3.1.1 Periodicity Flashcards
First ionisation energy definition.
Energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
How does atomic radius affect ionisation energy?
Greater the distance between nucleus and outer electrons, the less the nuclear attraction.
How does nuclear charge affect ionisation energy?
More protons in nucleus of an atom, greater the attraction between nucleus and outer electrons.
Therefore increased ionisation energy
How does electron shielding affect ionisation energy?
Shielding effect= inner shell electrons repel outer-shell electrons
Easier to lose outer electron, ionisation energy lower
Why would there be a large increase between 2 successive ionisation energies?
The second electron of the 2 is removed from an inner shell, which would take more energy.
What is the trend of first ionisation energies down a group and explain why.
Decreases down a group because
-atomic radius increases
-more inner shells, shielding increases
-nuclear attraction on outer electrons decreases
What is the trend of first ionisation energies across a period and explain why.
Increases because
-nuclear charge increases
-same shell, similar shielding
-nuclear attraction increases
-atomic radius decreases
Where is the s block on the periodic table
Two groups on the left
Where is the d block on the periodic table
Middle 10 groups
Where is the p block on the periodic table
Right 6 groups
Second ionisation energy definition
Energy required to remove one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
Metallic bonding definition
Strong electrostatic attraction between cations and delocalised electrons
Describe a giant metallic lattice
Billions of metal atoms held together by metallic bonding
Electrical conductivity of metals
Metals conduct electricity as when a voltage is applied across a metal, delocalised electrons can move through the structure and carry a charge
Melting and boiling points of metals
High
High temperatures needed to provide large amount of energy to overcome strong electrostatic attraction between cations and electrons
Solubility of metals
Do not dissolve
What is a giant covalent lattice
Billions of atoms held together by network of strong covalent bonds
Examples of giant covalent lattices
Carbon (diamond, graphene, graphite)
Silicon
Structure of diamond and silicon
Have 4 outer electrons that form covalent bonds with other atoms, tetrahedral structure formed
Structure of graphene
Single latter of graphite, hexagonally arranged carbon atoms linked by strong covalent bonds,
Bond angles of 120, 3 covalent bonds by each carbon
Structure of graphite
Parallel layers of hexagonally arranged carbon atoms, layers bonded by London forces
3 covalent bonds formed by each carbon atom, spare electron is delocalised between layers, so conducts electricity
Melting and boiling points of giant covalent lattices
High due to strong covalent bonds, lots of energy needed to break bonds
Solubility of giant covalent lattices
Insoluble in most solvents, covalent bonds holding together atoms in lattice are too strong to be broken
Electrical conductivity in giant covalent lattices
In diamond and silicon, all outer electrons are involved in covalent bonds, no mobile charge carriers
Graphene and graphite can conduct electricity
What’s the trend in melting points across period 2 and 3 and explanation
Melting point increases from group 1 to 14, sharp decrease in melting point between group 14 and 15, melting points comparatively low from group 15 to 18
Sharp decrease due to switch from giant to simple molecular substances
