Topic 1: Atomic Structure and the Periodic Table

Question 1

What is the relative mass of an electron?

Answer 1

0.0005

Question 2

With what type of properties will isotopes be different to the atoms?

Answer 2

Physical properties. Chemical properties are the same because the number of protons and electrons are the same.

Question 3

What is the definition of relative isotopic mass?

Answer 3

The mass of one atom of an isotope of an element relative to 1/12th of the mass of an atom of Carbon-12

Question 4

What is the definition of relative atomic mass?

Answer 4

The mean average mass of an atom of the isotopes of an element relative to 1/12th of the mass of an atom of Carbon-12

Question 5

What physical state does a substance need to be to undergo mass spectrometry

Answer 5

Gas

Question 6

How are atoms made charged when in a mass spectrometer?

Answer 6

They are bombarded with electrons, which remove electrons and ionise them, making them positively charged ions.

Question 7

After being ionised, what is the next step in mass spectrometry?

Answer 7

Acceleration of the ions by an electric field.

Question 8

How does the mass spectrometer separate the ions?

Answer 8

The ions are separated by their mass/charge ratio (m/z), by an electromagnetic field deflecting them.

Question 9

How are the deflected ions read off a display?

Answer 9

The detector gives an electric signal whch is converted to a spectrum.

Question 10

What are the axes of a mass spectrum?

Answer 10

x-axis: m/z
y-axis: abundance

Question 11

What does the number of peaks on a mass spectrum of an element indicate?

Answer 11

The number of isotopes in a sample.

Question 12

What extra peaks can be seen when the sample is a diatomic molecule?

Answer 12

The diatomic molecules may separate when the sample is bombarded, but some don’t, so peaks can be seen at the m/z ratios of a bonded molecule, or of the single ions

Question 13

How can mass spectrometry can be used to determine the relative molecular mass of a molecule with the ion M+

Answer 13

There will be a peak, with a higher m/z value than the other peaks, and the m/z value of this peak is the RFM of the M+ ion.

Question 14

What’s funny about the peaks of diatomic molecules

Answer 14

The diatomic molecules may be made from two of one ion, or one of each, so you get more peaks than there are ions (and the probabilities are all super cool!!!)

Question 15

What are the two elements where one electron jumps from an s subshell to a d subshell, becuase they’re more stable with half or full d orbitals?

Answer 15

Chromium and Copper

Question 16

How many electrons can an s, p, d and f subshell hold respectively?

Answer 16

2, 6, 10, 14

Question 17

What are the shapes of s and p subshells?

Answer 17

s: spherical

p: figure of 8

Question 18

Which subshells do electrons fill first?

Answer 18

Lower energy

Question 19

What does spin-pairing mean?

Answer 19

Electrons in the same orbital must spin in opposite directions.

Question 20

How do electrons fill up the orbitals in a subshell?

Answer 20

They fill up singly, before pairing up

Question 21

How many electrons can fill each of the first four quantum shells?

Answer 21

1st: 2
2nd: 8
3rd: 18
4th: 32

Question 22

Why can’t electrons exist ‘between’ energy levels

Answer 22

The energy levels are discrete energy levels and have a fixed energy

Question 23

How do atomic emission spectra provide proof for the existence of quantum energy levels?

Answer 23

The emission spectra have discrete lines of radiation emission, showing that electrons ‘jump’ between discrete energy levels, which are the quantum shells

Question 24

Why will the same element always have the same lines on an emission spectrum?

Answer 24

The frequency of the emission or absorption of EM radiation will always be the same because the energy levels of the shells containing the electrons is fixed.

Question 25

What is the definition of the first ionisation energy?

Answer 25

The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.

Question 26

Write a general equation for the first ionisation energy of an element X

Answer 26

X(g) -> X+(g) + e-

Question 27

What are the three factors that may influence ionisation energy?

Answer 27

1. Nuclear charge: More protons in nucleus = more positively charges = electrons attracted more strongly.
2. Distance of outer shell electron from nucleus: An electron in a shell close to the nucleus will be much more strongly attracted than one in a shell further away.
3. Shielding: Fewer electrons between the outer electron and the nucleus = more attraction between nucleus and electron.

Question 28

What is the trend in first ionisation energies going down a group?

Answer 28

Ionisation energies decrease, because the outer electron will be further away from the nucleus, and there will be more shielding.

Question 29

What is the general trend in first ionisation energies going across a period?

Answer 29

They increase generally, because there are the same number of quantum shells, but a more positively charged nucleus, so more attraction between the nucleus and outer electron.

Question 30

How do the successive ionisation energies of an element provide evidence for the existence of quantum shells, and also indicate the group an element is in?

Answer 30

There will be a larger increase between the successive ionisation energies of the final electron of a quantum shell, and the first electron of the next quantum shell, because there will be much less distance between it and the nucleus, as well as less shielding. This can also be used to work out which group an element is in, because it shows how many electrons there were in the outermost quantum shell.

Question 31

Write a general equation for the nth ionisation energy (successive ionisation energies :D)

Answer 31

X^(n-1)+ (g) -> X^n+ (g) + e-

Question 32

What does ‘periodicity’ mean?

Answer 32

The general repeating pattern or trend across periods in the periodic table (due to the number of quantum electron shells etc…)

Question 33

What decides the chemical properties of an element?

Answer 33

The number of outer shell electrons

Question 34

What is the trend in atomic radius across a period?

Answer 34

It decreases

Question 35

Why does the atomic radius decrease across a period?

Answer 35

• As the number of protons increases, there is a more positive charge in the nucleus, so the electrons are pulled closer to the nucleus, making the atomic radius smaller
• The extra electrons that the elements gain across a period are added to the outer quantum shell, so don’t impact the shielding

Question 36

There is a general trend of increasing ionisation energies across a period, apart from in two places. Where are these places in each period?

Answer 36

Between groups 2 and 3, and between groups 5 and 6.

Question 37

Why is there a drop in ionisation energies between groups 2 and 3?

(Despite the general trend of increase in it across a period)

Although the ionisation energy generally increases across the period, there are drops between group 2 and 3, and between 5 and 6.

Answer 37

Between group 2 and 3, the outer shell electron goes from the s-subshell to the p-subshell. The electron in the p-subshell is on average, further from the nucleus, and has slightly more shielding (from the two electrons in the s-subshell). These two factors both mean that despite the more positive nucleus charge, there is less attraction between the outer shell electron and the nucleus, therefore slightly decreasing the ionisation energy. Between group 3 and 4, the nuclear charge increases, and this increase in attraction exceeds the decrease in attraction from the increased shielding, so the ionisation continues increasing (and is more than the group 2 element).

These Qs are usually 3 markers!!

Question 38

Why is there a drop in ionisation energies between groups 5 and 6?

(Despite the general trend of increase in it across a period)

Although the ionisation energy generally increases across the period, there are drops between group 2 and 3, and between 5 and 6.

Answer 38

The elements in the same period in group 5 and 6 both have their outer shell electron in the p-subshell (which has 3 orbitals, and therefore can hold up to 6 electrons), but the electrons in the group 5 element will each be in an orbital that is singly-occupied, whereas in the group 6 element, the p-subshell will have 4 electrons, and therefore only one doubly-occupied subshell. The repulsion between the two electrons in this orbital exceeds the attraction from the more positively charged nucleus, therefore slightly decreasing the ionisation energy.

These Qs are usually 3 markers!!

Question 39

How do the first ionisation energies of successive elements (i.e. across a group) provide evidence for the existence of sub-shells?

Answer 39

The decrease in ionisation energies between elements in groups 2 and 3, and groups 5 and 6.

Question 40

Explain why the M.Ps and B.Ps increase for the first 3 elements in a period.

For periods 2 and 3

Answer 40

They are metals, so are metallically bonded. Going along the period, the atoms lose more electrons to bond metallically, so there is a greater negative charge in the sea of electrons, and more positively charged metal ions, so therefore a greater force of electrostatic attraction. This means more energy is needed to overcome them, so the M.Ps and B.Ps are higher.

Question 41

Explain why the M.Ps and B.Ps increase between the third and fourth elements in a period.

For periods 2 and 3

Answer 41

The 4th elements (Carbon and Silicon) have giant covalent structures. This means all the atoms have strong covalent bonds between them, so lots of energy is needed to overcome them, leading to high M.Ps and B.Ps.

Question 42

Explain why the M.Ps and B.Ps decrease after the 4th element in a period, and what will influence them in the elements in the groups after group 4.

For periods 2 and 3

Answer 42

• The elements after the 4th element bond as simple covalent molecules. This means the M.Ps and B.Ps depend on the strength of the London forces between the molecules, which are weak and easily overcome, giving the elements lower M.Ps and B.Ps.
• The strength of the London forces will vary depending on the number of electrons in a molecule, meaning that some of the group 5 - 7 elements will have higher/lower M.Ps and B.Ps than others (depending on how they bond)
• The noble gases at the end will have the lowest M.Ps and B.Ps, because they exist as individual atoms (they are monatomic), so have very weak London forces .