3.1 - The Periodic Table

Question 1

Describe the arrangement of the periodic table.

Answer 1

Ordered in increasing atomic (proton) number. Periods show repeating trends in physical and chemical properties. Grouped into elements with similar physical and chemical properties.

Question 2

Why is argon before potassium in the periodic table despite having a larger atomic mass?

Answer 2

Argon has a lower atomic number. Argon has the same chemical properties as other elements in group 8. Potassium has the same chemical properties as other elements in group 1.

Question 3

What is periodicity?

Answer 3

A repeating pattern across different periods.

Question 4

Why do elements in the same group exhibit similar properties?

Answer 4

Similar outer shell electron configurations. Outer electrons are what give substances their chemical properties. This is why isotopes of the same element have the same chemical properties despite having different masses.

Question 5

What can the periodic table tell you about electron configurations?

Answer 5

The period an element is in tells you the number of energy levels that element has. The group an element is in tells you the number of electrons an electron has in its outer shell. The block an element is in tells you the sub shell that the last electron is in.

Question 6

What is first ionisation energy?

Answer 6

The amount of energy required to remove an electron from each atom in a mole of gaseous atoms to form a mole of gaseous 1+ ions.

Question 7

What is successive ionisation energy?

Answer 7

A measure of the energy required to remove each electron in turn from each atom or ion in a mole of gaseous atoms or ions.

Question 8

Using chlorine, give examples of equations that represent ionisation energies.

Answer 8

First IE: Cl(g) –> Cl+(g) + e-
Second IE: Cl+(g) –> Cl2+(g) + e-
Third IE: Cl2+(g) –> Cl3+(g) + e-

Question 9

Describe and explain the influences on ionisation energy.

Answer 9

Nuclear charge - a higher nuclear charge increases the energy needed.
Shielding - more shielding (electrons at lower levels) decreases the amount of energy needed.
Distance - increasing the distance to the outermost electron decreases the amount of energy needed.
Shielding and distance outweigh nuclear charge.

Question 10

Analyse ionisation energies.

Answer 10

The first ionisation energy will be the outermost electron. As you move to successive ionisation energies, you are moving through outer subshells as the electrons get closer and closer to the nucleus.

Question 11

Describe and explain how you can predict when electrons have moved into a closer shell from the successive ionisation energies of an element.

Answer 11

There will be a big jump in the ionisation energy required. This indicates a new shell closer to the nucleus.

Question 12

Describe and explain the trend in ionisation energies down a group.

Answer 12

Ionisation energy decreases. Atomic radii increases so electron is further away from the nucleus. Electron experiences more shielding. These two effects outweigh the increasing nuclear charge. This results in a decreased attraction between the outer electron and the nucleus. Less energy will be needed to remove it.

Question 13

Describe and explain the general trend in ionisation energies across a period.

Answer 13

Ionisation energy increases. Electrons added to the same shell so similar shielding. Atomic radii decreases so electron is closer to nucleus. Increased nuclear charge. These factors make it more difficult for the electron to be removed. More energy is therefore required to remove the electron.

Question 14

Describe and explain the trend in atomic radii across a period.

Answer 14

Radii decrease across a period. Electrons are added to the same shell and experience similar shielding. Nuclear charge increases across the period which pulls the electrons in closer.

Question 15

Explain the dip in ionisation energy between group 2 and 3 as you go across a period (e.g. Be and B).

Answer 15

The electron added at group 3 goes into a p orbital. This is a higher energy orbital. This reduces the amount of energy needed to remove it.

Question 16

Explain the dip in ionisation energy between groups 5 and 6 (e.g. N and O).

Answer 16

At group 5 there is a single electron in each p-orbital. At group 6 the first p orbital gains a second electron. The repulsion between the 2 electrons in the p orbital reduces the energy needed to remove and electron.

Question 17

Describe the metallic bonding present in a giant metallic lattice structure.

Answer 17

A attraction between a lattice of positive metal ions surrounded by a sea of delocalised electrons.

Question 18

Describe the bonding and structure of metallic elements and explain some of their properties.

Answer 18

Giant structure with metallic bonding (ions attract to electrons).
Melting and boiling point: Moderate to high. There is a strong attraction between the ions and the electrons which requires a lot of energy to separate.
Electrical conductivity: Conducts due to the free electrons that can move through the metal.
Malleable: The metal ions can slide past each other due the ‘lubrication’ provided by the delocalised electrons.
Solubility in water: Metals are not soluble in water. Technically the electrons are still attracted to the ions and there is not enough energy for the water molecules to pull them apart. (Metal ions on their own without their associated electrons are soluble). Some very reactive metals will react with water.

Question 19

Describe the bonding in giant covalent lattices.

Answer 19

Giant covalent lattices are huge networks of covalently bonded atoms.

Question 20

Describe the bonding and structure of diamond and explain some of its properties.

Answer 20

Giant structure with covalent bonding.
Each carbon is bonded to 4 other carbons.
Very high melting point: the structure is made up of a huge number of covalent bonds which would require a large input of energy to break apart.
Electrical conductivity: Does not conduct electricity as no free electrons or ions (it has no ions).
Hardness: very hard due to the strength of the network of covalent bonds.
Solubility in water: not soluble. Not enough attraction with water to enable the water molecules to break up the lattice.
Silicon has the same type of structure with similar properties.

Question 21

Describe the bonding and structure of graphite and explain some of its properties.

Answer 21

Giant structure with covalent bonding within each sheet and Van der Waals’ between each sheet.
Layered structure with free electrons between layers.
Each carbon is bonded to 3 other carbons with 1 spare electron.
Melting and boiling point: Very high. The structure is made up of a huge number of covalent bonds which would require a large input of energy to break apart.
Electrical conductivity: Conducts electricity due to free electrons.
Hardness: Soft. The layers are weakly held together and can slide across each other.
Solubility in water: not soluble. Not enough attraction with water to enable the water molecules to break up the sheets.

Question 22

Describe the bonding and structure of graphene and explain some of its properties.

Answer 22

Graphene is one layer of graphite. Each carbon is bonded to 3 other carbon atoms with one delocalised electron per atom.
The sheet is one atom thick (two-dimensional compound).
Electrical conductivity: The best known electrical conductor. The delocalised electrons can move very quickly above and below the sheet as there as no layers.
Strength: Very strong as the delocalised electrons strengthen the covalent bonds.

Question 23

Describe and explain the trend in melting and boiling points across a period.

Answer 23

Metallic bonding: Strong bonding requiring lots of energy to break due to the attraction between the positive ions and sea of delocalised electrons. As you go across the period the boiling point increases for metallic bonding because: the charge on the metal ion increases, there are more delocalised electrons, and this increases the attraction between the positive ions and delocalised electrons.
Giant covalent: Vast network of covalent bonds that need lots of energy to be broken.
Simple covalent (simple molecular): Held together by weak induced dipole-dipole attractions between molecules. These need less energy to break. Phosphorous and sulphur are higher because they have more electrons per molecule (P4 and S8 – this is also why sulphur is higher then phosphorous)
Simple atomic: Held together by weak induced dipole-dipole attractions between molecules. These need less energy to break.

Question 24

Describe the redox reaction of a group 2 element with oxygen, water and steam.

Answer 24

Oxygen:
Ox states: 0 0 +2 -2
. 2Mg(s) + O2(g) → 2MgO(s)
Mg = oxidised from 0 to +2
O = reduced from 0 to -2
Water:
Ox states: 0 +1 -2 +2 -2 +1 0
. Mg(s) + 2H2O(l) → Mg(OH)2(s) +H2(g)
Mg = oxidised from 0 to +2
H = reduced from +1 to 0
Steam:
Ox states: 0 +1 -2 +2 -2 0
. Mg(s) + H2O(g) → MgO(s) +H2(g)
Mg = oxidised from 0 to +2
H = reduced from +1 to 0

Question 25

Describe and explain the trend in reactivity of group 2 elements.

Answer 25

Reactivity increases down the group. The ions are formed easier as you go down the group because atomic radii increases so electron is further away from the nucleus. Electron experiences more shielding. These 2 effects outweigh the increasing nuclear charge so the nuclear attraction increases. This results in less energy being needed to remove the electron.

Question 26

Give symbol equations including state symbols for the following reactions: Group 2 element with hydrochloric acid, group 2 oxide with hydrochloric acid, group 2 carbonate with hydrochloric acid, heating a group 2 carbonate and group 2 oxide and water and state the pH.

Answer 26

Group 2 element with hydrochloric acid:
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
Group 2 oxide with hydrochloric acid:
MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)
Group 2 carbonate with hydrochloric acid:
MgCO3 + 2HCl(aq) → MgCl2(aq) + H2O(l) + CO2(g)
Heating a group 2 carbonate:
CaCO3(s) → CaO(s) + CO2(g)
Group 2 oxide and water and state the pH:
CaO(s) + H2O(l) → Ca(OH)2(aq)
pH = 10-12 – Alkalinity increases down the group.

Question 27

What is calcium hydroxide used for?

Answer 27

Neutralising acidic soils in agriculture (lime).

Question 28

What is magnesium hydroxide used for?

Answer 28

Neutralising excess stomach acid (antacid tablets).

Question 29

Describe and explain the trend in boiling points of the halogens.

Answer 29

Boiling point increases down the group. More electrons in the atoms as you go down the group. Increased induced dipole-dipole forces. The attraction between the molecules increases. Results in more energy being needed to break the London forces.

Question 30

Describe and explain the relative reactivity of the halogens.

Answer 30

Reactivity decreases down the group. Halogens higher up the group capture electrons easier (i.e. have a greater oxidising power). This is due to the smaller radii higher up the group, therefore electron captured is closer to the nucleus with less shielding.

Question 31

Describe the displacement (redox) reactions of halogens including observations.

Answer 31

More reactive halogens displace less reactive halogens. You see the colour of the halogen that has been displaced. If Iodine is displaced you see purple. If bromine is displaced you see an orange/yellow colour. Reaction is often put in cyclohexane as halogens are more soluble in organic solvents so colour change is more apparent.
The actual reaction:
Cl2(aq) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq) Orange/yellow colour seen.
The sodium ions are removed from the equation as they are spectator ions.
In the exam you must write:
Cl2(aq) + 2Br-(aq) → 2Cl-(aq) + Br2(aq) Orange/yellow colour seen.
This is the same for the other reactions. The ones below already have had the metal ions removed:
Cl2(aq) + 2I-(aq) → 2Cl-(aq) + I2(aq) Purple colour seen.
Br2(aq) + 2I-(aq) → 2Br-(aq) + I2(aq) Purple colour seen.

Question 32

Describe the identification of halide ions including observations.

Answer 32

Add silver nitrate.
Ag+(aq) + Cl-(aq) → AgCl(s) White precipitate.
Ag=(aq) + Br-(aq) → AgBr(s) Cream precipitate.
Ag+(aq) + I-(aq) → AgI(s) Yellow precipitate.
Can then add ammonia as colour difficult to distinguish. Silver chloride dissolves in dilute ammonia. Silver bromide dissolves in concentrated ammonia. Silver iodide does not dissolve.

Question 33

Describe the reaction of chlorine with sodium hydroxide and give a use for the reaction.

Answer 33

Cl2(g) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)
Must be cold, dilute aqueous sodium hydroxide.
Cl is reduced from 0 to -1 (in NaCl) and oxidised from 0 to +1 in (NaClO).
Used to form bleach.

Question 34

Describe the reaction of chlorine with water and include a use.

Answer 34

Cl2(g) + H2O(l) → HCl(aq) + HClO(aq)
Cl is reduced from 0 to -1 (in NaCl) and oxidised from 0 to +1 in (NaClO).
Used to purify water.

Question 35

What is a disproportionation reaction?

Answer 35

A reaction where the same element is oxidised and reduced.

Question 36

What are the benefits and risks associated with the use of chlorine in water treatment?

Answer 36

Benefits – kills bacteria so can prevent the spread of disease (e.g. cholera).
Risks – chlorine gas is toxic (only give this answer is chlorine is mentioned specifically instead of chlorine compounds). You could form chlorinated hydrocarbons. These can cause cancer.

Question 37

Describe the qualitative tests for carbonate, sulphate, chloride ions and ammonium.

Answer 37

Carbonate (CO3 2-). React with acid and carbon dioxide will be released. CO3 2-(s) + H+(aq) → CO2(g) + H2O(l). The CO2 can be bubbled through limewater which will turn cloudy.
Sulphate (SO4 2-). React with barium nitrate – a white precipitate will form. Ba 2+(aq) + SO4 2-(aq) → BaSO4(s).
Chloride ions. Use the halide tests.
Ammonium (NH4 +). Add sodium hydroxide and warm the solution. Hold damp red litmus paper at the top of the test tube. Litmus paper will turn blue if ammonia is being produced. NH4 +(aq) + OH-(aq) → NH3(g) + H2O(l).