Topic 2 Atomic Structure

Subatomic particles

• Atoms consist of a positively charged, tiny, dense nucleus, made up of nucleons - positive protons and neutral neutrons.

• Negatively charged electrons surround the nucleus in energy levels called shells.

• Atoms are neutral as they contain the same number of positive protons and negative electrons.

Mass number, atomic number

Mass number = number of protons + neutrons

Atomic number = no. of protons

Isotopes

→ Isotopes are atoms of the same element that have the same number of protons but different number of neutrons.

• They have similar chemical properties, as they still have the same electron configuration.

• They have different properties (e.g. density and melting point) because they have different masses due to different no. of neutrons.

Relative atomic mass (Ar) + how to calculate it

The weighted average of the different isotopes of an element relative to 1/12th the mass of a carbon-12 atom.

Find percentage composition of an isotope

Subshells

= division of main energy levels

• Electrons are arranged in energy levels called electron shells, which have subshells.

• s, p, d, f are subshells with sucessively higher energies

• EXCEPTION: 4s has lower energy than 3d

• (number of energy level (n) = number of types of subshells, e.g. in the 3rd energy level there are 3 types of subshells, s, p, d→ notation = 3s, 3p, 3d)

Orbitals

1. An orbital is a fixed region of space where there is a high probability of finding an electron.

2. Each subshell consists of a fixed number of orbitals

• s=1, p=3, d=5, f=7 orbitals

3. Each orbital can hold 2 electrons with opposite spins.

Shapes of orbitals

s orbitals = spherical

p orbitals = dumbbell-shaped (can be Px, Py, Pz)

Max no. of electrons that each main energy level can hold

Rules and principles to follow when deducing the electon configuration

1. Aufbau Principle: electrons fill lower energy orbitals first(enter lowest energy orbital available)

2. Hund’s Rule: electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy levels are available

3. Pauli exclusion principle: Maximum 2 electrons per orbital, which must have opposite spins

Condensed electron configuration

→ replace part of full electron configuration by the previous noble gas, written in square brackets.

e.g. oxygen: 1s²2s²2p⁴ → [He] 2s² 2p⁴

The electromagnetic spectrum: relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum

High frequency → high energy, short wavelength

Low frequency → low energy, long wavelength

frequency is directly proportional to energy (since E = h x f, where h is the plank’s constant)

frequency is inversely proportional to wavelength (since c = f x λ , where c is speed, gcse physics equation)

Emission spectrum (plural = spectra)

The pattern of lights emitted that have different wavelengths, produced by atoms emitting photons ( releasing a small packet of electromagnetic energy) when electrons in excited states return to lower energy levels.

Difference between continuous spectrum and line spectrum

Continuous spectrum shows all the wavelengths/frequencies of visible light from red to violet.

(emission) Line spectrum only shows specific wavelengths or frequencies of light, which are shown as coloured lines on a black background.

How are emission spectra produced?

Electrons move/transition between energy levels when they absorb or emit energy, which is in the form of photons, a small packet of electromagnetic energy.

• If an electron absorbs an exact discrete amount of energy (possibly due to the atom being heated) , it transitions from a lower to a higher energy level, e.g. from n=2 to n=3

• The electron is now in an excited state, which is unstable relative to the ground state (lowest-energy state of an atom).

• Therefore the electron emits the same amount of energy it absorbed and transitions back to n= 2.

• The amount of energy emitted by the electron in the transition corresponds to the same wavelengths of visible light, e.g. red light.

• Thus a red line will appear on the emission line spectrum.

Hydrogen emission spectrum

→ shows 4 different coloured, discrete lines at specific wavelengths on a blackground, that are not evenly spaced and converge at higher energies

What does the hydrogen emission spectrum provide evidence for?

• Each electron transition corresponds to a specific line in the emission spectrum, and there are only 4 lines at specific wavelengths instead of a continouous spectrum showing all wavelengths.

  • → This suggests that electrons can only exist at certain energy levels.

• The lines are not evenly spaced and get closer to each other

  • → The energy levels converge at higher energies.

  • at n= infinity, the electron is completely removed from the atom and is no longer under its influence, it is ionised. (ionisation energy = the amount of energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of positive ions)

The complete hydrogen emission spectrum: electron transitions from higher energy levels to n= 1, n=2, n=3

1. electron transitions from higher energy levels to n=1 : → emit energy that corresponds to (the wavelength of) UV radiation = highest energy transition

2. electron transitions from higher energy levels to n=2 : → emit energy that corresponds to (the wavelength of) visible light

3. electron transitions from higher energy levels to n=3 : → emit energy that corresponds to (the wavelength of) infrared radiation = lowest energy transition