Periodicity Flashcards
What is ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
How do you write equations for ionisation energy
Atoms and ions must be gaseous.
Balance charge with one electron on RHS
Remember the ionisation number is equal to the charge on the ion formed
What is successive ionisation energy
Removing each electron in an atom in turn, from outer to inner electrons.
Why is there a general increase in ionisation
With each ionisation, ionic radius DECREASES (as nuclear charge on each electron increases).
As radius decreases, attraction between the nucleus and outer electron increases and more energy is needed to remove each electron.
What is the significance of a large increase with the table
Shows when electrons start being removed from a new inner shell.
The number of ionisations before the first large increase is equal to the number of outer shell electrons (group number
If there is a big increase between the 2nd and 3rd number what does this mean
There is a big increase at the 3rd ionisation.
This shows that it has two electrons in its outer shell.
Therefore it is in group 2.
How do you explain the general trend in first ionisation energy
Radius
Shielding
Charge
What happens with atomic radius
The greater the radius, the weaker the nuclear attraction, as the highest energy electron is further away from the nucleus.Pattern:
Atomic radius Decreases across the period
What happens with shielding
Shielding reduces the attractive force from the protons on the highest energy electrons.
The more inner shells there are, the larger the shielding effect, which will weaken the nuclear attraction.
Pattern: Shielding remains the same across the period (Electrons are added to the same shell)
What happens with charge
The more protons in the nucleus the greater the attraction on the highest energy electron.
Pattern:
The nuclear charge increase across the period
How do you explain the general trend of ionisation energy across a period
Across a period:
Atomic Radius decreases (R)
Shielding is similar (S)
Nuclear Charge increases (C)
Therefore there is a greater nuclear attraction on the highest energy electron, leading to an increase in ionisation energy.
How do you answer a 6 marker explain the trend of ionisation eenrgy across a period
Discuss the general trend (2 marks)
Discuss anomaly 1 - A fall group 2 to 3(2 marks)
Discuss anomaly 2 - A fall from group 5 to 6
(2 marks)
How would you explain the dip in the first ionisation energy for beryllium to boron or magnesium to aluminimum
2p sub-shell has a higher energy than 2s so the outer electron in boron is easier to remove than the 2s electron in beryllium.
How would you explain the dip in the first ionisation energy for nitrogen and oxygen
In oxygen, the outer electron is paired. Electrons in one of the 2p orbitals repel each other, making it easier to remove an electron from oxygen than nitrogen
Explain why theres a general increase in successive ionisation energy values
With each ionisation ionic radius decreases as the nuclear charge on each electron increases
As the radius decreases the attraction between the outer electron and nucleus increases and more energy is needed to remove each electron
Explain the significance of the large increase of ionisation energy
Electrons are being removed from next shield down so theres a dip in the radius and shielding leading to a jump in the nucleus
What is periodicity
A repeating pattern of properties across different periods
Why does first ionisation energy decrease down a group
Due to an increasing atomic radius and electron shielding which reduces the effect of the electrostatic forces of attraction
Explain why the first and second ionisation energies of strontium are less than those of calcium.
atomic radii of Sr > atomic radii of Ca/
Sr has electrons in shell further from nucleus than Ca/
Sr has electrons in a higher energy level/
Sr has more shells
Therefore less attraction
Sr has more shielding than Ca
(‘more’ is essential) 3
increased nuclear charge is outweighed / despite increased nuclear
charge …..by at least one of the factors above
What is the difference of the ionisation trend for period 2 and 3
All the first ionisation energies are higher than those for period 3 because of the smaller atomic radius less shielding so stronger nucleur attraction for electron removed
What are the structurs along period 2 and 3
Lithium Beryllium- Giant metallic
Boron and carbon- Giant covalent
Nitrogen, Oxygen Fluorine and Neon- Simple covalent
3- Sodium Magnesium and aluminimum- Giant metallic
Silicone- Giant covalent
Phosphorus Sulfur Chlorine Argon- Simple covalent
