted against time. It is
observed that model give good fits at temperatures around 500 °C. However, below that temperature
model suggests higher reaction rate like previous samples which was not actual case.
66
Table 4.9 Estimated Kinetic Parameters
Temperature
‘T’
Frequency
Factor
‘k’
Activation
Energy
‘EA’
Frequency
Factor
‘kg’
Activation
Energy
‘EA’(g)
Reaction
order
‘n’
°C Kcal/mol Kcal/mol
500-800 0.084 1320 0.0000065 20000 1.15
67
4.4 Charcoal (1% V + 1% Cu)
4.4.1 TG/DTG
Fig. 4.38 Fractional mass loss at various temperatures
Fig. 4.39 Conversion at various temperatures
In the above two figures fractional mass loss and conversion at various temperatures have been
plotted against time. It was observed that there is mass loss starting at 450 °C with a very slow pace.
It increases with a faster pace at 500 °C. After that the mass loss occurred in a more uniform way till
750 °C. However, above 750 °C it showed a little pronounced increase again till 850 °C These trends
showed that the combustion of charcoal impregnated with (1%V+1%Cu) follows complete combustion
68
pattern from 550 °C. At low temperatures till 450 °C the combustion is partial and followed by complex
mechanism.
Fig. 4.40 Time derivative of char mass fractions at low temperatures
Fig. 4.41 Time derivative of char mass fractions at high temperatures
In the above two figures time derivative of char mass fractions have been plotted as a function of
time. It was observed that the reactivity of char increases with increasing temperature. The char
oxidation reaction begins at around 450 °C and increases rapidly with time. After reaching a peak
value, the reactivity starts to decrease due to the combustion of the less reactive portion of the char.
Char reactivity was more or less same for temperatures 600 to 650 °C. However, at higher
69
temperatures it showed sudden remarkable increase. At 700 °C reactivity drops even below 600 °C
which is not understandable.
4.4.2 DSC
Fig. 4.42 Specific HF at low temperatures
In the above figure, specific heat flow data has been plotted as a function of time for three low
temperatures. It was observed that charcoal impregnated with (1%V+1%Cu) showed exothermic
reaction at temperature 450 °C. However, it shifted to endothermic in the midway. The pattern is
similar to the ones observed before. At little higher temperatures it was observed that CO2 formation
started to increase which resulted in higher exothermic reaction confirmed from the shape of higher
heat flow peaks and the following table. At 550 °C heat flow peak reaches to the maximum at 18mW.
Table 4.10 Gas mix composition at low temperature
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
450 -8.149 -30.883 -14.90 70
500 -8.017 -30.603 -22.18 37
550 -7.885 -30.311 -21.59 39
The above table explains the conversion of carbon to carbon monoxide. It is seen that CO formation
was at maximum at low temperature of 450 °C and then it decreases with the increase of
temperature. This trend suggests the temperature dependence of CO formation in comparison to
CO2. At low rate of reaction CO formation is more.
70
Fig. 4.43 Specific HF at high temperatures
In the above figure, heat flow data have been plotted against time for moderate to high temperatures.
It was observed from the figure that at temperatures of 600 °C, the heat flow peak is at the max i.e at
20mW. Further increase of temperature decreases the heat flow peak values and it was at the lowest
for temperature 850 °C. Further, having low peaks of specific heat flow also suggested increase in the
formation of CO in the gas mix which was validated through heat flow calculations discussed in
previous chapter and presented in the following table. It was also observed that with progressive
increase in temperature the isothermal segments decreases till 850 °C. However, they are same from
600 to 700 °C and from 750 to 850 °C. In these segments rate of reaction may be hindered by the
increased level of impurities.
Table. 4.11 Gas mix compostion at high temperature
Temperature
(°C)
Heat of Formation (Theoretical)
KJ/g
Experiment
KJ/g C → CO (%)
CO CO2
600 -7.750 -30.007 --23.68 28
650 -7.615 -29.692 -22.24 33
700 -7.478 29.365 -19.51 45
750 -7.340 -29.026 -14.79 66
800 -7.200 -28.675 -10.50 85
850 -7.06 -28.314 -6.22 100
The above table shows the conversion of carbon to carbon monoxide at high temperatures. It was
observed that CO formation decreases till 600 °C and it started to rise again; reaching to the
maximum at 850 °C. This trend of increase in CO with the increase of temperature suggests that
71
diffusion limitations started to develop after 600 °C and this phenomenon decreases the formation of
CO despite of increase in temperature. Diffusion limitations may be more pronounced due to high
levels of impurities in the sample.
Fig. 4.44 Percentage Carbon monoxide formation (1%V+1%Cu)
In the above figure, percentage of CO formation has been plotted against the temperature. It suggests
that with a given air flow rate the CO formation decreases rapidly in the begining and then followed by
a more or less constant segment it decrease with a slower pace. After reachig to a limiting value at
around 600 °C it started to increase uniformly and become maximum at 850 °C.
4.4.3 Kinetic Modelling
Fig.4.45 Model Fit at low Temperatures
72
Fig.4.46 Model Fit at Higher Temperatures
In the above two figures proposed model has been validated with the experimental data. The results
are again promising at moderate to higher temperatures.
Table 4.12 Estimated Kinetic Parameters
Temperature
‘T’
Frequency
Factor
‘k’
Activation
Energy
‘EA’
Frequency
Factor
‘kg’
Activation
Energy
‘EA’(g)
Reaction
order
‘n’
°C Kcal/mol Kcal/mol
500-800 0.082 1300 0.0000055 20500 1.13
73
5. Conclusions
5.1. TG
Raw charcoal conversion occurs around 500 o C
Impregnation of catalysts reduces the temperature for carbon conversion and it was
lowest for 1% V.
Charcoal impregnated with 1 % Cu and (1%V+1%Cu) showed conversion at 450 °C.
For low temperatures; inhibiting effect was observed after around 80 % conversion.
Least conversion time was shown by 1% Cu impregnation.
Raw charcoal and 1%V didn’t give sharp change in conversion at high temperatures.
5.2. DTG
Raw charcoal had the lowest reaction rate among others, also the rate increases
slowly with increase in temperature.
For Vanadium the rate was more pronounced at high temperatures.
Copper showed sharp increase in rate with temperature.
For catalyst mix; the rate first increases and then decreases at around 700 °C,
followed by an increase again.
5.3. DSC
Fig. 5.1 Percentage Carbon monoxide formation
For all samples the product was a gas mix.
Maximum CO2 in the gas mix was attained with raw charcoal.
Extreme low and high temperatures favored CO formation.
Moderate temperatures favored CO2 formation.
74
At high temperatures CO formation was attributed to the diffusion limitations.
Cu in the form of mix and alone found to promote diffusion limitations more effectively
in comparison to Vanadium.
5.4. Kinetic Modeling
Langmuir-Hinshelwood type kinetics seems appropriate to be considered as basis for
the charcoal oxidation reaction modeling.
For all samples the predicted kinetic model gave best fits at moderate to high
temperatures.
At low temperatures model suggested high rates which was not in accordance with
actual experimental data.
Model was found valid for complete oxidation reactions only as at low temperatures
the samples were following partial combustion.
Table 5.1 Estimated Kinetic Parameters
Sample
Temperature
‘T’
Frequency
Factor
‘k’
Activation
Energy
‘EA’
Frequency
Factor
‘kg’
Activation
Energy
‘EA’(g)
Reaction
order
‘n’
°C Kcal/mol Kcal/mol
Charcoal 600-800 0.068 1540 0.000005 21000 1.1
1%V 450-800 0.08 1380 0.000006 20000 1.16
1%Cu 500-800 0.08 1380 0.0000065 20000 1.16
1%V+1%Cu 500-800 0.082 1300 0.0000055 20500 1.13
For all samples impregnated with catalyst; the estimated kinetic parameters were found to be
very close. However, raw charcoal followed reaction mechanism with higher activation energy
and slightly lower order in comparison to charcoal impregnated with catalysts.
Predicted model gave good conformity with the experimental data from moderate to high
temperatures. However, the suggested model didn’t operate well at low temperatures where
conversion was partially followed by pyrolysis.
As recommendations; the model may be incorporated with some pyrolysis term to improve its
working at low temperatures.
75
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