Chemistry 101

Question 1

Identify the compound usually added to water to provide the Mn2+ ions.

Answer 1

MnSO4.2H2O (manganese(II) sulfate) / MnCl2.4H2O (manganese(II) chloride)

Question 2

What change is observed when concentrated sulphuric acid is added to a stoppered bottle, completely filled with a sample of the water, was shaken with a concentrated solution containing Mn2+ ions and then with a concentrated solution of alkaline potassium iodide (KOH/KI).

Answer 2

A brown solution is formed

Question 3

Describe how the conical flask was prepared for a titration and then used during a titration to help ensure that an accurate end point was reached.

Answer 3

It was first rinsed with deionised water.
swirl to mix the content of the flask.
Wash down the sides of the flasks with deionised water.

Question 4

Give one way of ensuring that the level of liquid in the burette was at eye level before taking a reading.

Answer 4

bend (crouch, kneel) until eye at level of liquid

Question 5

When preparing a solution of Sodium Thiosulfate and Standardising it by Titration against a Solution of Iodine,
what colour changes were observed at the titration stage of the experiment
(i) up to the point when the indicator was added,
(ii) at the end point?

Answer 5

i) Brown fades to pale yellow

ii) Blue-black to colourless

Question 6

The sodium thiosulfate solution contained 3.1 g Na2S2O3.5H2O per litre. On average, 6.0 cm3 of this solution were required for complete reaction with 200.0 cm3 portions of the iodine solution.
The balanced equation for the titration reaction is:
I2 + 2S2O32– = 2I– + S4O62–
Four moles of S2O32– were required for every one mole of dissolved O2 present originally.
Show by calculation:
(i) the average number of moles of sodium thiosulfate used in a titration,
(ii) the number of moles of O2 that released the I2 detected in each titration,
(iii) the concentration of dissolved oxygen in the water in moles per litre,
(iv) the concentration of dissolved oxygen in the water in p.p.m. (mg l–1).

Answer 6

i) moles= mass/ molar mass
M = 3.1/ 248 = 0.0125 (1/80) M
(6.0 x 0.0125) / 1000 = 0.000075 (7.5 × 10–5, 3/40000) moles thiosulfate.

ii) Thiosulfate : O2 = 4 : 1
0.000075 (7.5 × 10–5, 3/40000) ÷ 4
= 0.00001875 (1.875 × 10–5, 3/160000) moles O2 in 200.0 cm3.

iii) 0.00001875 (1.875 × 10–5, 3/160000) × 5 = 0.00009375 (9.375 × 10–5, 3/32000) moles /l (M) of O2

iv) 0.000009375 (9.375 × 10–5, 3/32000) × 32 = 0.003 g/l
0. 003 × 1000 = 3.0 p.p.m.

Question 7

A student placed a short strip of freshly‐sanded magnesium ribbon into a test tube containing a few drops of ethanoic acid in approximately 10 cm3 of water and swirled the test tube.

(i) What was observed?
(ii) Write formulae for both products of the reaction that occurred.

Answer 7

i) bubbling (fizzing, effervescence, foaming, gas evolves) / solid (magnesium, Mg) dissolves (disappears) / solution forms / white powder (solid, precipitate) formed.
ii) H2 (hydrogen) and (CH3COO)2Mg (magnesium ethanoate)

Question 8

Describe a chemical test to show that ethanal can be very easily oxidised.

Answer 8

Add Tollens’ reagent to ethanal. Then heat gently using a water bath. Silver (Ag, mirror, precipitate, solid) is positive test for aldehyde

Question 9

A student was asked to prepare a small quantity of ethene from ethanol using the reaction shown by the balanced equation below. A suitable catalyst and a heat source were provided.
C2H5OH = C2H4 + H2O
(i) Identify the catalyst and describe its appearance at the beginning of the preparation.
(ii) Explain clearly why a risk of hot glassware shattering, due to rapid cooling, is associated with this preparation.
How can this risk be minimised?
(iii) A student prepared ethene, starting with 2.9 cm3 of ethanol (density 0.8 g cm–3). Calculate the volume of ethene gas produced in 26% yield when measured
at room temperature and pressure.

Answer 9

i) Aluminium oxide. It is a white powder.
ii) Cold water enters hot test-tube (apparatus, glassware). A suck-back can occur. To minimise this risk, loosen stopper before reducing (removing) heat.

iii) Density = mass/volume
0.8= mass/2.9
m=0.80×2.9=2.32g
2.32/ 46 =0.0504 moles ethanol used
0.0504 × 0.26 = 0.013104 actual mol ethene
0.013104 × 24.0 = 0.300 - 0.315 litres ethene

Question 10

To determine the relative molecular mass (Mr) of compound X, a volatile liquid, a small volume of the pure liquid was vaporised in a suitable container. The atmospheric pressure at the location was obtained using a smart phone barometric sensor.

Explain
(i) why the pressure of the vapour at the end of the heating stage of the experiment was known to be equal to the atmospheric pressure
at your location,
(ii) how the volume of the vapour was measured.
(iii) Describe how the mass of the vapour in the container
at the end of the heating stage was found.
(iv) Calculate, correct to the nearest whole number, the relative molecular mass of X, given that 0.22 g of X occupied a volume of 76 cm3 at a pressure of 1.011 × 105 Pa and a temperature of 99C.
(v) Bromine (Br2) (Mr = 160) has a significantly greater relative molecular mass than water (Mr = 18). Account for the volatility of bromine (boiling point 58.8 C) compared to that of water (boiling point 100 C).

Answer 10

i) Some vapour (gas, air) escaped.
ii) Flask filled with water using (with) graduated cylinder.
iii) Subtract the initial mass of the flask, rubber band and foil from the mass of flask, rubber band, foil and its contents.

iv) n= pV/ RT
pV=nRT
n = 2.49 × 10-3moles
Mr = m/n
Mr = 0.22/ 2.49x 10-3
Mr = 88.4 = 88
v) hydrogen (H) bonds in water
van der Waals forces in bromine (Br2)
intermolecular forces in water (H2O) stronger.
Intermolecular forces in bromine (Br2) weaker.
So Bromine (Br2) non-polar (pure covalent) while water (H2O) is polar

Question 11

What are cathode rays?

Answer 11

beams of electrons

Question 12

Bananas contain small quantities of potassium‐40, a radioactive isotope. What is the daughter nucleus when K‐40 emits an electron in beta decay?

Answer 12

calcium-40 / Ca-40

Question 13

State Avogadro’s law.

Answer 13

Equal volumes of gases at the same temperature and pressure have equal numbers of particles.

Question 14

What is the oxidation number of sulfur in

(i) sulfur dioxide (SO2),
(ii) the sulfate ion (SO42–)?

Answer 14

i) SO2: 4 / +4

ii) SO42– :6/+6

Question 15

Give two differences between a sigma and a pi covalent bond.

Answer 15

Sigma bonds are formed from head-on overlap of orbitals while pi bonds are formed from sideways overlap of orbitals.
Sigma bonds are associated with single covalent bonds while pi bonds are associated with multiple covalent bonds.
Sigma bonds determine shape of molecule while Pi bonds do not determine shape of molecule.

Question 16

In the electrolysis of aqueous KI with inert electrodes one half‐equation is:
2I– (aq) = I2 (aq) + 2e–
Write a balanced half‐equation for the reaction that takes place at the negative electrode during this electrolysis.

Answer 16

2H2O +2e –→ H2 +2OH–

Question 17

Define bond energy.

Answer 17

This is the average amount of energy involved in breaking/ forming bonds of a particular type in a chemical species in the gaseous state.

Question 18

What is the principle of the separation of the components in a mixture using any type of chromatography?

Answer 18

The components have different attractions for stationary and mobile phases.

Question 19

A single 5 cm3 dose, suitable for children aged 1 to 3 years, contains 4.86 × 10–4 moles of ibuprofen.
How many milligrams of ibuprofen are in this single dose?

Answer 19

n = m/ Mr
m = Mr x n
4.86×10–4×206=0.100g
–→100.116mg

Question 20

Write a balanced equation for the reaction – between limestone and carbonic acid in rainwater – that gives rise to temporary hardness in the water.

Answer 20

H2CO3 + CaCO3 → Ca(HCO3)2

Question 21

Why is nitrogen gas unreactive?

Answer 21

very strong (high energy) triple bond

Question 22

What is meant by the periodic system in the context of Mendeleev’s 1869 periodic table of elements?

Answer 22

Elements listed according to relative atomic mass (weight) and whose chemical properties repeat periodically (at regular intervals).

Question 23

Comment on the positioning of tellurium (Te) and iodine (I) in the 1869 table.

Answer 23

Their chemical properties fitted better if their orders were reversed.

Question 24

Why did the 1869 table not include any noble gases?

Answer 24

They were undiscovered (unknown) in 1869.

Question 25

Define atomic number

Answer 25

This is the number of protons in the nucleus of an atom

Question 26

Define relative atomic mass

Answer 26

This is defined as the average of the mass numbers of the isotopes of the element, as they occur naturally. Taking their abundances into account relative to 1/12th of the carbon-12 isotope.

Question 27

Give an advantage of arranging the elements in order of atomic number in a periodic table.

Answer 27

There is no need to reverse order to force elements into correct groups.

Question 28

Explain why all the elements of Group 18 in the periodic table are chemically inert

Answer 28

Because they have a stable arrangement of electrons

Question 29

Explain how and why the reactivity of the halogens changes down Group 17.

Answer 29

As you go down the group, they become less reactive because there is an increasing atomic radius.

Question 30

The element francium (Fr), atomic number 87, was discovered by Marguerite Perey in 1939 but its physical and chemical properties had already been predicted
by the periodic system.
(i) How would you expect a small sample of francium to react in water? Justify your answer.
(ii) Predict the products of this reaction.

Answer 30

i) Vigorously because reactivity increases down the group.

ii) FrOH (francium hydroxide) and H2 (hydrogen)

Question 31

What term is used to describe compounds that have the same molecular formula but different structural formulae?

Answer 31

structural isomers

Question 32

(i) Name the oil refining process in which one molecule of alkane W was converted into one molecule of octane and two propene molecules.
(ii) Deduce the formula of W.

Answer 32

i) Catalytic cracking

ii) C14H30

Question 33

(i) What is the advantage of adding tetraethyllead to petrol?
(ii) Why was its use in car engines discontinued?

Answer 33

i) To reduce autoignition.

ii) It is toxic.

Question 34

Use equations to show that, when dissolved in water,
(i) HCl acts as a Brønsted‐Lowry acid,
(ii) NH3 acts as a Brønsted‐Lowry base.
Explain why
(iii) HCl has a weak conjugate base,
(iv) NH3 has a strong conjugate acid.

Answer 34

i) H2O + HCl ⇌ H3O+ + Cl–
ii) H2O + NH3 ⇌ NH4+ + OH–
iii) It is a good proton donor
iv) It is a poor proton acceptor

Question 35

Write the self‐ionisation constant (Kw) expression for water.

Answer 35

Kw = [H3O+][ OH–]

Question 36

What is a secondary alcohol?

Answer 36

Two carbon (C) atoms attached to carbon (C) to which the OH (alcohol group, hydroxyl group) is attached

Question 37

State and explain the trend in the boiling points of the four primary alcohols as their relative molecular masses increase.

Answer 37

Their boiling points increases because they have more intermolecular forces.

Question 38

Consider the oxidation of alcohols in which no carbon‐carbon bonds are broken.

(i) Give the systematic IUPAC names for the two possible organic products of such an oxidation of butan‐1‐ol.
(ii) Identify clearly which bonds in butan‐2‐ol are broken in this oxidation.

Answer 38

i) 1-butanal / butan-1-al and 1-butanoic acid / butan-1-oic acid.
ii) OH bond and the CH bond of carbon to which OH is attached

Question 39

The ester formed from methanol and propanoic acid is found in many fruits.

(i) How many carbon atoms in a molecule of this ester are tetrahedrally bonded?
(ii) What are the products of the hydrolysis of this ester by NaOH?

Answer 39

i) 3

ii) methanol / CH3OH and Sodium propanoate / sodium propionate / CH3CH2COONa / C2H5COONa

Question 40

Define rate of reaction.

Answer 40

This is the change in concentration by unit time.

Question 41

Explain the term heterogeneous catalysis.

Answer 41

This is when the catalysts is in a different phase from the reactants and products.

Question 42

Describe the surface adsorption theory of catalysis.

Answer 42

The reactants are adsorbed on the surface of the catalyst. The reaction takes place on the surface of the catalyst and the products desorb from the surface.

Question 43

What is activation energy?

Answer 43

Minimum combined energy of colliding particles required for reaction to take place between them.

Question 44

State and explain two ways of increasing the rate of this reaction, other than by using a catalyst.

Answer 44

Increase the concentration of the reactants. This way there will be a greater frequency of collisions between reactants.

Increase the temperature. This way the reactants have more energy and can move faster.

Question 45

What are free radicals?

Answer 45

Atoms that have an unpaired electron.