Define a Brønsted–Lowry acid.
Which species in the following reaction acts as a Brønsted-Lowry base.
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H2SO4 (aq) + H2O (l)
Which species in the following equation is acting as a Brønsted-Lowry acid.
CO32- (aq) + H+ (aq) HCO3- (aq)
Explain, using the Brønsted-Lowry theory, how water can act either as an acid or a base.
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Describe the difference between an amphiprotic and amphoteric species.
i) State what is meant by the term conjugate base.
[1]
ii) State the conjugate base of the hydroxide ion, OH–
[1]
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State an equation for the reaction of magnesium carbonate with dilute hydrochloric acid.
State an equation for the reaction of lithium oxide with dilute nitric acid.
Which acid and base would be required to produce ammonium sulfate, (NH4)2SO4
Nitric acid and calcium hydroxide react together.
i) State the type of reaction that takes place.
[1]
ii) State the formula of the products of the reaction.
[1]
iii) State the sign of the enthalpy change for this reaction.
[1]
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Using section 18 of the data booklet, identify an indicator that would show a yellow colour in ammonia solution?
Suggest two characteristics that make a good indicator for a titration.
A typical set of acid-base titration results in shown in the table.
| Rough | Run 1 | Run 2 |
|---|---|---|---|
Initial burette reading /+0.05 mL | 0.00 | 0.30 | 0.60 |
Final burette reading/ +0.05 mL | 24.15 | 22.55 | 22.95 |
Determine the mean volume from these results.
What is the recorded uncertainty on the mean volume calculated in part c)?
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State the relationship between pH and hydrogen ion concentration.
Determine the pH of 0.200 mol dm-3 hydrochloric acid.
Determine the hydrogen ion concentration in a sample of lake water of pH 5.60.
The table below shows the hydrogen ion concentration in three solutions:
| P | Q | R |
|---|---|---|---|
[H+] | 0.001 | 1 x 10-5 | 1.00 |
List the three solutions in order from low pH to high pH
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State one advantage and one disadvantage of using a pH meter instead of universal indicator to measure pH.
State the name and formula of a strong alkali and a weak alkali.
State the meaning of the term dissociation as applied to acids and bases
Write equations for the dissociation of:
Nitric acid, HNO3: ...................................
Methanoic acid, HCOOH: ............................
Identify the formula of the weakest conjugate base produced in the two acids in part d).
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Explain the difference between the terms strong acid and weak acid.
Other than measuring the pH, describe one laboratory method, other than pH measurement, to distinguish between the two solutions. State the expected result for each acid.
A solution of 2.00 mol dm-3 hydrochloric acid was added to marble chips and the volume of carbon dioxide recorded. A graph of the result is shown below:
On the same graph, sketch the result of repeating the experiment with 2.00 mol dm-3 ethanoic acid.
The same experiment in part c) can be carried out by measuring how the mass of the reaction flask changes with time.
Sketch a graph of the expected result.
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Malonic acid is a naturally occurring acid found in fruits and vegetables and is shown below.
The first dissociation of malonic acid is:
C3H4O4 (aq) + H2O (l) ⇌ C3H3O4- (aq) + H3O+ (aq)
Identify one conjugate acid-base pair from the equation.
The equilibrium constant for the first dissociation of malonic acid is 1.48 x 10-3.
State, with a reason, the strength of malonic acid.
The anion C3H3O4- may be classified as amphiprotic. Explain the meaning of amphiprotic and write equations, using C3H3O4-, to illustrate your answer.
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Salicylic acid has the structure shown below.
Draw the structure of the conjugate base of salicylic acid, showing all the atoms and all the bonds.
Predict what would be seen if a small amount of copper (II) oxide was added to an aqueous solution of salicylic acid, HOC6H4COOH, and warmed.
Write a balanced equation for the reaction.
Suggest, with a reason, whether salicylic acid is likely to be soluble in water.
Determine the relative molecular mass, Mr, of salicylic acid using Table 6 from the Data book.
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Glycolic acid, C2H4O3, is a colourless, odourless crystalline solid that is highly soluble in water and behaves as a Brønsted–Lowry acid.
i) Define the term Brønsted–Lowry acid. [1]
ii) State one difference between Brønsted–Lowry acids and the traditional theory of acids as substances that dissociate in water to form hydrogen ions. [1]
The systematic IUPAC name for glycolic acid is 2-hydoxyethanoic acid.
Draw the structural formula for its conjugate base, showing all the atoms and bonds.
Write an equation for the reaction between glycolic acid, C2H4O3, and limescale, CaCO3. State and explain one observation you would make.
State one reason why you would use glycolic acid to remove the limescale in a kettle at home, but not hydrochloric acid.
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Carbon dioxide gas dissolves in rainwater to form carbonic acid. State the formula of the conjugate base of carbonic acid.
Carbonic acid and sulfuric acid can be described as diprotic acids. Explain the meaning of diprotic.
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The equilibrium constant for the first dissociation of formic acid is 1.8 x 10-4 mol dm-3.
State, with a reason, the strength of formic acid.
Outline one laboratory method used to distinguish between equimolar solutions of formic acid and hydrochloric acid, giving the expected observations.
Formic acid has the chemical formula HCOOH. Identify the conjugate base of formic acid and state whether it is a weak or strong conjugate base.
Draw the structure of formic acid and give its systematic IUPAC name.
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The pH of an aqueous solution of salicylic acid at 298 K is 3.85. Determine the concentration of hydroxide ions in the solution, using section 2 of the data booklet.
A and B are two solutions of the same concentrations that have pH values of 3 and 6 respectively.
i) Identify which is the stronger acid and calculate the concentration of hydrogen ions in each solution.
[2]
ii) Calculate the ratio of the hydrogen ion concentrations in both A and B.
[1]
The variation of conductivity and concentration of a strong and weak monoprotic acid are shown below.
Identify the strong and weak acid from the information given and justify your choices.
For acid 1 and acid 2 in part (c) compare the volume of 0.2 mol dm-3 NaOH required to neutralise 20 cm3 of 0.1 mol dm-3 solutions of the acids.
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The concentrations of solutions of weak acids can be determined by titration against standard solutions of alkalis, such as sodium hydroxide.
i) Explain what is meant by the term standard solution.
[1]
ii) State the name of the indicator which should be used for this titration and what would be observed at the equivalence point of the reaction if the sodium hydroxide is placed in the burette.
[2]
A solution of 25.0 cm3 ethanoic acid was titrated against 0.150 mol dm-3 NaOH (aq) and it was found that 22.35 cm3 of the NaOH was needed for complete neutralisation.
Write an equation for the reaction and determine the concentration of the ethanoic acid.
A solution of 0.1 mol dm-3 ammonia has a pH of approximately 11. Predict how the pH value of 0.1 mol dm-3 sodium hydroxide solution would compare and calculate its value.
Write an equation for the reaction between ammonia and water and use it to classify each product as a Brønsted–Lowry acid or base.
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Four solutions of acids with identical concentrations are prepared. The equilibrium constants of these acids are given in the table below.
Acid | Kc at 298 K |
HCN | 4.9 x 10-10 |
HF | 6.8 x 10-4 |
CH3COOH | 1.7 x 10-5 |
HCl | 1.3 x 106 |
Write down the acid dissociation equation for HCN.
Use the information in part (a) to complete this question.
i) Write down the list of acids in part (a) in order of decreasing pH.
[1]
ii) Write down the list of acids in order of increasing concentration of molecules of the acid present in the solution.
[1]
State the name and formula of all the chemical species present in the solution of CH3COOH.
Write the name and formula of the conjugate base of HF.
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This question is about Brønsted-Lowry acids and bases.
i) Give the meaning of the term Brønsted-Lowry acid.
[1]
ii) Explain the term weak acid.
[2]
When an acid and a base react they produce a conjugate base and a conjugate acid.
acid + base ⇋ conjugate base + conjugate acid
Write an equation to show how hydrochloric acid behaves as a strong acid when it reacts with water, and state the role of water in this reaction.
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25.0 cm3 of 0.100 mol dm-3 propanoic acid, CH3CH2COOH (aq) , is titrated with of 0.100 mol dm-3 sodium hydroxide. The pH curve for this titration is shown below.
i) Label the equivalence point and half equivalence point on the curve.
[2]
ii) Explain what is meant by the half equivalence point.
[1]
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At 298K, water molecules dissociate into equal quantities of ions, and the pH is 7.
i) Write an equation to show the dissociation of water.
[1]
ii) At 313 K, the pH of water is 6.77. Explain why water is still neutral with a pH of 6.77.
[1]
The ionic product of water, Kw, can be used to find the pH of a strong base. Changing the temperature will affect the value for Kw.
i) Give the expression and units for the ionic product of water, Kw.
[2]
ii) As temperature increases, the value for Kw also increases. Explain why.
[3]
Determine the pH of pure water at 40 °.
Kw of pure water at 40 ° is 2.92 x 10-14 mol2 dm-3
Strong bases fully ionise in water, as shown by the equation of dissociation of sodium hydroxide:
NaOH (aq) → Na+ (aq) + OH- (aq)
At 298K, Kw is 1 x 10-14 mol2 dm-6.
Calculate the pH of a 0.05 mol dm-3 solution of NaOH at 298 K.
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Ethanoic acid (CH₃COOH) is a weak acid used in food preservation and as a cleaning agent.
Write balanced chemical equations for the reaction of aqueous ethanoic acid with:
i) solid calcium carbonate.
[1]
ii) aqueous sodium hydroxide.
[1]
Explain, with reference to the specific types and strengths of intermolecular forces, why the boiling point of ethanoic acid (118 °C) is significantly higher than that of ethanol (78 °C), despite their similar relative molecular masses.
Calculate the minimum volume, in cm3, of 0.500 mol dm-3 aqueous ethanoic acid required to completely neutralise 1.25 g of solid magnesium oxide (MgO).
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Explain why an ammonium ion can not behave as a Brønsted-Lowry base.
State and explain the acid-base character of aqueous ammonia at 298 K.
Acids can be classed as monoprotic, diprotic and triprotic. Sulfuric acid is a diprotic acid.
i) State the equation for the first ionisation step of sulfuric acid, including state symbols.
[1]
ii) Label the conjugate acid and base pairs in your answer to part i).
[1]
The second ionisation step is for the ionisation of sulfuric acid is as follows.
HSO4- (aq) + H2O (aq) ⇌ SO42- (aq) + H3O+ (aq)
Suggest why the second ionisation step reaches equilibrium.
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Sodium hydrogen carbonate solution, NaHCO3 (aq) , can act as an amphiprotic species. State the equation for the reaction fo NaHCO3 (aq) with the following compounds:
i) Sodium hydroxide solution.
[1]
ii) Hydrochloric acid.
[1]
Using your answer to part a) i) and ii), explain why NaHCO3 is amphiprotic.
Phosphine is usually prepared by heating white phosphorus, one of the allotropes of phosphorus, with concentrated aqueous sodium hydroxide.
The equation for the reaction is.
P4 (s) + 3OH− (aq) + 3H2O (l) → PH3 (g) + 3H2PO2− (aq)
Identify the amphiprotic species in this reaction giving the formulas of both species it is converted to when it behaves in this manner.
1.68 g of white phosphorus was used to make phosphine
i) Calculate the amount, in mol, of white phosphorus used.
[1]
ii) This phosphorus was reacted with 50.0 cm3 of 3.00 mol dm−3 aqueous sodium hydroxide. Deduce, showing your working, which was the limiting reagent.
[1]
iii) Determine the excess amount, in mol, of the other reagent.
[1]
iv) Using section 2 of the data booklet. Determine the volume of phosphine, measured in cm3 at standard temperature and pressure, that was produced.
[1]
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Oxalic acid, H2C2O4 , is a weak diprotic acid and can be used in titrations. State the equation for the reaction of oxalic acid with sodium hydroxide.
The ionisation of oxalic acid occurs in two steps. State equations for both of these steps.
Tartaric acid shown below behaves as a Brønsted-Lowry acid when it reacts with calcium hydroxide, Ca(OH)2. Sketch the structure of the salt formed from this reaction.
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Using ionic equations state how HPO42- can behave as an amphiprotic and amphoteric species.
Gallium oxide behaves as an amphoteric oxide. State two equations to show how gallium oxide reacts with a strong monoprotic acid and strong base.
Reaction with strong monoprotic acid .........................................................................................................................................
Reaction with strong base .........................................................................................................................................
Identify the Br∅nsted-Lowry acids in the following reaction.
CH3CH2O- (aq) + H2O (l) CH3CH2OH (aq) + OH- (aq)
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A solution of hydrochloric acid of concentration 0.001 mol dm-3 has a pH value of 3. Suggest, giving a reason, the pH of the following solutions of acids:
i) 0.01 mol dm-3 hydrochloric acid
[2]
ii) 0.01 mol dm-3 ethanoic acid
[2]
Two separate titrations are carried out using 25.00 cm3 of 0.01 mol dm-3 solutions of hydrochloric acid followed by ethanoic acid, against 0.01 mol dm-3 sodium hydroxide.
State what difference(s) would be observed in the two titrations.
Suggest a suitable indicator for the titration of hydrochloric acid and sodium hydroxide in part c), and state the colour changes observed.
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Show how the ionic product for water is derived from the dissociation of water and give it units.
Determine the pH of 0.001 mol dm-3 sodium hydroxide.
Suggest, with a reason, how the magnitude of Kw changes with increasing temperature.
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Malonic acid is a weak dibasic carboxylic acid with the formula C3H4O4. Draw the displayed structure of malonic acid.
Suggest, with a reason, which of the two equimolar aqueous solutions, ethanoic acid or malonic acid, would have a higher pH.
Apart from testing the pH, suggest how equimolar solutions of malonic acid and ethanoic acid may be distinguished.
Write the formulas of two conjugate bases that can be formed from malonic acid.
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Sketch the titration curve obtained when 0.1 mol dm⁻³ HNO3 (aq) is added from a burette to 25.0 cm3 of 0.1 mol dm-3 NH3 (aq).
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