Ionisation Energy Flashcards
What is the definition for the first ionisation energy?
- The amount of energy required to remove 1 mole of electrons from 1 mole of atoms in a gaseous state.
- (saying atoms makes this specific to FIRST ionisation energy.)
Equation for Na first ionisation energy?
INCLUDE STATE SYMBOLS!!
Na (g) –> Na⁺ (g) + e⁻
Definition of 2nd/3rd/ 4th etc (ie. successive) ionisation energies?
- The amount of energy needed to remove 1 mole of electrons from 1 mole of ions in a gaseous state.
- (ions is what makes definition specific to 2nd/3rd/ 4th ionisation energies.)
() - extra info for clarification.
What is equation for second ionisation energy of calcium?
Ca⁺(g) –> Ca²⁺(g) + e⁻
What is the general trend for ionisation energy going down a group?
- As you go down group, ionisation energy decreases.
What is the general trend for ionisation energy going across a period?
- As you go across a period, the ionisation energy increases.
Give 3 factors that affect ionisation energy.
- Number of protons in nucleus (ie. nuclear charge.)
- Distance of electron from the nucleus (ie. atomic radius.)
- (Electronic) shielding.
How does number of protons in nucleus affect ionisation energy?
- If atom is already ion, it will have more protons in nucleus (leading to greater attraction between positive nucleus and negative electron that needs to be removed - making it harder to remove the electron.)
Why does the ionisation energy increase as you go across periods?
- As number of protons increase, ionisation energy increases.
- As you go across periods, the number of protons in each element increases.
Why does ionisation energy decrease as you go down groups?
(in terms of distance between electron and nucleus.)
- As you go down groups, the distance between the negative electron/ positive nucleus increases.
- This means that the negative electron experiences less force from positive nucleus.
- So, the electron can be removed more easily/ less energy required.
Why does ionisation energy decrease as you go down groups?
(in terms of electronic shileding.)
- As you go down groups, the number of shells increase - meaning that number of inner electrons increases.
- The large number of inner electrons shields the outer electron from the force of attraction (from nucleus) and the electron is removed more easily.
How does the shielding of atoms change as you go across a period?
- Shielding = similar (for all atoms.)
Why is it that when you go across periods (from group 2 to group 3 horiontally), there is a small dip in the ionisation energy? (Ie. when you go from Berillyium to Boron.)
- (Outer) Electron in higher-energy 2p sub -orbital.
- Electron further away from the nucleus.
- Little more electronic shielding (from previous 2s sub-orbital.)
- Electrons are more easily lost from the higher- energy sub-orbital.
Why is there a slight dip in ionisation energy when going across periods from group 5 to group 6?
- At this point, 2 electrons need to be paired up in the p orbital.
- So the electrons in the same sub-orbital will repel each other (so the electron is more easily lost - even though the nuclear charge increases.)
- Shielding is THE EXACT SAME (because we HAVE NOT gone into another orbital!!)
Which has a higher ionisation energy AND why?
- Argon
- Potassium.
- Argon
- Potassium has more electron shells than argon so, due to electronic shielding, the electron in potassium is lost more easily.
- IMPORTANT NOTE: YES, THE NUCLEAR CHARGE INCREASES FROM ARGON TO POTASSIUM BUT THIS HAS LESS OF AN EFFECT THAN THE ELECTRON SHIELDING, SO IONISATION ENERGY FROM ARGON TO POTASSIUM DECREASES
