Chem topic 8.1 -3 Reactivity trends

Module 3

What is another name for group 2 metals?

alkaline earth metals

Why is the electron configuration of group 2 metals all ns² ?

Because all group 2 metals are located on the s block and in the second column in the periodic table.

The s orbital/ highest energy shell is always filled with 2 valency electrons.

In a redox reaction, what kind of agent are group 2 metals and what kind of reaction does it carry out?

Reducing agents, which cause the other species to gain electrons and reduced while group two elements themselves undergo oxidation reaction to lose two electrons for a stable, full outer shell.

Describe the reactivity and the trend of ionisation energies down group 2 metals. (6)

• Reactivity increases down the group

• Due to increasing atomic radius and shielding effect

• This cause an increasing distance between the nucleus and the electron, reducing the nuclear attraction

• Despite the nuclear charge is also increasing (outweighted by atomic radius and shielding effect)

• Reduced nuclear attraction/ less attraction means it would be easier for the outer electrons to lose down the group.

• Easier to lose outer electron down the group means a decreasing trend of ionisation energy (less energy required to remove outer electrons)- more reactive

Explain the overall melting point trend of group 2 metals down the group. (5)

• Melting point of group 2 metals decrease down the group except Mg due to a different crystalline structure

• Due to weaker metallic bonds because of increasing size of the positive ions

• A larger ion meaning lower charge density because the charges inside the ion are spread out more rather than densily packed together

• A lower charge density also means a weaker electrostatic attraction between the metal cations and the sea of delocalised electrons.

• Another reason is larger ionic radius increase distance between the outer electrons and the positive nucleus, weakening the metallic bonds so less energy are needed to break

Compare the melting point of group 1 and 2 in terms of their ionic radius, nuclear charge and the sea of delocalised electrons . (4)

• Group 2 have a higher melting point than group 1

• This is because group 2 has a smaller ionic radius due to a more positive nuclear charge (2+) comapre to group 1 (1+).

• Smaller ionic radius means smaller distance between outer electrons and positive nucleus in group 2 than group 1 so stronger metallic bonding means more energy is needed to overcome the bonds in group 2 metals.

• In group 1 metals, only one electron is donated by each metal ion compare to 2 electrons from each group 2 metal ion. (weaker electrostatic attraction in group 1)

Describe the reaction between group 2 metals and oxygen (give general equation and redox reaction- oxidation number) . (3)

Group 2 metals react with oxygen to form metal oxide which are white solids.

general equation: 2M (s) + O_{2 (g)} → 2MO (s)

Redox reaction: metal is oxidised- 0 → +2, oxygen is reduced- 0 → -2

What are the flame colours of each element from Mg to Ba in group 2 elements when the metals are burnt in air? (4)

Mg- Bright white light

Ca- Brick- red flame

Sr- Red flame

Ba- Pale green flame

Explain why metal ions emit coloured light when they are heated and why there are different colours. (5)

• Heat energy is transferred to the electrons and cause them to jump to a higher energy shell

• The electron is now unstable, it would fall back to the previous shell for stability

• It must release energy to go back to a lower energy shell

• The energy is emitted as coloured light

• Different colours are emitted depending on the sizes of the jumps

Describe the reaction between group 2 metals and (cold) water (give general equation and redox reaction- oxidation number) . (3)

Group 2 metals react with cold water to form alkaline hydroxide and hydrogen gas.

general equation: M _(s) + 2H₂O _(l) → M(OH)_{2 (aq)} + H_2(g)

Redox reaction: metal is oxidised 0 → +2, hydrogen in water is reduced to hydrogen gas +1 → 0

Explain how each element from Mg to Ba react with cold water (reactivity). (4)

Mg- react very slowly, forms a weak, thin layer of Mg(OH)2

Ca- react steadily

Sr- react rapidly

Ba- react vigorously

Describe the reaction between group 2 metals and steam (give general equation and redox reaction- oxidation number) . (3)

Group 2 metals react with steam to form alkane oxide and hydrogen gas.

General equation: M(s) + H₂O (g) → MO (s) + H₂ (g)

Redox reaction: metal is oxidised 0 → +2, hydrogen in water is reduced to hydrogen gas +1 → 0

Why does alkane metals react with steam faster than with cold water.

Because a higher temperature increase kenetic energy in particles.

Describe the reaction between group 2 metals and dilute acid (give general equation and redox reaction- oxidation number) . (3)

Group 2 metals react with dilute acid to form metal salt and hydrogen

General equation: M(s) + 2HCl (aq) → MCl₂ (aq) + H₂ (g)

Redox reaction: metal is oxidised 0 → +2, hydrogen in acid is reduced to hydrogen gas +1 → 0

Does Group 2 metals oxidise or reduce when react?

Oxidise, from 0 to +2 (lose two electrons)

Describe the bonding between group 2 oxides and compare it to the exception of Beryllium (Be). (3)

• All group 2 oxides have ionic bonding except Be that have covalent bonding

• Group 2 oxides (from Mg to Ba) have an increasing atomic radius down the group, so it is easier to lose electrons, making strong ionic bonds

• But Be at the top of group two is small, so it has high charge density and high electronegativity, being able to attract electrons more easily rather than losing electrons.

Describe the reaction of group 2 metal oxides with water. (give general equation) (2)

• Group 2 metal oxides reacts with water to form metal hydroxides

• general equation: MO(s) + H₂O (l) → M(OH)₂ (s)

Write the ionic equation of the G2 metal oxide dissolves/ react in water (slightly soluble) and when the solution saturated but with some ions left. (2)

• Ionic equation: MO (s) + H₂O (l) → M ²⁺ (aq) + 2OH^- (aq)

• Saturated solution with further ions: M ²⁺ (aq) + 2OH^- (aq) → M(OH)₂ (s)

Explain why BeO is insoluble in water.

BeO is insoluble in water because of the strong lattice structure from covalent bonds (high lattice enthalpy) that are unable to be seperated by the water molecules.

Describe the solubility trend down group 2 hydroxides in regards of their ionic bonds. (4)

• Solubility increases down the group

• Due to weaker ionic bonds down the group because metal cations get larger so the charge density decreases

• There is a lower attraction between the metal cations(2+) and the OH^- ions

• The lower attraction between the ions, the easier it can be seperated by water molecules

Explain why pH increases down G2 hydroxides. (2)

• Solubility also increases, more OH- ions will be released to the solution

• So it would be more alkaline

• The more alkaline a solution is, the higher the pH

Describe an experiment to test for the solubility and pH for the G2 hydroxides and explain the observation. (3)

1. Add a spactula of group 2 metal oxides in each sperate test tubes and add distilled water

2. Shake the mixture and there should be a cloudy solution with some white precipitate at the bottom of the test tube.

3. Measure the pH of each solution (should increase along with its solubility)

What is the group 2 compound that is used in agriculture? (give the balanced equation) (2)

Ca(OH)₂- calcium hydroxide is used to balance the pH of the (acidic) soil.

Ca(OH)_2(s) + 2HCl (aq) → CaCl₂ (aq) + 2H₂O (l)

What is the group 2 compounds that is used in medicine? (give the balanced equation)(4)

Mg(OH)_2, magnesium hydroxide(milk of magnesia) and CaCO₃, calcium carbonate are used as antacids to neutralise excess stomach acid

1. Mg(OH)_{2 (s)} + 2HCl (aq) → MgCl(aq) + 2H₂O (l)

2. CaCO_3(s)+ 2HCl (aq) → CaCl₂(aq) + H₂O (l) + CO_2(g)

What is the ionic equation for neutralisation reaction between milk of magnesia and acid.

H⁺ (aq) + OH^- (aq) → H₂O (l)

Why are antacids safe to swallow?

The antacids are insoluble so they cannot be absorbed into the blood stream.

Which group are halogens in and what kind of form do they naturally occur? (3)

Group 17 or group 7 non metals

They occur as diatomic molecules in pure elemental state

On earth, they stay as stable halide ions that dissolves in seawater or combine with sodium or potassium as solid deposits such as NaCl

What are the top 4 halogens state and appearance? (4)

1. Flourine (F₂)- Pale yellow gas

2. Chlorine (Cl₂)- Pale green gas

3. Bromine (Br₂)- Red- Brown liquid

4. Iodine (I₂)- Shiny grey, black solid

Explain the boiling point trend of halogens down the group. (6)

• Boiling point increases down the group

• This is because the molecules get larger down the group which also means there are more electrons

• The more electrons there are, the stronger the London forces

• Because there would be more induced moments and the induced dipoles are stronger

• The stronger the London force, the more energy require to break the intermolecular forces

• So it would be harder to change the state down the group

What kind of agent are halogens and how is it relavent to its electron configuration? (3)

Halogens are oxidasing agents which means they cause the other species to oxidise or lose electrons while they gain electrons themselves.

This is because they have 7 valency electrons so they need to gain one to make a full stable outershell.

Halogens have the electron configuration of ns² np⁵ so they need to have one more electron to get the same electron configuration as a stable noble gas. When they gain an electron, they become halide ions.

Explain the reactivity of halogens down the group along with the trend of electronegativity and oxidising power. (4)

• Reactivity decreases down the group of halogens

• This is because atomic radius and shielding effect increases down the group

• Despite the nuclear charge is increasing, the overall nuclear attraction gets weaker down the group.

• The smaller the attraction from the nucleus, the harder it is to attract or gain electrons which means it has a decreasing electronegativity and oxidising power.

What kind or reaction is carried out to show the reactivity of halogens down the group?

Displacement reaction.

Suggests an example of how this type of reaction is carried and what happens. (3)

During the displacement reaction between halogen and halide ions, for example: when clorine water is added to KBr aqueous solution, the clorine water is reactive enough to displace the Br^-, so Bromine (Br2) is formed in this reaction.

If the halogen is not reactive enough the dispalce the halide ions, then no reaction would occur.

Eg. Iodine water and KBr, iodine is not reactive enough so no displacement reaction occurs.

Give the full equations for reactions that occurs between halogens and halides. (3)

1. Cl₂ and Br^- : Cl₂ (aq) + 2KBr (aq) → 2KCl (aq) + Br₂ (aq)

2. Cl₂ and I^- : Cl₂ (aq) + 2KI (aq) → 2KCl (aq) + I₂ (aq)

3. Br₂ and I^- : Br₂ (aq) + 2KI (aq) → 2KBr (aq) + I_{2 (aq)}

Explain which agent halogens are, using oxidation numbers during displacement reactions.(2)

Halogens are oxidising agents as they need to gain electrons for a full outer shell.

0 → -1 (eg. Cl₂ → 2Cl^- )

Explain which agent Halides are in terms of their reactivity compare to halogens. Give examples through oxidation numbers during displacement reactions. (3)

Halides are reducing agents

because they are less reactive than halogens so are oxidised to lose electrons.

-1 → 0 (eg. Br^- → Br₂ )

Describe what colour is formed for the reactions occured during the displacement reactions and when cyclohexane is also added. (3)

1. Cl₂ and Br^- : Orange, when cyclohexane is added, orange layer

2. Cl₂ and I^- : Brown, when cyclohexane is added, purple layer

3. Br₂ and I^- : Brown, when cyclohexane is added, purple layer

Explain the fucntion of the cyclohexane? (3)

Cyclohexane are used to distinguish between similar colours of Bromide ions and Iodide ions before adding it.

Br₂ or I₂ that are displaced during the reaction dissolves into the cyclohexane, forming a coloured layer on top of the rest of the solution as it is less dense

The Br₂ or I₂ formed are non polar so they would rather to dissolve in the cyclohexane than in water.

They all form different colours when they dissolve in cyclohexane- (Br₂ = orange layer, I₂ = purple layer)

Define disproportionation and give two examples of chlorine reactions. (3)

Disproportionation is when the same element is both oxidised and reduced.

The two example: Chlorine react with water and chlorine react with sodium hydroxide

Give the word and chemical equation when chlorine reacts with water and show which chlorine in the products are oxidised or reduced through oxidation number. (4)

Chlorine + water → chloric (I) acid + Hydrochloric acid

Cl₂ (g) + H₂O (l) → HClO (aq) + HCl (aq)

Cl₂ → HClO = 0 → +1 oxidation

Cl₂ → HCl = 0 → -1 reduction

Why is mixing chlorine into water useful? (1)

It can be used to disinfect drinking water, killing harmful bacteria by the chloric (I) acid

Explain how you could test the formation of HClO, chloric (I) acid and HCl in the reaction? (3)

It can be test out by damp blue litmus paper, if it is present, it turns paper red (due to the acidic solutions) then white ( due to chloric(I) acid as it is a bleach).

Give the word and chemical equation when chlorine reacts with sodium hydroxide and show which clorine in the products are oxidised or reduced through oxidation number. (4)

Chlorine + sodium hydroxide → sodium chlorate(I) + sodium chloride + water

Cl_{2 (aq)} + 2NaOH (aq) → NaClO (aq) + NaCl (aq) + H₂O (l)

Cl₂ → NaClO = 0 → +1 oxidation

Cl₂ → NaCl = 0 → -1 reduction

Why is mixing chlorine with sodium hydroxide useful? (1)

It makes bleach, killing bacteria by chlorate(I) ions from sodium chlorate, NaClO

What are the benefits (2) and risks (2) of chlorine use?

Benefits:

1. Kills bacteria to prevent diseases

2. Prevents algae from growing

Risks:

1. Chlorine is a respiratory irritant

2. Clorine can react with orgainc hydrocarbons like methane to make chlorinated hydrocarbons that could possibly cause cancer

What are the two alternatives to chlorine and explain their disadvantages? (4)

1. Ozone (O3)- strong oxidising agent that kills microorganisms - But have a short half lide and is expensive

2. UV light- kills microorganisms by samageing DNA - But ineffective in cloudy water (preventing from killing microorganisms)

Describe how and why halide test is carried out. (6)

Halide ion solutions are all colourless so it is hard to distinguish which is which.

So silver nitrate (AgNO3) will be added to each solution to form different colour precipitates.

AgCl forms a white precipitate, AgBr forms a cream precipitate, AgI forms a yellow precipitate.

Then dilute ammonia will be added until there is no change- Chloride ions dissolves

Then conc ammonia will be added- Bromide ions dissolves and chloride ions also do.

Iodine does is insoluble.

Give the full word and chemical and ionic equation when potassium chloride react with silver nitrate. (3)

Silver nitrate + potassium chloride → silver chloride + potasssium nitrate

AgNO3 (aq)+ KCl (aq) → AgCl (s) + KNO3 (aq)

Ag⁺ (aq) + X^- (aq) → AgX (s)

Describe the trend of reducing power of halides as a reducing agent down the group. (5)

• The reducing power of halide ions increases down the group

• Halide ions further down the group have a greater ability to donate electrons

• Due to the increasing ionic radius and shielding effect

• There is a weaker electrostatic attraction/ nuclear attraction

• So electrons are easier to be removed

Define qualitative analysis.

A technique that is used to identify unknow ions, elements or functional groups based on non- numerical observations like colour changes, precipitate formation or effervescence through test- tube reactions.

What are the tests for anions? (3)

1. Carbonate

2. Sulfate

3. Halide

Describe the process of carbonate test. (3)

1. Dilute nitric acid will be added sodium carbonate

2. Bubbles should be produced

3. To prove if CO2 gas is produced, the gas will be delivered to another test tube with lime water which then turns milky if CO2 is present.

Give the balanced chemical equation of the carbonate test reaction.

NaCO₃ (aq) + 2HNO₃ (aq) → CO₂ (g) + 2NaNO₃ (aq) + H₂O (l)

Give the balanced chemical equation of between lime water and carbon dioxide

Ca(OH)₂ (aq) + CO₂ (aq) → CaCO₃ (s) + H₂O (l)

Describe what reaction you expect to see in a sulfate test (include word, balanced and ionic equations) (4)

Barium Chloride or Barium nitrate can be added to sodium sulfate to make white precipitate as barium sulfate that is made is very insoluble.

Sodium sulfate + Barium nitrate → Barium sulfate + Sodium nitrate

Na₂SO₄ (aq) + Ba(NO₃)₂ (aq) → BaSO₄ (s) + 2NaNO₃ (aq)

Ba²⁺ (aq) + SO4 ^2- (aq) → BaSO4 (s)

What is involved in the halide test and describe what do you expect to see. (2)

Silver nitrate reacts with halide ions to form different colour precipitates.

Then dilute and conc ammonia will be used to test the solubility of the halide ions.

What is the order of the anion tests to be carried out and why does it have to be in this sequence? (3)

1. Carbonate test will always be carried out first as neither do sulfate nor halide ions produce bubbles when react with dilute acid.

2. Sulfate test will always be carried out after the carbonate test as BaCO₃ is also a white precipitate which gives us the false positive result. It would be hard to distinguish if there are sulfate ions present in the unknown solution.

3. Halide test will always be carried out last because when Ag+ ions are introduced into the solution, Ag₂CO₃ and Ag₂SO₄ are both insoluble so are both precipitates. This would give the false positive result again so it is important to carry halide test after all of the elimination.

How would you carry out the test with a mixture of ions (in one test tube) ? (3)

1. Carry carbonate test first with adding dilute nitric acid until bubbling stops which means that all of the carbonate ions have been removed

2. Carry the sulfate test after that, adding excess barium nitrate. Remove any barium sulfate precipitate formed using filter paper.

3. Carry the halide test last and add the silver nitrate. Add NH₃ to confirm that halide ions presents.

Describe the test for cations (NH4⁺) and give its word, balanced and ionic equations. (7)

1. NaOH, sodium hydroxide would be added to a mixture containing ammonium ions (eg. NH₄NO_3, ammonium nitrate)

2. Ammonia gas would be produced but dissolves in water

3. The solution needs to be heated up in a warm water bath so ammonia gas can be released

4. To test if the gas produced is ammonia gas, a red damp litmus paper would turn blue (alkaline nature) along with a pungent smell if ammonia gas is produced

Sodium hydroxide + Ammonium nitrate → Ammonia + water + Sodium nitrate

NaOH(aq) + NH₄NO₃(aq) → NH₃ (g) + H₂O (l) + NaNO₃ (aq)

NH4⁺ (aq) + OH^- (aq) → NH₃ (g) + H₂O (l)