UNIT 8
eactivity 3.1—Proton transfer reactions
R
.1.1- 3.1.3: Brønsted–Lowry theory
3
- A Brønsted–Lowry acid is a proton donor
- A Brønsted–Lowry base is a proton acceptor
● Deduce the Brønsted–Lowry acid and base in a reaction
- A proton in aqueous solution can be represented as both H+(aq) and H3O +
(aq)
- The distinction between the terms “base” and “alkali” should be understood.
- An alkali is a soluble base
- All alkalis are bases but not all bases are alkali
- A pair of species differing by a single proton is called a conjugate acid–base pair.
● Deduce the formula of the conjugate acid or base of any Brønsted–Lowry base or acid.
- Some species can act as both Brønsted–Lowry acids and bases.
- These species are called amphiprotic, they can act as both an acid and a base by donating or accepting a proton
- Not to be confused with amphoteric which does not involve proton donation but in which the substance can act as both an
acid and base
● Interpret and formulate equations to show acid–base reactions of these species.
3.1.4: The pH scale
- The pH scale can be used to describe the [H+] of a solution:
- pH = –log10[H+]
- [H+] = 10–pH
● Perform calculations involving the logarithmic relationship between pH and [H+].
T he concentration of hydrogen ions increases by a factor of 100 between pH5 and pH3
-
● Include the estimation of pH using universal indicator, and the precise measurement of pH using a pH meter/probe.
- Red,orange,yellow,green,blue,violet
- Strong acid, weaker acid, weak acid, neutral, weak base, strong base
- 1,3,5,7,11,14
● What is the shape of a sketch graph of pH against [H+]?
- Inversely proportional with each decrease in pH translating to a large increase in hydrogen ion concentration.
-
● When are digital sensors (e.g. pH probes) more suitable than analogue methods (e.g. pH paper/solution)?
- Universal indicator: dip paper in solution and compare color to strip to measure pH. Universal indicator is a mix of dyes.
- Advantages: easier to read, quick, accurate, no calibration, no power
, - Disadvantages: imprecise (only measures to the nearest whole number), human decision.
- pH probe: rinse probe in water and put in solution, read measurement off screen. Device contains two electrodes that measure
the H+ activity in solution
- Advantages: more precise, can use data logging
- Disadvantages: hard to get a consistent reading, inaccurate calibration
- The equations for pH are given in the data booklet.
3.1.5: The ionic product constant of water, Kw
- The ion product constant of water, Kw, shows an inverse relationship between [H+] and [OH–] .
- Kw = [H+] x [OH–]
● Recognize solutions as acidic, neutral and basic from the relative values of [H+] and[OH–] .
- [H+] > [OH-] acidic
- [H+] = [OH-] neutral
- [H+] < [OH-] basic
- [OH-] = 10-2, [H+] =10-12, pH=12 and is basic as there are more [OH-] ions than [H+] ions
- Kw=10-14
- The equation for Kw and its value at 298 K are given in the data booklet.
3.1.6: Strong and weak acids and bases
- Strong and weak acids and bases differ in the extent of ionization.
- If an acid is strong its conjugate base will be weak and vice versa
- Recognize that acid–base equilibrialie in the direction of the weaker conjugate.
- HCl, HBr, HI, HNO3 and H2SO4 are strong acids, and Group 1 hydroxides are strong bases.
- Organic acids and other -COOH compounds as well as HF are weak acids, and ammonia, other amines and other hydroxides such as
Ca(OH)2 are weak bases
● The distinction between strong and weak acids or bases and concentrated and dilute reagents should be covered.
- Concentration is a measure of the number of moles of the substance/acid per unit volume, whereas the strength is a measure
of how likely these acid molecules are to donate the H+ and dissociate into their ions.
- pH depends on both: you can have a high concentration of a weak acid and a low concentration of a strong acid with a similar
pH value
● How would you expect the equilibrium constants of strong and weak acids to compare?
- Strong acids would have very high Kc values, whereas weak acids have low Kc values
- The position of equilibrium will lie further to the left/reactant side for very weak acids.
● Why does the acid strength of the hydrogen halides increase down group 17?
- Down group 17, the halogens become less reactive, and therefore the strength of the H-Halide bond is weaker down the
group. It is more likely for H-I to donate an H+ ionthan for H-F to break apart and donate an H+ ion. This greater likelihood
to dissociate means acid strength increases down the group.
● What physical and chemical properties can be observed to distinguish between weak and strong acids or bases of the same
concentration?
- Physical methods
- pH measurements either by probe or paper: The stronger the acid the lower the pH as more H+ ions will be in solution
- Conductivity: the stronger the acid or base, the greater the conductivity as more ions will be present in solution
- Chemical methods:
eactivity 3.1—Proton transfer reactions
R
.1.1- 3.1.3: Brønsted–Lowry theory
3
- A Brønsted–Lowry acid is a proton donor
- A Brønsted–Lowry base is a proton acceptor
● Deduce the Brønsted–Lowry acid and base in a reaction
- A proton in aqueous solution can be represented as both H+(aq) and H3O +
(aq)
- The distinction between the terms “base” and “alkali” should be understood.
- An alkali is a soluble base
- All alkalis are bases but not all bases are alkali
- A pair of species differing by a single proton is called a conjugate acid–base pair.
● Deduce the formula of the conjugate acid or base of any Brønsted–Lowry base or acid.
- Some species can act as both Brønsted–Lowry acids and bases.
- These species are called amphiprotic, they can act as both an acid and a base by donating or accepting a proton
- Not to be confused with amphoteric which does not involve proton donation but in which the substance can act as both an
acid and base
● Interpret and formulate equations to show acid–base reactions of these species.
3.1.4: The pH scale
- The pH scale can be used to describe the [H+] of a solution:
- pH = –log10[H+]
- [H+] = 10–pH
● Perform calculations involving the logarithmic relationship between pH and [H+].
T he concentration of hydrogen ions increases by a factor of 100 between pH5 and pH3
-
● Include the estimation of pH using universal indicator, and the precise measurement of pH using a pH meter/probe.
- Red,orange,yellow,green,blue,violet
- Strong acid, weaker acid, weak acid, neutral, weak base, strong base
- 1,3,5,7,11,14
● What is the shape of a sketch graph of pH against [H+]?
- Inversely proportional with each decrease in pH translating to a large increase in hydrogen ion concentration.
-
● When are digital sensors (e.g. pH probes) more suitable than analogue methods (e.g. pH paper/solution)?
- Universal indicator: dip paper in solution and compare color to strip to measure pH. Universal indicator is a mix of dyes.
- Advantages: easier to read, quick, accurate, no calibration, no power
, - Disadvantages: imprecise (only measures to the nearest whole number), human decision.
- pH probe: rinse probe in water and put in solution, read measurement off screen. Device contains two electrodes that measure
the H+ activity in solution
- Advantages: more precise, can use data logging
- Disadvantages: hard to get a consistent reading, inaccurate calibration
- The equations for pH are given in the data booklet.
3.1.5: The ionic product constant of water, Kw
- The ion product constant of water, Kw, shows an inverse relationship between [H+] and [OH–] .
- Kw = [H+] x [OH–]
● Recognize solutions as acidic, neutral and basic from the relative values of [H+] and[OH–] .
- [H+] > [OH-] acidic
- [H+] = [OH-] neutral
- [H+] < [OH-] basic
- [OH-] = 10-2, [H+] =10-12, pH=12 and is basic as there are more [OH-] ions than [H+] ions
- Kw=10-14
- The equation for Kw and its value at 298 K are given in the data booklet.
3.1.6: Strong and weak acids and bases
- Strong and weak acids and bases differ in the extent of ionization.
- If an acid is strong its conjugate base will be weak and vice versa
- Recognize that acid–base equilibrialie in the direction of the weaker conjugate.
- HCl, HBr, HI, HNO3 and H2SO4 are strong acids, and Group 1 hydroxides are strong bases.
- Organic acids and other -COOH compounds as well as HF are weak acids, and ammonia, other amines and other hydroxides such as
Ca(OH)2 are weak bases
● The distinction between strong and weak acids or bases and concentrated and dilute reagents should be covered.
- Concentration is a measure of the number of moles of the substance/acid per unit volume, whereas the strength is a measure
of how likely these acid molecules are to donate the H+ and dissociate into their ions.
- pH depends on both: you can have a high concentration of a weak acid and a low concentration of a strong acid with a similar
pH value
● How would you expect the equilibrium constants of strong and weak acids to compare?
- Strong acids would have very high Kc values, whereas weak acids have low Kc values
- The position of equilibrium will lie further to the left/reactant side for very weak acids.
● Why does the acid strength of the hydrogen halides increase down group 17?
- Down group 17, the halogens become less reactive, and therefore the strength of the H-Halide bond is weaker down the
group. It is more likely for H-I to donate an H+ ionthan for H-F to break apart and donate an H+ ion. This greater likelihood
to dissociate means acid strength increases down the group.
● What physical and chemical properties can be observed to distinguish between weak and strong acids or bases of the same
concentration?
- Physical methods
- pH measurements either by probe or paper: The stronger the acid the lower the pH as more H+ ions will be in solution
- Conductivity: the stronger the acid or base, the greater the conductivity as more ions will be present in solution
- Chemical methods:
