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Contents Section 1 – Indicators were identified with the observation that the colour of some flowers depends
on soil composition ................................................................................................................................. 4
1.1 Classify common substances as acidic, basic or neutral ............................................................... 4
1.2 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 range is
identified by change in colour. ........................................................................................................... 4
1.3 Identify and describe some everyday uses of indicators including the testing of soil
acidity/basicity .................................................................................................................................... 5
1.1 Perform a first‐hand investigation to prepare and test a natural indicator ................................. 6
1.2 Identify data and choose resources to gather information about the colour changes of a range
of indicators ........................................................................................................................................ 6
1.3 Solve problems by applying information about the colour changes of indicators to classify
some household substances as acidic, neutral or basic ..................................................................... 6
Section 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 .............................................................................................. 7
2.1 Identify oxides of non‐metals which act as acids and describe the conditions under which they
act as acids .......................................................................................................................................... 7
2.2 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 .......................... 7
2.3 Define Le Chatelier’s principle ...................................................................................................... 8
2.4 Identify factors which can affect the equilibrium in a reversible reaction ................................... 8
2.5 Describe the solubility of carbon dioxide in water under various conditions as an equilibrium
process and explain in terms of Le Chatelier’s principle .................................................................... 9
2.6 Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen ....................... 11
2.7 Describe, using equations, examples of chemical reactions which release sulfur dioxide and
chemical reactions which release oxides of nitrogen ....................................................................... 11
2.8 Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur
and nitrogen ...................................................................................................................................... 13
2.9 Calculate volumes of gases given masses of some substances in reactions, and calculate
masses of substances given gaseous volumes, in reactions involving gases at 0˚C and 100kPa or
25˚C and 100kPa ............................................................................................................................... 13
2.10 Explain the formation and effects of acid rain .......................................................................... 13
2.1 Identify data, plan and perform a first‐hand investigation to decarbonate soft drink and gather
data to measure the mass changes involved and calculate the volume of gas released at 25˚C and
100kPa ............................................................................................................................................... 14
2.2 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 ..................................................................................................................................... 14
Section 3 – Acids occur in many foods, drinks and even within our stomachs .................................... 15
3.1 Define acids as proton donors and describe the ionisation of acids in water ............................ 15
3.2 identify acids including acetic (ethanoic), citric (2‐hydroxypropane‐1,2,3‐tricarboxylic),
hydrochloric and sulphuric acid ........................................................................................................ 15
3.3 Describe the use of the pH scale in comparing acids and bases................................................. 16
3.4 Describe acids and their solutions with the appropriate use of the terms strong, weak,
concentrated and dilute.................................................................................................................... 16
3.5 identify pH as ‐log10 [H+] and explain that a change in pH of 1 means a ten‐fold change in [H+]
.......................................................................................................................................................... 16
3.6 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 ................................................ 17
3.7 Describe the difference between a strong and weak acid in terms of an equilibrium between
the intact molecule and its ions ........................................................................................................ 17
3.1 Solve problems and perform a first‐hand investigation to use pH meters/probes and indicators
to distinguish between acidic, basic and neutral chemicals ............................................................. 18
3.2 Plan and perform a first‐hand investigation to measure the pH of identical concentrations of
strong and weak acids ....................................................................................................................... 18
3.3 Gather and process information from secondary sources to write ionic equations to represent
the ionisation of acids ....................................................................................................................... 18
3.4 Use available evidence to model the molecular nature of acids and simulate the ionisation of
strong and weak acids ....................................................................................................................... 18
3.5 Gather and process information from secondary sources to explain the use of acids as food
additives ............................................................................................................................................ 19
3.6 Identify data, gather and process information from secondary sources to identify examples of
naturally occurring acids and bases and their chemical composition .............................................. 19
3.7 process information from secondary sources to calculate pH of strong acids given appropriate
hydrogen ion concentrations ............................................................................................................ 20
Section 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. ........................... 21
4.1 Outline the historical development of ideas about acids including those of: ............................ 21
‐ Lavoisier .......................................................................................................................................... 21
‐ Davy ................................................................................................................................................ 21
‐ Arrhenius......................................................................................................................................... 21
4.2 Outline the Brönsted‐Lowry theory of acids and bases .............................................................. 21
4.3 Describe the relationship between an acid and its conjugate base and a base and its conjugate
acid .................................................................................................................................................... 22
4.4 Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic,
neutral or basic nature ...................................................................................................................... 23
4.5 Identify conjugate acid/base pairs .............................................................................................. 24
4.6 Identify amphiprotic substances and construct equations to describe their behaviour in acidic
and basic solutions ............................................................................................................................ 24
4.7 Identify neutralisation as a proton transfer reaction which is exothermic ................................ 25
4.8 Describe the correct technique for conducting titrations and preparation of standard solutions
.......................................................................................................................................................... 25
4.9 Qualitatively describe the effect of buffers with reference to a specific example in a natural
system ............................................................................................................................................... 28
4.1 Gather and process information from secondary sources to trace developments in
understanding and describing acid/base reactions .......................................................................... 30
4.2 Choose equipment and perform a first‐hand investigation to identify the pH of a range of salt
solutions ............................................................................................................................................ 30
4.3 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 ................................................................ 30
4.4 Perform a first‐hand investigation to determine the concentration of a domestic acidic
substance using computer‐based technologies ................................................................................ 30
4.5 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 ........................................... 30
Section 5 – Esterification is a naturally occurring process which can be performed in the laboratory31
5.1 Describe the differences between the alkanol and alkanoic acid functional groups in carbon
compounds ....................................................................................................................................... 31
5.2 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 .......... 31
5.3 Explain the difference in melting point and boiling point caused by straight‐chained alkanoic
acid and straight‐chained primary alkanol structures ...................................................................... 32
5.4 Identify esterification as the reaction between an acid and an alkanol and describe, using
equations, examples of esterification ............................................................................................... 33
5.5 Describe the purpose of using acid in esterification for catalysis .............................................. 33
5.6 Explain the need for refluxing during esterification ................................................................... 34
5.7 Outline some examples of the occurrence, production and uses of esters ............................... 34
5.2 process information from secondary sources to identify and describe the uses of esters as
flavours and perfumes in processed foods and cosmetics ............................................................... 34
5.1 Identify data, plan, select equipment and perform a firsthand investigation to prepare an ester
using reflux ........................................................................................................................................ 35
Section 1 – Indicators were identified with the observation that the
colour of some flowers depends on soil composition
1.1 Classify common substances as acidic, basic or neutral
Acids (<7)
Common substances (acids): vinegar (acetic acid), lemon juice (citric acid), aspirin, vitamin C,
car batteries (sulphuric acid), carbonated drinks, rain can be slightly acidic
An acid is a substance which in solution produces hydrogen ions, H+ or more specifically H3O+
(hydronium ions)
Common properties:
o Sour taste
o Sting/burn the skin
o In solution, they conduct electricity
o Acids turn blue litmus red (Blue in Acid goes Red – BAR)
Bases (>7)
Common substances (bases): cloudy ammonia (for cleaning), washing soda (sodium
carbonate), antacid tablets (neutralise acids in our stomach – contain calcium carbonate
and/or magnesium hydroxide), oven and drain cleaners (sodium hydroxide), garden lime
(calcium carbonate), bleach, detergent
A base is a substance which either contains the oxide O2‐ or hydroxide ion OH‐ or which in
solution produces the hydroxide ion.
o A soluble base is called an alkali (contain, or in aqueous solution produce hydroxide
ions)
Common bases: sodium hydroxide (NaOH), barium hydroxide (Ba(OH)2), potassium oxide
(K2O), magnesium oxide (MgO), iron (III) oxide (Fe2O3), copper hydroxide (Cu(OH)2), and
ammonia (NH3)
o Only the first three and ammonia are alkalis
Common properties:
o Have a soapy feel
o Bitter taste
o In solution, alkalis are good conductors of electricity
o Turn red litmus blue
Neutral (=7) – blood (also slightly basic), water, sea/salt water, distilled water, alcoholic drinks,
sucrose solutions (sugar)
1.2 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 range is
identified by change in colour.
An indicator is a substance (usually a vegetable dye) that in solution changes colour
depending on whether the solution is acidic or basic
E.g. red cabbage indicator (universal) – red (acid) ‐> purple ‐> green (neutral) ‐> yellow
(base)
Litmus or bromothymol blue can be used to decide whether a solution is acidic, neutral or
basic and then methyl orange or phenolphthalein can be used to decide whether the
solution is highly/slightly acidic/basic
Methyl orange – 3.2‐4.4
Litmus – 5‐8
Bromothymol blue – 6.2‐7.6
Phenolphthalein – 8.2‐10
1.3 Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity
Everyday uses:
Testing the acidity/alkalinity of soils – some plants need an acidic soil (e.g. azaleas and
camellias) while other need an alkaline or near neutral soil (most annual flowers and
vegetables)
o Soil is partially dissolved in distilled water, and white barium sulphate powder is
added (so colour can be seen) because it is neutral and insoluble
o Universal indicator is added to test pH
o Can also use pH probe/meter
Testing home swimming pools (need to be approx. neutral, though adding chemicals to
sanitise the water can change its acid‐alkali balance) – use universal indicator
o Too acidic – corrosive to structural metals, causes excessive release of chlorine into
air (water is no longer disinfected), damages swimwear, stings/burns eyes and skin
o Too basic – causes scale build up, reduces capacity of chlorine to disinfect,
stings/burns eyes and skin
Monitoring wastes from laboratories that process photographic film (discharges to the
sewerage system must be nearly neutral: photographic solutions are often highly alkaline)
Testing aquarium water – pH of water must be suitable to variety of fish
1.1 Perform a first‐hand investigation to prepare and test a natural indicator
1.2 Identify data and choose resources to gather information about the colour changes of a range of
indicators
1.3 Solve problems by applying information about the colour changes of indicators to classify some
household substances as acidic, neutral or basic
Section 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
2.1 Identify oxides of non‐metals which act as acids and describe the conditions under which they act
as acids
What are oxides?
A class of compounds (all covalent) that contain oxygen and another element
Often display either acidic or basic properties
A small few are neutral
A small few are amphoteric ‐ (don't fit into any one category, so they get their own)
Acidic Oxides
Most oxides of non‐metals (with the exception of CO, NO and N2O) act as acids
Only do so when dissolved in water (gas dissolves into the water molecules, forming an acid)
– moist litmus paper can be used to test the acidity of gases
Or reacts with bases to form salts (or does both)
E.g. CO2 (g) + H2O (l) H2CO3 (aq)
Water + carbon dioxide carbonic acid
E.g. P2O5 (s) + 6NaOH (aq) 2Na3PO4 (aq) + 3H2O (l)
Diphosphorous pentoxide + sodium Hydroxide water + sodium phosphate
Other acidic oxides – NO2, P2O3, SO2, SO3, Cl2O
2.2 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
Basic oxides are oxides of metals (formed when group I or group II metals react with oxygen)
– they become more and more basic as you move down the periods
Most non‐metals produce acidic oxides – as the non‐metals become more metallic in
character, they decrease in acidity
2.3 Define Le Chatelier’s principle
A system at equilibrium, when introduced to a change, will shift in equilibrium to minimise that
change.
2.4 Identify factors which can affect the equilibrium in a reversible reaction
Changing concentration (or pressure with a gas) 1. Increasing concentration of A: This promotes the reaction that will decrease the concentration
of A, i.e. by reacting it with B and turning it into C + D. The position of equilibrium moves to the right.
e.g. CH3COOH + H2O ⇌ H3O+ + CH3OO‐ diluting/adding more water, will drive the equation to the right as it wants to achieve an equilibrium, thus the concentration of hydronium ions will go up and the pH will decrease.
2. Decreasing the concentration of C (you can achieve this by removing C as it forms): This will drive the forward reaction. It is an effective means to forcing an equilibrium to assume a one‐way reaction.
Changing the pressure of the system ‐ ONLY APPLIES TO GAS REACTIONS 1. Increasing the pressure: This will promote the reaction that will decrease the pressure ‐ i.e. the
direction that results in less molecules. A + 2B = 3 molecules, C + D = 2 molecules
NOTE: a single mole of gas will always occupy the same volume
2. Decreasing the pressure:
NOTE: reactions with equal numbers of molecules on each side of the equilibrium will be unaffected by pressure changes.
Changing the temperature 1. Increasing the temperature: This promotes the endothermic reaction that will decrease the
temperature as energy is absorbed.
2. Decreasing the temperature: This promotes the exothermic reaction that will increase the temperature as energy is given out.
2.5 Describe the solubility of carbon dioxide in water under various conditions as an equilibrium
process and explain in terms of Le Chatelier’s principle
Chemical Equilibrium Some reactions do not go to completion ‐ rather the end result is a mixture of reactants and
products Equilibrium reactions are DYNAMIC: the rate of the forward and reverse reactions are identical
in all condition Conditions for establishing and maintaining equilibrium The system is closed: nothing can leave or enter the reaction vessel No observable changes will occur at the macroscopic level: e.g. no colour change will be seen,
despite the fact that there are reaction occurring at the microscopic level The reactant and product concentrations remain constant (once equilibrium is established)
Carbon dioxide is only slightly soluble in water ‐ and its solubility varies with pressure and temperature Removal of carbon dioxide from the body Transport of CO2 for photosynthesis Carbonating drinks ‐ Carbon dioxide is the gas that makes carbonated drinks fizzy (reacts with
water to make carbonic acid) CO2 (g) + H2O (l) ⇌ H2CO3 (aq) ∆H = ‐19.4 kJ/mol 1 : 0 ‐‐> moles of gas 2 : 1 ‐‐> moles of whole system
CO2 (g) ⇌ CO2 (aq)
H2O (l) + CO2 (aq) ⇌ H2CO3 (aq)
H2CO3 (aq) ⇌ H+ (aq) + HCO3‐ (aq)
HCO3‐ (aq) ⇌ H+ (aq) + CO3
2‐
The position of the equilibrium can be shifted to the left or right depending on the conditions: 1. Increase pressure of CO2 (e.g. by pumping more CO2 into the reaction vessel))
o The system tries to decrease the pressure of CO2 o The equilibrium is driven to the RHS
Decrease pressure of CO2 o The system tried to increase the pressure of CO2 o The equilibrium is driven to the LHS
2. Increase pressure of the system (compression of the gas above) o The system tries to decrease the pressure o The equilibrium is driven to the right where there are fewer molecules o More carbon dioxide gas is dissolved
*this is why fizzy drinks are under pressure ‐ to get more CO2 into solution 3. Increase temperature of system
o The system tries to decrease the temperature of the system o The equilibrium is driven in the direction of the endothermic reaction ‐ i.e. to the LHS o Creates more carbon dioxide gas
*we keep fizzy drinks cold so the drink stays fizzy 4. Increasing acid concentration (adding H+)
o Increase in acid concentration will cause equilibrium to shift to the left – to ‘use up’ the added acid
o More carbon dioxide gas will be produced Increasing base concentration (adding OH‐) o H2CO3 (aq) + 2OH‐ (aq) ‐‐> 2H2O (l) + CO3
2‐ (aq) o The above reaction occurs, which has the effect of removing H2CO3 to ‘use up’ the
added base o The system tries to replace H2CO3 o The equilibrium is driven to the right
NOTES:
Shaking the bottle of drink when the cap is removed speeds up the release of carbon dioxide ‐> equilibrium is reached more rapidly
Carbon dioxide has a very low solubility in solid ice ‐> ice freezes, CO2 originally dissolved is forced as gas into the small gas space ‐> builds excessive pressure and bottle may explode
2.6 Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen
2.7 Describe, using equations, examples of chemical reactions which release sulfur dioxide and
chemical reactions which release oxides of nitrogen
Sulfur Dioxide
Naturally, sulphur dioxide is released from natural sources such as geothermal hot springs
and volcanoes
Burning organic matter (as in bushfires)
Oxidation of hydrogen sulphide produced during decay of organic matter
Main human activities that release sulphur dioxide to the atmosphere are
Processing or burning fossil fuels
o Coal contains some sulphur – mainly as metallic sulphides and sulphur in carbon‐
containing compounds
o Not practical to extract sulphur from coal before burning so in the combustion
process sulphur is converted to sulphur dioxide
o S (in compounds) + O2 (g) SO2 (g)
o Most coal‐burning power stations remove most of the sulphur dioxide, but low‐
sulfur coal is widely favoured/in demand
o Crude oil and sometimes natural gas contain a variety of sulphur compounds –
extracted at oil refineries or gas processing plants, some release of sulphur dioxide
o E.g. Sulphur compounds are oxidised to produce sulphur dioxide:
4FeS2 (s) + 11O2 (g) 2Fe2O3 (s) + 8SO2 (g)
Extracting metals from sulphide ores
o Many metals occur as sulphides (e.g. copper, lead, zinc, silver, nickel)
o First step is to roast the sulphide ore in air
o Produces sulphur dioxide along with the metal or metal oxide
o E.g. zinc
2ZnS (s) + 3O2 (g) 2ZnO (s) + 2SO2 (g)
o Sulphur dioxide is released to the atmosphere
o E.g. Smelting of chalcopyrite:
2CuFeS2 (s) + 5O2 (g) + 2SiO2 2Cu (l) + 2FeSiO3 (l) + 4SO2 (g)
Nitrous oxides
Major natural source of nitric oxide and nitrogen dioxide is lightning
o High localised temperatures generated causes nitrogen in the atmosphere to react
with oxygen
o O2 (g) + N2(g) 2NO (g)
o Nitric oxide slowly reacts with oxygen to form nitrogen dioxide
o 2NO (g) + O2(g) 2NO2 (g)
Nitrous oxide is formed naturally by action of certain bacteria or nitrogenous material in
soils
Also produced in combustion ‐ stationary sources (power stations) and moving ones (cars
and trucks)
o High temperatures in the combustion chambers
o Same reactions as caused by lightning
Human activity has increased release of nitrous oxide through the increased use of
nitrogenous fertiliser which provides more raw material for the bacteria ‐> enhances
greenhouse effect
2.8 Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur
and nitrogen
Analysis of air bubbles in ice cores taken from the Antarctic ice show evidence of an increase
of over 25% in atmospheric acidic oxide levels over the past two centuries
Carbon isotopes in trees and calcium carbonate in coral
Largely attributed to the burning of fossil fuels (in motor vehicles and electricity generators)
with sulphur and nitrous impurities
Systems to control fossil fuel emissions have stabilized air quality, however there is still a
clear increase as dependence on fossil fuel energy increases
Lack of reliable data for periods before about 1950 – hard to tell increase in SO2 and NO2
http://www.easychem.com.au/the‐acidic‐environment/acidic‐oxides/atmospheric‐concentrations‐
of‐oxides
The emissions of acidic oxides due to increasing industrialisation and growing population –
increase in combustion of fossil fuels
Increasing vehicles on roads add to sulphur + nitrogen oxide emissions
As NO2 and SO2 are regularly flushed out of the atmosphere by rain, there has not been any
appreciable build‐up during the 20th Century.
However, there is a lack of reliable data – technology that could measure concentrations
below 0.1 ppm (parts per million) was only developed in the 70s. (Their current
concentrations are around 0.001 ppm in clean air and 0.01 ppm in urban areas)
N2O has shown a steady accumulation over this time – over last 150 years, N2O has
increased by 15%.
2.9 Calculate volumes of gases given masses of some substances in reactions, and calculate masses
of substances given gaseous volumes, in reactions involving gases at 0˚C and 100kPa or 25˚C and
100kPa
Calculating Gas Volumes
One mole of any gas has the same volume as a mole of any other gas as long as they are at the same temp and pressure (MOLAR VOLUME)
If temp or pressure changes, volume will change
We must always state the conditions
Temp ‐ either 0oC or 25oC
Pressure ‐ 100 kPa (approx. 1 atmosphere)
n(gas) = V/VM
2.10 Explain the formation and effects of acid rain
high emissions of sulphur and nitrogen dioxides to the atmosphere lead to acid rain
these pollutants are converted to sulphuric and nitric acids which dissolve in the rain
o sulphur dioxide reacts to form sulphurous acid
SO2 (g) + H2O (l) H2SO3 (aq)
o various other impurities in air catalyse the conversion of sulphurous acid to
sulphuric acid
2H2SO3 (aq) + O2 (g) 2H2SO4 (aq) (with catalyst)
o nitrogen dioxide reacts with water to form a mixture of nitrous and nitric acids
2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)
o nitrous acid in solution is catalytically oxidise to nitric acid
2HNO2 (aq) + O2 (g) 2HNO3 (aq) (with catalyst)
acid rain has a higher hydrogen ion concentration that normal (higher than 10‐5 mol/L)
rain that is unaffected by any pollutants is slightly acidic because of dissolved carbon dioxide
which forms carbonic acid
Effects
increasing acidity of lakes and waterways ‐> poison human and marine life
damage to pine forests
erosion of marble and limestone of building surfaces and decorations – carbonates react
with acids
severe damage to vegetation – especially around mine and smelter sites
2.1 Identify data, plan and perform a first‐hand investigation to decarbonate soft drink and gather
data to measure the mass changes involved and calculate the volume of gas released at 25˚C and
100kPa
Soda experiment (one note)
2.2 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
Section 3 – Acids occur in many foods, drinks and even within our
stomachs
3.1 Define acids as proton donors and describe the ionisation of acids in water
acids are described as ‘proton donors’ – they transfer protons to other substances
water is not only made up of H2O molecules – there are small amounts of hydroxide and
hydrogen ions
H2O (l) ⇌ H+ (aq) + OH‐ (aq)
When acids are placed in water, they ionize by donating protons ‐> ionisation reaction
This forms a negative ion, and the H+ ion (which is just a proton) joins the water molecules,
forming a hydronium ion (H3O+)
HA + H2O H3O+ + A‐
E.g. nitric acid
HNO3(l) + H2O (l) H3O+ (aq) + NO3‐ (aq) or simply as HNO3 H+ + NO3
‐
3.2 identify acids including acetic (ethanoic), citric (2‐hydroxypropane‐1,2,3‐tricarboxylic),
hydrochloric and sulphuric acid
Hydrochloric Acid
HCl Found in the stomach to dissolve proteins ‐ monoprotic
Sulfuric Acid H2SO4 Found in manufacturing fertilizers ‐ diprotic
Citric Acid aka 2‐hydroxypropane‐1,2,3‐tricarboxylic acid
C6H8O7 Found in all citrus fruits ‐ triprotic
Acetic Acid aka ethanoic acid
CH3COOH Found in vinegar – monoprotic
Strong acids: HCl, H2SO4, HNO3
Weak acids: HF, CH3COOH, citric acid, carbonic acid: H2CO3
3.3 Describe the use of the pH scale in comparing acids and bases
pH of a solution is defined as the negative of the logarithm of the hydrogen ion
concentration
pH = ‐log10[H3O+]
this (logarithmic) scale goes from ‐1‐>14 and is used to compare acids and bases
as concentration of hydronium ions drops down, the pH goes up
as concentration of OH ions goes up, the pH goes up
3.4 Describe acids and their solutions with the appropriate use of the terms strong, weak,
concentrated and dilute
Classification Degree of H+ Donation Concentration
Slightly acidic Weak <100% Dilute e.g. 0.1 mol HCl
Highly acidic Strong 100 % Concentrated e.g. 10 mol HCl
strong – an acid or base that completely ionizes in water
weak ‐ an acid or base that does not completely ionize in water
concentrated – an acid or base that is concentrated has a high concentration of
hydrogen/hydroxide ions
dilute – an acid or base that is dilute has a low concentration of hydrogen ions or hydroxide
ions
Concentrated v dilute – simply referring to how much solute is dissolved in the water (solvent) – a
strong acid can be dilute and a weak acid can be concentrated
*the only strong acids dealt in the course: HCl (monoprotic) + H2O ‐‐> H3O+ + Cl‐ HNO3 (monoprotic) + H2O ‐‐> H3O+ + NO3
‐ H2SO4 (diprotic) + H2O ‐‐> 2H3O+ + SO4
2‐
3.5 identify pH as ‐log10 [H+] and explain that a change in pH of 1 means a ten‐fold change in [H+]
pH = ‐log10[H3O+]
Square brackets around a species means concentration of
Nitric acid HNO3 Monoprotic
Number of decimal place in the value for pH should equal the number of significant figures
in the value for [H3O+]
[H3O+] = 10‐pH
E.g. [H3O+] = 10‐7 mol/L – pH = 7.0
When the pH changes by 1, there has been a change of 10x the concentration of hydrogen
ions
3.6 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
3.7 Describe the difference between a strong and weak acid in terms of an equilibrium between the
intact molecule and its ions
Strong acid – e.g. HCl, H2SO4, HNO3
HCl is a strong acid ‐ almost all HCl molecules react with water
o HCl (g) + H2O (l) H3O+ (aq) + Cl‐ (aq)
In terms of its equilibrium, this means that for strong acids and bases the equilibrium lies
almost entirely to the right (products)
There are no (or few) neutral acid molecules present in the solution
The [H3O+] is the same as the molarity of the acid
0.1 mol/L acid has [H3O+] = 0.1 mol/L (10‐1 mol/L)
The ionisation reaction goes to completion
Weak acid – e.g. HF, CH3COOH, citric acid, H2CO3
Citric acid and acetic acid are both weak acids – reaction does not go to completion (only
some of the molecules react with water)
o E.g. CH3COOH (aq) + H2O (l) ⇌ H3O+ (aq) + CH3COO‐ (aq)
In terms of their equilibrium, this means that for weak acids and bases the equilibrium lies
almost entirely to the left (reactants)
Only SOME of the molecules of the acid ionise
There are many neutral acid molecules present in the solution
The [H3O+] is much less than the molarity of the acid
The ionisation reaction is an equilibrium ⇌
Degree of ionisation – fraction of the molecules that has ionised
Degree of weakness of acid is determined by degree of ionisation – however, there is no degree for
strong acids (an acid is strong if degree of ionisation is 100%)
*Calculating degree of ionisation as a %
Degree of ionisation = x 100
[H+] will be given to you or you can calculate it from [H+] = 10‐pH
3.1 Solve problems and perform a first‐hand investigation to use pH meters/probes and indicators to
distinguish between acidic, basic and neutral chemicals
3.2 Plan and perform a first‐hand investigation to measure the pH of identical concentrations of
strong and weak acids
3.3 Gather and process information from secondary sources to write ionic equations to represent the
ionisation of acids
Ionisation Reactions of Common Acids Monoprotic acids ‐ each molecule can donate one proton
o Hydrochloric acid HCl (g) + H2O (l) ‐> H+ + H2O + Cl‐ ‐> H3O+ (aq) + Cl‐ (aq)
o Hydrobromic acids HBr + H2O (l) ‐> H+ + H2O + Br‐ ‐> H3O+ (aq) + Br‐ (aq)
o Nitric acids HNO3 (l) + H2O (l) ‐> H+ + H2O + NO3‐ ‐> H3O+ (aq) + NO3‐ (aq)
o Acetic acid is a bit different ‐ not all the H atoms in the molecule are donated as H+; only the one attached to the ‐O‐
CH3COOH (l) + H2O (l) ⇌ H3O+ (aq) + CH3COO‐ (aq)
Diprotic and triprotic acids ‐ each molecule can donate 2 and 3 protons respectively o Sulfuric and phosphoric acids can donate more than one proton ‐ predict their ionisation
reactions H2SO4 + 2H2O ‐> 2H+ + 2H2O + + SO4
2‐ ‐> 2H3O+ + SO42‐
H3PO4 + 3H2O ‐> 3H3O+ + PO43‐
3.4 Use available evidence to model the molecular nature of acids and simulate the ionisation of
strong and weak acids
3.5 Gather and process information from secondary sources to explain the use of acids as food
additives
Acids are frequently added to food in order to
o Improve the taste (e.g. in drinks and sweets)
o Preserve the food by increasing its acidity (e.g. canned fruits, vegetables in pickles),
because bacteria cannot survive in acid conditions
Common acids added are acetic acid (vinegar), citric acid, and occasionally phosphoric acid
Propanoic acid is often used as a preservative in bread
Ascorbic acid (vitamin C) is often added to foods because it is a vitamin
Potassium hydrogen tartrate is often added to food along with sodium hydrogen carbonate
as baking powder to make cakes rise as they cook
3.6 Identify data, gather and process information from secondary sources to identify examples of
naturally occurring acids and bases and their chemical composition
Naturally occurring acids
Hydrochloric acid (HCl)
o produced in our stomachs to break down food molecules into easily transportable
small molecules absorbed into the blood stream
o also made commercially – uses such as cleaning metals, brickwork, neutralising
bases in manufacturing processes and adjusting acidity in swimming pools
acetic acid (CH3COOH)/ethanoic acid
o carboxylic acid functional group
o present in vinegar which is commonly made from wine by oxidation of ethanol
o produced industrially to make a wide range of organic chemicals
Citric acid
o Occurs in citrus fruit
o Made industrially and used as a food additive
Ascorbic acid (C6H8O6)/Vitamin C
o Occurs widely in fruit and vegetables
o Essential to our health and wellbeing
Synthetic acids
Sulphuric acid (H2SO4)
o Used to make fertilisers, synthetic fibres, industrial ethanol, detergents and car
batteries
Nitric acid (HNO3)
o Used to manufacture fertilisers and explosives
Naturally occurring bases
Ammonia (NH3)
o Present in the stale urine of humans and animals
o Formed by anaerobic decay (decay without oxygen) of organic matter
o Manufactured for the production of fertilisers and nitric acid
o Used in the home as a cleaning agent
Amines
o Compounds with some other group such as an alkyl replacing one or more of the H
atoms of ammonia
o Formed during anaerobic decomposition of organic matter
Metallic oxides
o E.g. iron (III) oxide, copper oxide, titanium oxide
o Extracted from insoluble bases
Carbonates
o E.g. sodium carbonate, calcium carbonate, magnesium carbonate
o Synthesised industrially for use in making glass, paper, soaps, detergents, treating
water, and a cheap base for neutralising acids
Synthetic bases
sodium hydroxide (NaOH)
o made from common salt
o used for making soap and detergents, rayon and other synthetic fibres
o used in oven cleaners
Calcium oxide (CaO) and hydroxide Ca(OH)2
o Calcium oxide is made by heating limestone and then hydroxide is made by reacting
the oxide with water
o CaO used to make cement while Ca(OH)2 is used in mortars and plasters
3.7 process information from secondary sources to calculate pH of strong acids given appropriate
hydrogen ion concentrations
Section 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.
4.1 Outline the historical development of ideas about acids including those of:
‐ Lavoisier
‐ Davy
‐ Arrhenius
Lavoisier (1770) – after experimenting on oxides of non‐metals, Lavoisier came to the conclusion
that acids were substances that contained oxygen atoms (e.g. CO2, SO2)
Davy (1810) – Davy noticed that HCl did not contain oxygen, however still acted as an acid. Davy
proposed that acids contained hydrogen atoms
Explains release of hydrogen when metals are reacted with acids
However, did not account for many compounds that contained hydrogen, but is not acidic
Arhenius (1884)
Acids ionize in water to produce hydrogen ions while bases produce hydroxide ions
Some acids/bases were strong or weak depending on the ionisation of hydrogen/hydroxide
ions
Acids react with bases to produce a salt and water (neutralization) – hydrogen ions and
hydroxide ions react to produce water
Problem is – only applies to aqueous solution (when water is the solvent) and did not explain
why some salts are acidic, basic, and neutral
Theory is still taught today – most situations are aqueous, and his definition recognises the
importance of the concentration of hydrogen/hydroxide ions in water solution ‐> led to the
development of the pH scale
4.2 Outline the Brönsted‐Lowry theory of acids and bases
Acids are PROTON DONORS o HA + H2O ‐‐> H3O+ + A‐ (<‐‐‐ H2O is a bronsted‐lowry base in this case) o B‐L Acid ‐ proton donor
bases are PROTON DONATORS o B + H2O ‐‐> HB+ + OH‐ (<‐‐‐ H2O is a B‐L acid in this case) o B‐L Base ‐ proton acceptor
Removed the need for ionization – could explain those acids that were acidic without introduction to water
An acid can't act as an acid unless another substance acts as a base ‐ there is an exchange of protons
Eq 1: HCl(g) + H2O(l) ‐‐> H3O+(aq) + Cl‐(aq) Eq 2: NH3(g) + H2O(l) ⇌NH4+(aq) + OH‐(aq)
The behaviour of the solvent is important Water can act as a B‐L base and a B‐L acid, depending on what it's with Therefore, acidity and basicity is related to the properties of the solute in relation to the
solvent
o If a substance A has a greater tendency to give up protons than the solvent, A will be an acid in that solvent
o If a substance B has a greater tendency to accept protons than the solvent, B will be a base in that solvent
Eq 1: HCl(g) + H2O(l) ‐‐> H3O+(aq) + Cl‐(aq) Eq 2: Nh3(g) + H2O(l) ⇌NH4+(aq) + OH‐(aq) Self‐ionization of water:
4.3 Describe the relationship between an acid and its conjugate base and a base and its conjugate
acid
When a B‐L acid has donated a proton (H+) to a base, it forms an anion called its conjugate
base
o The acid and anion is called a conjugate acid‐base pair
o The conjugate base is proton deficient so tends to accept a proton in solution – and
act as a base ‐> it will reform the original acid
When a B‐L base has accepted a proton from the acid, it forms a cation called its conjugate
acid
o The conjugate acid tends to donate a proton in solution – and act as an acid ‐> it will
reform the original base
Weak acids/bases have strong conjugate that react significantly with water in the solutions,
reforming the original weak acid/base
Strong acids/bases have very weak conjugates that barely react with water – the reverse
reaction is so insignificant it is considered as
4.4 Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic,
neutral or basic nature
A salt is an ionic chemical formed when an acid and base reacts
Although it is a neutralisation reaction, not all salts are neutral
o If there is a change in pH on dissolving, it is due to the reaction between the salt and
the water – hydrolysis
Acidic salts
Formed from the reaction between a strong acid and a weak base
E.g. HNO3 + NH3 NH4+NO3
‐
*no water with ammonia
The pH of an aqueous solution is <7 because the conjugate acid reacts to an extent with
water – it is a weak B‐L acid
The ammonium is a weak B‐L acid and donates a proton in solution
Due to the presence of the hydronium ion, ammonium salts are acidic in water
Other acidic salts: potassium hydrogen oxalate, sodium hydrogen sulfate
Basic salts
Formed from the reaction between a strong base and weak acid
E.g. H2CO3 + 2NaOH 2Na+CO32‐ + 2H2O
The pH is >7 because the conjugate base reacts to an extent with water – it is a weak B‐L
base
The carbonate ion is a weak B‐L base and accepts a proton in solution:
Due to the presence of the hydroxide ion, carbonate salts are basic in water
Other basic salts: potassium nitrate, sodium acetate, potassium cyanide
Neutral salts
Formed from the reaction between a strong acid and strong base
Their pH is (close to) 7 because their conjugates do not react significantly with water:
E.g. HCl + NaOH NaCl + H2O (the Na+ is a spectator)
o The conjugate base of a strong acid is weak and unlikely to accept a proton
o The conjugate acid of a strong base is weak and unlikely to donate a proton
Strong acids – hydrochloric, nitric, sulphuric
Strong bases – sodium hydroxide, potassium hydroxide
Neutral salts – sodium chloride, nitrate, sulfate; potassium chloride, nitrate, sulfate
Neutral salts may also be formed between a weak acid and weak base
Their pH is close to 7 because their conjugates cancel each other out
o E.g. NH3 + CH3COOH ⇌ NH4+ (aq) + CH3COO‐ (aq)
o NH4+ + H2O ⇌ NH3 + H3O+
o CH3COO‐ + H2O ⇌ CH3COOH + OH‐
o H3O+ + OH‐ H2O
o The conjugate base of a weak acid is weak and somewhat likely to accept a proton
o The conjugate acid of a weak base is weak and somewhat likely to donate a proton
4.5 Identify conjugate acid/base pairs
By adding or removing a hydrogen ion, the substances’ conjugate acid/base is produced.
E.g. HF/F‐, HCl/Cl‐, NH3/NH4+
4.6 Identify amphiprotic substances and construct equations to describe their behaviour in acidic and
basic solutions
Amphiprotic substances can act as either a B‐L acid (proton donor) or B‐L base (proton acceptor)
depending on the nature of the other chemical substances present.
This is why bicarbonate makes a good “go to” for neutralising both acid and base spills.
Example 2: hydrogen phosphate (HPO42‐)
Example 3: water
4.7 Identify neutralisation as a proton transfer reaction which is exothermic
Acid + base salt + water
Neutralisation reacts are PROTON TRANSFER reactions – acids are proton donors and bases
are proton acceptors
Normally from the H3O+ ion to the OH‐ ion
o H+ + OH‐ H2O
o H3O+ + OH‐ 2H2O
Neutralisation is an EXOTHERMIC REACTION
Weak acid/base give less energy/heat
E.g.
4.8 Describe the correct technique for conducting titrations and preparation of standard solutions
Titrations are a specific type of volumetric analysis – they are quantitative
The unknown concentration of a solution can be determined by titrating it against a solution
of known concentration – this is called the standard solution
Standard solutions – these are the “known” – we use the volume and concentration
Primary standards: These may not be used in the final analysis of the unknown, but are an essential part of the process. To make a standard solution, a known mass of a suitable chemical is added to a known volume of distilled water. The chemical used must satisfy these conditions: Highly pure Accurately known composition po reduce % error in weighing (Mr ‐ molecular mass) Reacts instantly and completely with the "unknown"
Unsuitable substance: Conc HCl ‐ fumes and loses HCl gas Conc H2SO4 ‐ absorbs water from atmosphere
NaOH ‐ absorbs moisture from air during weighing; reacts with CO2 to form NaHCO3 Hydrated Na2CO3 ‐ loses water to the atmosphere during weighing; composition cannot be
determined accurately These substances may be suitable as secondary standards… Suitable: Acid standards ‐ potassium hydrogen phthalate, benzoic acid, oxalic acid Base standards ‐ sodium carbonate (Na2CO3), sodium hydrogen carbonate, borax These may be used to determine the absolute concentration of a solution that can be used as
a secondary standard Titrate a primary standard against a secondary standard
Preparing a Standard Solution Weigh out some chemical and add it to a known volume of water Since concentration is measured in mol/L, we calculate c=n/V where n=m/M
To prepare 250 mL of 0.0500 mol/L sodium carbonate (aq) What mass of anhydrous sodium carbonate will you need to weigh?
C(Na2CO3) = 0.0500 mol/L n(Na2CO3) = 0.0500 x 0.250
= 0.0125 moles m(Na2CO3) = 0.0125 x ((2 x 22.99) + 12.01 +48)
= 1.32 g
1. Prepare and weigh solid ‐ 3dp 2. Dissolve solid in small beaker ‐ use distilled water and do not remove stirring rod 3. Transfer to clean 250ml volumetric flask ‐ pre‐cleaned and rinsed with distilled water 4. Use small glass funnel 5. Use rod to help pour 6. Rinse beaker and rod and put all washings in flask 7. Add water to the mark on the flask ‐ drop by drop for last 1cm.
*eye level to the bottom of meniscus 8. Add stopper, invert flask 4‐5 times and label
Titration Basics Acid‐base titrations ‐ i.e. neutralisation reactions
o Acid + base ‐‐> salt + water In a titration, we react exactly the number of moles of each reactant as determined by the
equation ‐ the stoichiometric ratio The point at which the reaction is complete is called the equivalence point ‐ i.e. the exact
stoichiometric ratio of n(acid):n(base) has been reached Experimentally, it is impossible to tell when this occurs without an indicator (or pH probe) ‐
this changes colour as the pH changes at neutralisation So we need
o A measured volume of the known concentration substance ‐ e.g. Na2CO3 o An indicator to show when neutralisation occurs o To know how much of the unknown we added ‐ e.g. HCl o To get reliable data o Accuracy and precision
Glassware Volumetric flask
Pipette Burette Conical flask
Purpose Make a standard solution
Extract known volumes of standard ‐ aliquot
Hold and dispense the titrant
To hold aliquot + carry out neutralisation reaction
Cleaning Washed in the same way as the burette. Volume of solution is called and aliquot.
Rinsed first with distilled water then with the solution (titrant) to be used. Washings are discarded before glassware is filled accurately.
Rinsed with distilled water and can be left wet**. Aliquot transferred into here.
Accuracy Electric balance = +/‐ 0.01 g Burette = +/‐ 0.05 mL 25 mL bulb pipette = +/‐ 0.03 mL
Performing a Titration 1. Fast/rough titre first
Prepare the burette and fill it with the "unknown" Prepare the bulb pipette and extract the standard solution (the aliquot) e.g. 25mL 0.05M
Na2CO3 Transfer the aliquot to a washed conical flask Add three drops of indicator to the flask and place on white tile. Repeat, but keep the
second as a reference With one hand open tap to dispense liquid slowly from burette, and swirl the flask with
the other hand (rinse any splashes from sides with distilled water) Stop at colour change and record the volume of titrant dispensed ‐ the titre
2. Accurate titre (based on rough one) Repeat the whole titration process, but slow down at least 1mL away from the end‐point Transfer the titrant drop by drop and swirl thoroughly until you reach the end‐point ‐
note: drops or part of drops can be washed into the flask using a wash bottle Repeat at least 3 times for close, reliable titres.
*the rough value is not used in the calculations because it is not accurate Titration calculations
1. Equation ‐ balanced 2. n(Na2CO3) = cV 3. n(HCl) from equation 4. C(HCl) = n(HCl)/V(HCl) <‐ average titre
4.9 Qualitatively describe the effect of buffers with reference to a specific example in a natural
system
A buffer solution resists changes to PH when small quantities of acid or base are added
Important in maintain pH in a variety of areas: e.g. biological systems, lakes, labs, cosmetics
and toiletries
A buffer solution has to contain species which will remove any H+ or OH‐ that you might add
to it – otherwise pH will change
Acidic and alkaline buffer solutions achieve this in different ways
Buffer solutions often contain weak acids and their conjugate bases (e.g. acetic acid and
sodium acetate)
Contains the SAME molar concentration of a weak acid and its conjugate base
Sodium salts of weak acids are good conjugate bases as they are all soluble (and Na is usually
a spectator ion) – e.g. NaHCO3, CH3COONa (sodium acetate)
1. Acid reaction: HA + H2O ⇌ H3O+ + A‐ Ionisation reaction e.g. acetic acid + acetate ‐ CH3COOH + H2O ⇌ CH3COO‐ + H3O+ carbonic acid + sodium hydrogen carbonate ‐ H2CO3 + H2O ⇌ HCO3
‐ + H3O+
2. Conjugate base reaction: H2O + A‐ ⇌ OH‐ + HA
Hydrolysis reaction e.g. acetic acid + acetate ‐ CH3COO‐ + H2O ⇌ CH3COOH + OH‐ carbonic acid + sodium hydrogen carbonate ‐ HCO3
‐ + H2O ⇌ H2CO3 + OH‐
By adding H+ (i.e. by adding an acid such as HCl) There is an increase in [H3O+] as the acid ionises in the water In response, the A‐ accepts H+ to form HA acid molecules (reaction 1) The position of the equilibrium 1 is shifted to the left No pH change as no change in [H3O+]
By adding OH‐ (e.g. by adding an alkali) There is an increase in [OH‐] as the base ionises in the water In response, the HA donates H+ to the OH‐ forming H2O and A‐ The position of the equilibrium 2 is shifted to the left No pH change as no change in [OH‐]
Natural Buffer systems
Biological reaction involve proteins called enzymes which are catalysts
They act best at their optimum pH and reasonably well within a narrow pH range
They are sensitive to changes in pH and are damaged when the pH is outside the narrow
range – and are rendered useless = death of the organism
4.1 Gather and process information from secondary sources to trace developments in understanding
and describing acid/base reactions
4.2 Choose equipment and perform a first‐hand investigation to identify the pH of a range of salt
solutions
4.3 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
4.4 Perform a first‐hand investigation to determine the concentration of a domestic acidic substance
using computer‐based technologies
4.5 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
Laboratory spills:
Never use concentrated acids/bases to neutralise spills due to the exothermic reaction that
can cause boiling and vapourise fumes
o Also, if added in excess, you have the same problem in reverse
Evacuate if fumes are present – e.g. concentrated HCl/NH3
Remove clothes and shower a person affected with lots of water – diluting reduces harmful
effects and washes the acid/base away
Isolate spills using sand and this can be neutralised later
Neutralise ACID with sodium carbonate:
e.g. Na2CO3 (s) + 2HCl (aq) 2NaCl (aq) + H2O (l) + CO2 (g)
Sodium hydrogen carbonate can neutralise both acid and base spills because it is
amphiprotic
e.g. NaHCO3 (s) + HCl (aq) NaCl (aq) + H2O (l) + CO2 (g) (carbonic acid decomposes: H2CO3
⇌ H2O + CO2)
o Very cheap and non‐toxic
o Very stable – can be stored long term
o The products are safe and unreactive
Large‐scale spills – isolate using sand, neutralise and dilute with loads of water
Important to dilute the products that enter the waste‐water system so as not to discharge
high concentrations
Sodium hydrogen phosphate is amphiprotic and can be used to neutralise
Multi‐step Ionisation Process
For acids that are diprotic and triprotic, the protons are transferred sequentially.
This is why you can get hydrogen carbonate and carbonate salts from carbonic acid, and hydrogen
sulfate and sulfate salts from sulfuric acid.
Ionisataion of Sulfuric acid
H2SO4 (l) + H2O (l) HSO4‐ (aq) + H3O+ (aq) Highly exothermic
HSO4‐ (aq) + H2O (l) ⇌ SO4‐ (aq) +H3O+ (aq) Not so exothermic and equilibrium
Section 5 – Esterification is a naturally occurring process which can
be performed in the laboratory
5.1 Describe the differences between the alkanol and alkanoic acid functional groups in carbon
compounds
It is the functional group on an organic compound that leads to its physical and chemical properties
Alkanols and alkanoic acids are involved in making esters Alkanols Contain the hydroxy functional group: ‐OH This leads to high melting points and boiling points due to the strength of intermolecular
forces General formula: CnH2n+1OH e.g. methanol: R=CH3; ethanol: R=C2H5
Alkanoic Acids Contain the carboxylic acid functional group, ‐COOH
They can lose a hydrogen ion from this group and behave as a weak acid o NOTE: no other H are released
General formula: Cn‐1H2n‐1COOH e.g. methanoic acid: R=H; ethanoic acid: R=CH3
5.2 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
Alkanol gives the 1st name, alkanoic acid gives the 2nd Count the no. of carbons in the alkyl group from the parent alkanol (the chain attached to the
SINGLE O) o Delete 'anol' and replace with 'y'
Count the no. of carbons in the alkanoate from the parent alkanoic acid (the chain attached to the double bonded O) o Delete 'anoic acid' and replace with 'oate'
Examples Methanol + hexanoic acid ‐‐> methyl hexanoate Butan‐2‐ol + propanoic acid ‐‐> 2‐butyl propanoate Propan‐1‐ol + formic acid (methanoic acid) ‐‐> 1‐propyl formate
5.3 Explain the difference in melting point and boiling point caused by straight‐chained alkanoic acid
and straight‐chained primary alkanol structures
C‐O and H‐O bonds are polar so alkanols are polar ‐ leads to dipole‐dipole forces C=O bonds are also polar so alkanoic acids are even more polar, and therefore have a higher
mp/bp O‐H bonds can also form H‐bonds Result: both alkanols and alkanoic acids have strong intermolecular forces and so much higher
mp/bp than alkanes of similar molecular weight, which only exhibit dispersion forces Look at the table of bps on page 173 Comparisons are made on molecular weight because dispersion forces will be similar Strength of dispersion forces increases as numbers of electrons increases What generalisations can we make from the table?
o Increasing bp as molecular weight increases o Alkanoic acids have higher bps
More explanation? Alkanol molecules can make hydrogen bonds between the H and the O of neighbouring
molecules
Each alkanoic acid molecule can make hydrogen bonds and dipole‐dipole bonds This leads to an even higher mp and bp, due to increased energy being required to overcome
the bonds Acetic acid molecules can form dimers as the diagram shows
Short chain alkanols and acids are soluble in water due to H bonding
5.4 Identify esterification as the reaction between an acid and an alkanol and describe, using
equations, examples of esterification
Formed when an alkanoic acid reacts with an alkanol in a condensation reaction o (although alkanols have an OH, they are not basic, so it's not an acid‐base reaction)
Alkanol + alkanoic acid ⇌ alkyl alkanoates (ester) + H2O (l) o e.g. ethanol + ethanoic acid ⇌ ethyl ethanoate + H2O (l) o Ethanol + propanoic acid ⇌ ethyl propanoate + H2O (l) o Propan‐1‐ol + ethanoic acid ⇌ propyl ethanoate + H2O (l)
OH comes from the alkanoic acid* H comes from the alkanol All esters have the ester functional group ‐ COO ‐
*OH from the longer chain – if it has more it has less to lose Often, esters are known by their common names (e.g. acetate, etc)
5.5 Describe the purpose of using acid in esterification for catalysis
Esterification is a slow reaction that does not go to completion ‐ it is an equilibrium reaction 1‐3 mL of concentrated H2SO4 used as a catalyst ‐ this gets the system to equilibrium faster NOTE: the acid catalyses both reactions, it allows the establishment of equilibrium from just
reactants to occur faster BUT: it also acts as a dehydrating agent ‐ absorbs the water product and forces equilibrium to
the right (not physically closed but the system is a closed system ‐ cold water cools the gases which
doesn’t allow them to escape) Extracting the ester Since the reaction is an equilibrium, our product is mixed with our reactants Some esters can be extracted by distillation (and other processes) For some esters, the BP of the alkanol and ester are too close. We then rely on the insolubility
of the ester in water and use a separating funnel (and other processes)
5.6 Explain the need for refluxing during esterification
Some alkanoic acid, alkanols and esters are volatile Reflux ‐ when gas rises but comes back down because it is condensed Low temperatures will cause the reaction to proceed too slowly…so we need to heat up the
reaction The higher temperatures allow volatile reactants and products to vapourise What will happen in an open vessel? ‐ the volatile substances will evaporate without
condensing them What effect would this have on the reaction? ‐‐> lose products (lose reactants, system will
replace and drive to LHS) Solution: reaction flask fitted with a condenser Vapours which rise from reaction mixture are condensed and run back into flask It is a closed system, without actually being physically closed, as the reactants and products
cannot escape This allows the equilibrium to be established and allows the reaction to occur at higher
temperatures than otherwise possible Why not put a lid on the reaction?
o This reduces the risk of explosion due to pressure build up (Esterification can occur for hours or days)
Yields Often use an excess of alkanoic acid to shift the equilibrium to the right and create more ester % yields are very variable ‐ the ester is lost mostly in the isolation and purification process
5.7 Outline some examples of the occurrence, production and uses of esters
5.2 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 and occur widely in nature as perfumes and flavouring
agents
E.g. methyl butanoate – apple flavour
Butyl ethanoate – raspberry
Pentyl ethanoate – pear
Benzyl ethanoate – jasmine
It is the odour of esters that contributes to flavours
There is a substantial industry for developing and manufacturing synthetic flavours and
perfumes
First step is to identify the constituents of the natural flavour and then to synthesise similar
esters which reproduce this flavour
Artificial flavours often cheaper than natural extracts
Ethyl acetate (ethanoate) is widely used as a solvent in industry, also in nail polish remover
o Manufactured from acetic acid and ethanol
High molecular weight (i.e. non‐volatile) esters such as dialkyl phthalates are used as
plasticisers in some plastics such as PVC – make the materials soft and pliable
A functional group is a centre of reactivity in a carbon compound
5.1 Identify data, plan, select equipment and perform a firsthand investigation to prepare an ester
using reflux