the acidic environment - acehsc · web viewtesting acidity or alkalinity of soils. optimum ph range...

THE ACIDIC ENVIRONMENT HSC CHEMISTRY

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

An acid is a substance which in a solution produces hydrogen ions, H+ or more strictly H3O+ - the hydronium ion.

- The H+ ions are available in solutions where the minute proton is stabilised by association with a solvent molecule

- Common laboratory acids are hydrochloric acid (HCl), sulfuric acid (H2SO4), nitric acid (HNO3). An organic acid is acetic acid (CH3COOH)

- Common properties of acids: a sour taste, sting/burn skin, conducts electricity in aqueous solutions, and turns blue litmus red.

A base is a substance which either contains the oxide O2- or hydroxide ion OH- or which in solution produces the hydroxide ion. A soluble base is called an alkali

- Common alkalis include sodium hydroxide (NaOH), ammonium hydroxide (NH4OH), and calcium hydroxide (Ca(OH)2).

- Other bases include magnesium oxide (MgO), ammonia (NH3) and copper hydroxide (Cu(OH2))- Common properties of alkalis: a soapy feel, a bitter taste, good conductors of electricity in solution

and turn red litmus blue

Neutralisation is the reaction of an acid with a base- Acid + Base = Water + Salt- A salt is an ionic compound formed when a base (alkali) reacts with an acid- For example, sodium hydroxide reacts with hydrochloric acid to form the salt sodium chloride and

water. HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

The NaOH reacts with the H+ ion in HCl. The Cl- anion remains unaltered, so it is called a spectator ion.

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 (usually a vegetable dye) that in solution changes colour depending on whether the solution is acidic or alkaline

- One of the first indicators was litmus, a dye extracted from lichens - An highly effective natural indicator is red cabbage solution (see practical below)- Litmus and bromothymol blue shows whether the solution is acidic, neutral or basic. Methyl orange

and phenolphthalein shows whether the solution is highly or slightly acidic or alkaline respectively.Colour 1 Colour 2HIn H+ + In-

- Increase H+, solution more acidic favours reverse (left) reaction to produce more HIn


- Decrease H+, solution more basic favours forward (right) reaction, hydroxide ions react with H+ to produce more In-

Indicator Acidic Transition Colour Basic pH range Methyl Orange Red Orange Yellow 3.2 – 4.4Bromothymol Blue Yellow Green Blue 6.0 – 7.6Litmus Red Purple Blue 5.0 - 8.0Phenolphthalein Colourless Light Pink Pink 8.2 – 10.0 Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity.

- Testing Acidity or alkalinity of soils. Optimum pH range of soil is 5.5 – 7.5. Add a thin layer of neutral white powder, such as barium sulfate, on top of moist soil. Add a few drops of universal indicator

- Testing home swimming pools. pH needs to be close to 7.4. Take sample of water and test with an indicator (phenol red, pH range 6.8 – 8.4)

- Monitor how acidic or alkaline waste solutions are from industries (discharges to sewerage system must be nearly neutral)

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

- Chop up one leaf of red cabbage and place into a beaker of 50ml distilled water and bring to boil.- Obtain 12 test tubes and put them in a test tube holder. The first 6 test tubes will be used for

different concentrations of HCl, and the last six for different concentrations of NaOH- Using a measuring cylinder, measure out 10ml of HCl and pour into the first test tube- Use a pipette to transfer 1 ml of HCl out of the first test tube into the second test tube- Using a measuring cylinder, pour 10ml of distilled water into the second test tube- Repeat up to test tube 6. Repeat for NaOH- Use a pipette to place 3 drops of beetroot solution to test substances of varying pH

Results: Acidity / (Substance) Colour

Strong Acid (HCl) RedWeak Acid (vinegar) Pink

Neutral (distilled water) PurpleWeak Base (baking soda) Blue/Green

Strong Base (NaCl) Yellow

Identify data and choose resources to gather information about the colour changes of a range of indicators.Solve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basic.

Substance Acid/Neutral/Base Substance Acid/Neutral/BaseDetol Antiseptic Base Lemon AcidBleach Base Antifreeze NeutralVinegar Acid Morning Fresh Washing AcidShampoo Acid Dishwasher rinse aid Base


Sodium hydroxide Base Ammonia Base Hydrochloric acid Acid Milk AcidHandwashing Liquid Base Block soap Base

2. While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulphur. The concentrations of these 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.

Oxides often displays acidic or basic properties

Acidic oxides - Reacts with water to form an acid and/or- Reacts with bases to form salts- Examples (are all gases) – carbon dioxide (CO2), diphosphrous oxide (P2O3), nitrite, sulfite- They are generally oxides of non-metals, hence they are also known as non-metallic oxides

Basic oxides - Reacts with acids to form salts but- Does not react with alkali solutions - If in a solution, it produces a basic solution. - Examples (are all solids) – copper oxide, iron (III) oxide, sodium oxide, magnesium oxide- They are generally oxides of metals, hence they are also known as metallic oxides

Amphoteric oxides react with both acid and base. Examples – ZnO, PbO, Al2O3

Neutral oxides do not change pH of water. Examples – CO, nitrous oxide, nitric oxide

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

- Basic oxides Amphoteric oxides Acidic oxides (going from left to right of periodic table) - As we move down the group the oxides become less acidic. - The oxide is more acidic if the element is more electronegative. - Therefore, acidic oxides occur towards the right and top. Basic oxides occur to the left.

Define Le Chatelier’s principle

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

Identify factors which can affect the equilibrium in a reversible reaction.Equilibrium is disturbed if:

- Concentration of one or more species involved is changed- Total pressure acting upon a reaction that involves gases is changed


- Temperature is altered

Note: adding a solid, increasing the pressure of a reaction that doesn’t involve gases, or using a catalyst does not affect the equilibrium

Increase Decrease Concentration Tries to reduce this concentration, so it

shifts of opposite side of the increased component

Tries to increase this concentration, so it shifts to side component was removed from

Pressure Concentration has increased, so it shifts to side with fewer moles of gas

Concentration has decreased so it shifts to side with more moles of gas

Temperature Favours endothermic reaction forward Favours exothermic reaction forward

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

CO2 (g) + H2O (l) H2CO3 (aq) + heat

Disturbance +/- Effect on EquilibriumPressure (total) Increase Favours reaction with few moles of gas, reducing pressure, shifts to right. CO2

dissolves more readily.Decrease Favours reaction with more moles of gas, increasing pressure, shifts to left.

More CO2 gas formed.Temperature Increase Favours endothermic reaction, shifts to left as it is the direction the liberates

heatDecrease Favours exothermic reaction, shifts to right as it is the direction that absorbs

heatConcentration (pressure of CO2)

Increase Reduce concentration of CO2, equilibrium shifts to right to counteract increaseDecrease Increase concentration of CO2, equilibrium shifts to left to counteract decrease

- Pumping gas besides carbon dioxide gas does not affect the equilibrium as the pressure on CO 2

remains unchanged- The solubility of carbon dioxide can also be increased by making the solution alkaline. The added OH -

reacts with carbonic acid H2CO3 (aq) + 2OH-

(aq) H2O (l) + CO32-

(aq)

Carbonic acid is removed from the equilibrium, shifting the equilibrium to the right so more carbon dioxide is formed.

- Freezing soft drink is dangerous because pure ice freezes out of the drink. The carbon dioxide originally dissolved is forced as gas into the small gas space, building up pressure and the bottle may explode. Carbon dioxide has very low solubility as solid ice.

Identify data, plan and perform first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 250C and 100kPa

- Weigh the bottle using an electronic balance (do not remove lid)


- Place the bottle of soda bottle into a beaker of 100 ml of hot (not boiling) water. Undo the lid but do not remove the lid (to prevent vapours from escaping)

- Once the water has cooled down and the bubbling is rapid, remove the lid completely- Take out of beaker when bubbling ceases.- Calculate the mass of carbon dioxide loss by subtracting the final mass from the initial mass.

Calculate the volume of carbon dioxide lost

Identify natural and industrial sources of sulphur dioxide and oxides of nitrogen.

Sulfur dioxide Oxides of NitrogenNatural - Volcanic eruption

- Vegetation decay- Geothermal hot springs

- Lightning strikes (nitric oxide)-Nitrogen fixing bacteria on nitrogenous material in soils (nitrous oxide)

Industrial - Combustion of fuel containing sulfur particles- Extracting metal from sulfide ores

- High temperature combustion eg. Internal combustion, power stations (nitric oxide)- Nitrogenous fertiliser which provides more raw material for the bacteria (nitrous oxide)

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

- Combustion of coal containing metallic sulfides or sulfur in carbon-containing compounds converts sulfur to sulfur dioxide

S (s) + O2 (g) SO2 (g)

- Extraction of metal from metallic sulfides, producing sulfur dioxide along with the metal or metal oxide

2ZnS (s) + 3O2 (g) 2ZnO (s) + 2SO2 (g)

- High temperature combustion (same process as how nitric oxide is formed from lightning)O2 (g) + N2 (g) 2NO (g)

2NO (g) + O2 (g) 2NO2 (g)

Assess the evidence which indicates increase in atmospheric concentration of oxides of sulphur and nitrogen.

- After the Industrial Revolution, there was a great increase in emission of sulfur dioxide in growing industrial cities. Pollution episodes in the ‘50s and ‘60s in USA and London cause many deaths. Regulations were then introduced to control SO2 and particulate emissions.

- Today, sulfur dioxide concentration in cities is significantly greater than in ‘pure’ air, but it is generally below the upper limits for safety recommended by health authorities.

- NOx pollution developed in the twentieth century as electricity generation and use of motor cars expanded dramatically. Photochemical smog became a problem in 1960s. Emission controls on cars were introduced.

- However, increasing city populations and increasing vehicle kilometres driven have cancelled out further benefits of increasingly stringent emission controls.


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

Explain the formation and effects of acid rain.

Acid rain is rain that has higher hydrogen ion concentration than normal – higher than about 10 -5 mol/L. Rain that is unaffected by pollutant is slightly acidic due to carbonic acid formed from dissolved carbon dioxide. Unpolluted rain has a pH between 10-5 mol/L and 10-6 mol/L

Formation - The pollutants SO2 (g) and NO2 (g) are released into the atmosphere. These pollutants mainly come

from the burning of fossil fuels, combustion and some other natural and industrial causes. - They readily dissolve and react with moisture in the atmosphere to produce acidic solutions of

sulphurous acid and a mixture of nitrous and nitric acid.SO2 (g) + H2O (l) H2SO3 (aq)

2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)

- SO2 (g) can also oxidise to form SO3 (g), which then reacts with water to form sulfuric acid2SO2 (g) + O2 (g) 2SO3 (g)

SO3 (g) + H2O (l) H2SO4 (aq)

- HNO2 (aq) is catalytically oxidised to nitric acidHNO2 (aq) + O2 (g) 2HNO3 (aq)

- Acid rain is usually formed in clouds where SO2 (g) and NO2 (g) react.- Precipitation i.e. Rain, fog etc. containing these solutions fall on the earth as acid rain (pH < 5)

Effects - Damage to trees, crops, plants

o decrease pH of soil reducing nutrients availableo removes waxy coating on leaves, affecting ability to photosynthesise

- Affects aquatic lifeo low pH reduced of lakes and ponds: most aquatic life will die at pH < 5o toxic ions released from soil are leaked into waterways

- Damage to buildings and statues, especially marble and limestone (can react with acids produced)CaCO3 (s) + H2SO4 (aq) CaSO4 (s) + H2O (l) + CO2 (g)

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.

Industrial origins:- Sulfur dioxide - Fuel combustion (especially power plants and motor vehicles), processing of metals,

metal processing, garbage incineration, petroleum refineries, food processing, sewage treatment.- Nitrogen dioxide – combutions of fossil fuel in motor vehicles and power stations- Dinitrogen monoxide – manufactured as fuel for racing cars- Nitrogen monoxide- burning of biomass, combustion of fossil fuel in motor vehicles and power

stations


Health effects: - Sulfur dioxide irritates the respiratory system and causes breathing difficulties. Asthma and

emphysema sufferers are particularly susceptible. The presence of particulates magnifies these effects.

- Nitrogen dioxide irritates the respiratory tract and causes breathing difficulties. At very high concentrations, it can cause extensive tissue damage. In the formation of photochemical smog, sunlight acts upon nitrogen dioxide in the presence of hydrocarbons and oxygen to form ozone, peroxyaclynitrates and haze (poor visibility)

Environmental effects:- Emissions of sulfur and nitrogen oxides can lead to the formation of acid rain (see effects of acid

rain)- Sulfur dioxide has an unpleasant odour and is detrimental to health- Oxides of nitrogen forms petrochemical smog

Sulfur dioxide and nitrogen dioxide are formed in localised industralised areas and mining towns, but they do not spread globally because natural processes remove them from the air. However, some mineral-processing areas frequently record sulfur dioxide levels exceeding the limit, so it is important to improve operating procedures, such as trapping sulfur dioxide in flue gases, to improve air quality.


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 proton donors. Although while H+ represents a proton, free H+ does not exist in aqueous solution

- For example, liquid nitric acid undergoes an ionisation reaction when reacted with water, forming an aqueous solution containing hydronium ions, H3O+, and nitrate ions, NO3

-

HNO3 (l) + H2O (l) H3O+ (aq) + NO3

- (aq)

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

- Acetic acid (CH3COOH) also known as ethanoic acid is a naturally occurring acid. It is a main ingredient of vinegar

- Citric acid (C6H8O7) is a natrually occurring acid present in citrus fruits, berries and vegetables

- Hydrochloric acid (HCl) is a naturally occurring acid produced by cells in the lining of stomach.

- Sulfuric acid (H2SO4) is the most manufactured chemical worldwide. It is used to manufacture fertilisers, synthetic fibres, industrial ethanol, detergents and car batteries.

Describe the use of the pH scale in comparing acids and bases.Identify pH as –log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+]

pH = -log10[H3O+][H3O+] = 10-pH

- The pH of a solution is defined as a negative of the logarithm (to base 10) of the hydrogen ion concentration

- A change in pH of one unit corresponds to a tenfold change in [H3O+]


- The number of decimal places in the value for pH equals the number of significant figures in the value for [H3O+]

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

- In a strong acid solution, each acid molecule is assumed to fully ionise. - The concentration of hydrogen ions and hence pH will depend on whether the acid is monoprotic,

diprotic (eg. H2SO4), or triprotic.- Use pH = -log10[H3O+] to calculate pH of hydrogen ion concentration.

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

Strong acid- Weak acid Organic acids contain –COOH groups. H can be attracted to water molecule as H+.

Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute.Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions.

A strong acid is one in which all the acid present in solution has ionised to hydrogen ions: there are no neutral acid molecules present

- Eg. All the hydrochloric acid is present as ionsHCl (g) + H2O (l) H3O+

(aq) + Cl- (aq)

- The ionisation reaction with water goes to completionA weak acid is one in which only some of the acid molecules present in the solution have ionised to form hydrogen ions

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.

In equation concentrations (eg. 1M) HCL is the strong acid - completely ionises

Citric and acetic acids are weak acids that only partial ionises. Typically only 1% ionised. In both weak acids, hydrogen ion concentration is much less than the concentration of dissolved acids.

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

Acids are added to- Improve the taste - Preserve the food by increasing its acidity, since bacteria cannot survive in acid conditions


Common acids added include acetic acid (vinegar), citric acid, with phosphoric acid used occasionally.

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

Substance Composition Acid or Base pH in naturally occurring form

Gastric acid Hydrochloric acid, potassium chloride, sodium chloride

Acid 1-2

Lemon Citric Acid Acid 2-3Vinegar Acetic acid Acid 2-3Cheese Lactic acid Acid 5-6Fruit/Vegetables Ascorbic acid Acid 2-3Putrefaction of animal and vegetable matter

Ammonia Base 11-12

Mineral deposits Sodium hydroxide Base 14

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

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

- Set up the datalogger and put the probe into a beaker of cold distilled water- Obtain the various acids of equal concentration in separate beakers. Place the probe into each

solution, one at a time and record the pH readings- Rank the acids according to their relative strength- HCl > HNO3 > H2SO4 > C6H8O7 > CH3COOH

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


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 - Acids were substances that contained oxygen- It attempted to define acids by their composition- This was disproved when it was found that many oxygen-containing substances were basic and some

were distinctly acids but contained no oxygen eg. (HCl)

Davy - Acids were substances that contained replaceable oxygen (hydrogen could be replaced by metals)- Bases were substances that reacted with acids to form salts and water- This reflected the composition of substances on that of HCl. - Since it classified substances without trying to interpret properties, it still does not explain why they

behave the way they do

Arrhenius - An acid is a substance that provides hydrogen ions when dissolved in an aqueous solution.- A base is a substance that provides hydroxide ions when dissolved in an aqueous solution.- Acids were strong if ionised completely, and weak if only ionised slightly- It interpreted acid properties in terms of the hydrogen ion they produced, and explained weak and

strong acids in terms of the extent to which the ionisation reaction proceeded- Definition was narrow in that it excludes metallic oxides that are distinctly basic

Limitations of Arrhenius theory- It does not explain the role of the solvent, as ionisation of an acid does not happen in isolation, but

rather is a reaction between the acid molecule and the solvent. The nature of the solvent is important

- Does not explain why many substances which behave as bases such as ammonia and sodium carbonate do not contain the hydroxide group in its formula.

- Does not explain how some substances can act both as an acid or base. It does not explain the role of the solvent.

- Does not explain why some combinations of acids and bases created solutions which were not neutral.

Outline the Bronsted-Lowry theory of acids and bases

An acid is a proton donor.A base is a proton acceptor.


Any molecule or ion that can act as a proton donor is a Bronsted-Lowry acid.Any molecule or ion that can act as a proton acceptor is a Bronsted-Lowry acid.

- If a substance has a greater tendency to give up protons, that substance in a solvent is an acidHA + H2O H3O+ + A-

- If a substance has a greater tendency to accept protons, that substance in a solvent in a baseB + H2O HB+ + OH-

- The Bronsted-Lowry concept showed that acidity depends on properties relative to those of the solvent of other reactant present in the solution

- Neutralisation did not need to involved transfer of H+ to OH- ions but could proceed directly by proton transfer

- Shows that the hydrolysis of salts to produce pH different from 7.0 = acid/base reaction- It also provided the basis for the quantitative treatment of acid-base equilibria and pH calculations

Hence, its advantages include that it can- explain why many substances are amphiphrotic (can act as acid and base)- explain why substances eg. carbonates and ammonia behave as bases and don’t contain hydroxide

ions- be applied to non-aqueous solvents and reaction that occur in the absence of solvent

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

Scientist DefinitionLavoisier Acid contains oxygen because non-metal oxides react with water producing acidic

solutionsDavy Acid contains hydrogen that can be replaced by reaction with a metal. Arrhenius Acid is a molecule or ion that contains hydrogen and dissolves in water to produce

hydronium ions. Base is a molecule or ion that dissolves in water to produce hydroxide ions.

Bronsted-Lowry

An acid is a proton donor. A base is a proton acceptor.

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

An acid gives up a proton to form a conjugate base.- Acid + water H3O+

+ conjugate base- The reaction product solution gains (accepts) a proton

Base accepts a proton to form a conjugate acid.- Base + water conjugate + OH-

- The reaction product solution gives up (donates) a proton

By the Bronsted-Lowry definition, anions, cations and neutral molecules can be acids or bases


A monoprotic acid is one that forms on proton (hydrogen ion) per molecule (eg. HCl, nitric acid)However, some acids, such as sulfuric and carbonic acid, are diprotic acids.

- H2SO4 (l) + H2O (l) H3O + (aq) + HSO4

- (aq)

HSO4-

(aq) + H2O (l) H3O+ (aq) + SO4

2- (aq)

- HSO4- is the conjugate base of H2SO4. However HSO4

- is itself an acid. SO42- is the conjugate base of

the acid HSO4-. H2SO4 is a strong acid, while HSO4

- is a weak acid- Thus, two series of salts are formed – the normal one such as sulfates or carbonates, and another set

known as hydrogen salts or acid saltsSimilarly, triprotic acids, such as phosphoric acid can form three series of salts: phosphates, hydrogen phosphates and dihydrogen phosphates Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature.

The reason why the pH of aqueous solutions of many salts is different from 7.0 is that many anions and cations can act as acids and bases

Basic salt (pH>7.0)- In a salt formed from a weak acid and a strong base (eg. sodium acetate from acetic acid and sodium

hydroxide), the anion is a weak base and so the salt in aqueous solution has a pH greater than 7- CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O(l)

CH3COONa (aq) + H2O (l) CH3COO – (aq) + Na+ (aq)

CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH-

(aq)

- The acetate ion accept a proton from water, forming hydroxide which makes the pH> 7.0Acidic salt (pH<7.0)

- In a salt formed from a weak base and strong acid (eg. ammonium chloride from ammonia and hydrochloric acid), the cation is a weak acid and so the salt in aqueous solution has a pH less than 7.

- NH4OH (aq) + HCl (aq) NH4Cl (aq) + H2O (l) NH4Cl (aq) + H2O (l) NH4

+ (aq) + Cl-

(aq)

NH4+

(aq) + H2O (l) NH3 (aq) + H3O+ (aq)

- The ammonia ion donate protons to water, forming hydronium which makes the pH<7.0Neutral salt (pH=7.0)

- A salt formed from a strong acid and strong base, or a weak acid and weak base forms a neutral salt- HCl (aq) + NaOH NaCl (aq) + H2O (l)

NaCl (aq) + H2O (l) Na+ (aq) + Cl-

(aq)

- Na+ is a metallic substance, so it stays as an ion

Cl- is a very weak base, so it does not accept protons- CH3COOH (aq) + NH3 (aq) (CH3COO)NH4 (aq)

(CH3COO)NH4 (aq) CH3COO- (aq) + NH4

+(aq)

CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH-

NH4+

(aq) + H2O (l) NH3 (aq) + H3O+ (aq)

- The hydroxide ions formed by CH3COO- and the hydronium ions formed by NH3 cancels each other, so the salt is neutral


Identify conjugate acid/base pairs

- The conjugate base of a strong acid is an extremely weak base (base reacts with water to a minute extent). Strong acids react completely with water, so consequently the conjugate bases do not react significantly with water

- The conjugate base of a weak acid is a weak base, so both acid/base react with water significantly- The conjugate acid of a strong base is an extremely weak acid (acid reacts with water to a minute

extent)

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

A substance that can act both as a proton donor and as a proton acceptor is called an amphiprotic substance

- For example, hydrogen carbonate ion is amphoproticHCO3

- (aq) + H2O (l) H3O+ (aq) + CO32- (aq) (HCO3

- is a proton donor – acid)HCO3

- (aq) + H2O (l) H2CO3-

(aq) + OH- (aq) (HCO3- is a proton acceptor – base)

- The reactions occur simultaneously when hydrogen carbonate is dissolved in water, but only to small extents

- If placed in alkaline (basic) solution, it acts as an acidHCO3

- (aq/s) + OH-

(aq) H2O (l) + CO32-

(aq)

- If placed in acid solution, it acts as a baseHCO3

- (aq/s) + H+

(aq) H2CO3 (aq)

- The reaction goes to completion in both cases

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

- Set up the data logger and obtain several small beakers, each containing 10 ml of various 0.1M solutions of salts (eg. NaCl, KNO3, NaHCO3, CH3COONa, Na2CO3)

- Place the pH probe into each salt solution, recording the pH reading. Universal indicator solution can also be used to determine pH.

- Identity species that are responsible for measured pH of salt solutions- Write chemical equations for the formation of salts during neutralisation, and for hydrolysis of

respective ions- Explain and acidic, basic or neutral nature of salts


Identity neutralisation as a proton transfer reaction which is exothermic

- The net ionic reactions for many neutralisation reactions show a proton transfer from H 3O+ ion to the OH- ion. The acidic and basic properties of the original base and acid are cancelled out.

H3O+ + OH- 2H2O- For example, acid transfers a proton to the base OH-

HCl (aq) + NaOH (aq) H2O (l) + NaCl (aq) - Equations tend to be written in terms of H3O+ when focusing on the proton transfer aspect of the

equation, and in terms of H+ when calculating quantities- Neutralisation is an exothermic reaction, meaning that heat is liberated and the enthalpy change is

negative (approx. -56 kJ/mol)- Enthalpy varies by a few kJ/mol depending on strength and concentrations of solutions involved.

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

- Volumetric analysis is a form of chemical analysis in which the concentration (or amount) of a substance A is determined by measuring the volume of a solution of known concentration of another substance B which is just sufficient to react with all of the sample of A.

- Titration is the process of determining the ‘just sufficient’ volume- The equivalence point is the point at which stoichiometrically equal amount of reactant solutions

have been added so that no acid or base is left in solution (complete neutralisation). This is when the ratio of moles of acid and base added is the same as mole ratio in balance equation. pH and equivalence point depends upon strength of acid and base

- The end point is where the indicator permanently changes colour (stop adding, stop titration)- An indicator is used to detect the equivalence point, where the colour will change sharply.

- A primary standard in volumetric analysis is a substance of sufficiently high purity and stability that a solution of it, of accurately known concentration, can be prepared by weighing out the desired mass, dissolving it in water, and making the volume up to an accurately known volume.

- Hence the primary standard should be:o solid – weighed accurately on an electronic balanceo easily be obtained in very pure form – analytical grade (a grade), known formula - e.g. HCl is

unsuitable because it is volatileo large molecular mass – more accurateo unaffected by exposure from air – e.g. NaOH is unsuitable because it is hygroscopic,

absorbing water from the air, and also reacts with carbon dioxide from the airo High water solubility – can make it into solution, must dissolve

- Sodium carbonate, sodium hydrogen carbonate and oxalic acid are suitable primary standards. - Samples to be analysed are often quantitatively diluted


- Need to ensure that standard solutions have precisely the concentrations we expect, and that the exact equivalence point is reached.

- Procedureo Use a volumetric flask (cleaned with distilled water) and note volume (common- 500.0ml)o Substance needs to be pure and free as possible. Dry in oven and cool in dessicator. o Calculate mass of solid required to make 500.0ml of the substances of a particular

concentrationo Mass is accurately weigh out using paper or plastic cup on an electronic balanceo Mass is carefully transferred to volumetric flask using wash bottle or funnelo The flask is one-third filled with distilled water and gently shaken to dissolve masso Distilled water is added until the bottom of the meniscus is on the 500.0ml line

Method- Fill a burette with solution of known concentration. This solution is called the titrant - Place solution of sample to be analysed in a flask under the burette.- Add several drops of indicator into flask- Place a white card under flask, allowing for easier detection of colour change- Slowly run solution from burette into flask. Continue swirling until colour change- Read volume delivered by burette. Repeat titration several times. (First titration overshoots the

equivalence point.)- Calculate required concentration of unknown. Ensure answer is given to the correct number of

significant figures.

Number of moles= massmolar mass

Molarity= number of molesvolume of solution∈litres

- Volumetric glassware needs to be thoroughly cleaned and rinsed with distilled water. - Pipettes and burettes needs to be rinsed with a portion of the solution to be measured. This

prevents dilution of the solution by water. - Pipette should be held vertically. Pipette filler is use to draw in solution until it is above the

graduation mark, and then run out until the meniscus sits on the mark. Hold at eye level to avoid parallax errors. Solution is let run out of pipette into required flask by gravity, with the pipette held with its tip in contact with wall of the receiving flask.

- A burette is overfilled, and the excess is run out. Once air bubbles are removed, clamp burette vertically and lower liquid level until meniscus sits on the zero mark. As burettes are marked with 0.1ml graduations, it can be read to +/- 0.03ml. Near the end-point, a portion or a drop can be allowed to build up on the tip, then washed into the flask using a wash bottle.


- Measuring cylinder is not accurate enough. For volumetric analysis, volumes need to be measured to +/- 0.5% or better.

- Indicator is selected to that it will change at the equivalence point. pH of equivalence point varies.- Strong Acid + Strong Base Neutral Salt + Water [pH=7] – bromothymol blue- Weak acid + strong base basic salt + water [pH>7] – phenolphthalein- Strong acid + weak base acidic salt + water [pH<7] – methyl orange - Weak acid-weak base titrations are not performed as pH change around equivalence point is not

large. There is no sharp colour change, so it is impossible to determine equivalence point accurately.

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

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

A buffer solution is a solution that contains comparable amounts of a week acid and its conjugate base and which is therefore able to maintain an approximately constant pH even when significant amounts of strong acid and strong base are added to it.

HA + H2O H3O+ + A-

- Adding hydrogen ion (more acidic) will shift the equilibrium to the left. The base A - combines with added H3O+ to form HA, minimising the change in hydrogen ion

- Adding hydroxide ion will shift the equilibrium to the right to minimise change as HA ionises to produce more A-

Blood- Complex processes in the human body involve enzymes that only operate over narrow pH ranges- Buffered solution contains carbonic acid and hydrogencarbonate ion.

H2CO3 (aq) + H2O (l) HCO3-

(aq) + H3O+ (aq)

- If an acid is added, the increase in H+ will increase the amount of hydronium ions. The equilibrium shifts to the left to form H2CO3, minimising the change in hydrogen ion

- In a base is added, the base combines with H2CO3. The equilibrium shifts to the right to increase the amount of HCO3

- ions, minimising the change.


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

- Many acids and alkalis are corrosive, so spills need to be neutralised quickly- Sewage authorities put strict limits on the pH of factory and laboratory effluents discharged- Sodium hydrogen carbonate is widely used because it is

o A stable solid which is easily and safely handled and storedo Cheapo Less danger is an excess is usedo Amphiprotic – can neutralise both acid and alkali

- Spills of acids or bases on skin should never be neutralised because neutralisation is exothermic, i.e could burn skin. Instead, rinse affected area with water.

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

Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds

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

Alkanols are compounds formed by replacing an H atom of an alkane with an OH group. - To name alkanols

o Identify longest chain that includes the carbon bearing the –OH groupo Suffix “-ol” is used to represent the –OH group, and the lowest possible number is assigned

to the carbon bearing this group.o Find whatever groups are not part of the longest continuous chain. Names these prefixes.

- Alkanols are a subset of a larger class of compounds called alcohols, which are compounds that contain an OH group attached to a C atom that only has C or H atoms attached to it, regardless of what else is in the molecule. [OH group commonly called the alcohol group]

Alkanoic acids have the general RCOOH, where R is an alkyl group or an H atom.- To name alkanols

o Identify longest chain includes carbon bearing –COOH groupo The suffix “oic acid” is used to represent the –COOH group and lower possible number

assigned to carbon bearing this group. –COOH is always number 1 and not stated.o Find whatever groups are not part of the longest continuous chain. Name these prefixes.

- Since the –COOH group has only on free carbon valence, it must always be at the end of a carbon chain. The –COOH group is called the carboxylic acid group. Carboxylic acids are compounds which contain this group, regardless of what is attached to it. CH3-COOH and Cl-CH2-COOH are both carboxylic acids; CH3COOH is an alkanoic acid, but Cl-CH2-COOH is not

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


- As C-O and O-H bonds are polar, alkanols are polar molecules- C=O bond is also polar, making alkanoic acids even more polar- O-H bonds can form hydrogen bonds, so there are strong intermolecular forces in alkanols and in

alkanoic acids. Hence, they have higher melting and boiling points than alkanes of similar molecular weight. They are also highly soluble

- Compounds with similar molecular weights are compared to show effect of polarity and hydrogen bonding. Strength of dispersion forces increases as the number of electrons in the molecule increases, which increases as molecular weight increases.

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

Esters are compounds formed when alkanoic acids react with alkanols, or more generally, when carboxylic acids combine with alcohols

- The C=O and C-O bonds make esters polar substances. However, esters have much lower boiling points than acids or alkanols of the same molecular weight because esters have no hydrogen bonding. They are also less soluble in water.

The reaction between an alkanol and an alkanoic acid is called esterification- When ethanoic (acetic) acid and ethanol are heated with a few drops of concentrated acid, ethyl

ethanoate (ethyl acetate) is formed:CH3COOH + C2H5OH CH3-C=O -O-CH2 – CH3 + H2O

- Esters contain the structural unit

- They’re named alkyl alkanoates – alkyl comes from the alcohol and alkanoate from the alkanoic acid

Describe the purpose of using acid in esterification for catalysis

- Esterification is moderately slow at room temperature and does not go to completion: it comes to equilibrium

- Concentrated sulfuric acid is used as a catalyst too Speed up the reactiono Absorb product water, forcing equilibrium to the right

Explain the need for refluxing during esterification

- Since the reaction is carried out at temperatures near boiling point, alcohol needs to be prevented from vaporisation. The water-cooled condenser allows alcohol vapour to be condensed and run back into the reaction vessel. This enables the reaction to occur at higher temperatures.

- Refluxing is the process of heating a reaction mixture in a vessel with a cooling condenser attached in order to prevent loss of any volatile reactant or product.

- Using a closed vessel could lead to a dangerous build-up of pressure and the possibility of explosion.

Outline some examples of the occurrence, production and uses of esters


Process information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmetics

- Esters have pleasant, fruity odours, occurring widely in nature as perfumes and flavouring agents - Natural flavours results from a complex mixture of many esters. For example

o apple – methyl butanoateo pear – pentyl ethanoateo jasmine – benzyl ethanoate

- Synthetic flavours and perfumes are produced by first identifying the constituents of the natural flavour, then synthesising similar mixtures of esters which reproduce this flavour. These are cheaper than natural extracts, and represent little health hazard.

- Ethyl acetate is widely used as a solvent in industry. It is also used as a solvent in nail polish remover- High molecular weight (i.e. non-volatile) esters e.g. dialkyl phthalates are used in plasticisers in PVC

Identity data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux

Risk Assessment: Alcohols and organic acids are flammable so keep away from flames or sparks and work in a well ventilated area. Concentrated sulfuric acid is highly corrosive to skin and eyes. Wear safety glasses and protective clothing, and immediately wipe away any spills.

MethodPreparation of Ester

- Add 10 ml of 1-pentanol, 12 ml of ethanoic acid and 5 drops of concentrated sulfuric acid into a round bottom flask. Excess alkanoic acid is used to increase the rate of forward reaction relative to the reverse, shifting the equilibrium to the right.

- Add 2 to 3 boiling chips of porous ceramic material to the reaction mixture in the flask. (for steady boiling)

- Attach a condenser to the flask and clamp the condenser vertically to a retort stand - Attach the cooling water supply to the bottom connector of the condenser and the outlet tube to

the upper connector. Ensure that the outlet tube exits into the sink. - Heat the mixture for 10 minutes by immersing the flask into a beaker of water heated by an electric

hotplate.- Allow to cool. Separate the condenser and flask.

Isolation of Ester- Transfer to separating funnel. Add 100ml water the funnel and shake, allowing the two layers to

separate. Discard the lower layer. The lower layer is the water layer, containing ethanoic acid and 1-pentanol. The organic layer at the top contains the ester – pentyl ethanoate and 1-pentanol (as ester is less dense) (1-pentanol is slightly soluble)

- Add 50ml of saturated solution of sodium carbonate (Na2CO3) and shake, allowing two layers to form. Discard the lower aqueous layer (remaining acid is remove)

- Pour the remaining layer into a clean round bottom flask, add 2 to 3 boiling chips and distil the mixture, collecting fractions into separate test tubes. Record the temperature at which each fraction boils off.


Diagram:

Equation:

Discussion- Yield is maximised by using a sulfuric acid catalyst and using excess alkanoic acid to shift the

equilibrium to the right. - The validity of the experiment is improved by adding water and sodium carbonate to remove all

acid. - Distillation – a thermometer is inserted in the distillation apparatus to record the temperature of the

avpur as it distils over. The condensed vapours are all collected into a test tube until the temperature reaches 138°C. A clean dry test tube collects the pure pentyl enthanoate, and the test tube is changed once the temperature of the vapour rises again.

o Boiling point of pentyl ethanoate: 149°Co Boiling point of 1-pentanol: 138°Co Boiling point of ethanoic acid: 118°C

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