f321 Flashcards
State and explain the trend in first ionisation energies across a period.
- 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.
State and explain the reactivity of elements down Groups 1&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.
Explain why water is polar.
- O is more electronegative than H.
- Creates a permanent dipole across O-H bonds.
- Water not symmetrical so dipoles do not cancel out.
State and explain the trend in atomic radii (size of atom) across a period.
- 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).
State and explain the trend in atomic radii (size of atom) down a group.
- 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.
State and explain the reactivity of elements down Group 7.
- 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.
Explain why there is an increase in melting/ boiling points from Na to Al.
- 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.
State and explain the trend in melting/ boiling points down Group 7.
- 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.
Describe the hydrogen bonding between two water molecules.
- 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.
Describe van der Waals’ (London) forces.
- 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.
In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Diamond
- 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.
In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Graphite
- 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.
In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Sodium chloride
- 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.
In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Iodine
- 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.
Outline how the ‘Electron Pair Repulsion Theory’ can be used to predict the shape of a molecule.
- 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.