Periodicity Flashcards
What is the periodic table arranged by?
Increasing atomic number
Periodicity
A repeating pattern (chemical or physical) across different periods
What happens along the period?
Nuclear charge increases/Atomic number increases
Shielding remains the same
Atomic radius decreases
Ionisation
When atoms gain or lose electrons to form positive and negative ions respectively
First Ionisation Energy
Energy required to remove one mole of electrons
From one mole of gaseous atoms of an element
To form one mole of gaseous 1+ ions
Why are ionisation energies endothermic reactions?
negative electrons are held in their electron shells (by
their attraction with the positively charged nucleus)
MORE energy must be SUPPLIED than is released in order overcome this attraction and therefore remove one electron
from a gaseous atom
Successive ionisation energies
The amount of
energy required to remove each electron in turn.
Second Ionisation Energy
the energy required to remove one mole of electrons
From one mole of gaseous X+ ions
To form one mole of gaseous X 2+ ions
Equation for the first ionisation energy of chlorine
Cl (g) 🡪 Cl+ (g) + e-
Equation for the fourth ionisation energy of chlorine
Cl3+ (g) 🡪 Cl4+ (g) + e-
How to differentiate an element’s group from a graph of successive ionisation energies
The point where there is the greatest difference in ionisation energies
Factors affecting ionisation energies
Atomic Radius
Nuclear Charge
Electron Shielding
How does ATOMIC RADIUS affect ionisation energy?
LARGE atomic radius = SMALL ELECTROSTATIC ATTRACTION between nucleus and outermost electron
Positivley charged NUCLEUS is FURTHER AWAY from outermost ELECTRONS
LOW IE because LESS ENERGY is REQUIRED to REMOVE OUTERMOST ELECTRON
How does NUCLEAR CHARGE affect ionisation energy?
HIGH Nuclear charge = LARGER ELECTROSTATIC ATTRACTION between nucleus and outermost electrons
HIGH IE because MORE ENERGY REQUIRED to REMOVE OUTERMOST ELECTRON
How does ELECTRON SHIELDING/SHIELDING EFFECT affect ionisation energy?
MORE electron shielding = REDUCED ATTRACTION/ MORE REPULSION between inner and outer electrons
LARGER the ATOM = more electron shells and GREATER ELECTRON SHIELDING
LOW IE because LESS ENERGY is REQUIRED
What happens to the ‘factors’ as you go down the group?
Atomic radius increases (IE decreases)
Electron shielding increases (IE decreases)
Nuclear charge increases (IE increases)
What is the trend in first ionisation energies across a period?
Nuclear charge increases (the number of proton increases) (IE increases)
Electron shielding stays the same (The number of electron shells stays the
same )
Atomic radius decreases (IE increases)
outer shell electrons experience an increase in nuclear attraction overall across a period
Overall INCREASE in IE
Why does Boron have a lower IE than Beryllium?
SHIELDING EFFECT
The outer electron is removed from a 2p sub-shell which is higher in energy than a 2s sub-shell.
The s electrons provide a slightly greater shielding of the p electron.
Therefore, the 2p electron is easier to remove than a 2s electron.
Why does Oxygen have a lower IE than Nitrogen? (Try to draw out the electronic configuration of both)
ELECTRON PAIR REPULSION
Nitrogen has a half-filled set of p orbitals.
In oxygen the paired 2p electrons repel one another
This repulsion makes it easier to remove an electron from oxygen compared to nitrogen.
Why do simple molecular structures have LOW MP?
Weak London forces exist between their molecules/atoms.
As a result, very little heat energy is needed to break these forces and therefore melt them
Why do giant covalent structures have HIGH MP?
Giant covalent lattices contain many strong
covalent bonds
A large amount of heat energy is needed to break these strong covalent bonds
Why do giant covalent structures have HIGH MP?
