C1 - Atomic structure

Question 1

Define the term relative isotopic mass based 12c

Answer 1

the mass of the isotope on a scale on which the mass of a carbon 12 atom is exactly 12 units

Question 2

Describe the structure of an atom in terms of electrons, protons and neutrons

Answer 2

Protons and neutrons are located in the nucleus. Electrons are located in energy levels surrounding the nucleus.

Question 3

Describe the relative mass and relative charge of protons, neutrons and electrons

Answer 3

Protons have a relative mass of 1 and a relative charge of +1
Neutrons have a relative mass of 1 and a relative charge of 1
Electrons have a relative mass of 1/1836 and a relative charge of -1

Question 4

Describe what is meant by the terms ‘atomic (proton) number’ and ‘mass number’

Answer 4

Atomic (proton) number is the number of protons in the nucleus of an atom. Mass (nucleon) number is the total number of protons and neutrons in the nucleus of an atom

Question 5

Describe how to determine the number of each type of sub-atomic particle in an atom, molecule or ion from the atomic (proton) number and mass number

Answer 5

number of protons = atomic number.
number of electrons = atomic number.
number of neutrons = mass number - atomic number.

Question 6

Define the term ‘isotopes’

Answer 6

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

Question 7

Define the term ‘relative isotopic mass’ based on the 12c scale

Answer 7

The mass of the isotope on a scale on which the mass of a carbon-12 atom is exactly 12 units.

Question 8

Define the term ‘relative atomic mass’, based on the 12c scale

Answer 8

The weighted average of the masses of the isotopes on a scale on which the mass of a carbon-12 atom is exactly 12 units.

Question 9

Define the term ‘relative molecular mass’

Answer 9

The ratio of the average mass of one molecule of an element or compound to one twelfth of the mass of an atom of carbon-12.

Question 10

Define the term ‘relative formula mass’,

Answer 10

The relative formula mass of a substance made up of molecules is the sum of the relative atomic masses of the atoms in the numbers shown in the formula . Relative formula mass has the symbol, M r.

Question 11

Describe how to calculate relative formula mass from relative atomic masses

Answer 11

Add up all of the relative atomic masses () of all elements in the formula.

Question 12

What is produced by a mass spectrometer?

Answer 12

A mass spectrometer produces a mass spectrum which shows lines at m/z where ions of that mass are present.

Question 13

What does m/z represent on a mass spectrum?

Answer 13

mass / charge ratio

Question 14

Describe how to identify the molecular ion on a mass spectrum

Answer 14

In the mass spectrum, the heaviest ion (the one with the greatest m/z value) is likely to be the molecular ion.

Question 15

What is shown by the relative heights of the peaks on a mass spectrum?

Answer 15

The relative heights of the peaks on the mass spectrum show the relative abundance of the different ions present.

Question 16

Describe how to calculate the Ar of an element from a mass spectrum

Answer 16

∑ (isotope mass x isotope abundance) of all isotopes / 100

Question 17

Predict the mass spectra, including relative peak heights, for a diatomic chlorine molecule

Answer 17

Chlorine has two isotopes 35 and 37. The mass spectrum for a chlorine molecule will have peaks at m/z 35 and 37 for the two Cl+ ions.There will be three peaks for the possible Cl2+ ions - 35 + 35 = 70, 35 + 37 = 72, 37 + 37 = 74. The relative heights of the 70, 72 and 74 lines are in the ratio 9:6:1 due to the ratio of the Cl35 and Cl37 isotopes

Question 18

Describe how mass spectrometry can be used to determine the relative molecular mass of a molecule

Answer 18

The m/z value for the heaviest ion, the molecular ion (m+), is the relative molecular mass of the molecule

Question 19

Define the term ‘first ionisation energy’

Answer 19

The energy involved in removing one mole of electrons from one mole of atoms in the gaseous state.

Question 20

Define the term ‘successive ionisation energies’

Answer 20

The energy that is required to remove the electron one after the other. Successive ionization energy will depend upon the number of electrons present in the outermost shell.

Question 21

Describe how ionisation energies are influenced by the number of protons, the electron shielding and the electron sub-shell from which the electron is removed

Answer 21

Nuclear charge increases with increasing atomic number, which means that there are greater attractive forces between the nucleus and outer electrons, so more energy is required to overcome these attractive forces when removing an electron.
Electrons in shells that are further away from the nucleus are less attracted to the nucleus so the further the outer electron shell is from the nucleus, the lower the ionisation energy.
The shielding effect is when the electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge so the greater the shielding of outer electrons by inner electron shells, the lower the ionisation energy.
Spin-pair repulsion: paired electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals; this makes it easier to remove an electron (which is why the first ionization energy is always the lowest).

Question 22

Explain why there is a general increase in first ionisation energy across a period

Answer 22

Across a period the nuclear charge increases, the distance between the nucleus and outer electron remains reasonably constant and the shielding by inner shell electrons remains the same.
There is a slight decrease in 1st I.E. between beryllium and boron as the fifth electron in boron is in the 2p subshell which is further away from the nucleus than the 2s subshell of beryllium. There is a slight decrease in 1st I.E. between nitrogen and oxygen due to spin-pair repulsion in the 2p subshell of oxygen.

Question 23

Explain why first ionisation energy decrease down a group

Answer 23

Although going down a group the nuclear charge increases, the ionisation energy down a group decreases and it is due to the following factors:
The distance between the nucleus and outer electron increases, the shielding by inner shell electrons increases and so the effective nuclear charge is decreasing as shielding increases.

Question 24

Describe how atomic emission spectra provide evidence for the existence of quantum shells

Answer 24

Spectral lines give evidence of electrons moving from one energy level to another within the atom.

Question 25

Explain how successive ionisation energies provide evidence for the existence of quantum shells and the group to which the element belongs

Answer 25

Successive ionisations of an atom suggest that there are energy shells with large energy differences between them.

Question 26

Explain how the first ionisation energy of successive elements provides evidence for electron sub-shells

Answer 26

There is in general an increase in ionisation energy across a period. This is because as one goes across a period , the number of protons increases making the effective attraction of the nucleus greater. Each successive ionisation energy is bigger than the previous one for the same reason. Some of the increases are much bigger, however, and these big jumps gives us evidence for the main principle electron shells.

Question 27

State the number of electrons that can fill the first four quantum shells

Answer 27

The first shell only has the s subshell ⟹ 2 electrons. The second shell has the s and p subshells ⟹ 2 + 6 = 8 electrons.
The third shell has the s, p, and d subshells ⟹ 2 + 6 + 10 = 18 electrons.
The fourth shell has the s, p, d, and f subshells ⟹ 2 + 6 + 10 + 14 = 32 electrons.

Question 28

What is an orbital?

Answer 28

An orbital is a region within an atom that can hold up to two electrons with opposite spins

Question 29

What are the shapes of s-orbitals?

Answer 29

The s orbital is a spherical shape.

Question 30

What are the shapes of p-orbitals?

Answer 30

A p orbital has the approximate shape of a pair of lobes on opposite sides of the nucleus, or a somewhat dumbbell shape. An electron in a p orbital has equal probability of being in either half.

Question 31

State the number of electrons that occupy s, p and d-subshells

Answer 31

SubShell Electrons
s 2
p 6
d 10
f 14

Question 32

Describe the filling of electron suborbitals

Answer 32

Electrons with similar spin repel each other which is also called spin-pair repulsion. Electrons will therefore occupy separate orbitals in the same subshell where possible, to minimize this repulsion and have their spin in the same direction.
E.g. if there are three electrons in a p subshell, one electron will go into each px, py and pz orbital. Electrons are only paired when there are no more empty orbitals available within a subshell, in which case the spins are the opposite spins to minimize repulsion.

Question 33

Describe standard notation used to indicate electron configurations.

Answer 33

For atoms, the notation consists of a sequence of atomic subshell labels with the number of electrons assigned to each subshell placed as a superscript. e.g, Phosphorus (atomic number 15) is as follows: 1s2 2s2 2p6 3s2 3p3.

Question 34

Descibe how to represent electron configurations using box notation

Answer 34

The electron configuration can be represented using the electrons in boxes notation. Each box represents an atomic orbital. The boxes are arranged in order of increasing energy from bottom to top. The electrons are represented by opposite arrows to show the spin of the electrons

Question 35

Give the subshells in order of increasing energy. What is the exception to this rule?

Answer 35

The subshells increase in energy as follows: s ＜ p ＜ d ＜ f .The only exception to these rules is the 3d orbital which has slightly higher energy than the 4s orbital Because of this, the 4s orbital is filled before the 3d orbital.

Question 36

Describe how ions are formed

Answer 36

Ions are formed when atoms lose or gain electrons. Negative ions are formed by adding electrons to the outer subshell

Question 37

Know that elements can be classified as s, p and d-block elements

Answer 37

Elements are classified as either s,p or d- block dependent upon the obribital their valence electron(s) are in.

Question 38

What determines the chemical properties of an atom?

Answer 38

The specific arrangement of electrons in orbitals.

Question 39

What is periodicity?

Answer 39

Periodicity is a repeating pattern across different periods.

Question 40

What properties are discussed in periodicity?

Answer 40

Various properties such as atomic radius, melting points, boiling points,electronegativity, ionization energy, electron affinity and metallic character.

Question 41

Explain the general trend in ionisation energies across a period

Answer 41

The general trend is for first ionisation energies to increase across a period. An increasing number of protons in the nucleus as you go from sodium across to argon. Increase in the number of protons in the nucleus increases nuclear charge.
That causes greater attraction between the nucleus and the electrons and pulls the electreons closer to the nucleus therefre increasing the ionisation energies.

Question 42

Explain why first ionisation energies fall between magnesium and aluminium and between phosphorous and sulfur.

Answer 42

Despite the increase in nuclear charge, aluminium’s outer electron is in a 3p orbital rather than a 3s.
The 3p electron is slightly more distant from the nucleus than the 3s, and partially screened by the 3s electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.

Question 43

Explain the trend in melting and boiling points across period 3

Answer 43

Melting and boiling points rise across the three metals because of the increasing strength of the metallic bonds. The number of electrons which each atom can contribute to the delocalised “sea of electrons” increases. The atoms also get smaller and have more protons as you go from sodium to magnesium to aluminium.Silicon has high melting and boiling points because it is a giant covalent structure. Phosphorus, sulphur, chlorine and argon are simple molecular substances with only van der Waals attractions between the molecules. Their melting or boiling points will be lower than those of the first four members of the period which have giant structures.

Question 44

Why is argon not included in the trend in electronegativites of period three?

Answer 44

Electronegativity is about the tendency of an atom to attract a bonding pair of electrons. Since argon doesn’t form covalent bonds, you obviously can’t assign it an electronegativity.