1 - Atomic structure and the periodic table

Question 1

What are isotopes?

Answer 1

Isotopes are atoms of the same element with the same number of protons and electrons but a different number of neutrons.

Question 2

What are ions?

Answer 2

Ions are charged particles formed when atoms lose or gain electrons to gain stability.

Question 3

What is relative isotopic mass?

Answer 3

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

Question 4

What is relative atomic mass?

Answer 4

The weighted average mass of an atom of an element compared with 1/12th the mass of an atom of carbon 12.

Question 5

What is a mass spectrometer used for?

Answer 5

It is an analytical piece of equipment. It is used to find the relative abundances of isotopes in a sample of an element.

Question 6

What type of molecules can be read by a mass spectrometer?

Answer 6

Positive ions only.

Question 7

what is a molecule?

Answer 7

2 or more atoms bonded together.

Question 8

If a question asks you to identify the species responsible for a peak on a mass spectrum, you would write:

Answer 8

e.g 35Cl+

Question 9

in mass spectroscopy, how are does a molecule become a positively charged ion?

Answer 9

The molecule is bombarded with electrons.

Question 10

what is an atomic orbital?

Answer 10

a region of space around the nucleus that can hold up to 2 electrons with opposite spins.

Question 11

Main features of the s-orbital

Answer 11

• each shell contains 1
• spherical in shape
• each s subshell holds 2 electrons

Question 12

Main features of the p-orbital

Answer 12

• 3 in every shell from 2nd shell
• dumbbell shape
• each p subshell holds 6 electrons

Question 13

d-orbital

Answer 13

• each shell from 3rd shell contains 5

- D orbitals can hold 10 electrons.

Question 14

f-orbital

Answer 14

• each shell from 4th shell contains 7

- F subshells can hold 14 electrons

Question 15

why is 4s filled before 3d?

Answer 15

Sub shells fill in order of increasing energy. 4s has less energy than 3d, so 4s is filled in first. However, when electrons are lost, they are lost from 4s first.

Question 16

how are electron configurations drawn?

Answer 16

• 2 electrons of opposite spin in each orbital.

- each orbital gains 1 electron before the sub shell finishes filling up.

Question 17

why are copper and chromium exceptions to the normal electron structure?

Answer 17

in both cases, is it a more stable arrangement to move 1 electron from the 4s sub shell to ‘fill out’ the 3d sub shell.

Question 18

what do you need to remember about ions when writing electron configuration?

Answer 18

remember that 4s is part of the outermost shell so electrons are lost from there first (even though it was filled up first).

Question 19

when writing electron configuration and electrons filled in 4s and 3d, what do you write first?

Answer 19

even though 4s fills in first, you write it in number order (e.g …3p6 3d9 4s2)

Question 20

What is first ionisation energy?

Answer 20

First ionisation energy is the energy required to remove one electron from each atom in one mole of atoms of an element to form one mole of positive 1+ ions in the gaseous state.

Question 21

what is second ionisation energy?

Answer 21

Second ionisation energy is the energy required to remove one electron from each 1+ ion in one mole of 1+ ions of an element to form one mole of 2+ ions in the gaseous state.

Question 22

what are the factors that affect ionisation energy?

Answer 22

• attraction of the nucleus
• atomic radius (distance of outermost electrons from nucleus)
• shielding effect

Question 23

explain the trend in first ionisation energy down a group.

Answer 23

• although nuclear charge increases
• atomic radius increases (distance between outermost shells and nucleus increases.)
• and electron shielding effect increases
• so the force of attraction between the outermost electrons and nucleus decreases.
• so first ionisation energy decreases.

Question 24

explain the trend in first ionisation energy across a period.

Answer 24

• nuclear charge increases
• atomic radius decreases (number of shells stays the same, number of electrons and nuclear charge increases, force of attraction between them increases).
• shielding effect stays the same (same number of shells).
• force of attraction between nucleus and outermost electrons increases
• first ionisation energy increases.

Question 25

explain the trend in first ionisation energy from end of period to start of new period.

Answer 25

• although nuclear charge increases
• atomic radius increases due to 1 more shell.
• electron shielding effect increases
• force of attraction between outermost electrons and nucleus decreases
• first ionisation energy decreases.

Question 26

why is second ionisation energy always greater than first?

Answer 26

• when the first electron is removed, a positive ion is formed.
• force of attraction between nucleus and remaining electrons increases.
• energy required to remove another electron increases.

Question 27

why does atomic radii decrease across a period?

Answer 27

• nuclear charge increases
• number of shells stay the same
• shielding effect stays the same
• force of attraction between outer electrons and nucleus increases.
• outer electrons are pulled in closer towards the nucleus.
• atomic radii decreases across a period.

Question 28

why is there a dip between group 2 and group 3 in ionisation energy?

Answer 28

In group 3, the 3p sub shell is beginning to fill. Electrons in the 3p subshell are higher in energy than 3s electrons, and are also slightly shielded by the 3s electrons. Hence it is easier to remove an electron from group 3 hence lower first ionisation energy.

Question 29

Why is there a dip between phosphorus and sulfur? (group 5 and 6).

Answer 29

Sulfur has 4 electrons in the 3p sub shell. The fourth electron fully fills the first 3p orbital. The electron pair in this orbital slightly repel each other due to opposite spin, making it easier to remove an electron. Hence first ionisation energy decreases between them.

Question 30

Trend in melting points for Na to Al

Answer 30

Metallic bonding increases as proton number increases. This means that more electrons are released into the sea of electrons, and with decrease in ionic radii, the metallic bonding gets stronger. Hence melting points increase.

Question 31

Why does Si have a greater boiling point than Na to Al?

Answer 31

Si is a giant covalent structure with many covalent bonds between the atoms. These covalent bonds a very high amount of energy to break these bonds, hence they have a very high mp, bp.

Question 32

P, S, Cl, Ar boiling points are low. Why?

Answer 32

Only weak london forces are between the molecules, which only require a little amount of energy to break, so low mp and bp.

Question 33

Why does Sulfur have a higher boiling point than Chlorine?

Answer 33

Sulfur has more electrons, and therefore has stronger London forces between molecules. So, more energy is required to overcome these forces.

Question 34

What is atomic (proton) number?

Answer 34

The number of protons in the nucleus of an atom.

Question 35

What is mass number?

Answer 35

the sum of the number of protons + number of neutrons in the nucleus of an atom.