3.1.1.3 Electron Configuration

Question 1

What is a principal quantum shell/energy level?

Answer 1

-the shells in which electrons are arranged around the nucleus
N=1 is the first shell closest to the nucleus etc.

Question 2

How many electrons can each principal quantum shell hold?

Answer 2

N=1 can hold up to 2
N=2 can hold up to 8
N=3 can hold up to 18
N=4 can hold up to 32

Question 3

What are subshells?

Answer 3

The principal quantum shells can be split into subshells, s p d and f. The energy of the electrons in these subshells increases from s to f.

Question 4

What is the shape of the s orbital?

Answer 4

Spherical

Question 4

How many electrons can each subshells hold

Answer 4

S= max of 2 electrons
P= max of 6 electrons
D= max of 10 electrons
F= max of 14 electrons

Question 5

What is an orbital?

Answer 5

Subshells contain orbitals, which are spaces that are occupied by a maximum of 2 electrons with opposite spins.

Question 5

How many orbitals does each subshells contain?

Answer 5

S= one orbital
P= three orbitals
D= five orbitals
F= seven orbitals

Question 6

Describe how shells are filled by electrons

Answer 6

-electrons are added one at a time
-the lowest available energy level is filled first
-each energy level must be full before the next one fills up
-orbitals fill singly before pairing up (Hundt’s rule)

Question 6

Describe how to simplify electronic configurations

Answer 6

Simplify to the nearest noble gas
Eg. Calcium= [Ar]4s2

Question 7

Why does the 4s subshells fill up before the 3d subshell?

Answer 7

-4s has a lower energy than 3d, so will always be filled before the 3d
-however, when written, 3d is always written first

Question 8

Describe how to write the electronic structure of an ion

Answer 8

Simply add or remove electrons from the highest energy subshell.

Question 8

Why are electrons lost from the 4s subshell before the 3d subshell?

Answer 8

4s=FIRST IN FIRST OUT
-electrons are removed from 4s first
-upon filling 3d, the d orbitals become lower in energy than 4s
-therefore on ionisation, electrons are lost from 4s first

Question 9

Define isoelectronic

Answer 9

Two atoms, ions or molecules that have the same electronic configuration as each other.

Question 10

What is ionisation energy?

Answer 10

A measure of the energy required to completely remove an electron from an atom of an element, to form an ion.

Question 11

How is ionisation energy measured?

Answer 11

It is measured in KJmol-1, as the energy required to remove one electron from each atom in 1 mol of gaseous atoms to form 1 mol of gaseous ions.
-it is measured at 298k and 101KPa

Question 12

Define first ionisation energy

Answer 12

First ionisation energy is the energy required to remove 1 mol of electrons from 1 mol of gaseous atoms, to form 1 mol of 1+ gaseous ions.

Question 13

Define second ionisation energy

Answer 13

Second ionisation energy is the energy required to remove 1 mol of electrons from 1 mol of gaseous 1+ ions, to form 1 mol of gaseous 2+ ions.

Question 14

What is periodicity?

Answer 14

A trend across the periodic table

Question 15

What is the periodicity of ionisation energy?

Answer 15

Ionisation energy increases across a period, but decreases down a group.

Question 16

What are the 3 factors that affect ionisation energy?

Answer 16

-nuclear charge
-distance from nucleus
-shielding

Question 17

Explain why IE increases across a period

Answer 17

-the nuclear charge increases
-this causes the atomic radius to decrease, as the outer shell is pulled closer to the nucleus
-the distance between the nucleus and outer electrons decreases
-the shielding by inner shell electrons remains reasonably constant as electrons are added to the same shell
-it becomes harder to remove an electron across a period, so ionisation energy increases

Question 18

Explain why ionisation energy decreases down a group

Answer 18

-down a group, the nuclear charge increases
-however, the number of shells also increases
-more inner shells increases shielding
-this makes it easier to remove an electron, which decreases the ionisation energy

Question 19

What is the trend between successive ionisation energies?

Answer 19

The energy required to remove an electron (ionisation energy) increases, as the ionisation number increases
-this is seen as large jumps on a graph when electrons are removed from a lower energy level

Question 20

How can you determine group number from a graph of ionisation energy and ionisation number?

Answer 20

The number of electrons in the outer shell/ group number, is equal to the number of electrons removed before a sudden jump in ionisation energy

Question 20

Why does IE increase as ionisation number increases?

Answer 20

-as more electrons are removed, the attractive forces increase due to decreased shielding and increased proton to electron ratio -this distance from the nucleus to the outer electron also decreases

Question 21

Describe why beryllium has a higher IE than boron

Answer 21

-boron’s outer electron is in a 2p-subshell rather than a 2s-subshell, so it is in a higher energy level -this means less energy is required to remove this electron -therefore boron has a lower first ionisation energy than beryllium

Question 22

Describe why aluminium has a lower first ionisation energy than magnesium

Answer 22

-aluminium’s outer electron is in a 3p-subshell rather than a 3s-subshell, so it is in a higher energy level -this means less energy is required to remove this electron -therefore aluminium has a lower first ionisation energy than magnesium

Question 23

Describe why sulphur has a lower first ionisation energy than phosphorous

Answer 23

-sulphur has 4 electrons in the 3p-subshell, so one of them must pair in an orbital -this increases repulsion between the electrons -this means the outer electron is more easily removed -therefore sulphur has a lower first ionisation energy than phosphorous

Question 24

Why do the changes in trend in group 2 and 3 provide evidence for atomic structures?

Answer 24

It provides evidence for electron subshells as the change in the usual pattern of ionisation energies can only be explained or proven by the presence of subshells