Essential idea: Many reactions involve the transfer of a proton from an acid to a base.
8.1 Theories of acids and bases
UNDERSTANDINGS:
U8.1.1 A Brønsted–Lowry acid is a proton/H+ donor and a Brønsted–Lowry base is a proton/H+ acceptor.
U8.1.3 A pair of species differing by a single proton is called a conjugate acid-base pair.
Arrhenius and Bronsted-Lowry theories:
Arrhenius:
- Acid = substance that produces H+ ions in a solution
- Base = substance that produces OH= ions in a solution
The combination of an acid and a base is well known as a neutralisation reaction involving the combination of the hydrogen ion and the hydroxide ion.
H+ (aq) + OH- (aq) --> H2O (l)
An example of this type of neutralisation is the reaction of hydrochloric acid in the stomach with aluminium hydroxide contained in an antacid tablet:
3HCl (aq) + Al(OH)3 (s) --> AlCl3 (aq) + 2H2O (l)
Arrhenius' theory had its limitations. The reaction between the weak base ammonia and hydrogen chloride gas could not be explained, as ammonia does not contain hydroxide ions:
NH3 (g) + HCl (g) --> NH4Cl (s)
Bronsted-Lowry:
- Acid = substance that donates a proton (remember as B.A.D - Bronsted Acid Donates)
- Base = substance that accepts a proton
Common acids are referred to as being monoprotic e.g. hydrochloric acid, diprotic e.g. sulphuric acid or triprotic e.g. phosphoric acid. Hydrochloric and sulphuric acid are strong acids whereas phosphoric acid is weak.
For a compound to act as a BL acid, it must have a hydrogen atom in it that it is capable of losing (as a proton). A BL base must be capable of accepting a hydrogen ion. Some compounds such as water may act as both.
UNDERSTANDINGS:
U8.1.1 A Brønsted–Lowry acid is a proton/H+ donor and a Brønsted–Lowry base is a proton/H+ acceptor.
U8.1.3 A pair of species differing by a single proton is called a conjugate acid-base pair.
Arrhenius and Bronsted-Lowry theories:
Arrhenius:
- Acid = substance that produces H+ ions in a solution
- Base = substance that produces OH= ions in a solution
The combination of an acid and a base is well known as a neutralisation reaction involving the combination of the hydrogen ion and the hydroxide ion.
H+ (aq) + OH- (aq) --> H2O (l)
An example of this type of neutralisation is the reaction of hydrochloric acid in the stomach with aluminium hydroxide contained in an antacid tablet:
3HCl (aq) + Al(OH)3 (s) --> AlCl3 (aq) + 2H2O (l)
Arrhenius' theory had its limitations. The reaction between the weak base ammonia and hydrogen chloride gas could not be explained, as ammonia does not contain hydroxide ions:
NH3 (g) + HCl (g) --> NH4Cl (s)
Bronsted-Lowry:
- Acid = substance that donates a proton (remember as B.A.D - Bronsted Acid Donates)
- Base = substance that accepts a proton
Common acids are referred to as being monoprotic e.g. hydrochloric acid, diprotic e.g. sulphuric acid or triprotic e.g. phosphoric acid. Hydrochloric and sulphuric acid are strong acids whereas phosphoric acid is weak.
For a compound to act as a BL acid, it must have a hydrogen atom in it that it is capable of losing (as a proton). A BL base must be capable of accepting a hydrogen ion. Some compounds such as water may act as both.
BL acid-base reactions always involve an acid-base conjugate pair, one is the acid, the other is its conjugate base.
e.g. HCl/Cl-, CH3COOH/CH3COO-, NH4+/NH3
The conjugate base will always have one less H atom than the acid (and a more negative charge). In compounds where there are many hydrogen atoms, the one that is held the weakest is generally the one which is lost.
When a BL acid donates a proton, it forms a conjugate base. When a BL base accepts a proton, it forms a conjugate acid.
NOTE: a strong BL acid gives a weak conjugate base.
U8.1.2 Amphiprotic species can act as both Brønsted–Lowry acids and bases.
Some substances have the ability to act as either a BL acid or a BL base depending on the reaction in which they are taking part. These species are said to be amphiprotic e.g. the water molecule can donate a proton in a reaction (acting as a BL acid) or accept a proton (acting as a BL base). Amino acids also act as amphiprotic species.
Polyprotic species are frequently involved in reactions in which they behave amphiprotically for example:
e.g. HCl/Cl-, CH3COOH/CH3COO-, NH4+/NH3
The conjugate base will always have one less H atom than the acid (and a more negative charge). In compounds where there are many hydrogen atoms, the one that is held the weakest is generally the one which is lost.
When a BL acid donates a proton, it forms a conjugate base. When a BL base accepts a proton, it forms a conjugate acid.
NOTE: a strong BL acid gives a weak conjugate base.
U8.1.2 Amphiprotic species can act as both Brønsted–Lowry acids and bases.
Some substances have the ability to act as either a BL acid or a BL base depending on the reaction in which they are taking part. These species are said to be amphiprotic e.g. the water molecule can donate a proton in a reaction (acting as a BL acid) or accept a proton (acting as a BL base). Amino acids also act as amphiprotic species.
Polyprotic species are frequently involved in reactions in which they behave amphiprotically for example:
NOTE: To determine whether a substance is an acid or a base, count the hydrogens on each substance before and after the reaction. If the number of hydrogens has decreased that substance is the acid (donates hydrogen ions). If the number of hydrogens has increased that substance is the base (accepts hydrogen ions).
APPLICATION AND SKILLS:
AS8.1.1 Deduction of the Brønsted–Lowry acid and base in a chemical reaction.
AS8.1.2 Deduction of the conjugate acid or conjugate base in a chemical reaction.
See above points.
NOTE: Amphiprotic means the substance can both donate and accept a proton (H+), while amphoteric is a more general term meaning the substance can act as both an acid and a base. In a sense, water (H2O) can be considered both amphoteric and amphiprotic
H3O+ <------ H2O -----> OH-
Here we see that water can potentially do both, gain a proton while acting as a base, and lose a proton while acting as an acid.
So, in a manner, all amphiprotic substances are amphoteric - since when they donate a proton they are acting as an acid, and when they accept a proton they are acting as base.
On the other hand, not all amphoteric substances are amphiprotic, because only in the Bronstead and Lowry sense do acids and bases only accept and donate protons.
NOTE: a proton in aqueous solution can be represented as both H+ (aq) and H3O+ (aq).
AS8.1.1 Deduction of the Brønsted–Lowry acid and base in a chemical reaction.
AS8.1.2 Deduction of the conjugate acid or conjugate base in a chemical reaction.
See above points.
NOTE: Amphiprotic means the substance can both donate and accept a proton (H+), while amphoteric is a more general term meaning the substance can act as both an acid and a base. In a sense, water (H2O) can be considered both amphoteric and amphiprotic
H3O+ <------ H2O -----> OH-
Here we see that water can potentially do both, gain a proton while acting as a base, and lose a proton while acting as an acid.
So, in a manner, all amphiprotic substances are amphoteric - since when they donate a proton they are acting as an acid, and when they accept a proton they are acting as base.
On the other hand, not all amphoteric substances are amphiprotic, because only in the Bronstead and Lowry sense do acids and bases only accept and donate protons.
NOTE: a proton in aqueous solution can be represented as both H+ (aq) and H3O+ (aq).
Essential idea: The characterization of an acid depends on empirical evidence such as the production of gases in reactions with metals, the colour changes of indicators or the release of heat in reactions with metal oxides and hydroxides.
8.2 Properties of acids and bases
UNDERSTANDINGS:
U8.2.1 Most acids have observable characteristic chemical reactions with reactive metals, metal oxides, metal hydroxides, hydrogen carbonates and carbonates.
NOTE: acid + metal hydrocarbonate --> salt + water + carbon dioxide
Salts are composed of an anion and a cation e.g. sodium chloride.
Metals that are found above hydrogen in the activity series react with acids to form a salt and hydrogen gas. These reactions give off hydrogen at different rates according to the reactivity of the metal and the strength and concentration of acid.
UNDERSTANDINGS:
U8.2.1 Most acids have observable characteristic chemical reactions with reactive metals, metal oxides, metal hydroxides, hydrogen carbonates and carbonates.
NOTE: acid + metal hydrocarbonate --> salt + water + carbon dioxide
Salts are composed of an anion and a cation e.g. sodium chloride.
Metals that are found above hydrogen in the activity series react with acids to form a salt and hydrogen gas. These reactions give off hydrogen at different rates according to the reactivity of the metal and the strength and concentration of acid.
U8.2.2 Salt and water are produced in exothermic neutralization reactions.
Acid + base --> salt + water = exothermic neutralisation reaction
e.g. NaOH + HCl --> NaCl + H2O
Enthalpy of neutralisation is the enthalpy change that occurs when acid + base react to form 1 mole of water under standard conditions. This enthalpy change has a negative value.
Uses of neutralisation:
- heart burn (acidic)
- bee string (acidic)
- wasp sting (alkaline)
- soil acidity (acidic)
NOTE: alkali = forms OH- ions
base - accepts H+ ions
So all alkalis are bases, but not all bases are alkalis.
APPLICATION AND SKILLS:
AS8.2.1 Balancing chemical equations for the reaction of acids.
See above points.
AS8.2.2 Identification of the acid and base needed to make different salts.
NOTE: examples of bases include:
- metal oxides/hydroxides
- metal carbonates/hydrocarbonates
- ammonia/amines
AS8.2.3 Candidates should have experience of acid-base titrations with different indicators.
NOTE: The colour changes of different indicators are given in the data booklet in section 22.
Acid + base --> salt + water = exothermic neutralisation reaction
e.g. NaOH + HCl --> NaCl + H2O
Enthalpy of neutralisation is the enthalpy change that occurs when acid + base react to form 1 mole of water under standard conditions. This enthalpy change has a negative value.
Uses of neutralisation:
- heart burn (acidic)
- bee string (acidic)
- wasp sting (alkaline)
- soil acidity (acidic)
NOTE: alkali = forms OH- ions
base - accepts H+ ions
So all alkalis are bases, but not all bases are alkalis.
APPLICATION AND SKILLS:
AS8.2.1 Balancing chemical equations for the reaction of acids.
See above points.
AS8.2.2 Identification of the acid and base needed to make different salts.
NOTE: examples of bases include:
- metal oxides/hydroxides
- metal carbonates/hydrocarbonates
- ammonia/amines
AS8.2.3 Candidates should have experience of acid-base titrations with different indicators.
NOTE: The colour changes of different indicators are given in the data booklet in section 22.
Essential idea: The pH scale is an artificial scale used to distinguish between acid, neutral and basic/alkaline solutions.
8.3 The pH scale
UNDERSTANDINGS:
U8.3.2 A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [𝐻+].
U8.3.3 pH values distinguish between acidic, neutral and alkaline solutions.
The pH scale ranges from 0 to 14.
Acidic solutions have a pH < 7 and alkaline solutions have a pH > 7. Pure water has a pH of 7 at 25 degrees Celcius at 1 atm.
e.g. if the H+ ion concentration of a solution is 0.1, the pH is simply -log(base10)0.1 = pH 1.
If an acid is diluted 10 times its pH number is INCREASED by 1.
If an alkaline solution is diluted by 10 times its pH number is DECREASED by 1.
Each change of 1 pH unit represent a 10-fold change in the hydrogen ion concentration H+.
If the pH of a solution changes by 2 pH units, there is a 100-fold change in the hydrogen ion concentration.
If the pH of a solution changes by 3 pH units, there is a 1000-fold change in the hydrogen ion concentration.
If pH = 2 and the solution is changed to pH = 4, then there is a 100-fold decrease in the H+ concentration.
If pH = 12 and the solution is changed to pH = 9, then there is a 1000-fold increase in the H+ concentration OR a 1000-fold decrease in OH- concentration.
THUS, pH is inversely related to H+ concentration.
UNDERSTANDINGS:
U8.3.2 A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [𝐻+].
U8.3.3 pH values distinguish between acidic, neutral and alkaline solutions.
The pH scale ranges from 0 to 14.
Acidic solutions have a pH < 7 and alkaline solutions have a pH > 7. Pure water has a pH of 7 at 25 degrees Celcius at 1 atm.
e.g. if the H+ ion concentration of a solution is 0.1, the pH is simply -log(base10)0.1 = pH 1.
If an acid is diluted 10 times its pH number is INCREASED by 1.
If an alkaline solution is diluted by 10 times its pH number is DECREASED by 1.
Each change of 1 pH unit represent a 10-fold change in the hydrogen ion concentration H+.
If the pH of a solution changes by 2 pH units, there is a 100-fold change in the hydrogen ion concentration.
If the pH of a solution changes by 3 pH units, there is a 1000-fold change in the hydrogen ion concentration.
If pH = 2 and the solution is changed to pH = 4, then there is a 100-fold decrease in the H+ concentration.
If pH = 12 and the solution is changed to pH = 9, then there is a 1000-fold increase in the H+ concentration OR a 1000-fold decrease in OH- concentration.
THUS, pH is inversely related to H+ concentration.
Ionisation of water:
Water is slightly ionised:
H2O <--> H+ (aq) + OH- (aq)
Since [H2O] is constant, Kc = [H+][OH-] ÷ [H2O] - this is known as the ion product constant for water.
<-- APPLICATION AND SKILLS
Essential idea: The pH depends on the concentration of the solution. The strength of acids or bases depends on the extent to which they dissociate in aqueous solution.
8.4 Strong and weak acids and bases
UNDERSTANDINGS:
U8.4.1 Strong and weak acids and bases differ in the extent of ionization.
U8.4.3 A strong acid is a good proton donor and has a weak conjugate base.
U8.4.4 A strong base is a good proton acceptor and has a weak conjugate acid.
The strength of an acid or base depends on the degree to which it ionises or dissociates in water. A strong acid is an effective proton donor that is assumed to completely dissociate in water. Examples include hydrochloric acid, HCl, sulphuric, H2SO4 and nitric acid, HNO3:
These reactions are represented by chemical equations that are assumed to go to completion. The conjugate base of a strong acid is a very weak base. For HCl, the conjugate base is the chloride ion, Cl- which has almost no affinity for a proton in aqueous solutions.
A weak acid dissociates only partially in water; it is a poor proton donor. The dissociation of a weak acid is a reversible reaction that reaches equilibrium. At equilibrium only a small proportion of the acid molecules have dissociated. The conjugate base of a weak acid has a higher affinity for a proton than does the conjugate base of a strong acid.
A strong base also completely dissociates in water. The group 1 metal hydroxides are all soluble in water and are good examples of strong bases, e.g. NaOH and KOH.
A metal hydroxide does not act as a BL base because it does not have the capacity to accept a proton. However, in solution the hydroxide ion acts as a BL base, accepting a base.
Ammonia is an example of a weak base. In the reaction with water, ammonia accepts a proton and effectively undergoes ionisation.
UNDERSTANDINGS:
U8.4.1 Strong and weak acids and bases differ in the extent of ionization.
U8.4.3 A strong acid is a good proton donor and has a weak conjugate base.
U8.4.4 A strong base is a good proton acceptor and has a weak conjugate acid.
The strength of an acid or base depends on the degree to which it ionises or dissociates in water. A strong acid is an effective proton donor that is assumed to completely dissociate in water. Examples include hydrochloric acid, HCl, sulphuric, H2SO4 and nitric acid, HNO3:
These reactions are represented by chemical equations that are assumed to go to completion. The conjugate base of a strong acid is a very weak base. For HCl, the conjugate base is the chloride ion, Cl- which has almost no affinity for a proton in aqueous solutions.
A weak acid dissociates only partially in water; it is a poor proton donor. The dissociation of a weak acid is a reversible reaction that reaches equilibrium. At equilibrium only a small proportion of the acid molecules have dissociated. The conjugate base of a weak acid has a higher affinity for a proton than does the conjugate base of a strong acid.
A strong base also completely dissociates in water. The group 1 metal hydroxides are all soluble in water and are good examples of strong bases, e.g. NaOH and KOH.
A metal hydroxide does not act as a BL base because it does not have the capacity to accept a proton. However, in solution the hydroxide ion acts as a BL base, accepting a base.
Ammonia is an example of a weak base. In the reaction with water, ammonia accepts a proton and effectively undergoes ionisation.
In this reaction water displays its amphoteric nature by acting as a BL acid, donating a proton.
U8.4.2 Strong acids and bases of equal concentrations have higher conductivities than weak acids and bases.
All acids and bases dissociate to a degree in water and create ions. The conductivity of an aqueous solution depends on the concentration of ions present. This can be measured in a simple experiment using a power pack and graphite electrodes connected to an ammeter. The voltage applied must be identical for ach solution so that any difference in current passing through aqueous solutions of different acids or bases reflects the concentration of ions.
Strong acids and bases are strong electrolytes, so they display a higher conductivity than weak acids and bases. For example, a comparison of the conductivity of equimolar solutions of HCl and ethanoic acid would demonstrate that the HCl gives a higher ammeter reading and so has a higher degree of dissociation than ethanoic acid, which is a weak acid.
APPLICATION AND SKILLS:
AS8.4.1 Distinction between strong and weak acids and bases in terms of the rates of their reactions with metals, metal oxides, metal hydroxides, metal hydrogen carbonates and metal carbonates and their electrical conductivities for solutions of equal concentrations.
The reactions of strong and weak acids with metals, metal hydroencarbonates and metal carbonations all produce a gas. The rate of the reaction can be determined through observation followed by analysis and quantitatively by monitoring the rate at which gas is evolved through loss of mass.
All acids and bases dissociate to a degree in water and create ions. The conductivity of an aqueous solution depends on the concentration of ions present. This can be measured in a simple experiment using a power pack and graphite electrodes connected to an ammeter. The voltage applied must be identical for ach solution so that any difference in current passing through aqueous solutions of different acids or bases reflects the concentration of ions.
Strong acids and bases are strong electrolytes, so they display a higher conductivity than weak acids and bases. For example, a comparison of the conductivity of equimolar solutions of HCl and ethanoic acid would demonstrate that the HCl gives a higher ammeter reading and so has a higher degree of dissociation than ethanoic acid, which is a weak acid.
APPLICATION AND SKILLS:
AS8.4.1 Distinction between strong and weak acids and bases in terms of the rates of their reactions with metals, metal oxides, metal hydroxides, metal hydrogen carbonates and metal carbonates and their electrical conductivities for solutions of equal concentrations.
The reactions of strong and weak acids with metals, metal hydroencarbonates and metal carbonations all produce a gas. The rate of the reaction can be determined through observation followed by analysis and quantitatively by monitoring the rate at which gas is evolved through loss of mass.
Essential idea: Increased industrialization has led to greater production of nitrogen and sulfur oxides leading to acid rain, which is damaging our environment. These problems can be reduced through collaboration with national and intergovernmental organizations.
8.5 Acid deposition
All understandings and applications are covered below.
Acid deposition:
Acid deposition is the process by which acid-forming pollutants are deposited on the Earth's surface. Increased industrialisation and economic development in many parts of the world have led to rapidly increasing emissions of the nitrogen and sulphur oxides that cause acid rain, the most prevalent form of acid deposition.
Acid deposition affects the environment in many ways. These include:
- deforestation
- the leaching of minerals from soils leading to elevated acid levels in lakes and rivers
- the uptake of toxic minerals from soil by plants
- reduction in the pH of lake and river systems
- increased uptake of toxic metals by shellfish and other marine life which can affect the fishing industry and ultimately people health
- corrosive effects on marble, limestone and metal buildings, bridges and vehicles
Acid rain:
Pure water has a pH of 7.0. Rainwater is naturally acidic due to the presence of dissolved carbon dioxide which forms weak carbonic acid, H2CO3. A typical pH value of rainwater is 5.6.
All understandings and applications are covered below.
Acid deposition:
Acid deposition is the process by which acid-forming pollutants are deposited on the Earth's surface. Increased industrialisation and economic development in many parts of the world have led to rapidly increasing emissions of the nitrogen and sulphur oxides that cause acid rain, the most prevalent form of acid deposition.
Acid deposition affects the environment in many ways. These include:
- deforestation
- the leaching of minerals from soils leading to elevated acid levels in lakes and rivers
- the uptake of toxic minerals from soil by plants
- reduction in the pH of lake and river systems
- increased uptake of toxic metals by shellfish and other marine life which can affect the fishing industry and ultimately people health
- corrosive effects on marble, limestone and metal buildings, bridges and vehicles
Acid rain:
Pure water has a pH of 7.0. Rainwater is naturally acidic due to the presence of dissolved carbon dioxide which forms weak carbonic acid, H2CO3. A typical pH value of rainwater is 5.6.
Acid rain has a pH less than 5.6. The major pollutants that cause acid deposition are sulphur dioxide SO2 and nitrogen oxides, NO and NO2. These are products of natural occurences such as volcanic eruptions and the decomposition of vegetation, as well as man-made primary pollutants from the combustion of fossil fuels containing high levels of sulphur impurities. Acid rain results principally from the formation of two strong acids, nitric acid and sulphuric acid and can be considered as a major global environmental problem.
For example, at high temperature in the internal combustion engine of a car or jet engine, nitrogen gas reacts with oxygen gas to form the oxide, nitrogen (II) oxide (nitrogen monoxide):
For example, at high temperature in the internal combustion engine of a car or jet engine, nitrogen gas reacts with oxygen gas to form the oxide, nitrogen (II) oxide (nitrogen monoxide):
Pre- and post-combustion technologies:
- pre-combustions methods to reduce sulphur emissions refer to techniques used on fuels before their conbustion
- plysical cleaning or mineral beneficiation involves crushing coal, followed by floatation that reduces the amounts of sulphur and other impurities
- combinations of different pre-combustion methods results in the removal of up to 80-90% of inorganic sulphur
____________________________
- post-combustion methods focus on several complementary technologies to remove sulphur dioxide, nitrogen oxides, heavy metals and dioxins from the combustion gases before they are released into the atmosphere
- e.g. calcium oxide or lime will react with sulphur dioxide and remove it from flue gases:
CaO (s) + SO2 --> CaSO3 (s)
- pre-combustions methods to reduce sulphur emissions refer to techniques used on fuels before their conbustion
- plysical cleaning or mineral beneficiation involves crushing coal, followed by floatation that reduces the amounts of sulphur and other impurities
- combinations of different pre-combustion methods results in the removal of up to 80-90% of inorganic sulphur
____________________________
- post-combustion methods focus on several complementary technologies to remove sulphur dioxide, nitrogen oxides, heavy metals and dioxins from the combustion gases before they are released into the atmosphere
- e.g. calcium oxide or lime will react with sulphur dioxide and remove it from flue gases:
CaO (s) + SO2 --> CaSO3 (s)