3.1.1 Periodicity Flashcards
what order is the periodic table arranged in
increasing atomic number
what are the vertical columns called
groups (1 to 18)
what are the horizontal rows called
periods
what are the blocks
correspond to electron found on highest subshell
what is periodicity
there are repeating trend in properties across each period
what is ionisation energy
how easily an atom loses electrons and forms positive ions
what is first ionisation energy
the energy needed to remove one electron from each atom in a mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
what are the ionisation energy equations
Na(g) ==> Na(g)1+ + e-
Na(g)1+ ==> Na(g)2+ e-
which 3 factors effect the first ionisation energy of an element
1) atomic radius
2) nuclear charge
3) electron shielding
what is atomic radius and how does it effect ionisation energy
-greater the distance between outer electrons
-less nuclear attraction
-electrons easily lost
-LOWEST AT TOP RIGHT OF PERIODIC TABLE:
- across period: more protons, means more positive charge of nucleus, means electrons are pulled closer together, reducing atomic radius
explain nuclear charge
-more protons in nucleus
- greater nuclear attraction
explain electron shielding
- inner shell electrons repel outer electrons (all negative)
- causes repulsion called shielding effect
- reduces nuclear attraction
what is nuclear attraction
attraction between positive nucleus and outer electrons
how many IE’s will an atom have
- as many electrons as it has
why does IE increase with successive energies
- less electrons, e- pull closer to nucleus and is nearer it, more IE
- as size of ion is smaller
what is the definition of 2nd ionisation energy
the amount of energy needed to remove one electron from from each 1+ ion in one mole of gaseous ions of an element and form one mole of 2+ gaseous ions
how can the ionisation energy graph tell you which element we are looking at
- number of ionisation energies = number of electrons = number of protons
- large jump between ionisation energies = moving between shells = tells you number of electrons on outer shell = group number
what do the periodic trends of ionisation energy show proof of
shells and subshells
what is the trend in ionisation energy
- decreases down a group
- increases across a period
-(sharp decrease from end of one period to start of the next one)
why does ionisation energy decrease down a group
1) atomic radius increases, so outer electrons further from nucleus, so lower nuclear attraction
2) more electron shielding, lowering NA
3) nuclear charge increases, but outweighed by other two factors
why does ionisation energy increase across a period
1) atomic radius decreases
2) electron shielding stays the same
3) nuclear charge increases
- so NA increases, and harder to lose electron
what are the 2 exceptions to the IE rule
- between group 2 and 3, and group 5 and 6
why does IE decrease from group 2 to 3
- electron in group 3 is in the p subshell
- has a higher energy state
- electron is easier to remove
- despite increased nuclear charge
why does IE decrease from group 5 to 6
- electron in group 5 is unpaired in the p orbital
- electron in group 6 is paired in a p orbital
- electron pair repels
- and is at higher energy state
- so easier to remove electron
- despite increased nuclear charge
what is metallic bonding
strong electrostatic forces of attraction between cations and delocalised electrons
what is the position of subatomic particles in metals
- cations are in fixed positions, maintaining structure and shape
- delocalised electrons are in a mobile pool
how are metals held
in a giant metallic lattice
what and why is the melting point of metals
HIGH: (most solid except mercury)
- large amount of energy needed to overcome the strong electrostatic forces of attraction between cations and delocalised electrons
- increases across a period: more delocalised electrons per each ion (e.g. 3- each), higher nuclear charge and smaller atomic radius means stronger electrostatic forces of attraction (referred to as high charge density metals) (still all have same, metallic bonding)
explain the conductivity of metals
VERY GOOD:
- have delocalised electrons that can move through the structure and act as mobile charge carriers
explain the solubility of metals
INSOLUBLE: (as atoms)
- metal would react with solvent before chance to dissolve
- also strong metallic bonding
which atoms form giant covalent lattices
B, Si, C
what is the structure of diamond and silicon
- formed of 4/4 available carbon bonds
- tetrahedral shape
-109.5 degree bond angle
explain melting point of giant covalent lattices
HIGH:
very strong covalent bonds need to be overcome
explain solubility of giant covalent lattices
INSOLUBLE:
- covalent bonds holding the lattice together are too strong to be broken down by interactions with solvent
explain the conductivity of giant covalent lattices
DO NOT:
- no ions or delocalised electrons to move and carry charge
explain the exceptions to the inconductivity of giant covalent lattices
GRAPHENE AND GRAPHITE:
- made of hexagonal sheets
- have delocalised electron
- can conduct electricity
GRAPHITE: single layer
GRAPHENE: sheets joined by weak london forces, so slippery
which element make up simple molecular lattices
non metals
explain the boiling point of simple molecular lattices
LOW:
- only need to overcome the weak london dispersion forces between molecules, which don’t require a lot of energy to overcome
- more electrons= more LDF = higher boiling point, so generally increases across period (NOT S8)
explain the conductivity of simple molecular lattices
NONE:
no delocalised electrons or ions
( all electrons are in localised bonds)
explain the solubility of simple molecular lattices
ONLY IN NON-POLAR:
- not in water, as polar
- LF can form between molecule and non-polar molecules like hydrocarbons
explain the particles and forces in types of structures
METALLIC= metallic bonding, cations, delocalised electrons
GIANT COVALENT= atoms, covalent bonds
SIMPLE MOLECULAR= molecules, london forces
