acids,bases and buffers Flashcards
What is meant by a weak acid and Brønsted–Lowry acid
Proton/H+ donor
AND
Partially dissociates/ionises
Calculate the pH of 0.500 mol dm–3 potassium hydroxide.
H+] = 1.00 ×10–14 0.5(00) OR 2(.00) × 10–14 (mol dm–3) pH = –log 2(.00) × 10–14 = 13.7(0)
Write the full equation for the reaction of aqueous propanoic acid with sodium carbonate.
2C2H5COOH + Na2CO3 → 2C2H5COONa + CO2 + H2O
Write the ionic equation for the reaction of aqueous propanoic acid with aqueous potassium hydroxide.
H+ + OH– → H2O
A small amount of aqueous ammonia, NH3(aq), is added to the buffer solution.Explain, in terms of equilibrium, how the buffer solution would respond to the added NH3(aq).
Added ammonia C2H5COOH removes added NH3/alkali/base OR C2H5COOH + NH3 / OH– → OR NH3/alkali reacts with/accepts H+ OR H+ + NH3 → OR H+ + OH– → Equlibrium → C2H5COO– OR Equilibrium → right
How can an aqueous solution of an acid contain hydroxide ions?
Water dissociates/ionises OR H2O ⇌ H \+ \+ OH – OR 2H2O ⇌ H3O \+ \+ OH–
Write a full equation for the reaction between ethanoic acid and solid calcium carbonate.
2CH3COOH + CaCO3 → (CH3COO)2Ca + CO2 + H2O
Explain why this buffer solution has formed.
solution contains CH3COOH AND CH3COO–
Explain how this buffer solution controls pH when either an acid or an alkali is added.
Quality of written communication, QWC
2 marks are available for explaining how the equilibrium
system allows the buffer solution to control the pH on addition
of H+
and OH-
(see below)
——————————————————
CH3COOH ⇌ H
+
+ CH3COO–
—————————————————–
CH3COOH reacts with added alkali
OR CH3COOH + OH–
OR added alkali reacts with H+
OR H
+
+ OH–
Equilibrium right OR Equilibrium CH3COO– (QWC)
CH3COO–
reacts with added acid
Equilibrium left OR Equilibrium CH3COOH (QWC)
0.14 mol dm–3 solutions of hydrochloric acid, HCl, and chloric(I) acid, HClO (pKa = 7.43), have different pH values.Explain why the pH values are different and calculate the pH of 0.14 mol dm–3 solutions of HCl and HClO to two decimal places.
Show any working in calculations.
HCl is a strong acid AND HClO is a weak acid HCl: pH = –log 0.14 = 0.85 (2 DP required) HClO: CHECK THE ANSWER ON ANSWER LINE IF answer = 4.14, award all three calculation marks -------------------------------------------------------------------- Ka = 10–7.43 OR 3.7 x 10–8 (mol dm–3) [H+] = Ka [HClO] OR Ka [HA] OR Ka 0.14 OR 3.7 x 108 x 0.14 pH = 4.14 (2 DP required)
Aluminium powder is added to aqueous ethanoic acid, CH3COOH.Write full and ionic equations for the reaction that takes place.
2Al + 6CH3COOH 2(CH3COO)3Al + 3H2
2Al + 6H+ 2Al3+ + 3H2
Calculate the pH of a 0.40 mol dm–3 solution of NaOH.
[H+] = Kw [OH ] OR 1.0 1014 [OH ] OR 1.0 1014 0.4(0) OR 2.5 x 10–14 (mol dm–3) Correctly calculates pH = –log 2.5 x 10–14 = 13.6(0)
Explain what is meant by the term buffer solution.
Describe how a buffer solution based on methanoic acid can act as a buffer.
A buffer solution minimises pH changes on addition of small amounts of acid/H+ or alkali/OH–/base -------------------------------------------------------------------------------- HCOOH ⇌ H+ + HCOO– Equilibrium sign essential Added alkali HCOOH reacts with added alkali/base/OH– OR added alkali/OH– reacts with H+ QWC: Equilibrium shifts forming HCOO– OR H+ OR (HCOOH) Equilibrium right Added acid HCOO– reacts with added acid/H+ QWC: Equilibrium shifts forming HCOOH OR (HCOOH) Equilibrium left
A chemist prepares a buffer solution by mixing together the following:200 cm3 of 3.20 mol dm–3 HCOOH (Ka = 1.70 × 10–4 mol dm–3) and800 cm3 of 0.500 mol dm–3 NaOH.The volume of the buffer solution is 1.00 dm3. • Explain why a buffer solution is formed when these two solutions are mixed together. • Calculate the pH of this buffer solution. Give your answer to two decimal places.
HCOOH reacts with NaOH forming HCOO– /HCOONa OR HCOOH + NaOH HCOONa + H2O Equilibrium sign allowed (Some) HCOOH/(weak) acid remains OR HCOOH/(weak) acid is in excess n(HCOOH) OR [HCOOH] = 0.24(0) (mol / mol dm–3) n(HCOO–) OR [HCOO– ] OR [HCOONa] = 0.4(00) (mol / mol dm–3) [H+] = K a [HCOOH] [HCOO– ] pH = –log [H+] = –log(1.70 10–4 0.24 0.4 ) = 3.99
Write the ionic equation for the reaction between aqueous butanoic acid and magnesium.
Mg + 2H+ Mg2+ + H2
Write the ionic equation for the reaction between aqueous butanoic acid and aqueous sodium carbonate.
CO32– + 2H+ H2O + CO2
The student adds 50.0 cm3 of 0.250 mol dm–3 butanoic acid to 50.0 cm3 of 0.0500 mol dm–3sodium hydroxide. A buffer solution forms.
Explain why a buffer solution forms.
CH3(CH2)2COONa OR CH3(CH2)2COO– forms
OR
CH3(CH2)2COOH + OH– CH3(CH2)2COO– + H2O
CH3(CH2)2COOH is in excess OR acid is in excess
OR some acid remains
What is the difference between a strong acid and a weak acid?
A strong acid completely dissociates
AND
a weak acid partially dissociates
Calculate the pH of 0.375 mol dm–3 nitrous acid, HNO2.Give your answer to two decimal places.
pH = –log 0.0129 = 1.89
OR
pH = –log 0.0129 = 1.9
Explain what is meant by the term Brønsted–Lowry base.
Proton acceptor
Calculate the pH of 0.0400 mol dm–3 Ca(OH)2.Give your answer to two decimal places.
[OH–] = 2 × 0.04(00) = 0.08(00) (mol dm–3) [H+] = 1.00 10–14 0.08(00) OR 1.25 × 10–13 (mol dm–3) pH = –log 1.25 × 10–13 = 12.90 --------------------------------------------- pOH variation (also worth 3 marks) [OH–] = 2 × 0.04(00) = 0.08(00) (mol dm–3) pOH –log 0.08(00) = 1.10 pH = 14.00 – 1.10 = 12.90
Aqueous calcium hydroxide is added to nitrous acid, HNO2.Write the overall equation and the ionic equation for the reaction that takes place.
overall: ……………………………………………………………………………………………………………………..
ionic: ……………………………………………………………………………………………………………………
Ca(OH)2 + 2HNO2 Ca(NO2)2 + 2H2O
H+ + OH– H2O
Carbonic acid, H2CO3, is a weak Brønsted–Lowry acid formed when carbon dioxide dissolves in water. Healthy blood is buffered to a pH of 7.40. The most important buffer solution in blood is a mixture of carbonic acid and hydrogencarbonate ions, HCO3–.
Explain how the carbonic acid–hydrogencarbonate mixture acts as a buffer in the control of blood pH.
In your answer you should explain how equilibrium allows the buffer solution to control the pH.
Equilibrium H2CO3 ⇌ H+ + HCO3 – Action of buffer Added alkali H2CO3 reacts with added alkali OR H2CO3 + OH– OR added alkali reacts with H+ OR H+ + OH– Equilibrium right OR equilibrium shifts forming H+ OR HCO3 – Added acid HCO3 – reacts with added acid Equilibrium left OR equilibrium shifts forming H2CO3
Healthy blood at a pH of 7.40 has a hydrogencarbonate : carbonic acid ratio of 10.5 : 1. A patient is admitted to hospital. The patient’s blood pH is measured as 7.20.Calculate the hydrogencarbonate : carbonic acid ratio in the patient’s blood.
n blood at pH 7.40, [H+] = 10–pH = 10–7.40 = 3.98 × 10–8 (mol dm–3) Ka = [H+ ] [HCO3 – ] [H2CO3 ] = 3.98 10–8 10.5 1 OR Ka = 4.18 × 10–7 (mol dm–3) In blood at pH 7.20, [H+] = 10–pH = 10–7.20 = 6.31 × 10–8 (mol dm–3) [HCO3 – ] [H2CO3 ] Ka [H+ ] OR 4.18 107 6.31 10–8 = 6.6 1 OR 6.6 : 1 (up to calc. value, see below)
What is meant by the term Brønsted–Lowry acid ?
proton donor
What is meant by the strength of an acid?
(the proportion of) dissociation
What important factor does the student need to consider when deciding on the most suitable indicator to use for this titration?
Vertical section matches the (pH) range (of the
indicator)
OR colour change (of the indicator)
OR end point (of the indicator)
explain whether the dissociation of water is an exothermic or endothermic process.
Endothermic because Kw increases with
temperature
Many experimental measurements use published data, such as Kw, measured at 25 °C. Often these measurements have been taken at different temperatures, especially in experimental work carried out at body temperature.What is the consequence of this for published scientific work?
(Work is) inaccurate OR inva
because K varies with temperature
lid
w
The reverse reaction of the dissociation of water is called neutralisation.Plan an experiment that a student could carry out to measure the enthalpy change of neutralisation.
Acid and alkali mixed Amounts of acid AND alkali stated Temperature taken at start AND finish energy, Q = mc∆T OR in words AND meaning of m, c AND ∆T given Energy scaled up to form 1 mol of water ∆H = –energy change
A student made a buffer solution by mixing an excess of propanoic acid to an aqueous solution of sodium hydroxide at 25 °C. This buffer solution contains an equilibrium system that minimises changes in pH when small amounts of acids and alkalis are added. • Explain why a buffer solution formed when an excess of propanoic acid was mixed with aqueous sodium hydroxide.• Explain how this buffer solution controls pH when an acid or an alkali is added.
Propanoic acid reacts with sodium hydroxide forming propanoate ions/sodium propanoate OR CH3CH2COOH + NaOH CH3CH2COONa + H2O Some propanoic acid remains OR propanoic acid AND propanoate (ions) / sodium propanoate present equilibrium: CH3CH2COOH ⇌ H+ + CH3CH2COO– Added alkali CH3CH2COOH reacts with added alkali OR CH3CH2COOH + OH– OR added alkali reacts with H+ OR H+ + OH– CH3CH2COO– OR Equilibrium right Added acid CH3CH2COO– reacts with added acid OR [H+] increases CH3CH2COOH OR Equilibrium left
Write an equation for the reaction of aqueous propanoic acid with magnesium.
2CH3CH2COOH + Mg
(CH3CH2COO)2Mg + H2
Write an ionic equation for the reaction of aqueous propanoic acid with aqueous sodium carbonate.
2H+ + CO32– H2O + CO2 OR 2H+ + CO32– H2CO3 OR H+ + CO32– HCO3 –
