1

William KuangThe Acidic Environment

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

classify common substances as acidic, basic or neutralAll of the substances we use are acidic, basic or neutral

AcidsBasesNeutral

VinegarAmmonia (NH3)Pure water

Lemon JuiceSodium carbonateNaCl solution (sea water)

Battery Acid (H2SO4)Caustic soda (NaOH)Ethanol

DetergentsGlucose solution

Slaked lime

Acids sour taste, corrosive, in solution conducts electricity, BARBases bitter taste, caustic, usually soapy feel, in solution conducts electricity,

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

IndicatorHighly AcidicSlightly AcidicNeutralSlightly AlkalineHighly AlkalinepH Range

Methyl Orange(tests for HIGH acidity)red yellowyellowyellowyellow3.1 4.4

Bromothymol Blueyellowyellow(green)blueblue6.0 7.6

Litmusredred(purple)blueblue5.0 7.6

Phenolphthalein(tests for HIGH alkalinity)colourlesscolourlesscolourlesscolourlessred(purple/pink)8.0 10

Methyl orange

Litmus

Bromothymol Blue

Phenolphthalein

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

An indicator is a substance which in solution changes colour depending on whether the solution is acidic or basic. The colour change indicates the concentration of hydrogen ions which is how substances are classified as being acidic, basic or neutral.

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

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

Indicators provide a cheap and convenient way of determining the acidity or alkalinity of substances. Some everyday uses of indicators are:

Testing pH of soil- as most plants have a preferred soil pH range. In the testing of soil, a sample is taken and placed in a petri dish with distilled water. A white neutral powder (e.g. barium sulfate) is added (as soil can hide the indicator colour change) and then the indicator is added and the colour change observed.

Testing pH in swimming pools - Water needs to be approximately neutral as to avoid irritation to the skin and eyes. Add HCl to increase acidity. Add sodium hydrogen carbonate to increase basicity.

Aquarium water testingSome aquarium animals are sensitive. Check using indicator

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

Aim: To prepare red cabbage as a natural indicator and test in a number of known acidic, basic, and neutral solutions.

Safety: Wear safety glasses to avoid substances coming into contact with eyes.

Method:

1) Equipment was set up as diagram above2) Red cabbage was shredded and 50g was placed into the beaker. 300mL water was also placed into the beaker3) Bunsen was ignited and the water was allowed to reach near boiling. Bunsen was put out and the solution was allowed to cool4) Solution was decanted into 5 test tubes placed in a test tube rack.5) 30mL of 0.01mol/L HCl was placed in a test tube. Colour change was recorded. (Dependant variable observed colour change)6) Step 5 was repeated with 0.01mol/L acetic acid, 0.01mol/L ammonia, 0.01mol/L sodium hydroxide, and ethanol. (Independent variable known acidic, basic, or neutral solution mixed with the natural indicator)7) Results from 5 other groups performing a similar investigation were collected and recorded to establish reliability.

Results:

Red Cabbage Solution

Original Colour: PurpleColour in Acidic solution (HCl, acetic acid): RedColour in Basic solution (ammonia, NaOH): GreenColour in neutral substance (water, ethanol): Purple

Analysis: Red cabbage is a very good natural indicator as it gives a distinct colour change in acidic (turns red), basic (turns green), or neutral (remains purple) substances. However it is not very good at determining the pH difference between two acidic or basic substances.

2. While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulfur. 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 are a class of compound that often displays acidic or basic properties

Acidic Oxides (Non-metal oxides) lie on the right hand side of the periodic table.Non-metal oxides act as acids when in contact with water (to form an acidic solution) or bases (to form salts and usually water)

They react with water to form acidic solutions (proton donors in solution)

CO2 (g) + H2O (l) H2CO3 (aq)Carbonic acidSO2 (g)+ H2O (l) H2SO3 (aq) Sulfurous acidSO3 (g) + H2O (l) H2SO4 (aq) Sulfuric acid2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)Nitrous acid + Nitric acidP2O5 (g) + 3H2O (l) 2H3PO4 (aq) Phosphoric acid

They react with bases to form salts (and usually water)

CO2(g) + 2NaOH(aq) H2O(l) + Na2CO3(aq) (sodium carbonate)

Basic Oxides (Metal oxides) lie on the left hand side and centre of the periodic table. Metal oxides act as bases when in contact with water (to form a basic solution, usually a hydroxide) or acids (to form salts and usually water)

They react with water to form basic solutions (proton acceptors in solution)

MgO (s) + H2O (l) Mg(OH)2 (aq)

They react with acids to form salts (and usually water)

CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)

[SKILL] Learn: Acid + base salt + waterAcidic oxide + water acid Basic oxide + water base 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

Left hand side (metals) Basic oxides e.g. CuO, MgO Right hand side (non-metals) Acidic oxides, e.g. SO2, CO2, NO2

Basicity increases down the group. Acidity decreases down the group.

Exceptions Group VIII (noble gases) do not form oxides Some non-metal oxides are neutral- CO, NO, N2O Some elements form amphoteric oxides - Be, Zn, Al, Sn, Pb

Amphoteric oxides react with both acidic and basic solutionse.g. Al2O3 (s) + 6HCl (aq) 2AlCl3 (aq) + 3H2O (l)Al2O3 (s) + 2NaOH (aq) 2NaAlO2 (s) + H2O (l)

define Le Chateliers principle

When a system in equilibrium is disturbed, the system adjusts itself as to minimise the disturbance and return to equilibrium.

- This means that either the forward or reverse reaction proceeds at a faster rate until equilibrium is re-established.

About equilibrium- A system is in equilibrium when the forward and reverse reactions are occurring at the same rate and therefore does not go to completion. The concentrations of products and reactants are always constant. - For equilibrium to be established the system must be closed. (No matter or energy enters or leaves the system)- Macroscopic properties (observable properties) remain constant (e.g. colour, temperature, pressure, state) as concentrations remain constant.- Microscopic changes continue to occur as the forward and reverse reactions continue to occur at the same opposing rate. Thus equilibrium is said to be dynamic.

identify factors which can affect the equilibrium in a reversible reaction [SKILL practice predicting the reaction shift]

Factors that can affect the equilibrium in a reversible reaction include changes in:

Concentrations of substances involved in the reaction. If the concentration of a certain substance is increased, the reaction will shift in the direction that decreases its concentration. Conversely, if the concentration is decreased, the reaction will shift in the direction that increases its concentration.

Total pressure on a reaction involving gases.

Increase pressure Increases concentration of all gases Reaction shifts in the direction which forms the least gas molecules Lower concentration of gas molecules = a reduction in pressure

The converse is also true: Decrease pressure decreases concentration of all gases Reaction shifts in the direction which forms the most gas molecules Higher concentration of gas molecules = a rise in pressure

Temperature

Increase temperature Reaction shifts in the endothermic direction (to reduce heat)Decrease temperature Reaction shifts in the exothermic direction (to increase heat)

Adding a catalyst lowers the activation energy for both the forward and reverse reactions. It has no effect on the position of equilibrium but it allows for equilibrium to be reached faster.

describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chateliers principle

CO2 dissolves in water producing H2CO3 (carbonic acid) in an exothermic reaction

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

According to Le Chateliers principle, when a system in equilibrium is disturbed (change in: pressure, temperature, concentration), the system adjusts itself as to minimise the disturbance. Thus in the carbon dioxide in water equilibrium

DisturbanceEffect on equilibrium (right means inc. solubility)

Change in pressureIncreasing the pressure causes the system to favour the reaction direction which forms the least amount of gas molecules which in this case is the forward direction. [i.e. reaction shifts right more CO2 dissolves]

Converse is also true [decrease pressure reaction shifts left to form more CO2 gas molecules]

Change in concentrationIn general increasing the concentration of a substance favours the direction which reduces it. Converse is also true.E.g.Increase H2O Reaction shifts rightIncrease CO2 Reaction shifts right (same as pressure inc.)Increasing acidity (e.g. add an acid) Reaction shifts leftReducing acidity (e.g. neutralisation) Reaction favours right

Change in temperatureIncreasing temperature causes the reaction to favour the endothermic direction i.e. the reaction shifts left. (CO2 is less soluble with a temp. increase)

Decreasing temperature causes the reaction to favour the exothermic direction i.e. the reaction shifts right. (CO2 is more soluble in temp. decrease)

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 25C and 100kPa

Aim: To decarbonate soft drink and calculate the volume of gas released at 25C and 100kPa

Safety: Wear safety glasses to avoid substances coming into contact with eyes.

Method:1. The following were weighed - Closed 350mL can of soft drink- 20g NaCl - Empty 2L bottle with cap2. Can was opened and poured into the 2L bottle along with the 20g NaCl. (NaCl is added as it is more soluble than CO2 in water; as NaCl dissolves CO2 is forced out as part of a physical reaction) 3. Bottle was shaken at 1 minute intervals and opened periodically to allow the CO2 to escape. This was done 30 times.4. Bottle including contents was weighed and mass of CO2 that escaped was calculated. Volume of CO2 was subsequently calculated. [Dependant variable volume of CO2 released]5. Steps 1-4 were repeated with 5 different types of soft drink [Independent variable different soft drinks used]

Results:

Closed can = 405g2L Bottle = 6.5gEmpty can = 6.8gSalt = 20g 2L Bottle + salt + soft drink (after shaken and opened 30 times) = 432.1g WRONG

Soft drink (originally) = 405 6.8 = 398.2gSoft drink (final) = 432.1 6.5 20 = 405.6g [FIGURES ARE WRONG! MASS WAS GAINED RATHER THAN LOST]

[LETS PRETEND]Mass of CO2 released = 7.4g

nCO2 = 7.4 / (12 + 16 +16) = 37/220 mol. 37/220 mol. CO2 = 37/220 x 24.79 = 4.169L

Therefore 4.169L CO2 was released

Controls: Ensure experiment is conducted at 25C and 100kPa

calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0C and 100kPa or 25C and 100kPa [SKILL mole calculations]

At 0oC and 100kPa, molar volume of gases is 22.71 LAt 25oC and 100kPa, molar volume of gas is 24.79 LV

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

SourcesSulfur dioxide (SO2)Oxides of nitrogen NOx: nitric oxide (NO), nitrogen dioxide (NO2)nitrous oxide (N2O)

Natural Volcanoes Geothermal hot springs S (s) + O2 (g) SO2 (g)

Bacteria decomposing organic matter, producing H2S, and oxidising to SO22H2S (g) + 3O2 (g) 2SO2 (g) +2H2O (g) Lightning Forrest fires N2 (g) + O2 (g) 2NO (g)

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

Activity of bacteria on nitrogenous matter in soil

Industrial Combustion of fossil fuels(sulfur is present as impurities) S (s) + O2 (g) SO2 (g)

Smelting sulfide ores to extract metals 2ZnS(s) + 3O2(g) 2ZnO(s) + 2SO2(g) Internal combustion engines (cars, power stations)

N2 (g) + O2 (g) 2NO (g)

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

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

Sulfur DioxideCombustion of Fossil Fuels in power plants and vehicles (sulfur impurities)The sulfur impurities in these fuels directly combust with oxygen to produce sulfur dioxide. S (s) + O2 (g) SO2 (g)

Smelting of Sulfide Ores to extract metals from their compounds.e.g. 2ZnS (s) + 3O2 (g) 2ZnO (s) + 2SO2 (g)

Volcanoes & Geothermal Hot SpringsRelease large amounts of SO2 into the atmosphereS (s) + O2 (g) SO2 (g)

Oxides of NitrogenLightning, Forrest fires and Combustion EnginesThe extreme temperatures of both lightning and in internal combustion engines provide sufficient energy to allow atmospheric nitrogen and oxygen to form oxides.

N2 (g) + O2 (g) 2NO (g)NO then oxides slowly to form NO2

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

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

Overall there is valid and reliable evidence which indicates the rise in oxides of sulfur and nitrogen since the industrial revolution. Such evidence includes the comparison of old air trapped in layers of ice in Antarctica with modern air. This comparison successfully demonstrates the increase in oxides of sulfur and nitrogen since the industrial revolution.

This increase is of particular concern in heavily industrialised societies such as the USA. The rise is a result of the heavy industrialisation since the industrial revolution with its reliance on fossil fuels as energy and metals (requires smelting). The large increase in motor vehicles has also had a significant contribution (production, combustion engines, etc.)

However whilst evidence has shown there has been a significant increase in oxides of sulfur and nitrogen since the industrial revolution, evidence has shown that in recent years there has been no significant increase. This is a result of both human intervention (increasingly stringent emission controls and regulations) and the soluble nature of SO2 and NO2 (These two gases do not remain in the atmosphere as they are largely washed out by rain, however N2O is insoluble and has increased by more than 30% in the past century).

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

Since the Industrial Revolution in the 1800s there has been an increase in SO2 and oxides of nitrogen (NO, NO2, N2O). The rise is a result of the heavy industrialisation during and since the industrial revolution with the: Heavy reliance on fossil fuels as energy [sulfur impurities in coal (SO2), combustion engines (NO, NO2) in transport and power stations] And Metals (requires smelting (SO2)).

Also the increased use in nitrogenous fertilisers has contributed to increased bacteria activity in the soils forming larger amounts of SO2)

The large increase in motor vehicles has also had a significant contribution (production, combustion engines,)

Industrial origins - SO2 (burning of coal which contains sulfur impurities and the smelting of sulfide ores) [remember equations]Oxides of nitrogen (high temperature combustion environments e.g. power stations and combustion engines) [remember equations]

Reasons for Concern:

There is significant reason for concern in the release of these gases into the environment. This is because:

Both NO2 and SO2 irritate the respiratory system causing breathing problems NO2 leads to the formation of ozone due to photochemical smog. This is of serious concern as ozone Readily oxidises bodily tissue is a toxic gas which causes severe respiratory problems even in low concentrations High emissions of NO2 and SO2 to the atmosphere lead to the formation of acid rain causing: Erosion of iron structures such as bridges Erosion of stone buildings (sandstone, marble, limestone) Increased acidity of lakes which damages marine life Damage to plants (e.g. pine forests) and inhibiting their growth (as the soil pH is lowered)

Thus the release of SO2 and oxides of nitrogen need to be restricted through stringent controls and regulations/

explain the formation and effects of acid rain

Acid rain is rain with a higher hydrogen ion concentration than normal i.e. [H+] > 10-5 mol/L. The pH is usually between 3 and 5.

Regular rain is slightly acidic due to CO2 in the air dissolving to form carbonic acidCO2 (g) + H2O (l) H2CO3 (aq)

High emissions of oxides of sulfur and nitrogen into the atmosphere lead to the formation of acid rain. This is explained by the following equations.SO2 (g) + H2O (l) H2SO3 (aq) (sulfurous acid)2H2SO3 (aq) + O2 (g) 2H2SO4 (aq) (sulfuric acid)SO3 (g) + H2O (l) H2SO4 (aq) (sulfuric acid)

2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq) (nitrous acid + nitric acid)2HNO2 (aq) + O2 (g) 2HNO3 (aq) (nitric acid)

Effects of acid rain Increases acidity of waterways has a detrimental effect on marine life Inhibits reproduction and growth of fish

Causes damage to plants Pine forests lose their waxy coating and cannot withstand the cold Soil pH is lowered which can inhibit the growth of plants

Causes damage (erosion) to iron structures

Fe (s) + H2SO4 (aq) FeSO4 (s) + H2 (g)

Causes damage (erosion) to stone structures (limestone, sandstone and marble) as they contain carbonates.

E.g.CaCO3(s) + H2SO4(aq) CaSO4(aq) + H2O(l) + CO2(g)

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 [POSSIBLE CLEANUP]

An acid is a proton donor. An acid in water ionises and donates an H+ (proton) to water to form hydronium ions (H3O+).

For example:HCl (g) + H2O (l) H3O+ (aq) +Cl (aq)HCl (g) H+ (aq) +Cl (aq)

In general: acid + water hydronium ion + negative ion A base can be defined as a proton acceptor. In water they ionise to form hydroxide ions (OH-) which can accept a proton to form water.

For example: NaOH (aq) Na+ (aq) + OH- (aq)OH- (aq) + H+ (aq) H2O (aq)

Gather and process information from secondary sources to write ionic equations to represent the ionisation of acidsMonoprotic acids HCl (g) H+ (aq) + Cl (aq)HNO3 (l) H+ (aq) + NO3 (aq)CH3COOH (aq) H+ (aq) + CH3COO (aq)

ConverselyMonoprotic bases = NaOH, KOH, etc. [note: strong hydroxides are usually the group 1 and 2 hydroxides]

Diprotic acidsH2SO4 (l) 2H+ (aq) + SO42 (aq)H2CO3 (aq) 2H+ (aq) + CO32 (aq)

ConverselyDiprotic bases = Mg(OH)2, Ba(OH)2 etc.

Triprotic acids H3PO4 (aq) 3H+ (aq) + PO43 (aq)

ConverselyTriprotic bases = Al(OH)3, Fe(OH)3 etc.

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

Aim: To use pH probes and indicators to distinguish between acidic, basic and neutral substances

Method: 1) 4 drops of 0.1mol/L HCl were placed into a test tube. 4 drops of universal indicator were then added and the colour change observed. The colour change was then matched to a chart to identify pH. 2) Step 1 was repeated with 0.1 mol/L NaOH, NaCl solution, baking soda solution, washing soda solution, sugar solution, lemonade, chlorine bleach solution, tap water, milk, cloudy ammonia, shampoo, vinegar3) The pH probe was placed in solutions with a known pH (buffers) to establish its validity. When the reading on the probe corresponded with the known pH of the solution the validity of the pH probe was established. 4) The pH probe was washed and dried. It was then immersed in a beaker of 200mL 0.1mol/L HCl and the reading on the probe observed and recorded. 5) Step 4 was repeated with 0.1 mol/L NaOH, NaCl solution, baking soda solution, washing soda solution, sugar solution, lemonade, chlorine bleach solution, tap water, milk, cloudy ammonia, shampoo, vinegar6) Results for 5 groups performing a similar investigation were collected and recorded to strengthen reliability.

Results:Universal Indicator

SolutionpHUniversal Indicator Change

0.1 mol/L HCl 10-7 mol/L (at 25C)

Neutral pH = 7 [H+] = [OH-] = 10-7 mol/L (at 25C)

Basic pH > 7 [OH-] > [H+] [OH-] > 10-7 mol/L (at 25C)

identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+] [SKILL-calculations]

pH = - log10 [H+]

(note: H+ = H3O+)

10 - pH = [H+]

[H+] [OH-] = 10-14(mol/L) 2

(at 25C) (Ionic product constant for water)Because pH is a logarithmic scale to the base 10, it follows that a pH difference of 1 means a ten-fold change in [H+].I.e. pH increased by 1 = [H+] decreased 10 timespH decreased by 2 = [H+] increased 102 timespH increased by 6 = [H+] decreased 106 times

Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations [SKILL-calculations]

1) Write the ionic equation (note whether the acid is monoprotic, diprotic, etc.)2) Given the molarity of the acid, determine [H+] (strong acids go to completion)3) Calculate the pH using: pH = - log10 [H+]

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

A strong acid is one which (virtually) completely ionises in water, releasing H+ (protons) to form H3O+. Strong acids include (HCl), (H2SO4) and (HNO3). They are good proton donors. E.g.HCl (aq) + H2O (l) H3O+ (aq) + Cl (aq)

(Note the reaction essentially goes to completion; all water molecules react with all acid molecules to produce H3O+)

(Conversely a strong base is one which completely ionises in aqueous solution to form OH e.g. group 1 and 2 metal hydroxides)

A weak acid is an acid which only partially ionises in water, releasing H+, to form H3O+. Weak acids include (CH3COOH), citric acid, and carbonic acid. Bad proton donors.E.g.CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO-(aq)

(Note the reaction does not go to completion instead establishing equilibrium; only some acid molecules react with the water molecules to produce H3O+)

(Conversely a weak base is one which does not completely ionise to form OH, instead establishes equilibrium e.g. carbonates (CO3 2) and ammonia (NH3))

A concentrated acid is one which has a high number of (solute) moles per litre. A dilute acid is one which has a low number of (solute) moles per litre.

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

Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acidsWe made up mixtures of ethanoic acid and hydrochloric acid that were exactly 0.1 M. We then used universal indicator to find out the pH of these solutions. We found that the pH of the strong acid was higher than that of the 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

The relative strengths of weak acids can be compared by their degree of ionisation.

A strong acid (like HCl) ionises completely and thus its degree of ionisation is 100%. However as weak acids (like citric and acetic) only partially ionise, if their concentrations are the same their strengths can be compared.

Some calculations: With 0.1 mol/L solutions of HCl, Acetic acid, and Citric acid their repective pHs are 1, 2.9, 2.1. Calculate their degree of ionisation.

As pH = - log10 [H+], Therefore for citric, 2.1 = - log10 [H+]Therefore for HCl [H+] = 10( 2.1) mol/LTherefore HCl degree of ionisation = 10( 2.1)/0.1 = 7.94%

As pH = - log10 [H+], Therefore for acetic acid, 2.9 = - log10 [H+]Therefore for acetic acid [H+] = 10( 2.9) mol/LTherefore acetic acid degree of ionisation = 10( 2.9)/0.1 = 1.25%

STRONGEST ------- Hydrochloric Citric Acetic ------------ WEAKEST

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

Strong acids (virtually) completely ionise and release hydrogen ions (protons) which join with water molecules to create hydronium ions (H3O+). The equilibrium very strongly favours the formation of products and is thus denoted with a forward arrow.

Weak acids only partially ionise and release hydrogen ions (protons)E.g. CH3COOH (aq)+ H2O (l)H3O+ (aq) + CH3COO(aq)

In the equilibrium of the weak acid, the formation of reactants is strongly favoured and hence very few acid molecules ionise. In solution weak acids contain most of its acid molecules along with few of its ions.

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

Weak acids are often added to food because:

They inhibit the growth of micro-organisms such as bacteria and moulds (by lowering the pH as micro-organisms cannot usually survive below a pH of 5) They can be used for food preservation (e.g. acetic acid in vinegar used in pickling) They can act as an antioxidant to prevent the spoilage of food through oxidation They add flavour to foods (as acids have a characteristically sour taste; e.g. acetic acid in vinegar used in sweet and sour sauces)

Vitamin C (ascorbic acid) is added in foods as it is an essential nutrient for humans as it promotes the growth of connective tissue. Can be found as an additive in juices.

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

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

outline the historical development of ideas about acids including those of: Lavoisier Davy Arrhenius

Historical development of ideas about acids:

Lavoisier (1743 1794) (Oxygen theory of acids) He defined an acid as a non-metal compound which contained oxygen. Observed that non-metal oxides, when dissolved in water, formed acids Limitation: Not all acids contain oxygen, e.g. HCl Could not explain why metal oxides were basic

Davy (1778 1829) (Hydrogen theory of acids) Proposed that all acids contain hydrogen [as opposed to oxygen] Demonstrated that HCl was an acid which did not contain oxygen Observed that metals reacted with acids to form H2 Limitation: Could not explain why other compounds containing hydrogen were not acidic (e.g. methane)

Arrhenius (1859 1927) Defined an acid as a substance which ionises in solution to produce H+ Defined a base as a substance which ionises in solution to produce OH Developed the concept of Strong (if ionises completely) and Weak (ionises only partially) acids/bases Suggested that reactions between acids and bases (neutralisation reactions) produce water and salts Limitations: Only applies to aqueous solutions

Does not give recognition to the role of the solvent (ionisation is a reaction between the acid and solvent) (acid strength is determined also by the solvent)

Only accounts for substances which already have H+ or OH in their structure (e.g. NH3 is basic)

Does not explain why some salts (ZnCl2: acidic, NaS: basic) act as acids or bases

Cannot explain how some substances can act as both an acid and a base (amphiprotic)

outline the Brnsted-Lowry theory of acids and bases

Brnsted-Lowry (1923) - Brnsted and Lowry proposed independent theories stating: An acid is a proton donor A base is a proton acceptor

Gave due recognition to the importance of the proton exchange in determining acidic/basic properties

Gave due recognition to the solvent

- A substance will act as an acid when it has a greater tendency to donate protons than the solvent- A substance will act as a base when it has a greater tendency to accept protons than the solvent

Allowed chemists to venture outside aqueous chemistry to non-aqueous solutions and in gas phases etc

i.e. A substance cannot behave as an acid without another behaving as a base

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

According to the Bronsted-Lowry theory of acids and bases, each acid has a conjugate base, and each base a conjugate acid.

An acid gives up a proton to form its conjugate base. A conjugate base has the tendency to accept a proton in solution

Acid + H2O H3O+ + conjugate base (Conjugate base = whats left over after proton is donated)

A base accepts a proton to form its conjugate acid. A conjugate acid has the tendency to donate a proton in solution

Base + H2O OH + conjugate acid (Conjugate acid = whats left over after proton is accepted)

Acid (in decreasing order of strength)Base (in decreasing order of strength)

HCl (hydrochloric)O2 (oxide ion)

H2SO4 (sulfuric)OH (hydroxide ion)

HNO3 (nitric)S2 (sulfide ion)

H3O+ (hydronium ion)CO32 (carbonate ion)

HSO4 (hydrogensulfate ion) NH3 (ammonia)

HF (hydrofluoric)HCO3 (hydrogencarbonate ion)

CH3COOH (ethanoic)HS (hydrogensulfide ion)

H2S (hydrogen sulfide)CH3COO (ethanoate ion)

H2CO3 (carbonic)F (fluoride)

NH4+ (ammonium ion)SO42 (sulfate ion)

HCO3 (hydrogencarbonate ion)H2O (water)

HS (hydrogensulfide ion) NO3 (nitrate ion)

H2O (water)HSO4 (hydrgensulfate ion)

OH (hydroxide ion)Cl (chloride ion)

(NOTE: strong acids have weak conjugate bases, and strong bases have weak conjugate acids [as seen in table above])

Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature [SKILL determining the acidic/basic nature of salts]

The Bronsted-Lowry theory accounts for the acidity or basicity of salts by water acting as the Bronsted-Lowry acid or base

The reaction of a salt with water to produce a change in pH is called hydrolysis

Acidic salts:

A strong acid & weak base react to form an acidic salt.

e.g. NH4Cl [possibly from NH4OH (weak base) and HCl (strong acid)]

Hydrolysis: NH4+ (aq) + H2O (l) NH3 (aq) + H3O+ (aq)>>>>>>[note the cation is acting as the acid]

Basic salts:

A strong base & weak acid react to form a basic salt.

e.g. KF [possibly from KOH (strong base) and HF (weak acid)]

Hydrolysis: F (aq) + H2O (l) HF (aq) + OH (aq)>>>>>>[note the anion is acting as the base]

e.g. NaCH3COO [possibly from NaOH (strong base) and CH3COOH (weak acid) Hydrolysis: CH3COO (aq) + H2O (l) CH3COOH (aq) + OH (aq)>>>>>>[note the anion is acting as the base]

e.g. NaCO3 [possibly from NaOH (strong base) and H2CO3 (weak acid)]Hydrolysis: CO32 (aq) + H2O (l) HCO3 (aq) + OH (aq)>>>>>>[note the anion is acting as the base]

Neutral salts:

If both the acid & base are strong, the resulting salt is neutral. The conjugates of strong acids and bases are so weak, that no reactions with the salt ions and the water molecules will take place.

e.g. NaCl [possibly from NaOH (strong base) and HCl (strong acid)]e.g. KBr [possibly from KOH (strong base) and HBr (strong acid)]

Similarly, if both the acid & base are weak, the resulting salt is usually neutral. This is because both anion and cation react to only small extents with water and so approximately cancel out each others effect on pH.

Identify conjugate acid/base pairs [SKILL]

AcidConjugate baseBaseConjugate acid

HClCl O2 OH

HNO3NO3OHH2O

H3O+H2OS2HS

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

Aim: To identify the pH range of salt solutions

Safety: Wear safety glasses to avoid substances coming into contact with eyes

Equipment: Spot dish, ammonium chloride (acidic salt), ammonium nitrate (acidic salt), sodium acetate (basic salt), sodium carbonate (basic salt), sodium chloride (neutral salt), potassium bromide (neutral salt)

Method: 1. 4 drops of 0.1mol/L NaCl were placed into a spot dish along with 3 drops of universal indicator. Any colour change was observed and recorded (Dependant variable indicator colour change). The colour change was matched to a calibrated pH chart to determine pH.

2. Step 1 was repeated with potassium bromide, ammonium chloride, ammonium nitrate, sodium acetate, and sodium carbonate. (Independent variable use of different salts)

3. Results from 5 other groups performing a similar investigation were collected and recorded to strengthen reliability.

Results: Ammonium chloride and ammonium nitrate were determined to be acidic.NH4+ (aq) + H2O (l) NH3 (aq) + H3O+ (aq)

Sodium acetate and sodium carbonate were determined to be basicCH3COO (aq) + H2O (l) CH3COOH (aq) + OH (aq)CO32 (aq) + H2O (l) HCO3 (aq) + OH

Sodium chloride and potassium bromide were determined to be neutral (lack of tendency to gain or accept protons)

Discussion: Testing with indicators is an example of destructive testing, as equilibrium is disturbed, and hence cannot be collected for reuse.

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

Amphiprotic substances can act as both a proton donor, and as a proton acceptor. (In basic conditions they tend to act as acids and vice-versa.)

According to the Brnsted-Lowry theory, if a substance has greater tendency to give up protons than the solvent, then it is an acid in the solvent. If the substance has greater tendency to accept protons than the solvent, it is a base in the solvent

So for some substances, if the solvent is a stronger proton donor, the substance will act as a base whilst if the solvent is a stronger proton acceptor, the substance will act as an acid. Such substances are said to be amphiprotic (either accepting or donating a proton). An example of this is the hydrogen carbonate ion:HCO3 (aq) + H3O+ (aq) H2CO3 (aq) + H2O (l)BASEACID conj. ACID conj. BASE

HCO3 (aq) + OH (aq) CO32 (aq) + H2O (l)ACIDBASE conj. BASE conj. ACID

[LESS IMPORTANT]Also whilst in water, both equations below occur simultaneouslyHCO3(aq)+H2O (l) H2CO3 (aq) +OH(aq)BASEACIDconj. ACIDconj.BASE

HCO3(aq)+H2O (l) CO32(aq)+H3O+ (aq)ACIDBASEconj. BASEconj.ACIDe.g Water is also amphiprotic HCl (aq) + H2O (l) H3O+ (aq) + Cl (aq) ACIDBASEconj. ACIDconj. BASE

NH3 (g) + H2O (l) NH4+ (aq) + OH (aq) BASEACIDconj. ACIDconj. BASE

Also, the self ionisation of water can be explained in terms of the B-L theory

H2O (l) + H2O (l) H3O+ (aq) + OH (aq)ACIDBASEconj. ACIDconj. BASE

HSO4 and HPO4 2- are also amphiprotic.

Identify neutralisation as a proton transfer reaction which is exothermic

Neutralisation is a proton transfer reaction (protons are transferred from the acid to the base to form a salt). All neutralisation reactions are exothermic.Acid + base water + salt + HeatAcid + carbonate carbon dioxide + water + saltIn the neutralisation of ammonia by nitric acid:

HNO3 (aq) + NH3 (aq) NH4+ (aq) + NO3-(aq) The acid transfers a proton to ammonia (to form the NH4+ ion)

Neutralisation is a proton transfer reaction in which an acid and a base react to form a salt and usually water. Neutralisation reactions are exothermic. The amount of heat liberated when neutralisation occurs depends on the strengths and concentrations of the acid and base.

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

Neutralisation reactions are an effective method to minimise the damage caused by hazardous acid/base spills by turning them into relatively harmless products such as salts, water, and possibly CO2.

Situations which require neutralisation Large scale outdoor acid/base spills Small scale indoor acid/base spills [isolate spill area, people evacuated, ventilate area if fumes present, spilt on clothing or people wash with copious amounts of water (as its weak and amphiprotic), sand can be used to soak up the acid, which can then be neutralised (e.g. with NaHCO3) and diluted (to reduce ion concentration) in a safe location] Industrial acidic/basic waste chemicals being discharged as effluent

Reasons for neutralisation Damage to waterways and aquatic life Excess heat caused by the exothermic reaction between acids and water can lead to thermal pollution, killing aquatic life and degrading water quality (reduced dissolved oxygen) Altered pH of water harmful to fish populations as fish die, or cannot reproduce Stunted plant growth some plant species require specific soil pH levels. Damage to built and metal structures Dangerous to people Chemical burns

Properties of suitable neutralising substances

Only weak acids/bases should be used for chemical spills:- As neutralisation reactions are exothermic, strong acids and bases would generate excessive and potentially dangerous levels of heat, as well as the possible production of noxious fumes. Best to dilute spilt acid first with water, then neutralise with powdered NaHCO3, (powder helps stop the spread of the liquid].

- Excess weak acids/bases would pose little problems to environment. However if strong acids/bases were used, besides the excess heat generated, it becomes easier to misjudge the correct quantities to neutralise, and both excess and insufficient quantities lead to detrimental environmental effects.

Powders are most suitable as they do not spread the spill further such as liquids. Amphiprotic substances are most suitable as they are effective in both acidic and basic spills- Useful when pH nature of substance is unknown - Convenience, as only the one substance needed in case of spillE.g. Sodium hydrogen carbonate (NaHCO3) is ideal as its a weak, amphiprotic power. It is also not very corrosive and hence easly handled.

HCO3- + H3O+ H2CO3 + H2O HCO3_ + OH- CO32- + H2O

Describe the correct technique for conducting titrations and preparation of standard solutions [SKILL volumetric analysis calculations]

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

Standard solution: A solution of accurately known concentration. There are two types of standard solutions; Primary standard solution A solution made up to an accurate concentration by combining a suitable chemical (the primary standard) with water. Standardised without a prior titration.

The solute (primary standard) needs to: Have a high level of purity Be stable in air (does not react with oxygen, water or carbon dioxide) Readily soluble in pure water Have an accurately known formula (Ideally) have a high molar weight to reduce percentage error during weighingBase standards: Anhydrous Na2CO3, NaHCO3, Acid standards: Oxalic acid (H2C2O4.2H2O)

Secondary standard solution A solution whose concentration is determined accurately by titrations with a primary standard. eg. Hydrochloric acid titrated against (a known concentration of) sodium carbonate solution becomes a secondary standard.

Practical

Aim: To prepare a primary standard solution from oxalic acid, and standardise sodium hydroxide as a secondary standard using titrations. Safety: Wear safety glasses to avoid substances coming into contact with eyes.

Method [PART 1 Preparing the primary standard]:

1. Volumetric flask and beaker were washed with distilled water. Required mass of oxalic acid to make up a 200mL 0.05mol/L standard was calculated.

2. The required mass of oxalic acid was accurately weighed using an electronic balance and transferred to a beaker

3. The oxalic acid was dissolved in just sufficient amounts of distilled water in the beaker.

4. The solution was transferred to the volumetric flask with a small, glass funnel. The beaker and funnel were washed with distilled water and the washings allowed to enter the volumetric flask. Distilled water was added until the bottom of the meniscus of solution reached the engraved line.

5. Flask was stoppered and inverted 5 times to mix completely

Method [PART 2 Standardising the secondary standard] NOTE: Either the acid or base involved needs to be strong, to ensure reaction goes to completion, or mole calculations would be invalid.

1. Burette, pipette, conical flask, and volumetric flask were washed with distilled water. Pipette was then washed with oxalic acid solution and burette was then washed with the unknown-concentration of NaOH.

2. Burette was clamped to a retort stand with a burette clamp. The burette was then filled (to the zero mark) with the NaOH. (Avoid air bubbles)

3. The (25mL) pipette was filled to the engraved line with the oxalic acid solution using a pipette filler. (Bottom of meniscus should reach engraved line)

4. The standard solution in the pipette (aliquot) was then transferred to the conical flask. (Small quantity left in the bottom of the pipette should not be forced out as the pipette has been calibrated to take this into account).

5. 3 drops of phenolphthalein indicator were added to the conical flask. The conical flask was then placed on a white tile under the burette. 6. The burette tap was opened slowly and allowed to run into the conical flask. The conical flask was gently swirled and the tap closed when the indicator just changes colour (the end point). The volume of unknown solution required to reach equivalence point was recorded, and used as a rough. 7. Steps 1-6 were repeated 10 times to strengthen reliability. The molarity of the NaOH was subsequently calculated.

H2C2O4.2H2O (aq) + 2NaOH (aq) Na2C2O4 (aq) + 4H2O (l)

# 25 mL of 0.05mol/L oxalic acid was used, n = c.V, therefore n oxalic acid = 0.00125mol # n oxalic acid : n NaOH = 1:2, therefore n NaOH = 0.0025 mol# Therefore, c NaOH = 0.0025 mol / 0.0311 L (titre) = 0.0804 mol/L

Choosing the right indicator: An appropriate indicator will give a sharp colour change at the equivalence point (the inflexion of the titration curve). The choice is predictable by the acidic, basic or neutral nature of the salt. # Strong acid + strong base = neutral salt, therefore use litmus or bromothymol blue # Strong acid + weak base = acidic salt, therefore use methyl orange or bromothymol blue # Strong base + weak acid = basic salt, therefore use phenolphthalein Weak acids are not titrated against weak bases because there is no sudden change in pH to indicate the equivalence point. Hence unknown concentration cannot be calculated.

End point = Point when the indicator changes colourEquivalence point = Point when the acid and base have reacted stoichiometrically If the indictor is chosen correctly, end point = equivalence point Aliquot = Volume measured by the pipetteTitre = Volume delivered by the buretteRough = The initial titration to locate the approximate end point.

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

Aim: To determine the concentration of acetic acid in commercial vinegar using computer based technologies.

Safety: Wear safety glasses to avoid substances coming into contact with eyes

Method:

1. 25mL of commercial vinegar was made up to a 250mL solution with distilled water using a volumetric flask (i.e. 1 in 10 dilution). 2. Pipette, burette, and beaker were washed with distilled water. Pipette was washed once more with dilute vinegar solution, and burette was washed once more with NaOH solution (the secondary standard)3. Pipette was used to measure and transfer 25mL of dilute vinegar solution into the beaker. 3 drops of phenolphthalein indicator are added.4. Burette was filled with NaOH secondary standard solution. 5. A data logger attached to a pH probe was set up and placed into the beaker. The equipment was then set up as diagram above. The initial pH was recorded.6. The burette tap was gently opened and 2mL of NaOH solution was allowed to enter the beaker. The pH was then recorded. (Independent variable volume of NaOH solution added) (Dependent variable pH reading)7. Step 6 was repeated 6 more times after a sharp colour change was observed. The pH of solution was plotted as a function of NaOH added with computer software. The volume of NaOH required to reach the equivalence point was subsequently determined, and the concentration of commercial vinegar calculated. 8. Steps 1-6 were repeated 5 times to strengthen reliability. Results:NaOH (aq) + CH3COOH (aq) NaCH3COO (aq) + H2O (l)

[example figures] 38mL of 0.1 mol/L NaOH used, n=cV, therefore n NaOH = 0.1 x 0.038 = 0.0038 moln NaOH : n CH3COOH = 1:1, therefore n CH3COOH = 0.0038 molTherefore c CH3COOH = 0.0038 mol / 0.025 L = 0.152mol/LBut this is the concentration of ethanoic acid in the diluted vinegarTherefore the concentration of ethanoic acid in the commercial vinegar was = 1.52mol/L

Validity: The method is valid in determining the concentration of acetic acid in commercial vinegar. This is because equipment was washed with correct substances (everything once with distilled water, burette once with NaOH and pipette once with dilute vinegar) to ensure lack of contamination and hence improved accuracy of pH measurements. Calibrated volumetric equipment (pipette, volumetric flask, burette) was used, and this increased the accuracy of the volumes measured, and hence the accuracy of the pH measurements.

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

A buffer is a solution that resists rapid changes in pH to maintain an approximately constant pH. It contains similar amounts of a weak acid and its conjugate base (eg. H2CO3 and HCO3 ).

By Le Chateliers principle, when small amounts of an acid or base are added, the system will minimise the change and thus maintain an approximately constant pH.

HA + H2O H3O+ + AIf an acid was introduced into the system, then by Le Chateliers principle, the equilibrium will shift left to minimise the increase in H3O+ concentration in order to return to equilibrium. A + H2O HA + OHIf a base was introduced into the system, then by Le Chateliers principle, the equilibrium above will shift left to use up the added OH in order to return to equilibrium.

Example: Blood is a buffered solution, containing carbonic acid (H2CO3) and sodium bicarbonate (NaHCO3). This allows its pH to remain at a constant 7.4. Carbon dioxide is dissolved in blood, forming carbonic acid, which is a weak acid that ionise to form the hydrogen carbonate ion (HCO3).

1 CO2 (g) +H2O (l) H2CO3 (aq)2H2CO3 (aq) + H2O (l) H3O+ (aq) + HCO3 (aq)3HCO3 (aq) + H2O (l) H2CO3 (aq) + OH (aq)

If more CO2 is dissolved in blood, more H3O+ ions will form. This would lower the pH, causing the equilibrium (2) to shift to the left to use up the added H3O+ ions, forming more H2CO3 molecules. The increase in H2CO3 molecules causes equilibrium (3) to shift left to use up some of the H2CO3 molecules, forming more HCO3 ions. This results in a fairly constant pH range.

If the pH is too high, equilibrium (3) shifts left use up some of the OH ions, forming more HCO3 ions. This increase in HCO3 ions causes equilibrium (2) to shift left, forming more H2CO3 molecules. This again results in a fairly constant pH range.

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

Alkanols (n=1, 2)[CAREFUL] Alkanoic acids (n = 0, 1)(count the number of carbons)

CnH2n+1OHCnH2n+1COOH

Both alkanols and alkanoic acids are organic compounds.

Alkanols are alcohols which consist of a hydroxyl functional group (-OH) attached to an alkyl group. The polar nature of the hydroxyl group (and CO bonds, and asymmetric shape) makes alkanoic acids polar Their polar nature allows them to be soluble in polar solvents (e.g. water) due to the formation of hydrogen bonds. However, as with alkanoic acids, the longer the carbon chain, the less soluble they are (due to the dominance of dispersion forces over dipole-dipole forces) Also due to the hydrogen bonding between molecules, alkanols have higher M.P and B.Ps than corresponding alkanes. This is due their ability to hydrogen bond between molecules, and their greater molar mass, causing greater dispersion forces. M.P and B.P is related to the strength of intermolecular forces.

Alkanols are essentially neutral molecules. The hydroxyl functional group is strongly bonded to the carbon chain, and there is no tendency for the loss of hydroxide ions when alkanols dissolve in water.

Alkanoic acids are carboxylic acids which consist of a carboxyl functional group (-COOH) attached to an alky group. The polar nature of the carboxyl group (and C=O, CO, and asymmetric shape) makes alkanoic acids polar Their polar nature allows them to be soluble in polar solvents (e.g. water) due to the formation of hydrogen bonds. However, as with alkanols acids, the longer the carbon chain, the less soluble they are. Also due to the hydrogen bonding between molecules, alkanoic acids have higher M.P and B.Ps than corresponding alkanes or alkanols. This is due to their more extensive hydrogen bonding (than alkanols) between molecules, and their greater molar mass, causing greater dispersion forces. M.P and B.P is related to the strength of intermolecular forces. The longer the carbon chain, the greater the molar mass, and the greater the dispersion forces, hence the greater the M.P and B.P.

Alkanoic acids are weak acids

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

The M.P and B.Ps of alkanoic acids are greater than that of corresponding alkanols (or alkanes). This is because

- Alkanoic acids form more extensive hydrogen bonds than alkanols. (They can hydrogen bond due to the polar nature of their OH and COOH functional groups) This is because the carboxyl group (-COOH) can hydrogen bond two times with another carboxyl group. In alkanols, only one hydrogen bond is possible between molecules due to its single -OH group.

Hydrogen bonding in Alkanols: H-bonding in Alkanoic Acid:- Also, alkanoic acids have greater dispersion forces between molecules than corresponding alkanols. This is because alkanoic acids have greater molar masses than corresponding alkanols, causing more random movements of electrons, and hence causing greater dispersion forces.

Hence intermolecular forces are greater in alkanoic acids, and as M.P and B.P relate to the strength of intermolecular forces, alkanoic acids have higher M.P and B.Ps.

Also as the carbon chain increases, there is a trend of increasing M.P and B.P with increasing molar weight due to the increase in dispersion forces.

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 [SKILL naming /drawing esters]

1) Count the number of carbon atoms in the alkyl chain (separated by (inc.) O atom). (This came from the alkanol). Name the alkyl group (e.g. methyl, ethyl, propyl etc.)

2) Count the number of carbon atoms in the alkanoate chain (other side of the O atom). (This came from the alkanoic acid). Name the alkanoate group (methanoate, ethanoate, propanoate) [i.e. drop the oic with an oate]

3) Name the alkyl (from alkanol) then the alkanoate (alkanoic acid).

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

Esterification is the acid-catalysed (condensation) reaction between an alkanol and acid to produce an ester

Esters are organic compounds formed by the acid-catalysed (condensation) reaction between an alkanol and acid (esterification).

Alkanol + alkanoic acid alkyl alkanoate (ester) + water[Condensation reaction OH comes from the alkanoic acid, H comes from alkanol][Ester linkage / ester functional group ( COO )]

Esters are named as alkyl alkanoates, where the alkyl is the abbreviated form of the alkanol used and alkanoate the form of the alkanoic acid. For example:

Ethanol + Butanoic Acid Ethyl Butanoate + WaterEg: propanol + propanoic acid

Describe the purpose of using acid in esterification for catalysis Concentrated H2SO4 is used as a catalyst for two reasons: Catalyst Speed up rate of reaction by lowering the activation energy Dehydrating agent - Absorbs water, shifts equilibrium right increasing the yield of ester.

Explain the need for refluxing during esterification

Refluxing = The process of heating a mixture of liquids in a flask with an attached condenser in order to prevent the loss of volatile reactants or products

Heating = (increase the kinetic energy of the molecules, allowing for more collisions)

The esterification process requires heating, as the reaction proceeds slowly at room temperature. However, as the reactants and products involved in esterification are volatile, they readily vaporise on heating. To overcome this, refluxing is done heating the mixture with an attached water condenser.

Hence refluxing during esterification is done to accelerate the reaction (through heating) as the reaction is slow at room temperatures, and to avoid the loss of volatile chemicals to the atmosphere by condensing them with an attached water condenser.

[As the mixture is heated, the volatile components evaporate and rise into the condenser, where they are cooled and condensed by flowing water, and fall back into the reaction mixture.][The refluxing apparatus is open to the atmosphere to avoid the build up of pressure due to the production of vapours] [Vapours are flammable avoid build up of pressure, use a heated water bath etc.]

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

Aim: To prepare an ester using reflux

Safety: # Wear safety glasses to avoid substances coming into contact with eyes. # Alkanols are volatile work in a ventilated area. # Alkanols and alkanoic acids are flammable have a fire blanket/extinguisher at hand and avoid direct contact with naked flame. # Glassware is breakable carefully handle with firm supporting hands (both of them)

Equipment: retort stand, bosshead and clamp, bunsen burner, condensation tube, pear shaped flask

Method:

1. The following were added to the pear shaped flask- 7mL 1-butanol - 10mL of ethanoic acid (excess is used to shift the equilibrium towards the formation of ester)- 5mL of concentrated H2SO4 (as a catalyst and dehydrating agent)- 4 ceramic boiling chips (to provide a large surface area for vaporisation in order to avoid sudden superheating)

2. Equipment was set up as diagram above with water flowing in the condenser at a mild, uniform rate

3. The Bunsen burner was ignited and allowed to heat with a quite, blue flame. This was done for 30 minutes

4. Equipment was allowed to cool, then the pear-shaped flask detached. The solution was poured into a separating funnel and 25mL distilled water added. The separating funnel was shaken, left standing to allow layers to separate, and then the bottom aqueous layer allowed to run out. This was repeated two more times.

5. 15mL distilled water and 15mL of NaHCO3 were then added to the separating funnel and shaken (to neutralise the acid), left standing to allow layers to separate, and then the bottom aqueous layer allowed to run out.

Results: 1 butyl ethanoate (used in production of plastics) was produced.

The volatility of the reagents of esterification means that the extended heating must be performed under reflux. As the flask is heated the components vaporise and are then condensed due to the cooling water provided by the condensing tube. The reagents thus return to the flask and eventually equilibrium is achieved without loss of material.

Outline some examples of the occurrence, production and uses of estersEstersNaturally - The fragrance and flavour of flowers and fruits - Animal fats and plant oils (triesters)

Industry - Solvents or thinners as they are able to dissolve many polar and non polar substances (butyl-ethanoate)- Plastic production (butyl-ethanoate)- Artificial flavourings of food (pentyl ethanoate banana ester) - Flavouring and scenting cosmetics. (Lip gloss and perfumes)- Nail polish remover (ethyl ethanoate) - Jet engine lubricant (due to low viscosity at low temperatures)

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

Esters occur widely in nature as perfumes or flavourings in flowers and fruits. Within nature, a particular scent or flavour is usually of a complex mixture, but sometimes one particular ester can be identified as the major ingredient, and thus can be reproduced artificially.

Artificially produced, these are much cheaper than natural extraction. In processed foods = confectionary and ice cream flavouring (e.g. ethyl methanoate for artificial run flavouring without the alcohol content),

Cosmetics = In perfumes (e.g pentyl ethanoate), flavoured lip gloss (butyl butanoate), creams and lotions to sooth sore muscles (methyl salicylate)

Sam Hall33