Chemistry Paper 1 Mistakes

Question 1

Difference between enthalpy of atomisation and bond dissociation enthalpy in born haber calculations

Answer 1

Enthalpy of atomisation = 1/2 x bond dissociation

Question 2

How to calculate enthalpy of formation from a born haber cycle

Answer 2

Enthalpy of formation = everything else added together

Question 3

Why does the EMF of the cell decreases to 0V

Answer 3

The concentrations of the ions in the two solutions become equal

Question 4

In terms of electrostatic forces, suggest why the electron affinity of fluorine has a negative value

Answer 4

There is an attraction between the electron and nucleus.
Energy is released when electron is gained.

Question 5

Define enthalpy of hydration of an ion

Answer 5

The enthalpy change when 1 mol of a gaseous ion forms aqueous ions

Question 6

Platinum acts as a heterogenous catalyst in the reaction between ammonia and oxygen. It provides an alternative reaction route with a lower activation energy.

Describe the stages of this alternative route

Answer 6

Ammonia and oxygen adsorb onto the surface of platinum.
Bonds weaken/break
Product desorbs from the surface of platinum

Question 7

Explain why the second ionisation energy of calcium is lower than the second ionisation energy of potassium

Answer 7

In potassium the electron is being removed from a new energy level closer to the nucleus.
Less shielding in K+

Question 8

Fumes of sulfur dioxide are formed when sodium bromide reacts with concentrated sulfuric acid.
For this reaction:

Give an equation
Give one other observation
State the role of sulfuric acid

Answer 8

2NaBr + 2H2SO4 → Na2SO4 + SO2 + Br2 + 2H2O

Orange fumes are formed
Sulfuric acid acts as an oxidising agent

Question 9

Suggest 2 ways the student can reduce percentage uncertainty in the measurement of the volume of sodium thiosulfate solution, using the same apparatus as the experiment

Answer 9

Use more of the alloy
Use a lower concentration of the thiosulfate solution

Question 10

Copper(I) iodide is a white solid. Explain why copper(I) iodide is white

Answer 10

Full (3)d (sub)shell
No (d-d) transitions possible/ cannot absorb visible/white light

Question 11

A buffer solution has a constant pH even when diluted.
Use a mathematical expression to explain this.

Answer 11

ration [X-] / [HX] remains constant

Question 12

Explain with the use of an equation, why a solution containing [Al(H2O)6]3+ has a pH < 7

Answer 12

[Al(H2O)6]3+ + H2O ⇌ [Al(OH)(H2O)5]2+ + H3O+

Al3+ has a high charge density and is polarising.
It weakens and breaks the O-H bond in water, releasing H+ ions

Question 13

State what is observed when aqueous chlorine is added to sodium bromide solution.
Give an ionic equation for the reaction.

Answer 13

Observation : Solution becomes yellow/orange

Ionic equation : Cl2 + 2Br- → 2Cl- + Br2

Question 14

State what must be done to maintain the EMF of the fuel cell when in use

Answer 14

Keep the concentration of reactants constant

Question 15

Suggest why the pH probe is washed with distilled water between each of the calibration measurements

Answer 15

Different solutions must not contaminate each other

Question 16

State the meaning of the term electrochemical series

Answer 16

List of electrode potentials in numerical order

Question 17

Suggest why an aqueous electrolyte is not used for a lithium cell

Answer 17

Lithium would react with water

Question 18

In lithium cells, lithium cobalt oxide electrodes and lithium electrodes are used.
Give an equation for the reaction that occurs at the positive lithium cobalt oxide electrode and at the negative lithium electrode

Answer 18

Positive electrode : Li+ +CoO2 + e- → LiCoO2

Negative electrode : Li → Li+ + e-

Question 19

State how the relative abundance of 185-Re+ is determined in a TOF mass spectrometer

Answer 19

At the detector, ions gain an electron.
Relative abundance depends on the size of the current

Question 20

Half equation for the reduction of MnO4-

Answer 20

MnO4- + 5e- + 8H+ = Mn2+ + 4H2O

Question 21

Why is an indicator not used in manganate titrations

Answer 21

KMnO4 (example) is self indicating

Question 22

Define the term mean bond enthalpy

Answer 22

Heat energy change/ enthalpy changed needed to break 1 mole of bonds averaged over a ranged of compounds

Question 23

State how a buffer can be made from solution of potassium hydroxide and ethanoic acid

Answer 23

Add excess ethanoic acid to KOH

Question 24

State how a buffer solution made from KOH and CH3COOH resists change in pH when a small amount of acid is added

Answer 24

CH3COO- reacts with H+

Question 25

The pH of a barium hydroxide solution is lower at 50 °C than at 10 °C At 50 °C a 25 cm3 sample of this barium hydroxide solution was neutralised by 22.45 cm3 of hydrochloric acid added from a burette. Deduce the volume of this hydrochloric acid that should be added from a burette to neutralise another 25 cm3 sample of this barium hydroxide solution at 10 °C

Answer 25

22.45cm^3 because the concentration of OH- ions is the same

Question 26

Describe a test to distinguish between aqueous solutions of sulfur dioxide and sulfur trioxide

Answer 26

Use a pH meter to measure the pH of the solutions. Sulfur dioxide solution gives a pH of 2 Sulfur trioxide solution gives pH of 1

Question 27

Give two equations for the formation of the two gaseous sulfur-containing products when concentrated sulfuric acid is added to solid sodium iodide

Answer 27

H2SO4 + 2H+ + 2I- = SO2 + 2H2O + I2 H2SO4 + 8H+ + 8I- = H2S + 4H20 + 4I2

Question 28

Explain why aqueous solution containing [CuCl4]2- ions are yellow

Answer 28

Visible light is absorbed to excite d-electrons to excited state. Only yellow light is transmitted

Question 29

State 3 ways in which a transition metal complex can be changed to alter its colour

Answer 29

Change ligand Change the coordination number Change the oxidation state

Question 30

Suggest why the EMF values of the acidic and alkaline hydrogen-oxygen fuel cells are the same

Answer 30

Same overall reaction

Question 31

Iodine reacts slowly with propanone in the presence of an acid catalyst. The rate of this reaction can be followed by preparing mixtures in which only the initial concentration of propanone is varied. State and explain why different volumes of water are added to each mixture

Answer 31

To make volumes constant for all reactions so the volume of propanone is proportional to concentration

Question 32

Describe how a calibration graph is produced and used to find the concentration of the iron(III) complex.

Answer 32

Measure absorbance for (a range of) known concentrations Plot graph absorbance v concentration Read value of concentration for the measured absorbance from this graph

Question 33

During the experiment the concentration of cisplatin is measured at one-minute intervals. Explain how graphical methods can be used to process the measured results, to confirm that the reaction is first order.

Answer 33

Plot concentration (y-axis) against time (x-axis) and take tangents / (calculate the) gradients (to calculate rates) Plot rate/gradients against conc Straight line through origin / directly proportional confirms first order

Question 34

A calculation of the enthalpy of lattice formation of silver iodide based on a perfect ionic model gives a smaller numerical value than the value calculated before. Explain this difference

Answer 34

Silver iodide has covalent character so the forces/bonds are stronger

Question 35

Predict whether the pH of Mg(OH)2 is larger than, smaller than or the same as the pH of Ca(OH)2 solution

Answer 35

Mg(OH)2 has a lower pH as it is less soluble

Question 36

Describe the test you could carry out in a test tube to distinguish between sodium oxide and phosphorous oxide

Answer 36

React with/add water to both oxides. Measure their pH H3PO4 has pH 1 whereas NaOH has pH 13

Question 37

Explain why colorimetry cannot be used to determine the concentration of solutions containing [CuCl2]- In your answer refer to the electron configuration of the metal ion

Answer 37

Cu+ has a complete/full d-sub shell It is colourless/cannot absorb visible light

Question 38

Describe how a student could distinguish between aqueous solutions of KNO3 and K2SO4, using one simple test tube reaction

Answer 38

Reagent - BaCl2 Observation with KNO3 - No reaction. No visible change Observation with K2SO4 - White precipitate formed

Question 39

Describe how a student could distinguish between aqueous solutions of MgCl2 and AlCl3 using one simple test tube reaction

Answer 39

Reagent - NaOH Observation with MgCl2 - White precipitate formed Observation with AlCl3 - White precipitate which dissolves in excess NaOH

Question 40

Suggest why the oxidation state of chromium is zero in Cr(PF3)6

Answer 40

PF3 is neutral and the complex is zero

Question 41

Deduce the oxidation state of chromium in [Cr(NH3)4Cl2]Cl

Answer 41

+3

Question 42

The Ba(OH)2 is dissolved in a beaker containing 150cm^3 of distilled water. Describe how this solution is used to make 250cm^3 of the 0.100 mol dm^3 Ba(OH)2 solution

Answer 42

Transfer solution to volumetric/graduated flask Transfer washings to measure cylinder and make up to 250cm^3 mark with distilled water Invert many times/shake to mix

Question 43

Suggest one advantage of using methanol, rather than hydrogen, in a fuel cell for use in cars

Answer 43

Methanol can be stored/ transported easily