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HSC Chemistry Module 9.3 Summary 1. Indicators were identified with the observation that the colour of some flowers depends on soil composition Classify common substances as acidic, basic or neutral Examples of common acidic, basic, or neutral substances Acidic Neutral Basic Vinegar Pure water Caustic soda solution Carbonated soft drinks Sugar solution Drain cleaner Citrus juices Pure alcohol solution Ammonia cleanser Lemon juice Sodium chloride solution Milk of magnesia Stomach acid Ethanol Washing soda Aspirin Oils Soap Car battery acid Lactose solutions Lime water Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, and that the range is identified by change in indicator colour An indicator is a substance that changes colour depending on the degree of acidity of alkalinity of the solution Most indicators change colour over a small pH range, and produce 2 colours: one for acidic solutions, and one for basic solutions Examples of indicators are litmus, phenolphthalein, methyl orange, and bromothymol blue Each of the above substances changes colour over a limited range: o Litmus: 4.5-8.5 o Phenolphthalein: 8.2-10.0 o Methyl orange: 3.1-4.4 o Bromothymol blue: 6.0-7.6

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HSC Chemistry Module 9.3 Summary

1. Indicators were identified with the observation that the colour of some flowers depends on soil composition

Classify common substances as acidic, basic or neutral

· Examples of common acidic, basic, or neutral substances

Acidic

Neutral

Basic

Vinegar

Pure water

Caustic soda solution

Carbonated soft drinks

Sugar solution

Drain cleaner

Citrus juices

Pure alcohol solution

Ammonia cleanser

Lemon juice

Sodium chloride solution

Milk of magnesia

Stomach acid

Ethanol

Washing soda

Aspirin

Oils

Soap

Car battery acid

Lactose solutions

Lime water

Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, and that the range is identified by change in indicator colour

· An indicator is a substance that changes colour depending on the degree of acidity of alkalinity of the solution

· Most indicators change colour over a small pH range, and produce 2 colours: one for acidic solutions, and one for basic solutions

· Examples of indicators are litmus, phenolphthalein, methyl orange, and bromothymol blue

· Each of the above substances changes colour over a limited range:

· Litmus: 4.5-8.5

· Phenolphthalein: 8.2-10.0

· Methyl orange: 3.1-4.4

· Bromothymol blue: 6.0-7.6

· Indicators can be used in combination to obtain a more exact pH value

Identify data and choose resources to gather information about the colour changes of a range of indicators

· Below are the pH colour change ranges for the above indicators. MEMORISE!!!

· An indicator changes colour according to the concentration of H3O+ in the solution, as it shifts the equilibrium of the indicator solution

, were an=anion

· In an indicator, Han and an- have different colours

· The H3O+ concentration shifts the equilibrium, and thereby producing more Han or an-, and thus the solution changes colour according to pH

Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity

· Soil testing

· Soil acidity is an important factor in determining what plants and crops can be grown, so knowing the pH of soil is vital

· A sample of soil is taken, and mixed in distilled water to make a slurry

· A neutral white insoluble powder (barium sulfate) is sprinkled over the sample, and a suitable indicator is added. The white powder allows the colour to be observed more clearly.

· If the soil is too acidic, chemicals such as lime (calcium oxide) are added to make the soil more basic

· Pool testing

· The acidity of swimming pools needs to be monitored and controlled to prevent the growth of microbes, whilst avoiding skin and eye irritation

· Water is sampled, and a few drops of indicator is added

· Pool chlorine or hydrochloric acid is added to achieve a suitable pH

· Monitoring pH of chemical wastes

· Wastes from laboratories or photographic film centres can be highly acidic or alkaline

· Discharges into the sewerage system need to be nearly neutral, so the pH of wastes are monitored, and neutralised if necessary

Solve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basic

EXAMPLE

A drain cleaner’s pH was tested using litmus paper and phenolphthalein. The litmus turned blue, and the phenolphthalein turned pink. What can be said about the pH of the drain cleaner?

· As the litmus paper turned blue, the drain cleaner is basic

· As the phenolphthalein turned pink, it must be strongly basic

Some dilute vinegar was also tested using methyl orange, which turned yellow, and phenolphthalein, which turned pink. What can be said about its pH?

· The test is inconclusive, as the ranges do not definitively determine whether the vinegar is acidic, basic, or neutral.

Perform a first-hand investigation to prepare and test a natural indicator

METHOD

· One large beetroot was peeled and chopped, then placed into a beaker of 100mL of distilled boiling water, heated with a Bunsen burner. 5mL of HCl, NaOH, distilled water, and NaCl each were added to four test tubes, and then 3mL of beetroot juice was added to each test tube. Any colour changes were observed.

· SAFETY: Acids are corrosive, bases are caustic => use low concentrations (1M), wear safety glasses

RESULTS

· The beetroot juice remained red in water, NaCl, and HCl

· The beetroot juice turned yellow in NaOH solution

· Thus beetroot juice can be used to distinguish basic substances by changing colour from red to yellow.

· A beetroot was used as its pigmentation can be easily extracted

VALIDITY/ACCURACY/RELIABILITY

· Results were compared to other class results, which were similar

· HCl and NaOH were used to test the aim, by measuring a wide range of the pH scale

· Distilled water minimised error in the preparation of the beetroot juice

· A fresh beetroot was used instead of a canned one as canned beetroots contain preservatives (such as acids) that may alter the results

· The experiment was limited because…

· The size of the beetroot was not controlled

· The exact pH transition was not measured => could be measured by titration with data loggers and sensors

2. While we usually think of the air around us as neutral, the atmosphere contains acidic oxides of carbon, nitrogen and sulfur. The concentration of the acidic oxides have been increasing since the Industrial Revolution

Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids

· ACIDIC OXIDES

· Compounds that contain oxygen that act as acids

· React with water to produce an acid

· Reacts with bases to produce salts

· Non-metals oxides mainly behave as acids

· EXAMPLES OF ACIDIC OXIDES

· Sulfur dioxide (SO2)

(sulfurous acid)

· Carbon dioxide (CO2)

(carbonic acid)

· Nitrogen dioxide (NO2)

(nitric and nitrous acid)

· Diphosphorus pentoxide (P2O5)

(phosphoric acid)

· BASIC OXIDES

· React with acids to form salts

· Do NOT react with alkali solutions (an alkali is a water-soluble base)

· Metal oxides mainly behave as bases

· EXAMPLES OF BASIC OXIDES

· Potassium oxide (K2O)

(potassium hydroxide)

· Magnesium oxide (MgO)

(magnesium hydroxide)

· AMPHOTERIC OXIDES

· Amphoteric oxides can react with both acids and bases

· For example, aluminium oxide (Al2O3) can react with hydrochloric acid

· BUT can also react with sodium hydroxide

· NOTE!!! => Amphoteric substances can REACT with both acids and bases, but do not necessarily ACT as both acids and bases. Such substances are called amphiprotic substances

· NEUTRAL OXIDES

· Some elements form neutral oxides, such as hydrogen (H2O), carbon monoxide (CO), dinitrogen oxide (N2O), and nitric oxide (NO)

Analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the Periodic Table and acidity/basicity of oxides

GENERALISATION/TRENDS

· Non-metal oxides are acidic

· The further to the right & top of the Periodic Table (except the noble gases), the more acidic its oxide is

· This is because the more electronegative the element is, the greater ionisation occurs when in contact with water (~100% for S, N, Cl)

· Metal oxides are basic

· The further to the left & bottom of the Periodic Table, the more basic its oxide is

· This is because the less electronegative elements ionise easily in water to form oxide ions, which react with water to form hydroxide ions

· Amphoteric oxides are between the acidic and basic oxides

· Noble gases have no oxides

· Some other elements (H, N, C) form neutral oxides

Define Le Chatelier’s principle

· Recall that not all chemical reactions go to completion, i.e. the reaction goes in one direction (reactants to products)

· Not all reactions are one-directional => many reactions consist of a forward reaction (left to right) and a reverse reaction (right to left)

· These reactions are reversible reactions

· One example of a reversible reaction is a saturated solution of a salt in water

· Another example is the reaction of carbon dioxide in water

· Reversible reactions do not go to completion, but reach a chemical equilibrium

· At chemical equilibrium, the concentration of reactants and products do not change

· The forward and reverse reactions still continue, but proceed at the same rate (the equilibrium is dynamic)

· Chemical equilibrium occur in closed systems, where the macroscopic (observable) properties are constant

· If the conditions of the system change, the system will no longer be at equilibrium

· Le Chatelier’s principle states that:

If a system at equilibrium is disturbed, then the system adjusts itself so as to minimise the disturbance

· A system at equilibrium will readjust to oppose any changes

Identify factors which can affect the equilibrium in a reversible reaction

· Some of the factors that can affect the equilibrium in a reversible reactions are the concentration of particular chemical species, the total pressure of the system (only for reactions that involve gases), and the temperature

CONCENTRATION

BlueGreen

· Consider the following four scenarios guided by Le Chatelier’s principle:

· If the concentration of Cl- ions are increased, the equilibrium shifts to the right (i.e. the rate of the forward reaction increases), so the system becomes more green as more CuCl42+ ions are produced

· If Cl- ions are removed, the equilibrium shifts to the left, and the system becomes more blue

· If water is added, the equilibrium shifts to the left, and the system becomes more blue

· If water is removed, the equilibrium shifts to the right, and the system becomes more green

· Generally…

· If the concentration of a particular species is INCREASED, the equilibrium point will shift to the OPPOSITE side of the equation to oppose the disturbance by reducing the concentration of that species.

· If the concentration of a particular species is DECREASED, the equilibrium poin till shift to the SAME side of the equation to oppose the disturbance by increasing the concentration of that species.

PRESSURE

BrownColourless

· Recall that equal number moles of gases occupy equal volumes

· The pressure of a closed system can be increased by reducing the volume of the system, and can be decreased by increasing the volume.

· If the pressure is increased, the equilibrium will shift to the right, as the right has less moles of gas (therefore occupies less volume), and the system becomes more colourless

· If the pressure is decreased, the equilibrium will shift to the left, as the left has more moles of gas (therefore occupies more volume), and the system becomes more brown

· Generally…

· If the pressure of a system is increased, the equilibrium will shift to the side with fewer moles to decrease the pressure of the system.

· If the pressure of a system is decreased, the equilibrium will shift to the side with more moles to increase the pressure of the system

TEMPERATURE

· Recall that an endothermic reaction absorbs heat, whilst an exothermic reaction releases heat.

· By using Le Chatelier’s Principle, the following predictions can be made about a closed system:

· An increase in temperature will increase the rate of the endothermic reaction

· A decrease in temperature will increase the rate of the exothermic reaction

· If the ΔH of a reversible reaction is given, it is referring to the forward reaction

· Remember that ΔH is positive for endothermic reactions, and negative for exothermic reactions

· EXAMPLE

Colourless White

· In the above reaction, the forward reaction is exothermic, and the reverse reaction is endothermic

· If the temperature of the system is increased, the equilibrium shifts to the left (favouring the endothermic reverse reaction), and the system becomes more colourless

· If the temperature of the system is decreased, the equilibrium shifts to the right (favouring the exothermic forward reaction), and the system becomes more white

Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle

· The reaction between carbon dioxide and water is a reversible reaction that goes to equilibrium

· The equilibrium between carbon dioxide and carbonic acid can be affected by concentration, pressure, and temperature.

EFFECT OF CONCENTRATION

· The effect of concentration on the equilibrium can be explained by Le Chatelier’s principle

· If the concentration of CO2 is INCREASED, the equilibrium will shift to the right to dissolve more CO2, thereby opposing the increase of CO2

· If the concentration of CO2 is DECREASED, the equilibrium will shift to the left to produce more CO2, thereby opposing the decrease of CO2

· The effect of acids (H3O+) and bases (OH-) can also be explained by considering Le Chatelier’s principle

· If an acid is added, the concentration of H3O+ increases, so the equilibrium shifts to the left to decrease the concentration of H3O+.

· If a base is added, the OH- ions react with the H3O+ to produce water. Consequently, the equilibrium shifts to the right to produce more H3O+ ions to counteract the change.

PRESSURE

· The left side of the equation contains one mole of gas (CO2), whilst the right side of the equation contains zero moles of gas. Hence:

· If the pressure is INCREASED, the equilibrium shifts to the right, thereby reducing pressure by dissolving CO2 gas (decreasing the volume of gas decreases the pressure)

· If the pressure is DECREASED, the equilibrium shifts to the left, thereby increasing the pressure by producing CO2 gas.

TEMPERATURE

· Note that because ΔH is negative in the above equation, the forward reaction is exothermic, and the reverse reaction is endothermic. Hence:

· If the temperature is INCREASED, the equilibrium will shift to the left (i.e. the reverse reaction will be favoured), as it is an endothermic reaction, which will oppose the change by cooling the system

· If the temperature is DECREASED, the equilibrium will shift to the right (i.e. the forward reaction will be favoured), as it is an exothermic reaction, which will oppose the change by warming up the system

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen

Sulfur dioxide (SO2)

· Most of the atmospheric sulfur dioxide comes from the oxidation of sulfur in compounds

· NATURAL SOURCES include volcanic activity/gases, geothermal hot springs, bushfires [H2S], bacterial decomposition [H2S]

· INDUSTRIAL SOURCES include the smelting of metallic sulfide ores (e.g. Cu2S), and the combustion of fossil fuels (particularly coal) containing metallic oxides (e.g. FeS2) or carbon compounds containing sulfur

Oxides of nitrogen (NOx)

· NOx is the general for oxides of nitrogen. These oxides include nitric oxide (or nitrogen monoxide, NO), nitrogen dioxide (NO2), and nitrous oxide (or dinitrogen monoxide, N2O)

Nitric oxide (NO)

· Nitric oxide is primarily produced through the reaction of nitrogen and oxygen gas in high-temperature combustion environments (particularly high voltage)

· NATURAL SOURCES include lightning, which is a localised, high-temperature, and high-voltage environment

· INDUSTRIAL SOURCES include the combustion of fossil fuels, particularly in car engines and at power stations

· Nitrogen dioxide (NO2)

· Much of the atmospheric nitrogen dioxide is produced through the slow reaction of nitric oxide and oxygen

· NATURAL SOURCES include the conversion of nitric oxide produced by lightning

· INDUSTRIAL SOURCES include the direct emission of NO2 from motor vehicles and power stations, and the conversion of NO released from vehicles and power stations.

Nitrous oxide (N2O)

· NATURAL SOURCES include the natural actions of nitrogen-fixing bacteria on nitrogenous materials in the soil

· INDUSTRIAL SOURCES include the increased use of nitrogenous fertilisers, which is more food for bacteria to convert to nitrous oxide

Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen

Sulfur dioxide

· As mentioned above, the general reaction that produces atmospheric SO2 is the oxidation of sulfur in sulfur-containing compounds

· For example, sulfur dioxide is produced when organic matter releases H2S during decomposition:

· The smelting of zinc/iron also releases sulfur dioxide into the atmosphere

Oxides of nitrogen

· As mentioned above, nitrogen and oxygen gas react in high-temperature environments (such as lightning or combustion engines) to produce nitric oxide

· Nitric oxide and oxygen can slowly react to produce nitrogen dioxide

· The reverse equation can proceed in the presence of sunlight

Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen

· There is extensive evidence to suggest that there has been an increase in the atmospheric concentration of oxides of sulfur and nitrogen.

· Qualitative evidence includes:

· The increased formation of photochemical smog in urban areas, which indicates an increased concentration of NO2

· e.g. in 1952, 4000 people were killed due to smog in London, leading to environmental controls in the U.K.

· Increased incidence of acid rain indicates an increase in atmospheric SO2 and NOx

· Increased use of motor vehicles, combustion of fossil fuels for electrical production, and smelting of sulfide ores since the Industrial Revolution would logically suggest an increased concentration of SO2 and NOx

· THIS EVIDENCE IS CIRCUMSTANTIAL HOWEVER => it does not prove with certainty that concentrations of oxides of sulfur and nitrogen have increased. Nonetheless, they do provide a strong indication of increased levels of oxides of sulfur and nitrogen.

· Quantitative evidence for increasing levels globally has been difficult to obtain for the following reasons:

· Oxides of sulfur and nitrogen occur in relatively low concentrations => the global average is 0.001ppm, whilst peak levels in polluted urban areas such as Los Angeles peak at 0.47ppm. For comparison, CO2 occurs at ~360ppm.

· Chemical instruments able to measure these very low concentrations have only been available since the 1970s, so not reliable data is available before then

· SO2 and NOx are dissolved by atmospheric water, so measuring their levels globally is difficult.

· Quantitative measurements for oxides of sulfur and nitrogen locally is available

· The NSW EPA monitors atmospheric pollutant levels across Sydney, and has shown that the levels of SO2 and NOx have stabilised over the past two decades, and vary seasonally.

· This is due to emission controls that have been progressively implemented since the 1950s.

· The stabilised trend is similar in other Western cities, though levels in industrialising areas (notably China) have nearly tripled over the past twenty years

Explain the formation and effects of acid rain

FORMATION OF ACID RAIN

· Recall that pure water has a pH of 7 (i.e. is neutral)

· Unpolluted rain water typically has a pH of 6.0-6.5 due to atmospheric CO2, which dissolves in water to form carbonic acid (as discussed above)

· Acid rain is rain that has a pH less than 5.0

· Acid rain in industrial areas typically occurs due to the presence of SO2 and NOx, which react with water to form acids

· Sulfur dioxide reacts with water to produce sulfurous acid

· Sulfurous acid reacts with oxygen (catalysed by air particles) to produce sulfuric acid

· Nitrogen dioxide reacts with water to produce nitric acid and nitrous acid

· Nitrous acid reacts with oxygen (catalysed by air particles) to produce nitric acid

· Both sulfuric and nitric acids are strong acids, thus both can significantly reduce the pH of rain water to unnatural levels (pH of 3-5)

EFFECTS OF ACID RAIN

· Increasing acidity of lakes, which can kill marine life such as fish by killing eggs, and irritating skin and gills

· Damage to forests through the leeching of toxic ions such as aluminium and mercury. Also, the decrease in the pH of soil makes it difficult for plants to absorb sufficient calcium, potassium or potassium, and can kill bacteria important micro-organisms in the soil

· Erosion of marble and limestone decorations, as such materials contain carbonates (primarily calcium carbonate), which readily with acids.

Analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment

· See above for the industrial origins of sulfur dioxide and oxides of nitrogen, and the relevant chemical equations.

· The reasons for their concern are significant, due to the resultant health problems, and environmental problems.

Health problems

· Sulfur dioxide is a severe respiratory irritant, and causes breathing difficulties at concentrations as low as 1ppm. People suffering from asthma and emphysema are particularly susceptible.

· Nitrogen dioxide irritates the respiratory tract, and causes breathing discomfort at 3-5ppm. At higher concentrations it can destroy tissue, as it forms nitric acid, which is a strong acid.

· Nitrogen dioxide also leads to the formation of photochemical smog, which leads to the production of ozone, and other particulates (haze) and pollutants such as PAN (peroxyacylnitrates). Ozone has harmful effects on the body at concentrations as low as 0.1ppm (see Module 9.4 section 4)

Environmental problems

· Both sulfur dioxide and oxides of nitrogen lead to the formation of acid rain, which leads to environmental destruction (see below)

Calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0°C and 100kPa or 25°C and 100kPa

***NOTE*** => ALWAYS note the number of significant figures given in a question. The final answer cannot have more significant figures than the data with the least significant figures.

CALCULATING VOLUMES GIVEN THE MASSES OF SOME SUBSTANCES

· Given the mass of a substance, the volume of a gas can be calculated by the following steps:

1. Write the relevant chemical equation

2. Convert mass to moles

3. Use the stoichiometry of the equation to work out the number of moles of each chemical species as required

4. Convert moles to volume

NOTE: The molar volume of gases at 0°C and 100kPa or 25°C and 100kPa is given on the chemistry data sheet in exams

EXAMPLE

Calculate the volume of sulfur dioxide produced when 50g of hydrogen sulfide is oxidised at 25°C and 100kPa.

First write the relevant chemical equation

Next calculate the number of moles of H2S

From the stoichiometry of the equation, we can see that

Thus

Now we can calculate the volume of SO2

· CALCULATING MASSES GIVEN GASEOUS VOLUMES

· Given the volume of a gaseous substance, the mass of a particular species can be calculated by the following steps

1. Write the relevant chemical equation

2. Convert volume to moles

3. Use the stoichiometry of the equation to work out the number of moles of each species as required

4. Convert moles to mass

· The steps involved are the reverse of the example given above.

Identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25°C and 100kPa

METHOD

· A 300mL bottle of soda water was weighed on electronic scales. The can was then vigorously shaken, and the cap was opened. The bottle was shaken periodically for one hour, whilst ensuring no drink spilled, and the final mass of the drink was recorded.

· Make sure no drink spills, because it reduces the validity of the experiment, and poses a safety hazard.

RESULTS

· Initial mass: 330.36g

· Final mass: 328.50g

· Mass loss: 1.86g

· Volume of CO2 released: 1.05L (at 25°C 100kPa)

· In the experiment, we assume that any mass lost is due to carbon dioxide loss

THEORY

· Carbon dioxide exists in soda water by the following equilibrium:

· By shaking the bottle, we introduce additional energy (kinetic) to the soda water.

· By Le Chatelier’s Principle, the equilibrium opposes the increase in energy by favouring the reverse reaction, as it is an endothermic reaction that absorbs heat, thus CO2 gas evolves.

RELIABILITY

· The experiment was not repeated, no average was taken, the range of data wasn’t identified, and no outliers were identified => unreliable method

· Taking repeated results would improve reliability, but they must be conducted UNDER IDENTICIAL CONDITIONS

VALIDITY/ACCURACY

· Variables that could affect the investigation (namely the conditions of the CO2/H2CO3 equilibrium) were controlled, ensuring a fair test

· The use of electronic scales increased the accuracy of results

· The validity of the experiment could have been improved by having a bottle of water open as a control

· A more accurate result could have been obtained through a titration

· Factors to consider in the validity…

· Record the original mass of the soft drink whilst it is in the bottle

· Ensure that no soft drink spills in the experiment

· Minimise the evaporation of water (do not heat, keep lid on top etc.), as this would alter the mass change and hence be an invalid test

· Attempt to evolve all the gas from the soft drink

· Use soda water as opposed to cola or other soft drinks, as most soft drinks contain other acids (e.g. H3PO4) for flavouring, which may also exist in an equilibrium, and thus alter the results

3. Acids occur in many foods, drinks and even within our stomachs

Define acids as proton donors and describe the ionisation of acids in water

· Acids are called proton donors, as they donate protons (H+ ions) in reactions

EXAMPLE

· In the above reaction, HCl has donated a proton (H+) to the OH- ion I NaOH to form H2O

· Acids are ionised in water, as the acid donates a proton to a water molecule to form a hydronium ion (H3O+)

EXAMPLE

· HCl donates a proton H2O to form the hydronium ion (H3O+) and a chloride ion, which is solvated by the water

Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic), hydrochloric and sulfuric acid

ACETIC ACID

· Systematic name: Ethanoic acid

· Molecular formula: CH3COOH

· Weak monoprotic acid (4.2% ionisation at 0.010molL-1)

· Present in vinegar

· Manufacture to make organic chemicals

CITRIC ACID

· Systematic name: 2-hydroxypropane-1,2,3-tricarboxylic acid

· Molecular formula: C6H8O7

· Weak triprotic acid (27.5% ionisation at 0.010molL-1)

· Present in citrus fruit

· Used as a food additive

HYDROCHLORIC ACID

· Molecular formula: HCl

· Very strong monoprotic acid (100% ionisation)

· Present as stomach acid

· Used in swimming pools

SULFURIC ACID

· Molecular formula: H2SO4

· Very diprotic strong acid

· Used to make fertiliser

Identify pH as –log10[H+] and explain that a change in pH of 1 means a ten-fold change in [H+]

· pH is defined by the following equation:

OR

· As pH is a logarithmic scale base 10, a change in pH of 1 means a ten-fold change in [H+]

· e.g. A pH of 3 means [H+] = 1.0x10-3molL-1, whilst pH of 4 means [H+]=1.0x10-4molL-1

· The logarithmic scale was chosen for convenience of calculations

· Another scale used is pOH, where

· The relationship between pH and pOH are

·

·

· NOTE: In calculations, make sure that you use the concentration of H+ ions, NOT the number of moles

Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations

· Calculating the pH of strong acids:

· Calculate the concentration of the acid from the given data (if in doubt, convert any chemical data to moles, then divide by volume)

· Write the acid’s ionisation equation

· Monoprotic

· Diprotic

· Triprotic

· Use the stoichiometry of the equation to calculate the concentration of hydrogen ions => multiply by the ratio

· Use the pH equation to calculate the pH of the substance

EXAMPLE

Calculate the pH of 0.50mol of sulfuric acid, if the volume of acid is 5.0L

The ionisation reaction is

Thus

REMEMBER

· ***VERY IMPORTANT NOTE*** SIGNIFICANT FIGURES!!! The rules for significant figures in pH is an EXCEPTION to other calculations => the number of significant figures in the hydrogen ion concentration should equal the number of decimal places in the pH value (i.e. the whole number doesn’t count in significant figures)

· For example, if [H+]=0.00100, then pH=3.000 not and I repeat NOT 3.00

· Use log10 (log base ten) on the calculator, NOT the natural logarithm (loge or ln)

· When using the pH formula, ensure that you are using the concentration of H+ ions, NOT the number of moles

· If the acid is polyprotic, and the concentration of the acid is given, you must multiply the concentration according to the number of ionised hydrogen ions to find the hydrogen ion concentration => writing the ionisation equation significantly helps

· If the concentration of hydroxide ions is given, calculate the pOH of the solution, and find the pH by using the relationship pH+pOH=14

· If the concentration of a weak acid is given, and the degree of ionisation, multiply the concentration of the acid by the degree of ionisation (as a percentage) to find the concentration of ionised acid particles => continue from step 2

Describe the use of the pH scale in comparing acids and bases

· The pH scale is a logarithmic scale to determine the acidity or basicity of a substance

· The pH scale is a measure of the hydrogen ion [H+] or hydronium ion [H3O+] in a solution

· A substance is classed as acidic or basic according to its pH in reference to the pH of pure water

· Pure water exists in an equilibrium with hydronium and hydroxide ions

· Careful measurements have shown that

· Thus neutral substances have a pH of 7

· Acidic substances have a pH < 7 => the lower the pH, the more acidic a substance is

· Basic substances have a pH > 7 => the higher the pH, the more basic a substance is

Describe acids and their solutions with appropriate use of the terms strong, weak, concentrated and dilute

STRENGTHS OF ACIDS => DEGREE OF IONISATION

· A strong acid is an acid that completely ionises in water solution

· Only acids that achieve 100% ionisation are classed as strong

· The ionisation reaction goes to completion

· EXAMPLE

· The following table provides a list of strong acids and bases. MOST OTHER ACIDS OR BASES ARE WEAK => MEMORISE THIS TABLE

ACIDS

BASES

Name

Formula

Name

Formula

Hydrochloric acid

HCl

Potassium hydroxide

KOH

Nitric acid

HNO3

Sodium hydroxide

NaOH

Sulfuric acid

H2SO4

Lithium hydroxide

LiOH

Hydrobromic acid

HBr

Rubidium hydroxide

RbOH

Hydroiodic acid

HI

Caesium hydroxide

CsOH

Perchloric acid

HClO4

· In a weak acid, only some of the acid molecules ionise to form hydronium ions => they only partially ionise

· A weak acid is releases less protons than a strong acid of the same concentration

· Weak acids ionise to variable extents => they have a degree of ionisation

· A weak acid reaches an equilibrium between ionised and intact molecules

EXAMPLE

STRENGTH IS NOT RELATED TO CONCENTRATION

CONCENTRATION OF ACIDS => CONCENTRATION OF PARTICLES

· A concentrated acid has a high concentration of acid particles (>~5molL-1)

· A dilute acid has a low concentration of acid particles (<~2molL-1)

· pH measures the concentration of an acid, hence a strong acid could still have a relatively high pH (i.e. a strong acid can be dilute)

· BE CAREFUL TO DISTINGUISH BETWEEN STRENGTH AND CONCENTRATION => a strong acid can be both concentrated and dilute, so can a weak acid

Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their molecules

· HCl => Citric => Acetic (decreasing strength)

· The following table shows the pH of each acid at 0.100molL-1

Acid

pH at 0.100mol-1

Hydrochloric acid (HCl)

1.0

Citric acid (C6H8O7)

2.1

Acetic acid (CH3COOH)

2.9

· By using the pH formula, we can calculate the concentration of hydrogen ions [H+] in each solution

· HCl: [H+] = 0.1molL-1

· Citric: [H+] = 0.0079molL-1

· Acetic: [H+] = 0.00126molL-1

· From the above data we can calculate the degree of ionisation of each acid dividing [H+] for each acid by the total concentration of substance

· HCl: 0.1/0.1 = 100% ionisation

· Citric: 0.0079/0.1 = 7.1% ionisation

· Acetic: 0.00126/0.1molL-1 = 1.26% ionisation

· Thus acetic acid is the weakest acid, as only 1.26% of the acid particles ionise (i.e. it has the lowest degree of ionisation) => it is a very weak acid

· Citric acid is slightly stronger than acetic acid, as a greater percentage (7.1%) of its molecules ionise in solution => it is a weak acid

· Hydrochloric acid by far the strongest acid, as it achieves 100% ionisation of its molecules => it is a strong acid

· NOTE => the degree of ionisation of a weak acid depends on the concentration of the acid => changing the concentration of a weak acid solution can change the acid’s strength.

Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions

· The strength of an acid depends on its degree (100%) ionisation

· A strong acid will completely ionise (100 in solution => the ionisation reaction goes to completion

· A weak acid will partially ionise in solution => the ionisation reaction goes to an equilibrium

· The greater proportion of an acid’s ions to intact molecules, the stronger the acid is

· In other words, the faster the rate of the forward reaction, the stronger the acid is

· In the diagram below, HA is a stronger acid than HB, as there is a greater ratio of A- ions to HA molecules than B- to HB in their respective equilibriums

· For example, a solution of acetic acid reaches an equilibrium between intact molecules and ions

· Typically acetic acid’s degree of ionisation is ~1.3%

· An equilibrium exists between intact acetic acid molecules and its ions, but the proportion of ions to intact molecules is very low ([CH3COOH]/[CH3COO-]), hence acetic acid is a weak acid

· Weak acid equilibriums adhere to Le Chatelier’s principle, thus the pH of an acid can change according to the factors that influence the equilibrium

Gather and process information from secondary sources to write ionic equations to represent the ionisation of acids

· Recall that strong acids go to completion. Generally…

EXAMPLE

· Weak acids on the other hand reach an equilibrium. Generally…

EXAMPLE

· NOTE: Polyprotic ionise in steps, and the degree of ionisation for each step significantly decrease

EXAMPLE

Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids

· Above is a model of the molecular nature of acids, and their subsequent ionisation

· The beaker on the left contains a strong acid. As can be seen, there are no intact molecules, as the acid has completely ionised, releasing one H+ ion per molecule

· The beaker on the right contains a weak acid of equal concentration. The acid has only slightly ionised, with most of the molecules remaining intact. The concentration of H+ ions is significantly less than in the left beaker, and hence its pH would be much higher.

· The ionisation of strong and weak acids can also be modelled using molecular model kits, which can demonstrate the breaking and formation of chemical bonds during the ionisation process

Gather and process information from secondary sources to explain the use of acids as food additives

· Weak/dilute acids are used as food additives for two primary reasons: as preservatives (as many microbes can’t survive in acidic conditions), and as flavouring (adds a certain tartness, or sharp/sour flavour)

· PRESERVATIVES

· Acetic/ethanoic acid (CH3COOH) => Used as vinegar (4% solution) to preserve food by ‘pickling’

· Propanoic acid => controls bacteria and mould growth, particularly in bread, potato crisps, and cake mixes

· Citric acid => natural preservative, often added to jams and conserves

· Tartaric acid => preservative in jams, fruits, pickles, and soft drinks

· FLAVOURINGS

· Phosphoric acid (H3PO4) => added to soft drinks for a tartness of flavour

· Carbonic acid (H2CO3) => adds the ‘fizz’ to soft drinks

· Acetic/ethanoic acid => also used as acidic flavouring, such as salad dressing

· Malic acid => flavour enhancer particularly in fruit fillings; adds a savoury taste

Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition

Naturally occurring acids

Naturally occurring bases

Name

Chemical formula

Name

Chemical formula

Stomach acid (hydrochloric acid)

HCl

Ammonia

NH3

Vinegar (acetic acid)

CH3COOH

Metallic oxides

CuO, Fe2O3

Citric acid

C6H8O7

Limestone (calcium carbonate)

CaCO3

Lactic acid

C3H6O3

Nicotine

C8H14N2

Acid rain (carbonic acid)

H2CO3

Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals

METHOD

· A range of substances dissolved in 20mL solutions were placed in separate beakers

· The substances tested were washing powder, an antacid tablet, table salt solution, milk, a vitamin C tablet, and aspirin

· Each substance was tested with a pH probe and data logger to determine its pH

· 2 drops of universal indicator was added to each test tube, and the colour was checked against a pH chart

· SAFETY: General risks of using acids => use low concentration, and use a dropper when transferring solutions

RESULTS

Substance

pH

Colour

Classification

Washing powder

8.0-9.0

Blue

Base

Antacid tablet

8.0-9.0

Blue

Base

Table salt

6.5-7.5

Green

Neutral

Milk

6.5-7.5

Green

Neutral

Vitamin C tablet

5.0-6.0

Yellow

Acid

Aspirin

5.0-6.0

Yellow

Acid

RELIABILITY/ACCURACY/VALIDITY

· Universal indicator was used as it could distinguish between acids, bases, and neutral substances

· The use of a pH probe provided more accurate quantitative results

· Equal quantities of each substance was dissolved in 20mL of water to make a fair test of the substance’s pH

· A wide variety of substances were tested, and results matched expected results and other results in the class

Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids

· Refer to practical 9.2.3j)

METHOD

· 50mL of 0.1M solutions of sulfuric acid, hydrochloric acid, citric acid, and acetic acid were placed into four separate small beakers. The pH of each solution was tested with a pH probe.

· NOTE: Ensure that the pH probe is always wet

· SAFETY: Acids are corrosive substances => use dilute concentrations, use a dropper to transfer acid, wear safety goggles

RESULTS

· Sulfuric acid (H2SO4): 2.20

· Hydrochloric acid (HCl): 2.37

· Citric acid (C6H8O7): 2.85

· Acetic acid (CH3COOH): 3.44

· Sulfuric acid had the lowest pH, because it is a strong diprotic acid, whilst hydrochloric acid had a higher pH because it is a strong monoprotic acid. Citric acid and acetic acid are both weak acids, but citric acid ionises to a greater extent than acetic acid, and so has a lower pH.

VALIDITY/ACCURACY/RELIABILITY

· Computer technologies were used to maximise accuracy and reliability.

· Results were compared to other results in the class and were found to be similar, and matched expected results

· Other variables (e.g. temperature, pressure) were controlled, ensuring a fair test

· The pH values did not match the expected theoretical results because...

· The theoretical values are taken at standard laboratory conditions

· The standard solutions prepared for the classroom are not accurate => taken to one significant figure

4. Because of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, the definitions of acid and base have been refined

Outline the historical development of ideas about acids including those of:

· Lavoisier

· Davy

· Arrhenius

LAVOISIER => proposed that acids contained OXYGEN

· Stated that acids are corrosive, sour-tasting substances containing oxygen

· Oxygen was only a recently-discovered element

· Since most known acids contained oxygen (e.g. CH3COOH, H2SO4, H2CO3), he believed all acids must contain oxygen

· Lavoisier only experimented with oxyacids

· He believed that oxygen was the source of acidity

DAVY => proposed that acids contained HYDROGEN

· Davy demonstrated that HCN and HCl did not contain oxygen, thus disproving Lavoisier’s hypothesis

· Other hydrohalic acids had recently been discovered (e.g. HBr, HF), leading Davy to propose that hydrogen gave acids their acidic properties

· This led to his hypothesis that all acids contain hydrogen

· German scientist von Liebig later extended Davy’s theory on acids to state that acids contained replaceable hydrogen, thus explaining why methane (CH4) was not acidic

ARRHENIUS => proposed that acids IONISED in WATER to produce H+ IONS

· Davy’s hypothesis was limited as it did not explain why many acidic properties, such as the production of NO2 instead of hydrogen gas when an acid reacted with a metal

· After Arrhenius’s extensive work on hydrolysis, he discovered that hydrogen gas evolved at the cathode during the electrolysis of water

· This lead him to propose that acids dissociated in water to produce H+ ions, which then reacted during hydrolysis to produce hydrogen gas

· For example, hydrogen chloride gas ionised in solution to produce H+ ions.

· He also proposed that bases dissociate in water to produce hydroxide ions (OH-)

· He recognised that some acids were weaker than others (e.g. acetic acid), and proposed that weaker acids only partially ionised in water => this led to the development of the pH scale

Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions

· The definition of acids/base reactions have changed throughout history, due to advances in technology and understanding have altered the direction of scientific thinking

· Both the Lavoisier and Davy model of acids were limited, and are not used today

· The Arrhenius model of acids/bases is used as a simple model today

· The current models of acid/base reactions are the Brӧnsted-Lowry model, and the Lewis model (not required for HSC)

Theory

Acid definition

Base definition

Development

Limitations

Lavoisier

Contained oxygen

No definition

Investigated non-metallic oxides that produced oxyacids in water (e.g. H2SO4). Oxygen had recently been discovered.

Substances not containing oxygen (e.g. HCl, HCN) were shown to be acidic

Davy

Contained hydrogen

No definition

Investigated hydrohalic acids (e.g. HCl, HBr)

Did not explain the production of other gases (e.g. NO2) in acid/metal reactions

Arrhenius

Ionised in water to produce H+ ions

Ionised in water to produce OH- ions

Investigated electrolytes and electrolysis of acids. Explained how acids (H+) and bases (OH-) react to produce water in neutralisation reactions.

Does not explain why metallic oxides and carbonates are basic. Does not explain acidic/basic salts. Does not explain anhydrous neutralisation reactions.

Brӧnsted-Lowry

Proton (H+) donor

Proton (H+) acceptor

Investigated the problems with Arrhenius’s definition of acids

Limited to reactions containing hydrogen => Lewis acids

· Each definition has had an impact on our understanding of the nature and properties of acids/bases and their reactions

· Lavoisier’s definition of acids was wrong, but did increase awareness of the need to define an acid

· Davy’s definition helped classify substances as acidic, and directed scientific thinking on acids towards the study of hydrogen’s role

· Arrhenius’s definition helped interpret acid properties in terms of the hydrogen ions produce, though his definition of bases was limited => this was a significant increase in our understanding of acids, and in the development of the concept of an acid

· Brӧnsted-Lowry’s definition further increased our understanding by allowing for more accurate quantitative analysis of acids (e.g. pH, treatment of acid/base equilibriums). It also changed scientific thinking by demonstrating the importance of the solvent in acid/base reactions, and showed that the acidic and basic salts were due to simple acid or base reactions.

Outline the Brӧnsted-Lowry theory of acids and bases

· The Arrhenius definition of acids/bases was limited for the following reasons:

· Did not explain why metallic oxides and carbonates were bases, despite their ability to neutralise acids

· Did not explain acidic salts (e.g. zinc chloride) or basic salts (e.g. sodium sulfide)

· Did not explain why anhydrous neutralisation reactions would occur (e.g. hydrochloric acid dissolved in benzene and ammonia reacting to produce ammonium chloride)

· Did not explain why some substances can act as both an acid and a base

· NEW UNDERSTANDINGS IN ACIDS AND BASES LED TO THE DEVELOPMENT OF THE BRÖNSTED-LOWRY THEORY

· The Brӧnsted-Lowry theory of acids states:

· Acids are proton (H+) DONORS

· Bases are proton (H+) ACCEPTORS

· A neutralisation reaction is a proton-transfer reaction

· The Brӧnsted-Lowry theory applies to non-aqueous environments (e.g. non-aqueous solvents, gas-phase reactions) => acids and bases can be solids, gases, or anhydrous solutions

EXAMPLE

· Under the Arrhenius definition of acids, the above reaction would not be classed as an acid/base reaction, as there are no free H+ ions present

· But under Brӧnsted-Lowry theory, the above reaction is classed as an acid/base reaction, as HCl had donated protons, whilst NH3 has accepted protons

· Thus HCl has acted as an acid in an anhydrous environment, and NH3 has acted as a base => IMPROVEMENT ON THE PREVIOUS DEFINITION ON ACIDS AND BASES

· The Brӧnsted-Lowry theory also has implications for the role of the solvent

· When an acid is ionised in water, water acts as a Brӧnsted-Lowry base, as it accepts protons. For example…

· When a base is ionised in water to produce an alkali solution, water acts as a Brӧnsted-Lowry acid, as it donates protons

· In the self-ionisation of water, water acts as both an acid and a base

·

Describe the relationship between an acid and its conjugate base and a base and its conjugate acid

· Under Brӧnsted-Lowry theory, an acid donates a proton in a reaction, resulting in the acid species becoming deprotonated

· The deprotonated acid can then accept protons, hence it can act as a base

· Similarly, a protonated base species can donate protons, hence act like an acid

· This concept is known as conjugate acids and bases

· A CONJUGATE BASE is the original acid with a hydrogen ion removed

· A CONJUGATE ACID is the original acid with a hydrogen ion added

· Conjugate means ‘linked with’ => under Brӧnsted-Lowry theory, each acid has a conjugate base, and each base has a conjugate acid

· By convention, conjugate acids and conjugate bases are on the right-hand side of an acid-base equation

EXAMPLES

· H2SO4 is the acid; its conjugate base is HSO4- (H2SO4 minus a proton)

· HCl is the acid; its conjugate base is Cl- (HCl minus a proton)

· H2O is the base; its conjugate acid is H3O+ (H2O plus a proton)

· NH3 is the base; its conjugate acid is NH4+ (NH3 plus a proton)

· H2O is the acid; its conjugate base is OH- (H2O minus a proton)

RELATIVE STRENGTHS

· Generally…

· The stronger the acid, the weaker its conjugate base

· The stronger the base, the weaker its conjugate base

· The relationship between an acid and its conjugate base can be considered in the equilibrium below

· If HA is a strong acid, then the equilibrium will strongly favour the forward reaction

· Thus the reverse reaction will only very weakly proceed (i.e. A- will not accept protons well) => A- is thus a weak conjugate base

· Similarly for a base and its conjugate acid…

· If A- is a strong base, the equilibrium favours the forward reaction

· Thus reverse reaction does not proceed well (i.e. HA does not deprotonate well) => HA is thus a weak conjugate base.

· The diagram below illustrates the relationship between the strengths of an acid or base and its conjugate

Identify conjugate acid/base pairs

· Conjugate acid/base pairs can be easily identified from a chemical reaction

· Remember…

· A conjugate base is the original acid minus a proton

· A conjugate acid is the original base plus a proton

· Conjugate acid/base pairs can be identified by the following method:

1. Identify the proton donor and proton acceptor in the forward reaction

2. Classifying accordingly the reactants as B-L acids and bases

3. Identify the conjugate acid and base

4. Record the acid/base pair

EXAMPLES

· Acid/base pairs are H2CO3/HCO3- and HS-/S2-

· Acid/base pairs are H2O/OH- and H3O+/H2O

· Below is a table of common acid/base pairs

Identify a range of salts which form acidic, basic and neutral solutions and explain their acidic, neutral or basic nature

· Definitions…

· A salt is an ionic compound produced by a neutralisation reaction, consisting of an anion and a cation

· The hydrolysis of a salt is the reaction of a salt and water, producing a pH change

· Under Brӧnsted-Lowry theory, any ion can act as an acid or a base, and react with water

· Consequently, many salts can form acidic or basic solutions in water

· The pH of a salt in solution depends on the relative strengths of the anion and cation as either acids or bases

· The acidity or alkalinity of a salt can be predicted by considering the reactants of the neutralisation reaction producing the salt

ACIDIC SALTS

· Occurs when the cation’s acidity is greater than the anion’s alkalinity

· The cation reacts with water to a greater extent than the anion, thus producing more H3O+ than OH- => the pH of the solution decreases, thus the salt undergoes hydrolysis

· Acidic salts result when a strong acid reacts with a weak base

· The cation from the base acts as a weak acid by reacting with water to produce H3O+, whilst the anion reacts to a much lesser extent, if at all

EXAMPLES

· NH4Cl (ammonium chloride) is an acidic salt, as the NH4+ ion acts as a weak acid, whilst the Cl- does not react

· ZnSO4 (zinc sulfate) results from the reaction of zinc hydroxide and sulfuric acid. The hydrated zinc ions act as a weak acid when dissolved in water, whilst the sulfate ions do not react.

BASIC SALTS

· Occurs when the anion’s acidity is greater than the cation’s alkalinity

· The anion reacts with water to a greater extent than the cation, thus producing more OH- than H3O+ => the pH of the solution increases, thus the salt undergoes hydrolysis

· Basic salts result when a weak acid reacts with a strong base

· The anion from the acid acts as a weak base by reacting with water to produce OH-, whilst the cation reacts to a much lesser extent, if at all

EXAMPLES

· KF (Potassium fluoride) is produced from potassium hydroxide and hydrofluoric acid. The fluoride ion acts as a weak base in water, whilst the potassium ion does not react

· NaCH3COO (sodium acetate) is produced from sodium hydroxide and acetic acid. The acetate ion acts a weak base in water, whilst the sodium ion does not react

NEUTRAL SALTS

· The pH of a neutral salt in water is 7

· Neutral salts can be produced from the reaction of…

· A strong base and a strong acid (e.g. NaCl, Ba(NO3)2) => neither the cation or anion of the salt reaction with water sufficiently to alter its pH => the ions of such salts do not hydrolysis

· The anions of strong acids are all the halide ions (except F-), and strong oxyanions such as NO3- and ClO4-

· The cations of strong bases are those from Group 1 and Ca2+, Sr2+, and Ba2+ from Group 2

· A weak base and a weak acid (e.g. NH4CH3COO) => the cation and anion react with water to approximately same extent, thus causing no net change in pH => the ions of such salts hydrolysis to the same extent

· NOTE: The pH of a weak acid/weak base salt MUST be tested to verify its neutrality, because the two ions must be hydrolysis to the same extent => any imbalance will cause the pH of the salt solution to change

· In summary…

· A strong acid and a strong base will produce a neutral salt

· A weak acid and a weak base will generally produce a neutral salt

· A strong acid and a weak base will produce an acidic salt

· A weak acid and a strong base will produce a basic salt

Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions

· An amphiprotic substance is one that can act as BOTH a proton donor and a proton acceptor, i.e. can act as both a Brӧnsted-Lowry acid and base

· They act as an acid or base depending on the conditions of the reaction, usually the acidity of other substances in the reaction

EXAMPLES

· Water (H2O) is amphiprotic, as seen in its self-ionisation

· If a substance (e.g. HCl) has a greater tendency to lose protons than water, then water acts like a base, producing an acidic solution

· If a substance (e.g. NH3) has a lesser tendency to lose protons than water, then water acts like an acid, producing a basic solution

· The hydrogen carbonate ion (HCO3-) is also amphiprotic

· When placed in an alkaline solution, HCO3- acts as an acid

· When placed in an acidic solution, HCO3- acts as a base

· In the above reaction, the H2CO3 decomposes to H2O and CO2, thus carbon dioxide gas evolves

· The hydrogen sulfite ion (HSO3-) acts in a similar way to HCO3-, releasing chocking sulfur dioxide gas when it reacts with a hydroxide ion

· Hydrogen phosphate (HPO4-) also acts similarly, except that it forms a slightly basic solution when added to water

Identify neutralisation as a proton transfer reaction which is exothermic

· Neutralisation reactions are reactions between acids and bases

· Under Brӧnsted-Lowry theory, and acid is a proton-donor and a bases is a proton-acceptor, thus in a reaction of acids and bases, a proton is transferred from and acid to a base

EXAMPLE

· The reaction of hydrochloric acid and sodium hydroxide is a neutralisation reaction:

· Whilst a proton transfer may not be immediately evident, by considering the net ionic equation…

· The above equation is the same for all Arrhenius neutralisation reactions (i.e. hydronium ion and hydroxide ions)

· We can see that a proton has been transferred from the hydronium ion to the hydroxide ion, thus it is a proton transfer reaction

· Generally, all neutralisations are proton transfer reactions (see below for more examples)

· Also note that the above reaction is exothermic => the change in enthalpy for all neutralisations is around -57kJmol-1, depending on the strength of the acid or base (a weak acid/base won’t fully ionise in solution)

OTHER EXAMPLES

· A proton is transferred from the hydronium ion to the ammonium molecule

· Note that the enthalpy change is less than the reaction above => acetic acid is a weak acid, so does not ionise 100% in solution

Analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills

· It is important to safely clean up any spills involving acids and bases, because they are corrosive, thus can damage equipment or the environment, and pose a serious safety risk

· If an acid or base is spilt on skin, first aid must be administered immediately by washing the skin with copious volumes of water

· If an acid or base is spilt on the floor, the following procedure should be followed:

· The area first needs to be isolated to minimise further damage and the possibility of inhaling toxic fumes => this involves spreading sand around the spill site to minimise the spill from spreading

· The spill should then be cleaned up first by neutralising the acid/base, then once safe, the spill should be cleaned up with paper towels and disposed

· The preferred chemicals used in neutralising spilt acids/bases are stable (safely handled), solid in a powdered form (easily transported and stored), cheap, and amphiprotic (so it is weak, and can clean up both acid and base spills)

· The substance most commonly used to spill up laboratory acid/base spills is sodium hydrogen carbonate (NaHCO3)

· The HCO3- ion is amphiprotic, cheap, and readily available

· Additionally, the neutralisation reaction with an acid produces CO2 gas, thus the fizzing can be used to monitor the progress of the neutralisation reaction

· Thus NaHCO3- has been very effective as a safety measure or to minimise damage in accidents or spills

· It is very important that the chemical used to neutralise spills is a weak acid/base, as neutralisation reactions are exothermic, so the quantity of heat produce needs to be minimal.

· In addition, if excess chemical is used, the spill will be weak, thus does not pose further safety risks => if excess of a strong acid/base is added, the spill will become dangerous again.

· NOTE: As proton transfer reactions are exothermic, special care must be taken when diluting acids => a small volume of concentrated acid should be added to a large volume of water, so that the water can dissipate the generated heat. If water is added to acid, the acid’s temperature can quickly reach boiling point, which could be ejected from a beaker and pose a very serious safety issue.

Qualitatively describe the effect of buffers with reference to a specific example in a natural system

· Buffers are solutions that resist changes in pH when small quantities of acids or bases are added to them

· A buffer solution contains comparable quantities of a weak acid and its conjugate base, which exist in the following equilibrium

· By considering Le Chatelier’s principle, we can see that the equilibrium can resist changes in pH

· If an acid is added, the concentration of H3O+ is added, then the concentration of H3O+ ions increases. The equilibrium resists the change by favouring the reverse reaction to reduce the concentration of H3O+, thereby maintaining the original pH.

· If a base is added, it reacts with the H3O+ ions to produce water, thus increasing the concentration of water and decreasing the concentration of H3O+. The equilibrium resists the change by favouring the forward reaction, thus the pH of the solution is maintained.

· As the buffer contains comparable quantities of the weak acid and its conjugate base, the system can resist large changes in pH.

· The action of buffers can be observed by considering the titration curve of a weak acid and a strong base, and the corresponding titration curve of a weak base and a strong acid, as the buffer zone lies in the flat region of the curve

EXAMPLE OF A NATURAL BUFFER

· System: The carbonic acid system in freshwater lakes and rivers

· The carbonic acid molecules are formed when carbon dioxide dissolves directly in the water, or from the dissolved carbon dioxide in rain water

· The hydrogen carbonate ions are leeched out of rocks and minerals in the lake

· The additional HCO3- pushes the above equilibrium to the left, thus raising the pH to between 6.5 and 7.5.

· The buffer resists the addition of acids (such as acid rain) or bases, thus a neutral pH is maintained.

· For example, the absence of the above buffer in Scandinavian lakes due to the absence of carbonate rocks (thus no additional HCO3-) led to these lakes being the first to detect a falling pH from acid rain

· The natural carbonic acid buffer in lakes and rivers is important as marine life requires a neutral pH to live => for example, fish eggs are destroyed if water is acidic.

· Other possible buffers are shown below:

Describe the correct technique for conducting titrations and preparation of standard solutions

BACKGROUND THEORY

· Titration, also known as volumetric analysis, is a chemical technique used to determine the concentration of an unknown solution through experimental first-hand data

· It is a technique of quantitative chemical analysis, which requires high degrees of accuracy

· The concentration of an acid or a base can be determined by titration, using a neutralisation reaction

· A certain volume of and acid/base of unknown concentration is slowly reacted with another solution of known concentration and volume, until the endpoint is reached

· By measuring the volume of the acid/base that was required to reach the endpoint, the concentration of the unknown solution can be calculated

· It is important to distinguish between the equivalence point and endpoint of an acid-base titration

· At the equivalence point, equal volumes of the two reactants have been reacted to cause the complete consumption of both reactions (which depends the stoichiometry of the reaction)

· At the endpoint, there is a permanent colour change in the indicator

· For titration, the equivalence point and endpoint should coincide

EXAMPLE

· Titration of sulfuric acid against a standardised 0.100molL-1 sodium hydrogen carbonate

· First write the relevant neutralisation reaction:

· Next determine the molar ratio of reactants

· For every 1 mole of acid, there are 2 moles of base

· Perform the titration with a methyl orange indicator, and determine the volume of NaHCO3 required to react with H2SO4

· The endpoint was reached after 25.00x10-3L of NaHCO3 was titrated against 28.35x10-3L of H2SO4

· Ideally, repeat the titration as many times as possible, discount outliers, and average the reliable results

· Calculate the unknown concentration

· There are many rules to follow to achieve a highly accurate titration, including the preparation of a standard solution, indicator choice, and the titration itself

PREPARATION OF STANDARD SOLUTION

· A standard solution (also known as titrant) is a solution of accurately known concentration

Selecting an appropriate standard solution

· We cannot accurately determine the concentration of many common acids or bases (e.g. HCl, H2SO4, NaOH), as they’re concentration changes over time for various reasons

· We can accurately determine the concentration of such acids or bases through a titration against a primary standard (a solution prepared from dissolution, which has an accurately known concentration)

· To prepare a primary standard, we first need to use a suitable substance that will not cause error in the titration

· The criteria for choosing a primary standard are…

· High level of purity

· Accurately known composition

· High molecular (to reduce errors associated with weighing)

· Free from moisture

· Stable and unaffected by air in weighting

· Readily soluble in pure (distilled) water

· Reacts instantaneously and completely

· Non-hydroscopic (does not absorb water from surroundings) and non-efflorescent (does not release water to surroundings)

· The table below lists some suitable substance for use as a primary standard

Acid standard

Base standard

Potassium hydrogen phthalate (KHC8H4O4)

Sodium carbonate (Na2CO3)

Benzoic acid (C7H5O2)

Sodium hydrogen carbonate (NaHCO3)

Oxalic acid (H2C2O4.2H2O)

Borax (Na2B4O7.10H2O)

· Other common acids/bases are unsuitable, as they don’t fit the above criteria

· HCl is efflorescent

· H2SO4 is severely hydroscopic

· NaOH reacts with gases in the air

· Na2CO3.nH2O contains moisture

Preparing the standard solution

· To prepare a primary standard, the procedure must be followed with high accuracy and precision

· Ensure that all glassware is clean, and rinsed with distilled unless indicated otherwise

· In the method below, analytical-grade anhydrous sodium carbonate (Na2CO3) is used (99.9% purity after roasting for 30 mins at 150°-180°C in a drying oven, then cooled to form crystals in desiccator)

· Calculate the mass of primary standard required to achieve a desired molarity. The result is a guide only; it is the final concentration that it is important.

· Weigh the anhydrous sodium carbonate (1.325g) in a dry 50mL beaker on an electronic scale. Remember to first place the beaker on the scale, and then zero it.

· Use a plastic wash bottle to add a little distilled water to the beaker. Stir the solution with a fine, short glass rod to dissolve the sodium carbonate, making sure that the glass rod does not touch foreign substances. Ensure all solid has dissolved.

· Transfer the sodium carbonate solution from the beaker into a clean 250mL volumetric flask by placing a small, clean glass funnel over the neck of the flask. Rinse the beaker, glass rod, and funnel into the volumetric flask with the wash bottle to ensure that all the sodium carbonate has been transferred.

· Fill the flask with distilled water until it reaches about 1cm below the graduation mark. Continue to fill the flask up to the graduation mark drop-by-drop using a dropper or Pasteur pipette until the bottom of the meniscus is aligned with the graduation mark. Avoid parallax error (error can be minimised by marking a piece of paper with a black line and placing it at the graduation mark to aid locating the graduation mark.

· Stopper the flask, and invert the flask 10 times to ensure a homogenous solution. Label the flask as 0.0500molL-1 Na2CO3.

TITRATION

· First we’ll define relevant terms and equipment

· Aliquot – A known volume of a liquid

· Titrant – A solution that is added from the burette (typically standard solution)

· Burette – A piece of cylindrical glassware, held vertically, with volumetric divisions on its full length and a precision tap (stopcock) on the bottom. It is used to dispense precise volumes of the liquid reagent, with the volume dispensed being the difference between volumetric divisions. Burettes are very precise, with accuracy up to ± 0.05mL.

· Pipette – A glass tube used to transfer precise volumes of liquid reagants. They are usually dsigned to trnasfer a set volume, such as 25mL. The reagant is drawn up the pipette using a pipette filler (e.g. a rubber bulb)

· Volumetric flask – A glass flask with a long neck with a graduation mark to measure specific volumes. Volumetric flasks are used to prepare and hold standard solutions

· Conical flask – Typically used to hold the reactants during titration. Its shape prevents the reactants from spilling as they are swirled together.

· The diagram below shows the above glassware. The flask with the yellow stopper is the volumetric flask, and then clockwise is the burette, pipette and bulb, and the conical flask.

· Next we must choose an appropriate indicator

· Recall that not all salts are neutral

· A strong acid and a weak base produce an acidic salt

· A weak acid and a strong base produce a basic salt

· Hence at the equivalence point, the pH of the solution may not be neutral

· To ensure the equivalence point and endpoint align, we must choose an appropriate indicator

· The table below summarises the commonly used indicators

Indicator

Initial Colour

Range

Final colour

Diagram

Methyl orange

Red

3.1-4.4

Yellow

Bromothymol blue

Yellow

6.0-7.6

Blue

Phenolphthalein

Colourless

8.2-10.0

Pink

· Consider the titration curve for a strong acid/ strong base titration

· There is a steep increase in pH from 3 to 11 within a range of about a drop. Thus any of the above indicators could be used, as they would all register the pH change. The equivalence point lies at around a pH of 7, so use bromothymol blue if available, but methyl orange or phenolphthalein can be used too.

Strong acid/weak base

· The equivalence point lies at around 5.3 for HCl/NH3, so the most suitable indicator would be methyl orange. Bromothymol blue could also be use, but phenolphthalein cannot be used, as its pH transition range is outside the steep increase in pH shown above

Weak acid/strong base

· The equivalence point lies at around 8.7 for NaOH/CH3COOH, so phenolphthalein is the most suitable indicator, but bromothymol blue could also be used too. Methyl orange cannot be used, as its pH transition range lies outside the steep increase in pH shown above.

Weak acid/weak base

· As shown above, there is no sharp change in pH for a weak acid/weak base reaction, thus these reactions should be avoided in titrations.

· pH probes and data loggers can also be used in titrations to determine the equivalence point by analysing the titration curves similar to those above.

· Now we must rinse the glassware appropriately

· The volumetric flask (including stopper) should be rinsed thoroughly with distilled water. The flask can be left wet, as water is going to be added to it later.

· The conical flask should be rinsed thoroughly with distilled water, and left wet for similar reasons

· The burette should first be rinsed with distilled water by filling the burette, then opening the tap to clean the tip. Fill the burette again with distilled water, swirl the water to rinse the sides, then poured out the top. Repeat this THREE times, and then do it once more with the solution it is going to contain.

· The pipette should be rinsed three times with distilled water, then once with the solution it is going to contain.

· Now we can set up the titration

· Using a funnel, pour the titrant into the burette until above a suitable marking (e.g. 30mL) to facilitate measurements. Open the tap slightly until the meniscus sits just above the mark. Holding a white card with a black line makes the meniscus clearer.

· Using the pipette, transfer a fixed volume (typically 25mL) of the other reagent into the conical flask. UNDER NO CIRCUMSTANCES SHOULD THE FINAL DROP OF THE PIPETTE BE EXPELLED => pipettes are designed with this feature in mind.

· Add three drops of the appropriate indicator to the solution in the conical flask

· Set up the burette above the conical flask using a retort stand and clamp, and place the conical flask on a white tile to make the indicator colour clear

· Now we can conduct the titration

· First conduct a trial run, where the volume required to reach the endpoint is roughly determined by running the burette at a steady speed until the endpoint is reached. Keep this flask so that the final colour can be compared to in subsequent titrations

· With one hand, open the tap so that the titrant runs steadily into the conical flask. With your other hand, swirl the flask to homogenise the solution.

· Once within 10mL of the expected volume, close the stopcock. Wash down the conical flask with a wash bottle as the solution may have splashed. Open the stopcock slightly, so that the titrant slowly drips into the conical flask.

· Stop as soon as the endpoint is reached, and record the final volume reading on the burette. If unsure if the indicator has changed colour, record the volume, then add drop-by-drop, recording each reading until the endpoint is definitely reached. The diagram below shows methyl orange just before, at, and just after the endpoint has been reached.

· Repeat the titration as many times as possible, keeping each flask for colour comparison. For increased accuracy, add partial drops from the burette by quickly turning the stopcock on and off so that a drop hangs from the tip, then resting the drop on the conical flask.

· We now calculate the unknown volume

· Compare the readings from each titration => the volume of titrant used is calculated by subtracting the final reading of the burette from the initial reading.

· Discount any outliers, and ideally average the three closest results.

· Use the average value in subsequent calculations. The following formula can be used if the molar ratio is 1:1

Choose equipment and perform a first-hand investigation to identify the pH of a range of salts

METHOD

· The pH of various salts was determined by using a pH meter, and then confirmed by using universal indicator

· The salt solutions tested were sodium carbonate, sodium hydrogen carbonate, sodium chloride, ammonium chloride, and ammonium sulfate.

· All salt solutions were 0.1M, and 20mL of each salt was tested

· SAFETY: Ammonium chloride releases toxic fumes if heated => keep away from flame. Ammonium chloride, sodium acetate, and sodium carbonate are slightly toxic if ingested and are skin irritants => wear eye protection, use a dropper when transferring solutions

RESULTS

Salt (0.1M)

pH

Colour

Na2CO3

11-12

Dark blue

NaHCO3

8-9

Blue

NaCl

6.5-7.5

Green

NH4Cl

5-6

Orange

(NH4)2SO4

4-5

Light pink

VALIDITY/RELIABILITY/ACCURACY

· The results confirmed expected results

· The use of technology (pH probes) provided more accurate quantitative analysis, and more reliable results

· The results may have been inaccurate due to the inaccuracies of the concentration of solutions prepared in the school lab.

Perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases

METHOD

· See above for extensive notes on preparing standard solutions and conducting titrations

· A standard solution of sodium carbonate (Na2CO3) was prepared, and was titrated against an HCl solution to determine the concentration of HCl

· SAFETY: The concentration of HCl is unknown, thus it could be of high concentration => clean up spills immediately, wear safety glasses, use a pipette to transfer HCl solutions

RESULTS

· The following reaction occurred in the titration

· m(Na2CO3)=14.56g

· [Na2CO3] = 0.549molL-1 (Na2CO3 was dissolved in a volumetric flask)

· 22.3mL of Na2CO3 was required for 25mL of HCl solution to reach endpoint (methyl orange indicator was used)

· Thus the concentration of HCl solution was 0.98molL-1

VALIDITY/RELIABILITY/ACCURACY

· See titration notes above for a detailed account of conducting valid titrations

· All glassware was rinsed appropriately prior to use

· Electronic scales were used for accuracy, providing mass up to 2 decimal places

· Multiple titration runs were conducted, the first one discounted as a rough guide, and any outlier results were discounted before taking an average => increases reliability

Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies

METHOD

· The concentration of household vinegar was determined through titrations involving computer-based technologies

· See above for extensive notes on titration

· A primary standard of potassium hydrogen phthalate was prepared accurately

· The potassium hydrogen phthalate was titrated against a prepared solution of NaOH to produce a secondary standard solution

· The NaOH secondary standard was titrated against the vinegar solution to determine the concentration of vinegar.

· A pH probe and data logger were used in both titrations to determine the equivalence point, and the pH probes were calibrated using pH 4 and 7 buffers

RESULTS

· The data logger produced titration curves similar to those in the titration notes above (weak acid/strong base), from which the equivalence point could be determined

· The concentration of vinegar was calculated to be 0.79M (see above for notes on calculations)

JUSTIFICAITON OF METHOD

· A secondary standard of NaOH was produced to increase the reliability of results, as weak acid/ weak base titrations provide unreliable results

· All glassware was rinsed appropriately prior to use

· Trial runs were conducted for both titrations, and average results were taken, discounting outliers

5. Esterification is a naturally occurring process which can be performed in the laboratory

Describe the difference between the alkanol and alkanoic acid functional group in carbon compounds

· Recall that a functional group is a group of atoms responsible for particular characteristics of a compound

· Alkanols contain the hydroxyl group (-OH) attached to a carbon atom

· For example, below is the structural formula of ethanol

· Alkanoic acids contain the carboxyl group (-COOH) attached to a carbon atom

· One of the oxygen molecules is double-bonded to the central carbon, whilst the OH group is single-bonded to the carbon (shown below)

· NOTE: The –OH group in the carboxyl group is NOT a hydroxyl group

· The hydrogen atom of the carboxyl group can ionise in solution, hence the group is acidic

· For example, below is the condensed structural formula of ethanoic acid

Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures

· The melting/boiling point of a molecular substance reflects the strength of the intermolecular bonds within a given substance

· The stronger the intermolecular bonds, the more energy is required to weaken (melting) or break (boiling) the bonds, thus the melting/boiling point is higher

ALKANOLS

· Alkanols contain the hydroxyl (-OH) group an C-O bonds, both of which are polar

· This allows for strong hydrogen bonds to be formed between alkanol molecules, in addition to dispersion forces

· Alkanols thus have a higher melting/boiling point than many carbon compounds of similar mass, such as alkanes

ALKANOIC ACIDS

· Alkanoic acids have a higher molecular mass than their corresponding alkanol, since the carboxyl group has a higher mass than the hydroxyl group

· Thus alkanoic acids have stronger dispersion forces between molecules than the corresponding alkanol

· Alkanoic acids also have more extensive hydrogen bonding than alkanols

· The C=O bond in alkanoic acids is also polar in addition to the C-O and O-H bonds

· This allows two hydrogen bonds to be formed per molecule, compared to 1 hydrogen bond per alkanol molecule

· Thus alkanoic acids have stronger intermolecular forces between molecules than alkanols of similar molecular mass

· As a result, alkanoic acids have higher melting/boiling points than alkanols of similar size, which in turn have higher melting/boiling points than similarly-sized alkanes and alkenes

· In summary, alkanoic acid > alkanol > ester > alkane

Identify esterification as the reaction between an acid and an alkanol and describe, using equations, examples of esterification

· Generally, esterification is the reaction between an acid and an alcohol to produce an ester

· In the HSC course however, we only consider the reactions between straight-chained alkanoic acids, and straight-chained primary alkanols => these produce carboxylate esters

· Note that esterification is NOT an acid-base reaction => the –OH group of the acid reacts with the hydrogen atom of the hydroxyl group of the alkanol.

· Carboxylate esters contain the -COOC- structural unit:

· Generally…

· The equilibrium at room temperature lies much to the left at room temperature

· See below for an example of esterification, and naming of alkyl alkanoates

Identify the IUPAC nomenclature for describing esters produced by reactions of straight-chained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8

· First let’s review the IUPAC nomenclature for describing alkanoic acids and alkanols

ALKANOIC ACIDS

· Count the number of carbons, take the parent alkane name, drop the final –e, and add on the suffix ‘-oic acid’

· Since the –COOH contains only one free carbon valence, no number is needed to signify its location

EXAMPLE

· There are three carbon atoms, thus the parent alkane is propane. By dropping the suffix –e and adding –oic acid, we have ropanoic acid

ALKANOLS

· Count the number of carbons, take the parent alkane name, drop the final –e, and add the suffix ‘-ol’

· Remember to number the carbon where the hydroxyl group is by counting from the end that gives the smallest number

EXAMPLE

· There are three carbons, thus the parent alkane is propane. By dropping the ‘-e’, adding –ol, and number the carbon atom where –OH lies (2), we have 2-proponol

ESTERS

· Esters from alkanols and alkanoic acids are named as alkyl alkanoates

· The alkyl comes from the alkanol, and the alkanoate comes from the alkanoic acid => hence the alkanol is named first, then the alkanoic acid

· Esters can be named by considering the reaction that produces the reaction, or the given structural formula of an ester

· If the alkanol and alkanoic acid are given, name the ester by the following procedure:

· Replace the –anol of the alkanol with –yl (i.e. the alkyl is named according to the corresponding alkanol)

· Replace the ‘-oic acid’ of the alkanoic acid with ‘-oate’

· Place the two words together

EXAMPLE

· From the 1-proponol we get ‘propyl’, and from the pentanoic acid we get ‘pentanoate’, thus the name of the resulting ester is propyl pentanoate

· NOTE: The preferred IUPAC name for methanoic acid and ethanoic acid are formic acid and acetic acid respectively. Thus methanoate may sometimes be named formate, and ethanoate may sometimes be named acetate.

· If the structural formula is given, follow the following method:

· Split the ester across the C-O-C bond => the side with the C=O bond is the alkanoate, the side without is the alkyl group

· Count the number of carbons in the alkyl and alkanoate groups respectively, and name similarly to above

EXAMPLE

· We can see that the alkyl group is to the right of the C-O-C bond, whilst the alkanoate is to the left

· There are three carbons in the alkyl group, hence it is named propyl

· There are five carbons in the alkanoate group, hence it is named pentanoate

· Thus the name of the ester is propyl pentanoate

· NOTE: As the dot point only asks for the reactions involving straight-chained primary alkanols, numbering is not important for esters encountered in the HSC course

Describe the purpose of using acid in esterification for catalysis

· Concentrated sulfuric acid is often added as a catalyst in esterification (1-3mL of concentrated acid)

· As a catalyst, concentrated sulfuric acid is not consumed in the reaction, but allows the equilibrium point to be reached faster (esterification is naturally a slow reaction)

· Concentrated sulfuric acid is added because…

· It speeds up the rate of reaction by lowering the activation energy of the forward reaction, thus the equilibrium point is reached faster

· Sulfuric acid is a dehydrating agents, thus the equilibrium shifts to the right, and this increases the yield of ester

Explain the need for refluxing during esterification

· Refluxing involves heating up a mixture with a cooling condenser placed above, so any volatile reactants and products are condensed and returned to the reaction mixture => this reduces loss of substance

· A refluxing apparatus consists of a water condenser mounted on top of a reaction vessel (normally a round-bottomed flask). The circulating cold water cools any vapours, causing them to condense and run back to the reaction vessel.

· The refluxing apparatus is open to the atmosphere, so there is no pressure build-up inside the apparatus

· Heating is desirable in esterification because…

· The higher temperatures increase the rate of reaction, thus the equilibrium is reached faster

· Esterification is endothermic in the forward reaction, so a higher yield of ester is obtained if the mixture is heated

· The reactants and products of esterification are volatile however, so heatin