Topic 1 Atomic Structure & the Periodic Table

Question 1

Relative Atomic Mass

Answer 1

The weighted mean mass of an atom of an element relative to 1/12 the mass of an atom of carbon 12.

Question 2

Relative Isotopic Mass

Answer 2

The mean mass of an atom of an isotope relative to 1/12 of the mass of an atom of carbon 12.

Question 3

Formula for RAM

Answer 3

Sum of (Mass of isotope x abundance of isotope) / total abundance of all isotopes.

Question 4

Periodicity

Answer 4

Regularly repeating patterns across different periods/with increasing atomic number, e.g., in chemical , physical and/or atomic properties.

Question 5

Describe the trend in melting point across period 3.

Answer 5

Increases: Na to Mg to Al. Metallic bonding: ionic radius decreases, protons & e- increase and attraction between the ions & e- increases.
Big jump to Si– giant molecular, lots of strong covalent bonds.
Big drop down to P4, slight rise to S8 and decreasing thereafter to Cl2 & Ar. All simple covalent molecules except Ar which is monoatomic with weak intermolecular forces.

Question 6

d-block element

Answer 6

The last electron to fill an orbital does so in the d-subshell.

Question 7

s-block & p-block elements

Answer 7

The highest-energy electron is found in an s-/p-orbital respectively.

Question 8

Why don’t the d-block elements fit the same definition structure as p-block and s-block elements?

Answer 8

The 4s orbital e- has more energy than the 3d orbital e-.

Question 9

Orbital

Answer 9

A region within an atom that can hold up to 2 e- with opposite spins.

Question 10

In what order are e- lost?

Answer 10

e- further from the nucleus are lost first.

Question 11

Filling of subshells

Answer 11

Singly first, then doubling up, so that e- in each orbital have opposite spins.

Question 12

What determines the chemical properties of an element?

Answer 12

The number of electrons in the outer shell.

Question 13

What’s so special about Cu and Cr?

Answer 13

They only fill the 4s1 orbital, not 4s2 like all the other elements in their period. This pattern is repeated periodically for Mo & Ag too!

Question 14

Filling diagram

Answer 14

1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s etc.

Question 15

Graph of 1st IE against atomic number for periods 1 & 2

Answer 15

Big jump from H to He then a huge drop down below H to Li.
1st IE increases from Li to Ne: the shielding remains the same, but the number of protons increases, so the attraction between the nucleus & outer electrons increases.

Question 16

What are the exceptions to the trend in increasing 1st IE across period 2?

Answer 16

B & O.
B: less energy to remove the e- as it’s alone in the p-orbital, which has more energy than an s-orbital and is shielded by e- in the full s-subshell. A similar trend applies to Al concerning 3p.
O: 2p4, the 2 e- in the same orbital repel, so less energy is needed to remove one. Similar applies to S.

Question 17

Why does He have the highest 1st IE?

Answer 17

No shielding. Smallest atomic radius.

Question 18

What does the big jumps on a 1st IE graph of an element indicate?

Answer 18

A shell change. e- are removed from a shell closer to the nucleus that experiences less shielding.

Question 19

Why do successive IEs increase for an element?

Answer 19

The number of protons is the same, but the e- get closer to the nucleus.

Question 20

Isotopes of the same elements have identical chemical properties. Why?

Answer 20

They have identical electronic configurations.

Question 21

What is the difference between mass number & relative isotopic mass?

Answer 21

Mass number is always a whole number as it’s the sum of protons & neutrons in the nucleus) whereas relative isotopic mass is relative to the 1/12 the mass of an atom of carbon-12, so not likely to be an integer.

Question 22

Describe the mass spectrum of chlorine.

Answer 22

2 peaks at 35 & 37 with relative heights of 3:1. 3 peaks at 70, 72 & 74 with relative heights 9:6:1. Peak heights won’t be exactly 9:6:1 (as relative isotopic masses are not integers).

Question 23

What is the M + 1 peak?

Answer 23

In organic compounds, there is always a small % of the carbon-13 isotope present. the peak is only present/significant in molecules with large masses where the % of C-13 becomes significant.

Question 24

Molecular ion peak

Answer 24

The peak with the highest m/z ratio in the mass spectrum, the M peak.

Question 25

Quantum shell

Answer 25

Defines the energy level of an electron.

Question 26

Orbital

Answer 26

Region within an atom that can hold up to 2 electrons with opposite spins.

Question 27

Hund’s rule

Answer 27

Electrons will occupy orbitals singly before pairing takes place.

Question 28

Pauli Exclusion Principle

Answer 28

2 electrons cannot occupy the same orbital unless they have opposite spins.

Question 29

1st quantum shell

Answer 29

Electrons here have the lowest energy for that element. Contains only one subshell: 1s.

Question 30

2nd quantum shell

Answer 30

Two subshells: 2s & 2p.

Question 31

3rd quantum shell

Answer 31

3 subshells: 3s, 3p & 3d.

Question 32

4th quantum shell

Answer 32

4 subshells: 4s, 4p, 4d & 4f.

Question 33

How many orbitals in the s subshell?

Answer 33

1 (,so holds 2 e-).

Question 34

How many orbitals in a p subshell?

Answer 34

3 (, so can hold 6 electrons). The 3 orbitals are arranged at mutual right angles.

Question 35

Describe the shape of p orbitals?

Answer 35

Elongated dumbbell shape with variable charge density. The only difference between the 3 orbitals is their orientation in space.

Question 36

How many orbitals in a d subshell?

Answer 36

5, so can hold 10 electrons.

Question 37

How does atomic emission spectra provide evidence for the different energy levels in which electrons can exist within atoms?

Answer 37

When electrons in atoms are excited then return to ground state, they release electromagnetic radiation. All atoms of a particular element radiate the same specific set of frequencies, hence electrons can only have certain, fixed values.

Question 38

First ionisation energy

Answer 38

The energy required to remove one electron from each atom in one mole of atoms in the gaseous state.

Question 39

Second ionisation energy

Answer 39

The energy required to remove one electron from each singly-charged, positive ion in one mole of positive ions in the gaseous state.

Question 40

3 factors affecting ionisation energy

Answer 40

• The orbital in which the electron exists
• The nuclear charge of the atom.
• The repulsion/shielding experienced by the electron.

Question 41

Trend in ionisation energy down group 1.

Answer 41

Down the group, nuclear charge increases with the increasing no. of protons, so the attraction between the nucleus and outer electrons increases. However, quantum shells are added down the group, which means there is more shielding and shells further from the nucleus are at higher energy levels than shells closer to the nucleus. The increasing nuclear charge is outweighed. First IE decreases down the group.

Question 42

Atomic radius

Answer 42

Distance between 2 nuclei / 2. The largest of these measures is Van der Waals’s radius; you can also measure covalent radius and metallic radius. Always compare like for like.

Question 43

Why does atomic radius decrease across periods 2 & 3?

Answer 43

No. of protons increases, so does nuclear charge. This increases the attractive force between the nucleus and outer electrons, which outweighs the slight increase in shielding due to additional e-.