Chapter 7 - Periodicity Flashcards
(31 cards)
How are the elements arranged in the periodic table?
In order of increasing atomic (proton) number
In vertical columns called GROUPS
In horizontal rows called PERIODS
What is meant by periodicity?
A repeating trend in properties of elements across each period
e.g. electron configuration, ionisation energy, structure, melting points
What does the number of the period tell us?
Gives the number of the highest energy electron shell in an element’s atom
What does the group number tell us and how does this lead to elements in the same group having simlar chemical properties?
The number of electrons in their outer shell
Elements in each group also have the same number of electrons in each sub-shell
This similarity in electron configuration gives elements in the same group similar chemistry
How are the elements of the periodic table divided?
Into blocks corresponding to the highest energy sub-shell filled:
What is meant by the first ionisation energy?
The energy required to remove one electron from each atom in one mole of gaseous atoms of an element
Units: kJ mol-1
What is the first ionisation energy of sodium?
Na(g) -> Na+(g) +e-
How does atomic radius affect ionisation energy?
The greater the distance between the nucleus and the outer shell electrons, the weaker the nuclear attraction between the nucleus and the outer electrons
How does nuclear charge affect ionisation energy?
The more protons there are in the nucleus of an atom, the greater the nuclear attraction between the nucleus and the outer shell electrons
How does inner shell shielding affect ionisation energy?
Inner shell-electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons
What is the trend in first ionisation energies as you go down a group?
Decreases
Atomic radius increases
More inner shells, so shielding increases
Nuclear attraction decreases
What is the trend in first ionisation energies as you go across a period?
Increases
Nuclear charge increases
Same number of inner shells, shielding stays the same
Atomic radius decreases
Nuclear attraction increases
Ionisation energy example:
A rise from lithium to beryllium
A fall to boron followed by a rise to carbon and nitrogen
A fall to oxygen followed by a rise to fluorine and neon
Describe the fall in first ionisation energy from Be to B
The 2p sub-shell has a higher energy than the 2s sub-shell
It is easier to remove the outer 2p electron in B, than the 2s electron in Be
The first ionisation energy of boron is less than the first ionisation energy of beryllium
What is important about an element in terms of ionisation energies?
It has as many ionisation energies as it has electrons
e.g. helium (2 electrons)
Why is the second ionisation energy of helium greater than the first?
After the 1st electron is lost, the 2nd electron will be removed from a +ve ion, in which the remaining electrons have been pulled closer to the nucleus
Nuclear attraction on remaining electron increases
More energy is needed to remove this second electron
Describe the fall in first ionisation energy from N to O
Nitrogen contains three 2p electrons
One electron in each 2p orbital
Spins are at right angle - equal repulsion as far apart as possible
Oxygen contains four 2p electrons
Two electrons in one 2p orbital
2p electrons start to pair
The paired electrons repel
What is metallic bonding?
The strong electrostatic attraction between metal cations and delocalised electrons
Describe the metallic structure:
‘Sea’ of delocalised electrons (from metal atoms) - these are free to move over the whole structure
Metal cations - held in fixed positions in the lattice maintaining the structure and shape of the metal
What are the physical properties of metals?
Metals have high electrical conductivity
Metals conduct when solid and in molten form
The delocalised electrons can move throughout the structure, carrying charge
Most metals have high melting, boiling points
Is dependent on the strength of the metallic bonds
Therefore large amounts of energy required to break them
Metals are insoluble and do not dissolve
What do most non-metals form?
Simple molecular lattice structures:
Strong covalent bonds between atoms within molecules
Weak intermolecular forces between molecules
What non-metals form giant covalent lattice structures?
Boron, carbon and silicon
What is a giant covalent lattice?
Atoms are held together by a network of strong covalent bonds
Describe the carbon (diamond form) and silicon structure:
Uses the four outer-shell electrons to form covalent bonds to other carbon, or silicon atoms
Results in a tetrahedral structure (Bond angles are all 109.5)
