Ionisation Energy

Question 1

Describe first ionisation

Answer 1

First ionisation energy is the minimum energy required to completely remove one mole of electrons from the outer shell of one mole of atoms in the gas phase (it needs to be in the gas phase otherwise you can’t pull it)

Question 2

What’s the equations to show the first ionisation of an atom?

Answer 2

X(g) —> X+(g) + e-

Question 3

When will it be difficult to remove the outer electron from an atom? When will it be easy?

Answer 3

If is difficult to remove the outer electron from an atom, the ionisation energy for that atom will be high
If it is easy to remove the outer electron from an atom, the ionisation energy for that atom will be low

Question 4

What’s the equation to show the second ionisation of an atom?

Answer 4

X^+(g) —> X^2+(g) + e^-

Question 5

What are the three most important factors that affect ionisation energy?

Answer 5

The nuclear charge (No. of protons in the nucleus).

Shielding - The protection an electron has from the pull of the nucleus by the repulsion of inner shell electrons. Otherwise know as screening.

The distance the outer electron is from the nucleus.

Question 6

What’s the short hand for nuclear charge and electron shielding?

Answer 6

Nuclear charge = Z
Electron shielding = n

Question 7

How does nuclear charge affect I.E?

Answer 7

As the number of protons in the nucleus increases the strength of the nuclear charge gets greater and thus electrons experience a greater attractive force. This means that the Ionisation energy goes up with nuclear charge.

Question 8

How does electron shielding affect I.E?

Answer 8

The inner electron shells shield the outer electrons from the attractive force of the nucleus. Because more inner shells mean more shielding, the ionisation energy will be lower

Question 9

How does effective nuclear charge affect I.E?

Answer 9

It is the attraction that the electron experiences from the nucleus when shielding is taken into account

Question 10

What would the electrons experience if there was no shielding at all

Answer 10

If there was no shielding at all, the electron would experience the actual positive nuclear charge (Z) from the nucleus. However, shielding causes electrons to experience a reduced positive charge

Question 11

Explain the general trend in 1st ionisation energy across a row

Answer 11

Generally the effective nuclear charge increases across a period since the nuclear charge increases and electrons are added to the same shell (no change in shielding)

Z increases
N stays same
Overall I.E goes up

Question 12

Explain when and where there is a decrease in 1st ionisation energy across a row

Answer 12

The ionisation energy decreases from Be to B because the effective nuclear charge of Boron has decreased on the outer electron due to increased shielding
N increases
Z increases
I.E. goes down

And N to O, oxygen has a pair of repelling 2p electrons and nitrogen has a half full stable 2p orbital
N stays same
Z goes up

Question 13

What group always has the lowest 1st ionisation energy?

Answer 13

Group 1

Question 14

What group always has the highest 1st ionisation energy?

Answer 14

Group 8

Question 15

Why is group 3 1st ionisation lower than group 2?

Answer 15

Because we are removing the outer electron from atoms that have a whole shell more of shielding

Question 16

What’s is greater, 1st or 2nd ionisation energies? Why?

Answer 16

2nd because…

they involve removing an electron from a positive ion rather than a neutral atom = forces of attraction are greater on valence electrons

Sometimes the 2nd is removed from a shell closer to the nucleus.
I.e. the radius of an ion is smaller than an atom sot eh effective nuclear charge is greater on this electron

Removing an electron from an atom to form an ion, means that the net positive charge on the remaining electrons is going to be greater since there are less electrons to spread the attractive forces over. I.e. the effective nuclear charge has increased.

Question 17

Why does I.E. go down a group in the periodic table?

Answer 17

As we go down a group valence electrons enter new shells which are further from the nucleus and are more shielded by additional completed inner shells of electrons. Therefore, the effective nuclear charge on the outer electron decreases down a group due to the increased shielding

Question 18

Why do successive ionisation energies always increase?

Answer 18

There is a greater effective nuclear charge in the remaining electrons

The distance that the electron is from the nucleus decreases as we move between shells

As each electron is removed, there is less electron-electron repulsion and each shell will be drawn in slightly closer to the nucleus