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Question 1

State and explain the trend in first ionisation energies across a period.

Answer 1

• First ionisation energies increase.
• Outermost electrons in same shell.
• Nuclear charge increases (more protons).
• Atomic radius decreases (outermost electrons pulled closer to nucleus).
• Nuclear attraction on outermost electron increases.
• More energy needed to remove outermost electron.

Question 2

State and explain the reactivity of elements down Groups 1&2.

Answer 2

• First ionisation energy decreases down the group.
• Reactivity increases.
• Outermost electron is in a new shell further from the nucleus with more shielding effect from inner electrons.
• Atomic radius increases (more shells).
• Nuclear attraction on outermost electron decreases.
• Although nuclear charge increases this is ‘far outweighed’ by more distance and shielding.

Question 3

Explain why water is polar.

Answer 3

• O is more electronegative than H.
• Creates a permanent dipole across O-H bonds.
• Water not symmetrical so dipoles do not cancel out.

Question 4

State and explain the trend in atomic radii (size of atom) across a period.

Answer 4

• Atomic radius (size of atom) decreases.
• Outermost electron in the same shell.
• Nuclear charge increases (more protons).
• Nuclear attraction on electrons increases (electrons are pulled closer to the nucleus).

Question 5

State and explain the trend in atomic radii (size of atom) down a group.

Answer 5

• Atomic radii (size of atom) increases (more shells).
• Greater distance of outermost electron to nucleus with more shielding from inner electrons.
• Nuclear attraction on electrons decreases (electrons are not pulled as close to nucleus).
• Although nuclear charge increases this is ‘far outweighed’ by more distance and shielding.

Question 6

State and explain the reactivity of elements down Group 7.

Answer 6

• Reactivity decreases.
• More shells.
• Greater distance to nucleus with more shielding from inner electrons.
• Nuclear attraction (on electrons to be gained) decreases.
• Atom is less likely to gain
  an electron to form the 1- ion.

Question 7

Explain why there is an increase in melting/ boiling points from Na to Al.

Answer 7

• Ionic charge increase: Na+, Mg2+ and Al3+ across period.
• More delocalised electrons across period.
• Attraction between positive metal ions and delocalised electrons (metallic bonding) increases.
• More energy is needed to break the stronger metallic bonds.

Question 8

State and explain the trend in melting/ boiling points down Group 7.

Answer 8

• Melting/ boiling points increase.
• More electrons down Group 7.
• Strength of van der Waals’ forces between molecules increase.
• More energy is needed to break the forces.

Question 9

Describe the hydrogen bonding between two water molecules.

Answer 9

• Well labelled diagram*
• Water is polar
• Permanent dipole across O-H bonds.
• Hδ+ on one water molecule attracts lone pair of electrons on highly electronegative O atom of another water molecule.

Question 10

Describe van der Waals’ (London) forces.

Answer 10

• Occur in non-polar molecules (overall).
• Temporary dipole forms in one molecule due to uneven distribution of electrons.
• Induces a dipole in a neighbouring molecule.
• Creates van der Waals’ forces between molecules.

Question 11

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Diamond

Answer 11

• Giant covalent.
• High melting/ boiling point as the strong covalent bonds need a lot of energy to break.
• Non-conductor of electricity as there are no charge carriers.
• Insoluble in water as strong covalent bonds cannot be broken.

Question 12

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Graphite

Answer 12

• Giant covalent.
• Weak van der Waals’ intermolecular forces between layers.
• High melting/ boiling point as the strong covalent bonds need a lot of energy to break.
• Good conductor of electricity as there are delocalised electrons free to move across the layers and carry charge.
• Insoluble in water as strong covalent bonds cannot be broken.

Question 13

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Sodium chloride

Answer 13

• Giant ionic
• High melting/ boiling point as the strong ionic bonds need a lot of energy to break.
• Good conductor when aqueous/ molten as IONS are free to move and carry charge.
• Non-conductor when solid as IONS are fixed in a lattice and cannot move or carry charge.
• Soluble in water as polar H2O is attracted to Na+ and Cl- and ionic lattice breaks up.

Question 14

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Iodine

Answer 14

• Simple covalent
• Weak van der Waals’ intermolecular forces between molecules.
• Low melting/ boiling point as the weak intermolecular forces between molecules need less energy to break.
• Non-conductor of electricity as there are no charge carriers.
• Partially soluble in polar water.
• Soluble in non-polar hexane (organic solvent) as hexane forms van der Waals’ with I2 molecules.

Question 15

Outline how the ‘Electron Pair Repulsion Theory’ can be used to predict the shape of a molecule.

Answer 15

• Electron pairs repel other electron pairs as far apart as possible.
• Lone pairs of electrons repel more than bond pairs of electrons.
• The number and type of electron pairs around the central atom determines the shape of a molecule.

Question 16

State and explain two anomalous properties of H2O resulting from hydrogen bonding.

Answer 16

• Ice is less dense than water: hydrogen bonds between the molecules are longer in ice than water holding H2O molecules further apart in an open lattice.
• ‘Higher than expected’ melting/ boiling points: Hydrogen bonds between molecules are the strongest intermolecular force and therefore more energy is needed to break them.