3.1.1 Periodicity Flashcards
Define group
columns of elements with the same number of electrons in their outer shells
- similar chemical properties
Define period
horizontal row of elements with the same number of electron shells
- Repeating trends in physical and chemical properties
Define periodicity
repeating trends in properties
What do the blocks of the periodic table tell us ?
the sub-shell with the highest energy electrons
- the WIDTH is the same as the number of electrons in the outer sub-shell
Define ionisation energy
the energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of gaseous ions
X (g) -> X+ (g) + e-
Define successive ionisation energy
removing additional electrons from each ion in a mole of gaseous ions
What happens to the ionisation energy as more electrons are lost ?
- the protons remain constant
- so there is a greater proton to electron ratio
- so there is increases nuclear attraction between the nucleus and outer electrons
- increasing the ionisation energy
What three factors affect ionisation energy ?
- Nuclear charge (number of protons)
- Atomic / ionic radius
- Extent of shielding in lower energy shells
How does ionisation energy change down the group ?
- DECREASES
- Down the group electrons are added to a new shell further from the nucleus so atomic radius increases
- Shielding (of electrons in lower energy shells) increases
- Nuclear charge increases
- Increased shielding outweighs increased nuclear charge
- so there is weaker attraction between the nucleus and the outermost electron
- less energy needed to remove it which lowers the ionisation energy
How does ionisation energy change across the period ?
- INCREASES
- Nuclear charge increases
- Atomic radius decreases (due to the increased nuclear charge)
- Shielding is constant (as electrons are added to the outer shells)
- increased attraction between outermost electron and nucleus
- so more energy is needed to remove it which increases the first ionisation energy
Explain the ionisation energy dip between magnesium-aluminum and berrylium-boron
- outer electron in the 3p orbital of aluminum is further away from the nucleus than the two electrons in the 3s orbital of magnesium
- so there is less nuclear attraction
- so less energy is required to remove it
- so aluminum has a lower ionisation energy than magnesium
Explain the ionisation energy dip between phosphorus-sulfur and nitrogen-oxygen
- sulfur has an extra electron so there is electron-electron repulsion in the 3p orbital of sulfur
- so the electron requires less energy to remove
- so sulfur has a lower first ionisation energy
How can you establish the number of shells from an ionisation energy graph ?
number of jumps shows the gaps between shells (so three jumps = four shells)
How can you establish the number of outer electrons in an ionisation energy graph ?
count how many electrons are removed before the first big jump (This also tells us the group number)
Define metallic bond
What structure do metals have ?
the strong electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons (delocalised electrons are in the outermost shell)
- Giant metallic lattice
Describe and explain electrical conductivity in metals
- good electrical conductors
- as delocalised electrons (in the outer shell) can move and carry charge
Describe and explain solubility in metals
- insoluble
- due to the strong metallic bonds
Describe and explain thermal conductivity in metals
- good thermal conductors
- as delocalised electrons can move and carry kinetic energy
Describe and explain the trend in melting point in metals
- more **delocalised electrons **(higher charge)
- stronger metallic bonding
- higher melting point
Size of metal ion and lattice structure also affect melting point - smaller radius = more attraction so higher melting point
Define giant covalent lattice
huge network of covalently bonded atoms
- very high melting and boiling points as a lot of energy is needed to break the strong covalent bonds
Describe the structure and properties of diamond
- what other element forms a very similar structure to diamond ?
What is the shape ?
Giant covalent lattice (each carbon is covalently bonded to four other carbon atoms)
* very high melting point
* very hard
* good thermal conductor due to the stiff lattice
* non-conductive as all four outer electrons are involved in covalent bonding
* will not dissolve (strong covalent bonds between carbon)
SILICON IS SIMILAR
tetrahedral
Describe the structure and properties of graphite
Giant covalent lattice (carbon atoms in flat hexagonal layers** covalently bonded** to three others - BUT weak intermolecular forces between layers - induced dipole dipole)
* HIGH melting point due to strong covalent bonds in hexagon sheets
* Slippy due to weak intermolecular forces between layers
* Conductive due to delocalised electrons
* Insoluble
* Less dense than diamond as layers are far apart
Describe the structure and properties of graphene
Single sheet of carbon atoms that is one atom thick
* conductive as delocalised electrons are free to move and carry charge
* Very strong as delocalised electrons strengthen bonds between carbon atoms
* transparent and light as it is only one layer
* insoluble due to strong covalent bonds
Define simple molecular lattice
a few atoms covalently bonded togethet
Describe and explain melting and boiling points in simple molecular structures
What about noble gases ?
- low melting/boiling point
- as you only have to overcome the induced dipole-dipole interactions
- which need minimal energy to overcome
- Larger molecule = stronger induced dipole-dipole forces = higher melting and boiling point
Noble gases
- exist as single atoms so have very weak induced dipole-dipole forces
- so have very low melting and boiling points
Describe the what affects the melting and boiling point across the period of:
- metals
- giant covalent lattice
- simple molecular structures
- noble gases
- Metals - more delocalised electrons and smaller atomic radius = stronger metallic bonds = increased melting and boiling point
- Giant covalent lattice - more covalent bonds = increased melting and boiling point
- Simple molecular structures - larger molecule = stronger induced dipole-dipole forces
- Noble gases - all exist as single atoms so have very weak intermolecular forces and low melting/boiling points
