Topic 1

Question 1

What is the atomic number of an element?

Answer 1

The atomic number is the number of protons in the nucleus of an atom. It defines the element.

Question 2

What is the mass number of an atom?

Answer 2

The mass number is the total number of protons and neutrons in the nucleus of an atom.

Question 3

Define isotopes.

Answer 3

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers.

Question 4

What is an ion?

Answer 4

An ion is a charged particle formed when an atom gains or loses one or more electrons.

Question 5

How do you calculate relative atomic mass?

Answer 5

Relative atomic mass is the weighted average mass of the isotopes of an element, taking into account their relative abundance.

Question 6

What does the term ‘electron configuration’ mean?

Answer 6

Electron configuration describes the arrangement of electrons in the orbitals of an atom.

Question 7

What is the electronic configuration of carbon (C)?

Answer 7

The electronic configuration of carbon is 1s² 2s² 2p².

Question 8

What is the shape and bond angle of a molecule with 2 bonding pairs of electrons?

Answer 8

The shape is linear, with a bond angle of 180°.

Question 9

What is the definition of ionisation energy?

Answer 9

Ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Question 10

What is the trend in ionisation energy across a period?

Answer 10

Ionisation energy increases across a period because the nuclear charge increases, attracting electrons more strongly.

Question 11

What is the trend in ionisation energy down a group?

Answer 11

Ionisation energy decreases down a group because the outer electrons are farther from the nucleus and experience more shielding.

Question 12

Define electronegativity.

Answer 12

Electronegativity is the ability of an atom to attract bonding electrons in a covalent bond.

Question 13

What is the trend in electronegativity across a period?

Answer 13

Electronegativity increases across a period as the atomic radius decreases and the nuclear charge increases.

Question 14

What is the trend in electronegativity down a group?

Answer 14

Electronegativity decreases down a group as the atomic radius increases, and the bonding electrons are further from the nucleus.

Question 15

What are the groups in the periodic table?

Answer 15

Groups are vertical columns in the periodic table that contain elements with similar chemical properties. They are numbered 1 to 18.

Question 16

What are periods in the periodic table?

Answer 16

Periods are horizontal rows in the periodic table that represent elements with the same number of electron shells.

Question 17

What is the periodic trend in atomic radius across a period?

Answer 17

Atomic radius decreases across a period as the number of protons increases, pulling electrons closer to the nucleus.

Question 18

What is the periodic trend in atomic radius down a group?

Answer 18

Atomic radius increases down a group as additional electron shells are added, causing the outer electrons to be farther from the nucleus.

Question 19

What is the structure of a simple covalent substance?

Answer 19

Simple covalent substances have discrete molecules held together by covalent bonds, with weak intermolecular forces (Van der Waals forces) between molecules.

Question 20

What is the structure of a giant ionic lattice?

Answer 20

A giant ionic lattice consists of positively and negatively charged ions held together by strong electrostatic forces, forming a regular, repeating structure.

Question 21

What is the difference between a molecular and a giant covalent structure?

Answer 21

Molecular covalent structures consist of discrete molecules, whereas giant covalent structures have a continuous network of atoms bonded by covalent bonds (e.g., diamond, graphite).

Question 22

What are the key properties of metals?

Answer 22

Metals are typically good conductors of electricity and heat, have high melting and boiling points, and are malleable and ductile due to the delocalized electrons in their structure.

Question 23

What are the key properties of non-metals?

Answer 23

Non-metals are usually poor conductors of electricity and heat, have low melting and boiling points, and are brittle in solid form.

Question 24

How do metals bond in a metallic structure?

Answer 24

Metals bond through metallic bonding, where positive metal ions are surrounded by a ‘sea’ of delocalized electrons that allow for electrical conductivity and malleability.

Question 25

Define atomic orbital.

Answer 25

An atomic orbital is a region of space around the nucleus of an atom where there is a high probability of finding an electron.

Question 26

How many orbitals are in the s, p, and d sublevels?

Answer 26

The s sublevel has 1 orbital. The p sublevel has 3 orbitals. The d sublevel has 5 orbitals.

Question 27

What are the main differences between the s, p, d, and f orbitals?

Answer 27

s orbitals are spherical. p orbitals are dumbbell-shaped. d orbitals have a cloverleaf shape. f orbitals are more complex with multiple lobes.

Question 28

What is Hund's rule?

Answer 28

Hund's rule states that electrons will occupy degenerate orbitals singly before pairing up, in order to minimize electron-electron repulsion.

Question 29

What is the Pauli exclusion principle?

Answer 29

The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers. In other words, an orbital can hold a maximum of two electrons, and they must have opposite spins.

Question 30

What is the trend in atomic radius across a period?

Answer 30

Atomic radius decreases across a period due to increased nuclear charge, which pulls the electrons closer to the nucleus.

Question 31

What is the trend in atomic radius down a group?

Answer 31

Atomic radius increases down a group as additional electron shells are added, increasing the distance between the nucleus and the outermost electrons.

Question 32

What are transition metals and where are they located on the periodic table?

Answer 32

Transition metals are elements found in the central block (groups 3-12) of the periodic table. They have incomplete d sublevels and often form colored compounds.

Question 33

What are the general properties of transition metals?

Answer 33

Transition metals have high melting and boiling points, good conductivity, form multiple oxidation states, and often form colored compounds.

Question 34

Why do transition metals form more than one ion?

Answer 34

Transition metals can lose different numbers of electrons from both the 4s and 3d orbitals, allowing them to form multiple oxidation states.

Question 35

What is the trend in electronegativity across a period?

Answer 35

Electronegativity increases across a period because the number of protons increases, attracting the bonding electrons more strongly.

Question 36

What is the trend in electronegativity down a group?

Answer 36

Electronegativity decreases down a group as the atomic radius increases and the outer electrons are further from the nucleus.

Question 37

What is the relationship between ionisation energy and electronegativity?

Answer 37

Both ionisation energy and electronegativity increase across a period due to increased nuclear charge, but ionisation energy also tends to decrease down a group, while electronegativity decreases slightly.

Question 38

What is the first ionisation energy of an element?

Answer 38

The first ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state, forming one mole of positive ions.

Question 39

What happens to the ionisation energy as you move across a period?

Answer 39

Ionisation energy generally increases across a period as the atomic radius decreases and the nuclear charge increases.

Question 40

What happens to the ionisation energy as you move down a group?

Answer 40

Ionisation energy generally decreases down a group due to the increased distance between the outer electron and the nucleus and greater shielding from inner electron shells.

Question 41

Why do noble gases have very high ionisation energies?

Answer 41

Noble gases have full outer electron shells, meaning their electrons are very stable and tightly bound, requiring a lot of energy to remove.

Question 42

What is the relationship between the group number and the number of valence electrons?

Answer 42

The group number of an element corresponds to the number of valence electrons in its outer shell. For example, elements in Group 1 have 1 valence electron.

Question 43

What are the characteristics of Group 1 elements (alkali metals)?

Answer 43

Alkali metals are highly reactive, have low melting and boiling points, are soft, and react vigorously with water to form hydroxides and hydrogen gas.

Question 44

What is the trend in reactivity of alkali metals down the group?

Answer 44

Reactivity increases down Group 1 because the outer electron is more easily lost due to increased atomic radius and electron shielding.

Question 45

What are the characteristics of Group 7 elements (halogens)?

Answer 45

Halogens are highly reactive non-metals, have low melting and boiling points, and form salts when combined with metals. They exist as diatomic molecules (e.g., Cl₂, Br₂).

Question 46

What is the trend in reactivity of halogens down the group?

Answer 46

Reactivity decreases down Group 7 because the ability to attract electrons decreases as the atomic radius increases and shielding increases.