Atomic Structure and the Periodic Table

Question 1

Define atomic number:

Answer 1

The number of protons in the nucleus of an atom of that element

Question 2

Define relative atomic mass:

Answer 2

The weighted average mass of an atom of a particular element divided by 1/12th the mass of a C-12 atom

Question 2

Define isotope:

Answer 2

Atoms of the same element that have the same number of protons but different numbers of neutrons

Question 3

Define orbital:

Answer 3

A region of space around the nucleus of an atom which can be occupied by a maximum of 2 electrons

Question 4

Define isoelectronic:

Answer 4

The same number of electrons

Question 5

How many electrons can each orbital hold?

Answer 5

2 electrons with opposite spins

Question 6

Which electrons fills and empties first?

Answer 6

In ions, the electrons in the highest energy levels are lost first, electrons are lost from 4s before 3d

Question 7

What is the electron configuration for Chromium?

Answer 7

1s2 2s2 2p6 3s2 3p6 4s1 3d5

Question 8

What is the electron configuration for copper?

Answer 8

1s2 2s2 2p6 3s2 3p6 4s1 3d10

Question 9

How does an atomic spectrogram works:

Answer 9

1. A sample (containing the element being studied) is vaporised and passed through a flame
2. Inside the atoms in the sample the electrons are being excited (due to being given more energy by the flame) and can move up to a higher energy level/quantum shell
3. However, this new arrangement is unstable so the electron will quickly drop back down to the ground state (normal shell)
4. When doing this it emits light of a particular wavelength

Question 10

Define ionisation energy:

Answer 10

The amount of energy required to remove an electron from a sample of an atom (in the gaseous state)

Question 11

Define successive ionisation energy:

Answer 11

Energy required to remove one mole of electron at a time from one mole of ions to form increasingly positive gaseous ions under standard conditions

Question 12

Define first ionisation energy:

Answer 12

The energy per mole required to remove one electron from gaseous atom under standard conditions
Eg/ Na (g) → Na+ (g) + e-

Question 13

What happens when an electron is removed from an inner shell?

Answer 13

Removing an electron from an inner shell (closer to the nucleus) requires more energy. This electron will be more strongly attracted to the nucleus because it is closer to the nucleus’s positive charge

Question 14

How do you know which group the element is in?

Answer 14

Count how many electrons are removed before the largest jump occurs

Question 15

How do you know which group the element is in?

Answer 15

Count how many electrons are removed before the largest jump occurs

Question 16

What happens if electrons from the same shell get removed?

Answer 16

If electrons are removed from the same shell this takes a little more energy each time (an increase due to the fact that there is the same amount of positive charge in the nucleus, but now that has fewer electrons left to attract, so can attract them more strongly)

Question 17

What happens if the electron is being taken from a new shell?

Answer 17

The attraction will be much greater and so this will take a lot more energy. This is because the electron is now being removed from much closer to the positive charge of the nucleus so is now much more strongly attracted and therefore a lot harder to remove

Question 18

What is the general trend for ionisation energy across a period?

Answer 18

H → He for example. The ionisation energy increases.

Question 19

Why does the ionisation energy increase across a period?

Answer 19

This is because the atoms have greater nuclear charge (positive - more protons) so the electrons are being attracted more strongly to the nucleus and are harder to remove.

Question 20

Why does He, Ne and Ar have particularly high ionisation energy?

Answer 20

Within their period, these have the highest nuclear charge and are particularly stable as they already have full shells of electrons; it is going to be particularly difficult to remove an electron from these atoms

Question 21

What is the general trend from one period to another?

Answer 21

The energy required is lower, as we go down the periodic table. As you go down the group it is requiring a bit less energy each time to remove an electron.

Question 22

Why does the energy required from one period to another is less each time?

Answer 22

This is because the electron in the outer shell being removed, is further away from the nucleus in Ca (more shells of electrons) so less strongly attracted to the nucleus.

Question 23

What happens with the shielding in larger atoms?

Answer 23

In a larger atom the outer shell electron is more strongly shielded (blocked from the attraction of the nucleus by inner shells of electrons, which also makes it easier to remove.)

Question 24

What are ‘blips’/evidence for subshells?

Answer 24

Eg/ Li → Ne. Although the ionisation energy required is increasing overall because of the increased charge. Not always true because B requires less energy to ionise than Be, and O requires less energy to ionise than N.

Question 25

Describe the O and N example blip:

Answer 25

O is 1s2 2s2 2p4, in this configuration the p electrons in the same orbital are repelling each other - reducing this repulsion is favourable and results in lower ionisation energy. Taking an e- here leaves a half full p subshell and reduced the electron-electron repulsion in the p orbital the electron is being taken from (makes it more stable). However, in N it is 1s2 2s2 2p3, so taking an e- leaves a partly full p subshell (less stable, doesn’t help with repulsion, and therefore more difficult)

Question 26

What is the use of a mass spectrometer?

Answer 26

Allows us to determine the RAM and it gives information about the different isotopes present in a naturally occurring sample of an element

Question 27

What is the first process in the mass spectrometer?

Answer 27

Vaporisation. The sample particle is entered into the ionisation chamber. The sample has to be in a gaseous form. Na (s) → Na (g)

Question 28

What happens in the ionisation stage?

Answer 28

The electron gun will shoot out electrons. These electrons are attracted to the other side of the chamber due to the positive plate. The electrons bombard the sample causing it to form a positive ion. Na(g) → Na+ (g) + e-

Question 29

What happens in the acceleration stage?

Answer 29

Two negatively charged plates with an electric current causes the ionised particles to move quicker and be focused into a single beam. It will then enter into a magnetic field.

Question 30

What happens in the deflection stage?

Answer 30

The electromagnet will cause the particle to deflect. Depending on mass to charge ratio (m/z). A sample with a high mass and low charge will deflect less than a sample with a low mass with a high charge

Question 31

What happens in the detection stage?

Answer 31

A tiny current is produced when each ion reaches the detector. The amount of current depends on the abundance of that ion. Detects electrical charge when an ion hits it.

Question 32

What are the three main points of periodicity?

Answer 32

Nuclear charge
Distance from the nucleus
Shielding (inner shell of electron blocking the attraction to nucleus)

Question 33

Describe the trend of the atomic radius across a period:

Answer 33

The atomic radius decreases (gets smaller)
- Nuclear charge increases
- Distance from nucleus: roughly the same (same shell of electron)
- Shielding: roughly the same (same shell of electron)
Therefore attraction to nucleus increases overall

Question 34

Describe the trend down a group:

Answer 34

The atomic radius increases (gets bigger)
- Nuclear charge increases
- Distance from nucleus increases a lot because more shells of electron
- Shielding increases a lot
- Therefore attraction to nucleus: decreases overall. This is because an increase in distance + shielding weakens the attraction; this outweighs the increase in attraction from the greater nuclear charge

Question 35

Compare the ionic radius to their atomic radius in metals?

Answer 35

Metal ions are smaller than the same metal atom because they have the same number of protons (nuclear charge) but have lost electrons so the attraction of electrons to the nucleus increases. (same number of proton is shared between fewer electrons)

Question 36

Compare the ionic radius to their atomic radius in non-metals?

Answer 36

Non-metal ions are bigger than the same non-metal atom because they have the same number of protons (nuclear charge) but have gained electrons so the attraction of electrons to the nucleus decreases. (the nuclear charge is shared between more electrons)

Question 37

Compare the ionic radius to their atomic radius in non-metals?

Answer 37

Non-metal ions are bigger than the same non-metal atom because they have the same number of protons (nuclear charge) but have gained electrons so the attraction of electrons to the nucleus decreases. (the nuclear charge is shared between more electrons)

Question 38

Define electron affinity:

Answer 38

The ability to attract electrons (electronegativity)

Question 39

Define electron affinity:

Answer 39

The ability to attract electrons (electronegativity)

Question 40

Describe the trend of electron affinity across a period:

Answer 40

Electron affinity increases:
- Nuclear charge increases
- Distance from nucleus roughly the same
- Shielding: roughly the same
Therefore attraction to nucleus: increases

Question 41

Describe the trend of electron affinity down a group:

Answer 41

Electron affinity decreases
- Nuclear charge increases
- Distance from nucleus increases a lot (more shells of electron)
- Shielding increases a lot
Therefore attraction to nucleus decreases overall. Distance and shielding increase would weaken the attraction, this outweighs nuclear charge which would strengthen the attraction therefore overall, the attraction to electrons decreases.

Question 42

What happens to the boiling/melting points of metals as you go across the period?

Answer 42

Melting and boiling point increase across the period because metallic bonds get stronger. The bonds get stronger because the metal ions have an increasing number of delocalised electrons and a decreasing radius (higher charge density). This means there’s a stronger attraction between the metal ions and delocalised electrons, so stronger metallic bonding.

Question 43

Why do giant covalent lattices have high melting/boiling points?

Answer 43

They have strong covalent bonds linking all their atoms together. A lot of energy is needed to break all of these bonds.

Question 44

Why do simple molecular structures have low melting and boiling points?

Answer 44

Their melting points depend upon the strength of the London forces between their molecules. London forces are weak and easy to overcome, therefore have a low melting/boiling point.