Chemistry - Atomic structure

Describe the structure of an atom. [3]

• atoms are mostly empty space surrounding a very small, dense nucleus.

• the nucleus contains protons and neutrons.

• electrons are found in shells in the empty space around the nucleus.

State the charge and mass of protons, neutrons, and electrons. [3]

Particle, Relative Mass, Relative Charge

Proton, 1, +1

Neutron, 1, 0

Electron, 1/1836, -1

State what is meant by atomic number/proton number.

The number of protons in the nucleus of an atom.

State what is meant by mass/nucleon number.

The number of protons and neutrons (nucleons) in the nucleus of an atom.

Describe the distribution of mass and charge within an atom. [2]

• virtually all of an atom’s mass is concentrated in the nucleus, which contains the nucleons (protons and neutrons). Electrons have negligible mass.

• the nucleus carries a net positive charge (because of the presence of protons) , while the space surrounding it carries a negative charge as it is occupied by electrons.

Describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an electric field (two plates of opposite charge) in terms of direction and angle of deflection. [3]

• a beam of protons would be deflected towards the negatively charged plate at a lower angle than the beam of electrons (protons have higher mass).

• a beam of neutrons would remain undeflected.

• a beam of electrons would the deflected towards the negatively charged plate at a high angle as a result of the small mass of electrons.

State and explain qualitatively the variations in atomic radius across a period. [4]

Across a period,

1. charge density/nuclear charge increases.

2. shielding remains the same.

3. stronger forces of attraction between nucleus and outer electrons and electrons are pulled closer.

4. atomic radius decreases.

State and explain qualitatively the variations in ionic radius across a period for cations. [4]

Across a period for cations:

1. atoms lose electrons.

2. charge density to electron ratio increases.

3. stronger forces of attraction between nucleus and outer electrons and electrons are pulled closer.

4. ionic radius decreases.

State and explain qualitatively the variations in ionic radius across a period for anions. [4]

Across a period for anions:

1. atoms gain electrons.

2. inter electron repulsion increases.

3. distance between nucleus and outer electrons increases to decrease repulsion.

4. ionic radius increases.

State and explain qualitatively the variations in atomic radius down a group.[4]

Down a group:

1. shielding increases, more occupied energy levels.

2. distance between nucleus and valence electrons increases.

3. weaker forces of attraction between nucleus and outer electrons.

4. atomic radius increases.

State and explain qualitatively the variations in ionic radius down a group. [2]

Down a group:

1. more occupied energy levels.

2. ionic radius increases.

Define the term isotope in terms of numbers of protons and neutrons. [3]

Atoms of the same element with different number of neutrons but same number of protons. → ! - Start the definition with the word “Atoms”

State what is meant by A and Z in the following notation. [2]

A: Mass/nucleon number

Z: Atomic/proton number

Explain why isotopes of the same element have the same chemical properties. [1]

Same electronic configuration.

Explain why isotopes of the same element have different physical properties, and provide two examples. [2]

Different number of neutrons → different mass numbers → differences in mass and density.

Describe the meaning of a shell and a principal quantum number. [1]

Shell → the main energy level of an electron.

Principal quantum number (n) → the number representing the shell/energy level.

Describe the meaning of a sub-shell. [1]

Subdivisions within a main shell containing orbitals.

Describe the meaning of an orbital. [1]

A region in space in which an electron is most likely to be found.

Describe the meaning of ground state. [1]

The lowest energy arrangement of electrons in an atom or ion.

How many electrons can each orbital hold? [1]

2 electrons.

State the sub-shells, number of orbitals and total number of electrons in each sub-shell. [3]

Sub-shell, Number of orbitals, Total number of electrons in each sub-shell

s, 1, 2e-

p, 3, 6e-

d, 5, 10e-

Describe the order of energy in the sub-shells. [2]

Sub-shells fill in order of increasing energy, with lower-energy sub-shells (closer to the nucleus) filling completely before higher-energy ones.

1s →2s →2p →3s →3p →4s →3d → 4p

Note: The 4s sub-shell has a lower energy level than the 3d sub-shell, so 4s fills before 3d.

Exceptions:

Cr (Chromium) → 1s²2s²2p⁶3s²3p⁶4s¹3d⁵ → 3d is filled before 4s, half-filled, more stable.

Cu (Copper) → 1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰ → 3d is filled before 4s, fully filled, more stable.

Explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion. [3]

1. electrons occupy the lowest available energy levels closest to the nucleus to maximize energy stability.

2. electrons occupy orbitals singly before pairing up, minimizing inter-electron repulsion.

3. when 2 electrons occupy the same orbital, they must have opposite spins to minimize repulsion.

Describe what is meant by a free radical. [1]

A species with one or more unpaired electrons.

Define the term first ionisation energy (IE). [3]

The energy needed to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions.

Construct equations to represent first, second and third ionisation energies and state the term used to describe this sequence. [4]

1. X_(g) → X⁺_(g) + e^-

2. X⁺_(g) → X²⁺_(g) + e^-

3. X²⁺_(g) → X³⁺_(g) + 3e^-

We call this sequence of ionisation energies successive ionisation energies.

Explain the trends in ionisation energies across a period. [5]

Across a period:

1. charge density/nuclear charge increases → Atomic radius decreases.

2. shielding remains the same.

3. stronger forces of attraction between nucleus and valence electrons.

4. outer electrons are held more tightly, requiring more energy to remove them.

5. ionisation energy increases.

Explain the trends in ionisation energies down a group. [5]

Down a group:

1. more occupied energy levels.

2. atomic radius increases.

3. shielding increases.

4. outer electrons are held less tightly, requiring less energy to remove them.

5. ionisation energy decreases.

Explain the variation in successive ionisation energies of an element. [4]

Successive ionisation energies for any given element always increases,

• as each electron is removed, the remaining species becomes more positive, nuclear charge increases,

• stronger forces of attraction between the nucleus and the outer electron,

• more energy is needed to remove electrons.

How to deduce the position of an element in the Periodic Table and the electronic configuration using successive ionisation energy data? [3]

• when plotting successive ionisation energies, when dramatic jumps occur, the electron after was removed from a new principal quantum shell that is closer to the nucleus.

• to determine the period number → number of jumps + 1

• to determine the group number → number of electrons before the first jump

State the factors affecting the ionisation energies (IE) of elements. [4]

1. nuclear charge.

2. atomic/ionic radius.

3. shielding by inner shells and sub-shells.

4. spin-pair repulsion.

Note: Half-filled and fully filled shells are have higher ionisation energies, usually requiring more energy to remove electrons because of their high stability. This appears as an anomaly in patterns.

This is observed in Beryllium and Nitrogen.

Beryllium → complete orbital and close to the nucleus.

Nitrogen → half-filled orbital.

Explain how nuclear charge affects ionisation energy (IE). [4]

1. charge density/nuclear charge increases.

2. stronger forces of attraction between nucleus and valence electrons.

3. outer electrons are held more tightly, requiring more energy to remove them.

4. ionisation energy increases.

Directly proportional

Explain how atomic/ionic radius affects ionisation energy (IE). [5]

1. atomic/ionic radius increases.

2. distance between nucleus and outer electrons increases.

3. weaker forces of attraction between nucleus and valence electrons.

4. outer electrons are held less tightly, requiring less energy to remove them.

5. ionisation energy decreases.

Inversely proportional

Explain how shielding by inner shells and sub-shells affects ionisation energy (IE). [4]

1. more occupied shells and sub-shells.

2. the inner shells/sub-shells shield and protect the outer electrons from the full nuclear charge of the nucleus.

3. shielding increases.

4. outer electrons are held less tightly, requiring less energy to remove them.

5. ionisation energy decreases.

Inversely proportional

1. Explain how spin-pair repulsion affects ionisation energy (IE). [4]

1. electrons occupy orbitals singly before occupying them in pairs.

2. when they do occupy them in pairs, repulsion between the electrons destabilizes the pair.

3. the increased repulsion makes it easier to remove electrons, requiring less energy.

4. ionisation energy decreases.

Inversely proportional