CHAPTER 8: REACTIVITY TRENDS

Question 1

How was the periodic table developed?

Answer 1

• Mendeleev arranged elements in order of increasing atomic mass.
• There was no knowledge of subatomic particles.
• He lined up elements in groups with similar properties.
• If group properties didn’t fit, he swapped elements around & left gaps.
• This assumed elements were yet to be discovered.
• Henry Moseley arranged elements in order of increasing proton number, rather than mass.

Question 2

How is the periodic table arranged?

Answer 2

Atomic Number:
• Arranged in order of increasing atomic number.
• Each successive element has atoms with 1 more proton.

Periods:
• Elements are arranged in horizontal rows: Periods.
• Period number = number of energy levels for atom.
• Periodicity is repeating trends in properties across each period.

Groups:
• Elements are arranged in vertical columns: Groups.
• Atoms in the same group have the same number of valence electrons.
• Elements in the same group have similar properties.

Question 3

What is ionisation energy?

Answer 3

• Energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms & form 1 mole of gaseous 1+ ions.
• X (g) → X⁺ (g) + e⁻.
• Elements have successive ionisation energies.
• There are as many ionisation energies as electrons in the atom (e.g. helium has 2 ionisation energies).
• Each successive ionisation energy is greater.

Question 4

What are factors affecting ionisation energy?

Answer 4

1) Atomic Radius:
• Force of attraction decreases as distance increases.
• Greater distance between nucleus & outer electrons means less nuclear attraction.
• Greater atomic radius = Lower Ionisation Energy.

2) Shielding:
• Electrons are negatively charged.
• Inner-shell electrons repel outer-shell electrons.
• This repulsion is called shielding.
• Reduces attraction between nucleus & outer electrons.
• Greater Shielding = Lower Ionisation Energy.

3) Nuclear Charge:
• More protons in nucleus increase attraction between nucleus & outer electrons.
• Greater nuclear charge = Higher Ionisation Energy.

Question 5

How can you make predictions from successive ionisation energies?

Answer 5

• Successive ionisation energies allow predictions to be made about:

1) Number of electrons in the outer shell.
2) Group of the element.
3) The identity of an element.

Question 6

What are trends in electronic configuration?

Answer 6

• Elements are arranged into S-, P- & D-Blocks.
• S-Block elements have an outer shell electron configuration of s¹ or s².
• P-Block elements have an outer shell configuration of s²p¹ to s²p⁶.
• D-Block elements have electronic configurations in which D-Subshells are filled.

Question 7

What are the general trends in first ionisation energies?

Answer 7

Across Periods 2 & 3:
• Increase in Atomic Number.
• Shielding remains the same.
• Decreasing Atomic Radius.
• Increase in effect of nuclear charge.
• Increase in first Ionisation Energy.
• Increase in Electronegativity.

Down a Group:
• Increase in Atomic Number.
• Increase in Shielding.
• Increasing Atomic Radius.
• Decrease in effect of nuclear charge.
• Decrease in first Ionisation Energy.
• Decrease in Electronegativity.

Bohr Model of the Atom:
• These trends provide evidence that energy levels exist.
• Decrease in I.E down groups supports Bohr model.

Question 8

What is metallic bonding?

Answer 8

Metallic Bonding:
• Strong electrostatic attraction between cations & delocalised electrons.

Giant Metallic Lattice Structure:
• Each atom donates its negative valence electron.
• This forms a delocalised, shared pool of electrons.
• Cations are fixed in position.
• This maintains the structure & shape of the metal.
• Delocalised electrons are mobile charge carriers.

Question 9

What is the structure of lattices of silicon?

Answer 9

• Each Si atom forms 4 covalent bonds to other Si’s.
• This gives a tetrahedral shape (109.5°).
• This forms a giant covalent 3D lattice of solid Silicon.
• It’s very hard due to strong Si-Si covalent bonds.
• It has a high melting/boiling point.
• Si-Si covalent bonds require lots of energy to break.
• It can’t conduct electricity.
• All electrons are involved in bonding.
• It’s insoluble, due to strong Si-Si bonds.

Question 10

What is the structure of diamond?

Answer 10

• Each C atom forms 4 covalent bonds to other C atoms.
• This gives a tetrahedral shape (109.5°).
• This forms a giant covalent 3D lattice of solid carbon.
• It’s very hard due to strong C-C covalent bonds.
• It has a high melting/boiling point.
• C-C covalent bonds require lots of energy to break.
• It can’t conduct electricity.
• All electrons are involved in bonding.
• It’s insoluble, due to strong C-C bonds.

Question 11

What is the structure of graphite?

Answer 11

• Each C atom forms 3 covalent bonds to other C atoms.
• The 4th electron is delocalised.
• This gives a trigonal planar shape (120°).
• This forms a 2D hexagonal structure.
• There are weak London forces between layers.
• This makes it soft & layers can slide over each other.
• It has a high melting/boiling point.
• This is due to strong C-C covalent bonds, which require lots of energy to break.
• It can conduct electricity.
• Delocalised electrons are charge carriers.
• It’s insoluble, due to strong C-C bonds.

Question 12

What is the structure of graphene?

Answer 12

Properties:
• It’s one layer of graphite, one atom thick.
• Each C atom forms 3 covalent bonds to other C atoms.
• The 4th electron is delocalised.
• This gives a trigonal planar shape (120°).
• This forms a 2D hexagonal structure.
• It has a high melting/boiling point.
• This is due to strong C-C covalent bonds, which require lots of energy to break.
• A single layer is transparent & very light.
• It’s the best, known, electrical conductor.
• Delocalised electrons are charge carriers.
• Without layers, they move faster above/below a sheet.
• It’s insoluble, due to strong C-C bonds.

Applications:
• Due to high strength, low mass & great electrical conductivity, it has potential applications in high-speed electronics & aircraft technology.
• Its flexibility & transparency make it a potentially useful material for touchscreens on electronic devices.

Question 13

What are the properties of giant metallic & giant covalent lattices?

Answer 13

Giant Metallic Lattices:

Melting/Boiling Points:
• Depends on strength of metallic bonds.
• Most metals have high melting/boiling points.
• High temperatures provide large amounts of energy to overcome strong electrostatic attraction between cations & delocalised electrons.

Solublity:
• Interactions lead to reactions rather than dissolving.

Electrical Conduction:
• Electricity can conduct in giant metallic lattices.
• Delocalised electrons (mobile charge carriers).

Giant Covalent Lattices:

Melting/Boiling Points:
• High melting/boiling points.
• There are strong covalent bonds between atoms.
• High temperatures provide large amounts of energy to overcome strong electrostatic attraction between shared pairs of electrons & nuclei of bonded atoms.

Solubility:
• Strong covalent bonds can’t be overcome/dissolved.

Electrical Conduction:
• Electricity can’t conduct in giant covalent lattices.
• No delocalised electrons/mobile charge carriers.

Question 14

What happens in redox reactions of Group 2 elements?

Answer 14

• Atoms have 2 outer shell electrons in S sub-shell.
• Metal atoms are oxidised in redox reactions.
• They lose 2 electrons to form a 2+ ion.
• Another species gains these electrons & is reduced.
• The group 2 element is called a reducing agent.
• This is because it has reduced another species.

Question 15

How do Group 2 elements react with oxygen, water & dilute acids in redox reactions?

Answer 15

1) Redox Reactions with Oxygen:
• All combust to form a metal oxide.
• 2X (s) + O₂ (g) → 2XO (s).
• Reactions produce a bright white flame.

2) Redox Reactions with Water:
• All redox reactions form bases.
• X (s) + 2H₂O (l) → X(OH)₂ + H₂ (g).
• Solubility of X(OH)₂ increases down group 2.
• Observations for precipitates are different.
• Mg(OH)₂ & Ca(OH)₂ are insoluble = white precipitate.
• Sr(OH)₂ & Ba(OH)₂ dissolve = no precipitate.
• Mg (s) + H₂O (g) → MgO (s) + H₂ (g).
• MgO is an insoluble base.

3) Redox Reactions with Dilute Acids:
• All react to form salt & hydrogen gas.
• X (s) + 2HCL (aq) → XCl₂ (aq) + H₂ (g).
• X (s) + 2HNO₃ (aq) → X(NO₃)₂ (aq) + H₂ (g).
• X (s) + 2H₂SO₄ (aq) → XSO₄ (s/aq) + H₂ (g) (sulfates decrease in solubility down the group).
• X (s) + 2CH₃COOH (aq) → X(CH₃COO)₂ (aq) + H₂ (g).

Question 16

Why does reactivity increase down Group 2?

Answer 16

• Atomic Radius increases down a group.
• Full energy levels are added between each element.
• All outer electrons are in s-orbitals.
• Shielding increases.
• Effect of nuclear charge decreases.
• Weaker attraction to outer electrons.
• Less energy needed to remove valence electrons.
• 1st & 2nd I.E’s decrease.
• X (g) → X⁺ (g) + e⁻ = 1st I.E.
• X⁺ (g) → X²⁺ (g) + e⁻ = 2nd I.E.

Question 17

How does water react with Group 2 oxides?

Answer 17

• Group 2 elements release OH⁻ ions.
• Alkaline solutions of metal hydroxide are formed.
• X (s) + H₂O (l) → X(OH)₂.
• Mg (s) + H₂O (l) → Mg(OH)₂ (pH~ 10).
• Ca (s) + H₂O (l) → Ca(OH)₂ (pH ~ 11).
• Sr (s) + H₂O (l) → Sr(OH)₂ (pH ~ 12).
• Ba (s) + H₂O (l) → Ba(OH)₂ (pH ~ 13).
• Alkalinity increases down the group.
• This is because more OH⁻ ions are released.

Question 18

What are the uses of Group 2 compounds as bases?

Answer 18

Agriculture:
• Ca(OH)₂ is added to fields as lime.
• This increases the pH of acidic soils.
• Ca(OH)₂ neutralises acid in soil, forming neutral water.
• Ca(OH)₂ (s) + 2H⁺ (aq) → Ca²⁺ (aq) + 2H₂O (l).

Medicine:
• Used as antacids to treat acid indigestion.
• Indigestion tablets use Mg(OH)₂ & CaCO₃.
• Stomach acid is mainly HCl.
• Mg(OH)₂ & CaCO₃ neutralise HCl.
• Mg(OH)₂ (s) + 2HCl (aq) → MgCl₂ (aq) + 2H₂O (l).
• CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + H₂O (l) + CO₂ (g).

Question 19

What are the properties of halogens?

Answer 19

1) Melting & Boiling Points:
• Halogens exist as non-polar diatomic molecules, X₂.
• Weak I.D forces exist between the molecules.
• They’re overcome by high energy provided by heating.
• I.D. forces are stronger down a group as Mr increases.
• This is because more electrons are involved in the I.D.
• Melting/boiling points increase down the group.
• At RTP: F₂ (g), Cl₂ (g), Br₂ (l), I₂ (s).

2) Electronegativity:
• Electronegativity decreases down the group.
• This is due to increase in energy levels & shielding.
• Fluorine is the most electronegative element.
• It has the greatest effect of charge on nucleus.
• This is due to low shielding.
• Increases ability to attract bonding pair of electrons.

What happens in redox reactions of Halogens?
• Each halogen has 7 valence

Question 20

What happens in redox reactions of halogens?

Answer 20

• Each halogen has 7 valence electrons.
• 2 in the s sub-shell & 5 in the p sub-shell: s²p⁵.
• Gain 1 electron in redox reactions to form 1- ions.
• Each Halogen atom is reduced.
• Another species loses electron to halogens.
• Halogens are called oxidising agents.
• This is because it has oxidised another species.

Question 21

What is the trend in reactivity of Cl₂, Br₂ & I₂?

Answer 21

• Reactivity of halogens decreases down the group.
• This is shown in displacements reactions of halogens & halide ions.
• Solution of each halogen is added to aqueous solutions of other halides.
• If halogen added is more reactive than halide:
1) Halogen displaces halide from solution in reaction.
2) Solution changes colour.
• Organic non-polar solvent (cyclohexane) is added.
• Makes colour change more distinct (Iodine = violet).
• Cl₂ (aq) + 2Br⁻ (aq) → 2Cl⁻ (aq) + Br₂ (aq) (Orange).
• Cl₂ (aq) + 2I⁻ (aq) → 2Cl⁻ (aq) + I₂ (aq) (Violet).
• Br₂ (aq) + 2I⁻ (aq) → 2Br⁻ (aq) + I₂ (aq). (Violet).

Question 22

What is disproportionation?

Answer 22

• Oxidation & Reduction of the same element.
• Reaction of Cl₂ with H₂O & NaOH are 2 examples.
• Cl₂ (aq) + H₂O (l) → HClO (aq) + HCl (aq).
• Oxidation number of Cl₂ goes from 0 → +1 in HClO.
• Oxidation number of Cl₂ goes from 0 → -1 in HCl.
• Chlorine is oxidised in HClO & reduced in HCl.
• Reaction of Cl₂ with water, is used in water treatment.
• Bacteria are killed by chloric(I) acid & ClO⁻ ions.
• Cl₂(aq) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H₂O(l).
• Oxidation number of Cl₂ goes from 0 → +1 in NaClO.
• Oxidation number of Cl₂ goes from 0 → -1 in NaCl.
• Chlorine is oxidised in NaClO & reduced in NaCl.
• Reaction of Cl₂ with cold, dilute NaOH (aq), is used to form bleach.

Question 23

What are benefits & risks of chlorinating Water?

Answer 23

Benefits & Risks:
• Chlorine ensures bacteria in water are killed.
• But, it’s a respiratory irritant in small concentrations.
• Large concentrations can be fatal.
• Can react with organic hydrocarbons.
• Chlorinated hydrocarbons are suspected carcinogens.
• It’s an ethical dilema to not give people choice in whether they want to chlorinate water.
• Overall risk to health of not adding chlorine to water supply is greater than risk of chlorinated hydrocarbons.
• Quality of drinking water would be compromised & diseases such as typhoid and cholera might break out.

Alternatives:
• Ozone (O₃) is a strong oxidising agent.
• This means it’s effective at killing microorganisms.
• But, it’s expensive to produce & has a short half-life.
• This means water treatment isn’t permanent.
• UV light kills microorganisms by damaging their DNA.
• But, it’s ineffective in cloudy water.
• So, it won’t stop water contamination in the long-term.

Question 24

How do halide ions react with aqueous silver ions?

Answer 24

• Ag⁺ (aq) + X⁻ (aq) → AgX (s).
• Precipitates of silver halides are formed.
• This forms the basis for the Halide Test:
1) Add aqueous AgNO₃ to aqueous solution of halide.
2) The silver halide precipitates are different colours.
3) AgCl is white, AgBr is cream, AgI is yellow.
4) Aqueous ammonia is added to test solubility.
5) This is useful because the colours aren’t very distinct.
6) AgCl is soluble in dilute NH₃.
7) AgBr is soluble in concentrated NH₃.
8) AgI is insoluble in concentrated NH₃.

Question 25

What is the correct sequence of tests to analyse an unknown inorganic compound?

Answer 25

• For anions, the correct order for tests is:

1) Carbonate: CO₃²⁻
2) Sulfate: SO₄²⁻
3) Halides: Cl⁻, Br⁻ and I⁻.

Question 26

Why is there a correct sequence of tests to analyse an unknown inorganic compound?

Answer 26

1) Carbonate Test:
• Add dilute acid & look for effervescence from CO₂ (g).
• Sulfate/halide ions don’t give bubbles with dilute acid.
• Carbonate test can be carried out without possibility of an incorrect conclusion.
• No bubbles = you can move onto next test.

2) Sulfate Test:
• Add solution of Ba²⁺(aq) & look for white precipitate.
• Precipitate is of BaSO₄, but BaCO₃ is also white & insolube in water.
• Sulfate test on carbonates also gives white precipitate.
• This is why it’s important to carry out carbonate test first & only do the sulfate test if no carbonate is present.

3) Halide Test:
• Add solution of Ag⁺(aq), through AgNO₃(aq).
• Look for a precipitate.
• Ag₂CO₃ & Ag₂SO₄ are both insoluble in water.
• This means they both form as precipitates in this test.
• This is why it’s important to do the halide test last, after doing carbonate & sulfate tests to rule them out.

Question 27

What is the correct sequence of tests to analyse an unknown mixture of ions?

Answer 27

1) Add Dilute Acid:
• If you see bubbles, add dilute HNO₃ until bubbling stops & all carbonate ions are removed.
• If you intend to test for SO₄²⁻/halide ions, use diulte HNO₃, as H₂SO₄ contains SO₄²⁻ ions, & HCL contains Cl⁻ ions, which show up in sulfate & halide tests.

2) Add excess Ba(NO₃)₂(aq):
• Any SO₄²⁻ ions should precipitate out as BaSO₄.
• Remove BaSO₄ by filtering the solution.
• If you intend to test for halide ions, don’t use BaCl₂, because the Cl⁻ ions will appear in the halide test.

3) Add AgNO₃(aq):
• Any CO₃²⁻/SO₄²⁻ ions have already been removed.
• Any precipitate formed must involve halide ions.
• Add NH₃(aq) to confirm which halide you have.

Question 28

What is the test for cations (NH₄⁺)?

Answer 28

• If heated, NH₄⁺ (aq) & OH⁻ (aq) react, forming NH₃.
• NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l).
• This reaction forms the basis for a test of NH₄⁺:
1) NaOH(aq) is added to a solution of an NH₄⁺ ion.
2) Ammonia gas is produced.
3) Gas bubbles are unlikely as NH₃ is soluble in water.
4) Mixture is warmed & ammonia gas is released.
5) Test for NH₃ gas with moist pH indicator paper.
6) NH₃ is alkaline & its presence will turn the paper blue.