Topic 1 - Atomic Structure & the Periodic Table

Question 1

What are the relative masses and charges of electrons, protons & neutrons?

Answer 1

>Electrons = 1/1840 & -1
>Protons = 1 & +1
>Neutrons = 1 & 0

Question 2

What is meant by the term mass number?

Answer 2

The total number of protons & neutrons.

Question 3

What is meant by the term atomic number?

Answer 3

The number of protons.

Question 4

What is meant by the term isotope?

Answer 4

Atoms with the same number of protons but different numbers of neutrons.

Question 5

What is meant by the term relative isotopic mass?

Answer 5

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

Question 6

What is meant by the term relative atomic mass?

Answer 6

The average mass of 1 atom compared to 1/12th of the mass of 1 atom of carbon-12.

Question 7

How do you calculate the relative atomic mass?

Answer 7

∑ (isotopic mass x % abundance) / 100 OR ∑ (isotopic mass x abundance) / total abundance

Question 8

What does a mass spectrometer measure?

Answer 8

m/z (mass/charge ratio) & abundance

Question 9

How can you work out the Mᵣ from the mass spec?

Answer 9

The larges peak on the mass spec is due to the complete molecule (the M+ peak). The M+1 peak is cause by the carbon 13 isotope.

Question 10

What is meant by the term 1st ionisation energy?

Answer 10

The energy required when 1 mole of gaseous atoms forms 1 mole of gaseous ions with a single positive charge. (energy to take 1 electron)
H → H⁺ + e⁻

Question 11

What is meant by the term 2nd ionisation energy?

Answer 11

The energy required when 1 mole of gaseous atoms with a single positive charge forms 1 mole of gaseous ions with a double positive charge.
Ti⁺ → Ti²⁺ + e⁻

Question 12

How does the number of protons affect the ionisation energy?

Answer 12

The more protons there are in the nucleus, the greater the attraction will be ∴ the higher the ionisation energy.

Question 13

How does electron shielding affect ionisation energy?

Answer 13

The more electrons there are, the more they repel the outer electrons ∴ the lower the ionisation energy.

Question 14

How does the distance of the electrons from the nucleus affect the ionisation energy?

Answer 14

The further the outer electrons are from the nucleus, the weaker the attraction will be ∴ the ionisation energy will be lower.

Question 15

Why does the 1st ionisation energy increase across a period?

Answer 15

The number of protons increases, increasing the attraction to the nucleus, without extra shells for more shielding.

Question 16

Why does the 1st ionisation energy decrease down a group?

Answer 16

The outer electrons are in shells further away from the nucleus and have more shells shielding them.

Question 17

Why are succesive ionisation energies larger?

Answer 17

When the 1st electron is removed a positive ion is formed which increases the attraction of the remaining electrons, making them harder to remove.

Question 18

Why has helium got the largest ionisation energy?

Answer 18

It has no shells before the outer electron to shield it, and is bigger that hydrogen as it has 1 more proton.

Question 19

Why are there drops from Be → B & Mg → Al?

Answer 19

B & Al start to fill a new orbital so the outer electrons have more shielding & energy ∴ decreasing their ionisation energies.

Question 20

Why are there drops from N → O & P → S?

Answer 20

The outer electrons start to double up in the orbitals, causing slight repulsion & making them easier to remove.

Question 21

What is an orbital?

Answer 21

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

Question 22

In what order do electrons fill subshells?

Answer 22

Singly, before pairing up with an electron with an opposite spin.

Question 23

Which elements are in the s, p & d blocks?

Answer 23

s = 1,2
p = 3,4,5,6,7,0
d = transition metals

Question 24

What is the electronic structure for ions?

Answer 24

When a positive ion is formed electrons are lost, when a negative ion is formed electrons are gained.

Question 25

What do orbitals represent?

Answer 25

The mathematical probability of finding an electron at a certain point around the nucleus.

Question 26

What is meant by the term periodicity?

Answer 26

The repeating pattern of physical or chemical properties going across the periods.

Question 27

Why does the atomic radii decrease across a period?

Answer 27

The electrons have the same shielding as they are in the same shell, but have an increased attraction to the nucleus, as the number of protons increases.

Question 28

Why does metallic bond strength increase across a period?

Answer 28

The more electrons there are in the outer shell, the more electrons can be released into the sea. (Na, Mg, Al)

Question 29

Why does Si have a high melting point?

Answer 29

It is macromolecular (lots of strong covalent bonds) ∴ needs lots of energy to break them.

Question 30

Why are the mp & bp of molecular gases low?

Answer 30

They have weak london forces ∴ little energy is needed to break them.

Question 31

Why are the mp & bp of Ne & Ar so low?

Answer 31

They are monoatomic, so have very weak london forces.

Question 32

Why does S₈ have a high mp than P₄?

Answer 32

They are both low as they are simple molecular (weak london forces) but S₈ has more electrons ∴ slightly stronger london forces.

Question 33

In what order do transition metals loose electrons?

Answer 33

Loose electrons from 4s before 3d.