Chapter 7 Ionisation Energies Flashcards
Ionisation energy
Measure of how easily an atom loses electrons to form positive ions
1st ionisation energy
The energy required to rem0ove 1 electron from each atom in 1 mole of gaseous atom of an element to form 1 mole of gaseous 1+ ions.
Factors affecting ionisation energy
- Atomic radius
- Nuclear charge
- Electron shielding
Factors affecting ionisation energy: Atomic radius
The greater the distance between the nucleus and the outer electrons of an atom = less nuclear attraction
Factors affecting ionisation energy: Nuclear charge
The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons
Factors affecting ionisation energy: Electron shielding
- Shielding effect - reduces attraction between the nucleus and outer electrons
Shielding effect
Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons
How many ionisation energies does an element have?
As many electrons as the element has
2nd ionisation energy
The energy required to remove 1 electron from each ion in 1 mole of gaseous 1+ ions of an element to form 1 mole of gaeous 2+ ions
What do successive ionisation energies allow predictions to be made about?
- The number of electrons in the outer shell
- The group of the element in the periodic table
- The identity of an element
Trend of ionisation energies down a group?
Decreases
Why do ionisation energies decrease down a group?
- Atomic radius increases
- More inner shells so shielding increases
- Nuclear attraction on outer electrons decreases
- first ionisation energy decreases
Trend in 1st ionisation energy across a period
Increases
Why does 1st ionisation energy increase across a period
- Nuclear charge increases
- Same shell, so similar shielding
- Nuclear attraction increases
- Atomic radius decreases
- 1st ionisation energy increases
Metallic bonding
Strong electrostatic attraction between cations and delocalised electrons
- The cations are fixed in position -> maintains the structure and shape of the metal
- The delocalised electrons are mobile and are able to move throughout the structure
Solid metal structure
Each atom has donated its negative outer shell electron to a shared pool of electrons which are delocalised throughout the whole structure
Delocalised
Spread out
Properties of metals
- Strong metallic bonds - attraction between cations and delocalised electrons
- High electrical conductivity
- High melting and boiling points
Properties of metals: Electrical conductivity
- Conduct electricity in solid and liquid states
- When a voltage is applied across a metal, the delocalised electrons can move throughout the structure (carrying charge)
Properties of metals: Melting and boiling points
High melting and boiling points - high temperatures are required to provide large amount of energy needed to overcome the strong electrostatic attraction between cations and electrons
Properties of metals: Solubility
Insoluble
Giant covalent lattice
Many billions of atoms held together by a network of strong covalent bonds
Giant covalent lattice: Carbon and silicon
- Both have 4 bonds = tetrahedral (109.5 degrees)
Properties of giant covalent structures: Melting and boiling points
High melting and boiling points
Strong covalent bonds require high energy to be broken
Properties of giant covalent structures: Solubility
Insoluble -> covalent bonds in lattice are too strong to be broken by interaction with solvents
Properties of giant covalent structures: Electrical conductivity
Non-conductors of electricity
All electrons are in covalent bonds
Special cases of electrical conductivity of giant covalent structures
- Graphene and graphite
