Topic 1: Atomic Structure And The Periodic Table

Question 1

Define isotope

Answer 1

Elements with the same number of protons but different number of neutrons

Question 2

Define relative atomic mass

Answer 2

The weighted mean mass of an atom of an element, compared to 1/12th mass of an atom of carbon 12

Question 3

Define relative isotopic mass

Answer 3

The mass of an atom of an isotope, compared to 1/12th of the mass of an atom of carbon 12

Question 4

What must your abundance add up to?

Answer 4

100%

Question 5

Formula for relative atomic mass

Answer 5

(Abundance A x m/z of A) + (Abundance B x m/z of B) all over total abundance

Question 6

How do you calculate isotopic mass?

Answer 6

Do relative atomic mass and then solve for x

Question 7

Why are arrows in electronic figuration in different directions?

Answer 7

Spin pairing. When 2 electrons occupy one orbital, they ‘spin’ in opposite directions

Question 8

Why do you fill orbitals singly first then pair them up?

Answer 8

Electron repulsion

Question 9

Why would a transition metal with an electronic configuration of 3d5 4s1 form a unipositive cation with an electronic configuration of 3d5 4s1 and not 3d4 and 4s2?

Answer 9

Lose from the 4s orbital first then from 3d as 4s has higher energy when filled but lower energy when empty.

Question 10

What is n=1 in the atomic emission spectra?

Answer 10

Ground state

Question 11

Why are line spectra used?

Answer 11

To identify elements and it is evidence for quantum shells

Question 12

What is a series in atomic emission spectra?

Answer 12

A group of lines

Question 13

Why do lines in atomic emission spectra get closer together?

Answer 13

The energy and frequency increases

Question 14

What does the line spectrum show?

Answer 14

The frequency of light in coloured bands

Question 15

Where must an electron fall to for a line to be produced at uv?

Answer 15

Ground state, n=1

Question 16

Where must an electron fall for a line to appear at visible light in atomic emission spectra

Answer 16

N=2

Question 17

Where will the line in atomic emission spectra appear if n=3?

Answer 17

Infrared

Question 18

Define first ionisation energy

Answer 18

The minimum amount of energy required to remove 1mol of electrons from 1mol of atoms in the gaseous state to form one mole of unipositive gaseous cations

Question 19

Ionisation energy trends down a group

Answer 19

Decreases down groups. Atomic radius increases. Outer electrons are further from nucleus, so there is a weaker attractive force, therefore less energy is required. Shielding increases so less energy is required to remove electrons.

Question 20

What is successive ionisation

Answer 20

The removal of more than 1 electron from the same atom

Question 21

What is the general trend of successive ionisation

Answer 21

There is a general increases in energy as removing electrons causes the atom to become increasingly more positive

Question 22

What happens to the atomic radius across period 3?

Answer 22

It decreases. There is an increased nuclear charge because there is an increase in protons

Question 23

What happens to the atomic radius down a group?

Answer 23

Increases due to extra electron shells being added

Question 24

What is the trend of ionisation energies across periods

Answer 24

Increases due to increasing number of protons and nuclear attraction increases. Shielding is similar and distance marginally decreases. More energy is required to remove outer electrons.

Question 25

Why is there a decrease at the 3rd element in the ionisation energies across periods

Answer 25

The element sits in a higher energy subshell (p instead of s) so it is slightly further from the nucleus and there is some shielding from s, the p orbitals are higher in energy so less enegry needed to remove an electron

Question 26

Why is there a decrease at the 6th element in ionisation energies across periods

Answer 26

Electronic configuration p3 to p4 which means the outer electron is paired in its orbital. This means there is more repulsion between the electrons and so they are higher in energy and therefore require less energy to remove.

Question 27

What is the general trend of melting points across period 3 metals

Answer 27

Increases as metal ions have a positive charge, increasing the number of delocalised electrons and a smaller ionic radius

Question 28

Why does silicon have the highest melting point in period 3

Answer 28

Due to its giant covalent structure

Question 29

Why is argon the lowest melting point in period 3

Answer 29

Argon only exists as atoms