Acids, bases and indicators
Identify and describe two everyday uses of indicators.
1. The maintenance of particular acidity levels (pH 7.2–7.6) in swimming pools. A sample of pool water is collected in a small vial, to which several drops of phenol red are added. The colour of the resulting solution is then measured against a colour chart.
2. Many plants do not grow well if the soil is too
acidic. Therefore indicators allow horticulturalists and farmers to test the acidity of their soil and make adjustments as required. A small amount of universal indicator is added to a soil sample, barium sulfate (a white powder) is then added. Wait a few minutes for a colour change and compare to a colour chart.
Outline the Brönsted-Lowry theory of acids and bases
An acid–base reaction is one that involves the transfer of a proton from one species to another.
The substance that donates the proton is an acid, and the proton acceptor is the base.
A nitric acid solution had a pH of 2.
10 mL of the solution was diluted to 100 mL. The pH of the fi nal solution would be closest to:
A 0.2
B 2.5
C 1.0
D 3.0
D
Each one-unit change in the pH scale corresponds to a ten-fold change in hydrogen ion concentration.
Explain why neutralisation reactions are exothermic
The reaction is exothermic because it
results in the formation of a new covalent bond between the hydrogen (from the H+) and oxygen (from the OH–) as water molecules are formed.
Universal indicator is added to lemon juice - what colour will result?
The indicator is then added to household bleach - what colour will it turn?
The indicator is then added to water - what colour will it turn?
For lemon - orange/red
Bleach - Purple/dark blue
Water - green
For the reaction:
CH3COOH(aq) + OH–(aq) <--> H2O(l) + CH3COO–(aq)
Identify conjugate acid/base pairs
CH3COOH is acting as an acid (it is donating a proton) - its conjugate acid-base pair is CH3COO–
OH– is acting as a base (it is accepting a proton) - its conjugate acid-base pair is H2O
The concentration of hydrochloric acid in our stomachs is normally about 0.12 mol L–1.
Calculate the normal pH of the stomach.
As its a strong acid,
[H+] = 0.12 mol L–1
pH = –log10[H+] = –log10[0.12]
pH = 0.92
1.336 g of pure anhydrous Na2CO3 was dissolved in water and made up to 250.0 mL of solution in a volumetric flask.
Calculate the concentration of this standard solution.
Step 1 - Calculate the number of moles of Na2CO3 using n=m/mm
= 0.012605 molesStep 2 - Calculate the conc. in 250mL using n=cV
=0.05042 mol L-1
Give the ionisation equations for:
the strong acid - hydrochloric
the weak acid - carbonic
the strong base - barium hydroxide
HCl(g) → H+(aq) + Cl–(aq)
H2CO3(aq) <--> H+(aq) + HCO3-(aq)
Ba(OH)2(s) → Ba2+(aq) + 2OH–(aq)
Summarise the key ideas about acids of:
– Lavoisier
– Davy
– Arrhenius
Lavoisier put forward the hypothesis that all acids
contained oxygen. Lavoisier thought that oxides, when dissolved in water, formed acids.
Davy proposed that all acids contained hydrogen rather than oxygen.
Arrhenius proposed the theory that an acid is a substance that produces hydrogen ions in
water.
When ammonia gas is dissolved in water, the resulting solution is basic. Account for this observation using a chemical equation to support your answer.
NH3(g) + H2O(l) <--> NH4+(aq) + OH–(aq)
The OH– causes the solution to be basic.
Write the equation for the reaction between nitric acid and sodium hydroxide. What type of reaction is this and what would be the pH of the resulting solution?
HNO3(aq) + NaOH(aq) → NaNO3(aq) + H2O(l) + heat
Neutralisation
pH = 7 (results of a strong acid + strong base)
What is a 'triprotic acid' and give the successive ionisation equations for phosphoric acid (H3PO4)
Tripotic acid = an acid that contains three acidic protons.
H3PO4(aq) <--> H+(aq) + H2PO4–(aq)
H2PO4-(aq) <--> H+(aq) + HPO42–(aq)
HPO42–(aq) <--> H+(aq) + PO43–(aq)
A 5.00 mL sample of commercial vinegar was completely neutralised by 24.50mL of 0.150 mol L–1 NaOH solution. Calculate the acetic acid concentration of the vinegar.
NaOH(aq) + CH3COOH(aq) → H2O(l) + NaC2H3CO2(aq)
Step 1 Calc moles of 24.50mL NaOH using c=n/V = 3.675x10-3 moles
Step 2 Calc moles of CH3COOH = 3.675x10-3 moles
Step 3 Calculate the concentration of 5mL CH3COOH using c=n/V = 0.735 mol L-1
The solubility of limestone (CaCO3) depends on the following equilibrium system:
CaCO3(s)+H2O(l)+CO2(g) <-> Ca2+(aq) + 2HCO3–(aq)
ΔH = –42kJ
Deep within limestone caves the pressure on gases is greater than atmospheric pressure. As water leaves a cave system, the pressure is reduced and temperature is increased.
(i) What would a reduction of pressure have
on this equilibrium system?
(ii) How would an increase in the temperature cause a change to the concentration of calcium ions?
(i) Decreased pressure would cause the equilibrium to shift to the left. This would result in the formation of CaCO3(s) and CO2(g).
(ii) A shift to the left to counteract the change. This would cause a decrease in [Ca2+].
Copper metal can be extracted from the ore chalcopyrite (CuFeS2) as shown in the following equation:
2CuFeS2(s)+5O2(g)+2SiO2(s)-->2Cu(l)+4SO2(g)+2FeSiO3(l)
If 500 kg of chalcopyrite is roasted with air in the
presence of silica (SiO2), calculate the volume of sulfurdioxide produced at 100 kPa and 25°C.
Why is the production of SO2 a concern?
Step 1 = moles of CuFeS2 using n=m/mm
n of SO2 = 2 x moles of CuFeS2
Step 2 = Vol. SO2 = n of SO2 x 24.79L
= 135 066L
SO2 reacts with water in the atmosphere to produce acid rain.