Topic 4 Inorganic Chemistry And The Periodic Table

Question 1

Trend in ionisation energies down group 2 and reasons for it:

Answer 1

Down the group, ionisation energies decrease. This is because:
- Effective nuclear charge increases down the group (ionisation energy would increase)
- Increased shielding effect down the group (lessening of the pull of the nucleus to the electrons in energy levels) meaning the outermost electrons experience the increase in nuclear charge less strongly (causing ionisation energy to decrease).
- The energy of the outermost electrons increases as a new quantum shell is added - so less energy needed to be supplied to remove them (causing ionisation energies to decrease).
So the combined effect of the 2nd and 3rd factors outweighs the first so EI decreases down the group.

Question 2

What is the trend in reactivity for group 2 and why?

Answer 2

Reactivity increases down group 2 due to the increase in energy levels means the outer electrons are further from the nucleus, this means there’s a weaker attraction to the nucleus and so is more likely to react with another atom. It’s also due to the decrease in ionisation energy.

Question 3

What is calcium oxide (in terms of commercialism)?

Answer 3

Calcium oxide is called quicklime and is used in farming to counteract soil acidity. It’s made commercially by the thermal decomposition of limestone.

Question 4

What are the products of elements Mg, Ca, Sr, and Ba with water? And how do they react?

Answer 4

• Mg - MgO + H2. Reacts slowly with cold water but more vigorously with steam.

Ca - Ca(OH)2 + H2. Reacts with cold water. This is faster than the magnesium reaction.

Sr - Sr(OH)2 + H2. - Both react vigorously with cold water
Ba - Ba(OH)2 + H2.

Question 5

What are the products of elements Mg, Ca, Sr, and Ba with Cl? And how do they react?

Answer 5

• Mg - MgCl2. Reacts with chlorine gas when heated to form MgCl2.
• Ca - CaCl2 - All react more readily with chlorine, forming their respective chlorides at lower temperatures compared to magnesium.
• Sr - SrCl2
• Ba - BaCl2

Question 6

What are the products of elements Mg, Ca, Sr, and Ba with O? And how do they react?

Answer 6

Mg - 2MgO. Burns with a bright white flame to form MgO

Ca - 2CaO. Reacts more readily than magnesium. Is less intense but still vigorous.

Sr and Ba. 2SrO and 2BaO. Reacts even more readily with increasing vigour as you go down the group.

Question 7

Describe the general trend of reactions of group 2 oxides with water with examples

Answer 7

The oxides of Group 2 elements are basic and react with water to form alkaline hydroxides. The reactivity with water increases down the group. The solubility increases down the group too. For example, MgO with H20 forms Mg(OH)2. This pattern continues with the other group 2 elements.

Question 8

General reactions of group 2 oxides with dilute acid with example (using HCl as the dilute acid)

Answer 8

Group 2 oxides react with dilute acids to form a salt and water. For example, MgO + 2HCl = MgCl2 + H20. The group 2 elements follow a pattern forming their respective chlorides and water.

Question 9

General reaction of group 2 hydroxides with dilute acid (using HCl as the example)

Answer 9

Group 2 hydroxides reaction with dilute acids to produce a salt and water. The reaction proceeds with increasing vigour as you move down the group due to the higher solubility and basicity of the hydroxides.

Question 10

Describe the solubility trend and its reason of group2 hydroxides.

Answer 10

• The solubility of hydroxides increases as you go down the group.
• This is due to the increasing lattice energy and the decreasing hydration energy as the ionic size of the cation increases down the group. The decreasing lattice energy (energy needed to separate the ions) more than compensates for the decrease in hydration energy, making the hydroxides more soluble.

Question 11

Describe the solubility trend and reason of group 2 sulphates

Answer 11

• The solubility of group 2 sulphates decreases as you go down the group.
• This is because the lattice energy of sulphates remains relatively high, but the hydration energy decreases significantly as the ionic size increases down the group. Resulting in a net decrease in solubility.

Question 12

When does lattice energy and hydration energy increase?

Answer 12

• Both increase as the charge increases and size of ions decrease

Question 13

Describe the trend and reason for the thermal stability of group 1 nitrates (example NaNO3 or KNO3)

Answer 13

• Thermal stability increases down the group
• Because group 1 cations have a +1 charge, and as you move down the group, the ionic radius increases. Larger cations have a lower charge density and this polarise the nitrate anion less. This reduced polarisation means the anion is more stable and requires more energy to decompose.

Question 14

Describe the trend and reason for the thermal stability of group 2 nitrates (example Mg(NO3)2 or Ca(NO3)2))

Answer 14

• Thermal stability increases down the group
• Group 2 cations have a 2+ charge, resulting in a higher charge density compared to group 1. The smaller cations at the top (e.g. Mg2+ and Ca2+) have higher polarising power, which distorts the nitrate anion more and makes it less stable. As you move down the group, the ionic size increases, reducing the polarising effect and so increasing thermal stability.

Question 15

Describe the trend and reason for thermal stability of group 1 carbonates ( Na2CO3, K2CO3)

Answer 15

• Group 1 carbonates are generally very stable
• The cations in group 1 have a +1 charge and a relatively low charge density. This low charge density doesn’t polarise the carbonate ion (CO3 2-) significantly, making the carbonates thermal,y stable.

Question 16

Describe the trend and reason of thermal stability for group 2 carbonates (MgCO3, CaCO3)

Answer 16

• Thermal stability increases down group 2
• Group 2 cations have a +2 charge, giving them a higher charge density than group 1 cations. The smaller cations at the top (Mg2+ and Ca2+) polarise the carbonates ion more, destabilising it and causing easier decomposition upon heating, as you go down the group, the ionic size increases, reducing the polarising effect and so increasing thermal stability.

Question 17

Describe the group 1 flame colours:

Answer 17

• Lithium - Crimson red
• Sodium - Bright yellow
• Potassium - Lilac
• Rubdoum - Reddish-violet
• Caesium - Blue-violet

Question 18

Describe the group 2 flame colours

Answer 18

• Beryllium - No characteristic colour
• Magnesium - No visible colour, bright white when burning
• Calcium - Orange-red
• Strontium - Crimson red
• Barium - Pale green

Question 19

Explain how the flame colours are produced

Answer 19

• When heated, electrons in the metal cation absorb energy and move from a lower energy level (ground state) to a higher energy level (excited state). When the electrons return to their original lower energy level, they emit energy in the form of light.
• The specific wavelength of the light emitted depends on the energy difference between the excited and ground states. Larger energy differences result in light of shorter wavelengths (violet/blue), while smaller energy differences produce light of longer wavelengths (red/yellow).

Question 20

Describe the experimental procedure with observations to observe the decomposition of nitrates of group 1 and 2.

Answer 20

Procedure: Place a small amount of the nitrate sample in a test tube. Heat the test tube gently and slowly use a stronger heat using a Bunsen burner.

Observations: Group 1 nitrates: All produce no brown fumes except for lithium nitrate which decomposes more similarly to group 2 nitrates. Group 2 nitrates: All produce brown fumes.

Question 21

Describe the experimental procedure with observations to observe the decomposition of carbonates of group 1 and 2.

Answer 21

Procedure: Place a small amount of the carbonates sample in a test tube. Heat the sample with a Bunsen burner. Use a delivery tube to pass any gas evolved through limewater. If the limewater turns cloudy, CO2 is present, confirming decomposition.

Observations: Group 1 carbonates: generally don’t decompose except for lithium carbonate. Group 2 carbonates: decompose to produce carbon dioxide and metal oxide (with decomposition being easier for smaller cations so MgCO3 decomposes more readily than BaCO3).

Question 22

Describe the experimental procedure to show flame colours in compounds of group 1 and 2 elements

Answer 22

Procedure: Dip the wire loop in concentrated HCl to clean it and heat it in the Bunsen burner flame until it produces no colour (ensuring for no contaminants). Dip the cleaned wire loop into the sample of the metal salt solution or the dry salt. Place the loop in the blue part of the Bunsen flame.

Question 23

Describe the trend and reasons for melting and boiling points of group 7 elements.

Answer 23

• The m.p and b.p of group 7 elements increases as you go down the group
• Due to:
  Intermolecular forces - As the halogens are non-polar. The primary intermolecular forces between these molecules are London forces which become stronger as the number of electrons in the molecule increases.
  Electron cloud size - As you move down the group, the size of the electron cloud increases, which enhances the polarising ability of the molecules and strengthens the London forces.
  Energy requirement - More energy is needed to overcome these stronger intermolecular forces, resulting in higher m.p’s and b.p’s for the heavier halogens.

Question 24

Describe the trend and reason for physical state at room temperature for group 7 elements

Answer 24

• The states of halogens changes from gas (F and Cl) to liquid (Br) to solid (I and At)
• It’s directly related to the increasing strength of intermolecular forces. F and Cl have weaker London forces, so they exist as gases. Br has intermediate London forces so is a liquid. While I and At have the strongest London forces and so are solids at room temperature.

Question 25

Describe the trend and reason for electro negativity of the group 7 elements

Answer 25

— Electronegativity decreases as you go down the group
— Due to:
- Atomic size - atomic radius increases as you go down the group due to the addition of energy levels. This increases distance between the nucleus and bonding electrons reduces the attraction that the nucleus can exert in shared electrons in a bond.
- Shielding effect - The inner electrons provide a shielding effect that reduces the effective nuclear charge felt by the outer electrons.
- Nuclear charge - Although the number of protons increases down the group, the effect of increases electron shielding and atomic radius outweighs this, leading to a lower ability to attract electrons.

Question 26

Reasons for the trend in reactivity of group 7 elements down the group

Answer 26

Reactivity decreases down group 7 because:
- Atomic radius - atomic radius increases down the group so as it increases, the outermost electrons are farther from the nucleus. This greater distance reduces the nuclear attraction on the incoming electron, making it harder for the atom to attract and gain an electron
- Electron shielding - electron shielding increases down the group so the increased shielding weakens the effective nuclear charge felt by the outermost electrons and any incoming electron. This reduces attraction make it more difficult for the larger halogens (I and At) to attract and gain an electron compared to smaller halogens (F).
- Nuclear charge - should increase the attraction for electrons, the combined effect of atomic radius and shielding outweighs this.

Question 27

Describe the oxidising power of halogens and the general trend

Answer 27

Halogens act as oxidising agents by gaining electrons to form halide ions. Their ability to oxidise other substances decreases as you move down the group. The oxidising power decrease down the group

Question 28

Describe the redox reactions of Cl2, Br2 and I2 with Halide ions and their general reaction

Answer 28

General reaction: A more reactive halogen will oxidise the hoarder ion of a less reactive halogen, displacing it from solution, Chlorine with Bromide and Iodide ions for example: Cl2 + 2Br- = 2Cl- + Br2 (Cl oxidises bromide to bromine). Cl2 + 2I- = 2Cl- + I2 (Cl oxidises iodide to iodine), bromine can oxidise iodide to iodine however iodine can’t oxidise bromide or chloride ions as less reactive halogens can’t displace a more reactive halide ion.

Question 29

Purpose of organic solvents in accordance with halogens

Answer 29

An organic solvent, like hexane or cyclohexane, is used to enhance the visibility of halogen displacement reactions as the colours of halogens are quite similar and there’s variation in colour with concentration. Halogens are more soluble in organic solvents than in water, and they display distinct colours in these solvents, making it easier to identify the presence of a certain halogen.

Question 30

Colour changes of halogens in cyclohexane (organic solvent)

Answer 30

Chlorine - the pale green colour of chlorine doesn’t change much
Bromine - the orange colour of bromine is a bit darker
Iodine - the colour of iodine changes to purple or violet

Question 31

What is disproportionation?

Answer 31

When the halogen is simultaneously reduced and oxidised to form a halide ion and a compound with oxygen.

Question 32

Describe the trend of reducing power of halide down group 7

Answer 32

Reducing power increases

Question 33

Describe the general reaction of halogen oxidation reactions with group 1 metals and the oxidation number changes in the formation of sodium chloride.

Answer 33

• When a group 1 metal reacts with a halogen, the metal is oxidises and the halogen is reduced. Typically forming an ionic compound.
• Sodium’s oxidation number changes from 0 to +1, indicating that sodium is oxidised.
  Chlorine’s oxidation number changes from O to -1, indicating chlorine is reduced.

Question 34

Describe the general reaction of halogen oxidation reactions with group 2 metals and the oxidation number changes in the formation of Magnesium bromide.

Answer 34

• Group 2 metals react with halogens to form ionic halides.
• Magnesium’s oxidation number changes from O to 2+, indicating magnesium is oxidised. While Bromine’s oxidation number changes from O to -1, indicating that bromine is reduced.

Question 35

Describe the disproportionation reaction of chlorine with water

Answer 35

It involves chlorine reacting with water to produce a mixture of two different chlorine compounds: hypochlorite acid (HOCl) and hydrochloric acid (HCl). In this reaction, chlorine is both reduced to chloride ions and oxidised to hypochlorous acid.

Question 36

How to use chlorine in water treatment?

Answer 36

In water treatment, chlorine is added to water where it undergoes disproportionation to form hypochlorous acid (which is the active species responsible for the disinfection process). It then dissociates in water where the hypochlorite ion is the effective oxidising agent that kills bacteria and viruses by damaging their cell walls and proteins, disrupting their metabolic processes, leading to their death.

Question 37

What are the oxidation number changes in water treatment?

Answer 37

Chlorine: it’s elemental form has an oxidation number of O
Hypochlorous acid: Chlorine has an oxidation number of +1
Chloride: has an oxidation number of -1

Question 38

Key points involved in the disproportionation reaction of chlorine with cold, dilute aqueous sodium hydroxide:

Answer 38

• One chlorine atom is reduced to chloride (Cl-) and oxidised to hypochlorite (OCl-)
• Sodium hypochlorite is a bleach which is produced in this reaction and is widely used for cleaning and disinfecting.
• The oxidation number of chlorine changes from 0 in Cl2 to -1 in NaCl (reduction) and +1 in NaOCl (oxidation)

Question 39

Key points involved the disproportionation reaction of chlorine with hot alkali

Answer 39

• The reaction of chlorine with hot concentrated sodium hydroxide is an example of a disproportionation reaction in which sodium chloride and sodium chlorate (I) are produced.
• One atom of chlorine is reduced to form chloride ions (Cl-), the other is oxidised to chlorate ions (ClO3-)
• In sodium chloride (NaCl), chlorine has an oxidation number of -1. In sodium chlorate (NaClO3), chlorine has an oxidation number of +5.
• Sodium chlorate is used in the manufacturing of bleach and herbicides. It’s a powerful oxidiser.

Question 40

Prove that bromine and iodine react similarly to Chlorine

Answer 40

• Bromine with water: it undergoes disproportionation when reacting with water, forming hydrobromic acid and hypobromous acid. Bromine’s initial oxidation number is 0. In HBr, bromine has an oxidation number of -1 (reduction), in HOBr, bromine has an oxidation number of +1 (oxidation).
• Iodine with water: Iodine reacts with water to form hydriodic acid (HI) and hypoiodous acid (HOI). Iodine initially has an oxidation number of 0. In HI, iodine has an oxidation number of -1 (reduction), in HOI, iodine has an oxidation number of +1 (oxidation).

Question 41

Understanding solid group 1 halides in reaction with concentrated sulphuric acid

Answer 41

When group 1 halide (like sodium chloride, sodium bromide, sodium iodide) react with concentrated sulphuric acid, a redox reaction occurs. The halides act as reducing agents and reduce the sulphuric acid, while the hydrogen halides are released as gaseous products. In Sodium chloride with concentrated sulphuric acid: no reduction of sulphuric acid as the chloride ions aren’t strong enough reducing agents. In sodium bromide with concentrated sulphuric acid: the Bromide ion is a stronger reducing agent, so it reduces sulphuric acid to sulphur dioxide. In sodium iodide with concentrated sulphuric acid, Iodide is the strongest reducing agent and reduces sulphuric acid to hydrogen sulphide (H2S).

Question 42

Understanding precipitation reactions of the aqueous anions (Cl-,Br-,I-) with aqueous silver nitrate solution, followed by aqueous ammonia solution.

Answer 42

When halide ions are added to aqueous silver nitrate (AgNO3), precipitation reactions occur, forming silver halides.
- Chloride with silver nitrate, AgCl is a white precipitate and it’s soluble in dilute ammonia solution
- Bromide with silver nitrate, AgBr is a cream-coloured precipitate and is soluble in concentrated ammonia solution.
- Iodide with silver nitrate, AgI is a yellow precipitate and is insoluble in ammonia solution.
The trends in solubility of the silver halides in ammonia reflects the polarising abilities of the halide ion:
- I- is the most polarisable, so AgI is least soluble
- Cl- is the least polarisable, so AgCl is most soluble
- Br- is intermediate, leading AgBr being soluble in concentrated ammonia.

Question 43

Understanding hydrogen halides reacting with ammonia

Answer 43

When a hydrogen halides reacts with ammonia, the ammonia acts as a base, accepting a H+ ion from the hydrogen halide to form ammonium halide (NH4X). X Representing the respective halide in the reaction.

Question 44

Understanding hydrogen halides with water

Answer 44

When hydrogen halides dissolve in water, they ionise to form hydronium ions (H3O+) and halide ions, producing strong acids. All hydrogen halides form strong acids in water, with HI being the strongest due to the weakest bond strength between hydrogen and iodine, making it most likely to fully dissociate.

Question 45

How to make predictions of fluorine and astatine and their compounds in terms of knowledge of trends in halogen chemistry

Answer 45

• Fluorine is a very strong oxidising agent, reacts explosively with metals, forms highly stable compounds and has the highest electronegativity, making it highly reactive with hydrogen and metals.
• Astatine is much larger and less electronegative so will be a weaker oxidising agent, form less stable compounds, have weaker acid characteristics. Its chemistry is less studied due to it being highly radioactive and its rarity. But it behaves similarly to other halogens, just in a less extreme manner.

Question 46

How to identify carbonate ions using an aqueous acid using an example (using sodium carbonate and dilute acid)?

Answer 46

When carbonate ions react with aqueous acid, they undergo an acid-base reaction, producing carbon dioxide, water and a salt.
Ex: Na2CO3 + 2H+ = CO2 + H2O + 2Na+.
When you bubble the resulting gas through limewater, the presence of CO2 will turn the limewater milky due to the formation of CaCO3.

Question 47

How to identify hydrogencarbonate ions using an aqueous acid using an example (using sodium carbonate and dilute acid)?

Answer 47

Hydrogencarbonate ions react with acids similarly to carbonate ions but form CO2 directly without needing the extra proton.
Ex: NaHCO3 + H+ = CO2 + H2O + Na+
The presence of CO2 can be tested by passing the gas through limewater, which turns milky due to the formation of calcium carbonate.

Question 48

Identifying sulphate ions (SO4 2-) using acidified barium chloride solution

Answer 48

When surface ions are present, they react with barium ions to form barium sulphate, a white precipitate.
Ionic equation: Ba2+ (aq) + SO4 2- (aq) = BaSO4 (s)
The solution must be acidified to prevent the formation of insoluble barium carbonate or barium phosphate, which could interfere with the test.

Question 49

Identifying ammonium ions (NH4+) using sodium hydroxide and warming to form ammonia

Answer 49

Ammonium ions react with sodium hydroxide to produce ammonia gas (NH3), which can be detected by its characteristic smell or by turning damp red litmus paper blue in the presence of the alkaline ammonia.
Ionic equation: NH4 (aq) + OH- (aq) = NH3 (g) + H2O (l)