EL 7&8: Ionic bonding, precipitation reactions, ionisation enthalpy

Question 1

What is a valence electron?

Answer 1

An outer shell electron that can participate in the formation of a chemical bond (if the shell is not full).

Question 2

Answer 2

A

Elements so no ionic bonding involved, so skip s-block

Question 3

What is ionic bonding?

Answer 3

Overall attraction in a lattice, consisting of attraction between ions of opposite charge and repulsion between ions of like charge.

Question 4

Name 3 ways of making ionic salts.

Answer 4

• Acid + alkali/base → salt + water
• Acid + carbonate → salt + water + carbon dioxide
• Acid + metal → salt + hydrogen

Question 5

Draw a dot-and-cross diagram of sodium hydroxide.

Answer 5

Question 6

Draw a dot-and-cross diagram of magnesium fluoride.

Answer 6

Question 7

Give an example of a non-polar solvent.

Answer 7

Hexane.

Question 8

Many ionic substances dissolve readily in water. List the ones which are insoluble.

Answer 8

• Silver/lead/barium/calcium sulphates
• Silver/lead halides
• Metal carbonates
• Metal hydroxides, except group 1 metals

Question 9

Name 3 types of ionic substance which are always soluble in water.

Answer 9

• Group 1 compounds
• Ammonium compounds
• Nitrates

Question 10

What is a precipitate?

Answer 10

A suspension of solid particles produced by a chemical reaction in solution.

Question 11

How would you test for Cu²⁺ ions? Name the change observed and the product responsible.

Answer 11

Precipitation reaction

• Add NaOH solution to Cu²⁺ solution
• Blue copper(II) hydroxide, Cu(OH)₂, precipitate forms

Blue-green flame test

Question 12

How would you test for Fe²⁺ ions? Name the change observed and the product responsible.

Answer 12

Precipitation reaction

• Add NaOH solution to Fe²⁺ solution
• Dirty green iron(II) hydroxide, Fe(OH)₂, precipitate forms

​(No flame test)

Question 13

How would you test for Fe³⁺ ions? Name the change observed and the product responsible.

Answer 13

Precipitation reaction

• Add NaOH solution to Fe³⁺ solution
• Orange-brown iron(III) hydroxide, Fe(OH)₃, precipitate formed

Question 14

How would you test for Pb²⁺ ions? Name the change observed and the product responsible.

Answer 14

Precipitation reaction

• Add KI solution to Pb²⁺ solution
• Yellow lead iodide, PbI₂, precipitate forms

(No flame test)

Question 15

How would you test for Cl^- ions? Name the change observed and the product responsible.

Answer 15

Precipitation reaction

• Add AgNO₃ solution to Cl^- solution
• White silver chloride, AgCl, precipitate formed

Question 16

How would you test for Br^- ions? Name the change observed and the product responsible.

Answer 16

Precipitation reaction

• Add AgNO₃ solution to Br^- solution
• Cream silver bromide, AgBr, precipitate forms

Question 17

How would you test for I^- ions? Name the change observed and the product responsible.

Answer 17

Precipitation reaction

• Add AgNO₃ solution to I^- solution
• Yellow silver iodide, AgI, precipitate forms

Question 18

What are the colours of the respective precipitates formed when solutions of chloride, bromide and iodide are reacted with silver nitrate solution?

Answer 18

• Chloride = white
• Bromide = cream
• Iodide = yellow

More saturated colour down group

Question 19

How would you test for SO₄^2- ions? Name the change observed and the product responsible.

Answer 19

Precipitation reaction

• Add BaCl₂ solution to SO₄^2- solution
• White barium sulfate, BaSO₄, precipitate forms

Question 20

Answer 20

C

A: Ba(OH)₂ is insoluble so (s) not (aq). Barium too electropositive for its chloride to be hydrolysed

B: Ba(OH)₂ is formed

D: BaSO₄ is insoluble so (s) not (aq)

Question 21

Answer 21

C

Silver chloride, silver sulfate and silver carbonate are all insoluble

but c uses silver carbonate - what’s special about this one?

Question 22

Answer 22

B

Question 23

Some students place a small piece of calcium in water with a few drops of universal indicator solution. The indicator turns an alkaline colour. After a time, the solution goes cloudy.
The equation for the reaction is Ca + 2H₂O → Ca(OH)₂ + H₂

1. Use the equation to suggest two observations that the students made, other than the indicator turning an alkaline colour.
2. A student says that the cloudiness is caused by insoluble calcium hydroxide being formed slowly. Comment on this statement.
3. The students filter off the solid from the solution, dry it and find it weighs 4.05 g. They calculate that the percentage yield of their reaction is 55.0%. Calculate the mass of calcium they started with.
4. State an assumption you made in order to calculate (3).

Answer 23

1. Calcium dissolves / reacts away, + bubbling / fizzing of hydrogen gas
2. Ca(OH)₂ is sparingly soluble, so dissolves at first, then reaches its maximum solubility and forms a precipitate, causing cloudiness
3. 100 x 4.05 / theoretical = 55 → theoretical = 100 x 4.05 / 55 = 7.364 g → mol Ca(OH)₂ = mol Ca = 0.09937 mol → mass Ca = 0.09937 x 40.1 = 3.98 g
4. No Ca(OH)₂ was in solution

Question 24

Some students attempt to make pure copper sulphate crystals from copper carbonate and sulphuric acid. They plan to use the following methods:

• A: react excess copper carbonate with sulphuric acid. Evaporate. Allow to crystallise.
• B: React copper carbonate with excess sulphuric acid. Filter, wash + dry.
• C: react excess copper carbonate with sulphuric acid. Filter. Evaporate filtrate until dry.

None of these students would end up with pure crystals of copper sulphate. State what each one would end up with.

Answer 24

• A: anhydrous CuSO₄ and CuCO₃
  
  • CuCO₃ was in excess so some left over. Anhydrous since both are insoluble
• ​B: nothing would collect on filter paper
  
  • ​Wouldn’t get anhydrous CuSO₄ bc hasn’t been evaporated yet
• C: anhydrous CuSO₄
  
  • Get completely dry solid rather than crystals bc they evaporate fully, rather than until half volume then crystallising

Question 25

Define “first ionisation enthalpy”.

Answer 25

The energy needed to remove one electron from each of one mole of isolated gaseous atoms of an element, forming one mole of gaseous 1+ ions.

Question 26

Write the equation representing the first ionisation enthalpy of hydrogen.

Answer 26

H(g) → H⁺(g) + e^-

Question 27

Define “second ionisation enthalpy”.

Answer 27

The energy needed to remove one electron from each of one mole of isolated gaseous 1+ ions of an element, forming one mole of gaseous 2+ ions.

Question 28

State and explain the trend in ionisation enthalpies across periods.

Answer 28

Ionisation enthalpy increases across periods:

• All outer electrons in a period are in same shell
• Proton number increases so greater electrostatic attraction between outer electrons + nucleus
• So more energy required to remove an electron

Question 29

State and explain the trend in ionisation enthalpies down groups.

Answer 29

Ionisation energies decrease down groups:

• More electron shells
• Distance between nucleus + outer shell increases
• Weaker electrostatic attraction
• Easier to remove an electron

Question 30

How does observed data of 1st ionisation enthalpies provide evidence for:

• electron shells?
• electron sub-shells?

Answer 30

Electron shells

• Ionisation enthalpy decreases down groups
• Supports idea that inner shells shield outer ones, causing less attraction to nucleus

Electron sub-shells

• Ionisation enthalpy increases across periods, with variations
• So energy needed to remove electrons depends on their location
• Supports idea of sub-shells with varied electron energies

Question 31

Why is there a decrease in 1st ionisation enthalpy between nitrogen and oxygen, despite the general trend of an increase across periods?

Answer 31

• N has electronic configuration 1s²2s²2p³ and O 1s²2s²2p⁴
• Extra repulsion from paired electron subshell in oxygen
• Less energy needed to remove one of these electrons, despite increased nuclear charge