Topic 1: Atomic Structure and the Periodic Table

Question 1

Structure of the atom

Answer 1

The structure of an atom is made up of protons, neutrons and electrons

Question 2

What is the relative mass and relative charge of a proton, neutron and electron ?

Answer 2

Protons: Have a relative charge of +1 and a relative mass of 1
Neutrons: Have a relative charge of 0 and a relative mass of 1
Electrons: Have a relative charge of -1 and a relative mass of 1/1840

Question 3

What is meant by atomic number and mass number?

Answer 3

Atomic number: The number of protons in an atom’s nucleus.
Mass number: The total number of protons and neutrons in an atom’s nucleus.

Question 4

How to determine the number of each type of sub-atomic particle in an atom?

Answer 4

Number of protons: The number of protons in an atom is equal to its atomic number.
Number of electrons: In a neutral atom, the number of electrons is equal to the number of protons.
Number of neutrons: The number of neutrons in an atom is equal to its mass number minus its atomic number.

Question 5

What is an isotope?

Answer 5

An isotope is an atom of an element that has the same number of protons but a different number of neutrons

Question 6

What is relative isotopic mass and relative atomic mass?

Answer 6

Relative isotopic mass is the mass of an isotope of an element compared to 1/12th of the mass of a carbon-12 atom.
Relative atomic mass is the average mass of an element’s atoms compared to one-twelfth the mass of a carbon-12 atom

Question 7

What is relative molecular mass and relative formula mass?

Answer 7

Relative molecular mass is the average mass of a molecule compared to one-twelfth of the mass of a carbon-12 atom.
Relative formula mass is the sum of the relative atomic masses of all the atoms in a substance’s formula.

Question 8

How to find relative molecular mass?

Answer 8

The sum of the relative atomic masses of all the atoms in a molecule. For example, to calculate the relative molecular mass of ethanol (C2H5OH), you can add up the relative atomic masses of each atom in the molecule: Mr(C2H5OH) = 46.0.

Question 9

How to find relative formula mass?

Answer 9

To calculate the relative formula mass of a substance, you determine how many atoms of each element are in the chemical formula, find the relative atomic masses of each element on the periodic table and add together the atomic masses.

Question 10

How do you analyse and interpret data from mass spectrometry?

Answer 10

We can determine the relative molecular mass by looking at the peak furthest to the right. The peaks to the left are peaks caused by fragments: For CO2, for example, the peak furthest to the right is due to the CO2 molecular ion. The peak at 28 is due to the CO molecular ion (an oxygen atom has been lost).

Question 11

How do you predict mass spectra of chlorine?

Answer 11

You can predict how a mass spectrum might appear for a given compound.
A compound containing one chlorine atom will therefore have two molecular ion peaks due to the two different isotopes it can contain
35Cl = M+ peak
37Cl = [M+2] peak

Question 12

How can mass spectrometry be used to determine relative molecular mass?

Answer 12

Mass spectrometry can be used to determine the relative molecular mass (RMM) of a compound by finding the peak with the highest mass to charge ratio (m/z) value on the mass spectrum.

Question 13

Define first ionisation energy and successive ionisations energies

Answer 13

First ionization energy is the amount of energy needed to remove the outermost electron from a neutral atom of an element in a gaseous state.
Successive ionization energies are the energies required to remove electrons one by one from a mole of gaseous ions or atoms.

Question 14

How are ionisation energies influenced by protons?

Answer 14

The more protons in the nucleus, the greater the attraction between the nucleus and the electrons. This means more energy is required to remove the electron, resulting in a higher ionization energy.

Question 15

How are ionisation energies influenced by the electron shielding?

Answer 15

Electron shielding affects ionization energy by reducing the nuclear attraction on the outer electrons. The more inner shells of electrons an atom has, the greater the shielding effect, and the lower the ionization energy.

Question 16

How are ionisation energies influenced by the electron sub shell when an electron is removed?

Answer 16

The energy of the orbitals increases from s to d, so the orbitals are filled in this order. For example, in magnesium, the large increase in ionization energy from the second to the third ionization energy indicates that the valence shell must contain only two electrons. This corresponds to moving from the 3s to the full 2p subshell.

Question 17

Why does first ionisation energy increase across a period?

Answer 17

First ionization energy increases across a period on the periodic table because of the increase in nuclear charge and the decrease in atomic radius.
Nuclear charge: As you move across a period, the atomic number increases, which means the number of protons in the nucleus increases. This increases the nuclear charge.
Atomic radius: As you move across a period, the atomic radius decreases slightly. This is because the increased nuclear charge pulls the outer electrons closer to the nucleus.

Question 18

Why does first ionisation energy decrease down a group?

Answer 18

First ionization energy decreases down a group on the periodic table because the distance between the nucleus and the outermost electron increases.
Increasing atomic radius: The atomic radius increases down a group, making the outermost electron farther from the nucleus.
Electron shielding: Shielding increases down a group, reducing the effect of the electrostatic forces of attraction.
Weaker attraction: The increase in distance and shielding outweighs the increase in the number of protons, making the attraction weaker.

Question 19

Process of mass spectrometry

Answer 19

Ionization: A gaseous sample is ionized by knocking electrons off to create positively charged ions
Acceleration: The ions are accelerated so they all have the same kinetic energy
Deflection: The ions are deflected by a magnetic field based on their mass and charge
Detection: The ions hit a negatively charged detector plate, generating a signal proportional to the number of ions

Question 20

How have the ideas of electronic configuration developed?

Answer 20

1. Atomic emission spectra provides evidence for the existence of quantum shells
2. Successive ionisation energies provide evidence for the existence of quantum shells and the group to which the element belongs
3. First ionisation energy of successive elements provides evidence for electron sub-shells

Question 21

How many electrons fill up the first four quantum shells?

Answer 21

2, 8, 18 and 32

Question 22

What is an orbital?

Answer 22

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

Question 23

Shape of an s-orbital?

Answer 23

Spherical

Question 24

Shape of a p-orbital?

Answer 24

Dumbbell

Question 25

How many electrons occupy the s, p and d sub shells?

Answer 25

s = 2 p = 6 d = 10

Question 26

How do electrons fill subshells?

Answer 26

Electrons fill subshells singly, before pairing up. Two electrons within the same orbital must have opposite spins

Question 27

How to predict electronic configuration?

Answer 27

Using 1s notation and electrons within boxes (one arrow going up while the other arrow goes down within a single box)

Question 28

How to classify elements?

Answer 28

Using s, p and d-block elements

Question 29

What determines the chemical properties of an element?

Answer 29

Electronic connfiguration

Question 30

What is periodicity?

Answer 30

A regularly repeating pattern of atomic, physical, and chemical properties with increasing atomic number.

Question 31

As you go along period 2, does the melting point and boiling point increase or decrease?

Answer 31

As you move across period 2 of the periodic table, the melting point of elements first increases and then decreases.

Question 32

As you go along period 3, does the melting point and boiling point increase or decrease?

Answer 32

For the metals in period 3, the melting and boiling points increase as you move from sodium to magnesium to aluminum. This is because the strength of the metallic bonds increases as the number of electrons and the charge on the nuclei increase. However, for the non-metallic elements in period 3, the melting points decrease from phosphorus to argon. This is because the melting points of these elements are low due to the weak van der Waals' forces between the molecules. The strength of these forces depends on the size and number of electrons in the molecules.

Question 33

As you go along period 2, does the ionisation energy increase or decrease?

Answer 33

Ionization energy increases as you move from left to right across period 2 on the periodic table. This is because the nuclear charge increases, which pulls the outer electrons closer to the nucleus. This makes it harder to remove an electron, so more energy is required. In period 2, all the outer electrons are in 2s or 2p orbitals, which are roughly the same distance from the nucleus. The 1s2 electrons screen these orbitals. Neon (Ne) has the highest ionization energy in period 2 because it is on the rightmost side of the period.

Question 34

As you go along period 3, does the ionisation energy increase or decrease?

Answer 34

The ionization energy generally increases as you move across period 3 of the periodic table. The nuclear charge increases as you move across period 3, but the shielding of the outer electrons remains relatively the same. This means that the attraction between the nucleus and the outer electron increases, requiring more energy to remove the electron. However, there are some exceptions to this general trend. Magnesium to aluminum: The ionization energy decreases between these two elements. Phosphorus to sulfur: The ionization energy decreases between these two elements. This is because two of the 3p electrons in sulfur are paired, which reduces the force of their attraction to the nucleus.