Atomic structure

Proton

• relative charge: +1
• relative mass: 1

Neutron

• relative charge: 0
• relative mass: 1

Electron

• relative charge: -1
• relative mass: 1/1840

Protons and neutrons are in the centre, held together by the strong nuclear force

The electrons are around the nucleus in orbit, attracted to the protons in the nucleus by electrostatic forces

1913 - Niels Bohr

• tiny positive nucleus surrounded by negatively-charged electrons to form an atom like a solar system
• electrons orbited in shells of fixed size and movement from one shell to the next explained how atoms absorbed and gave out light

1926 - Erwin Schrödinger - worked out an equation that used the idea that electrons had some of the properties of waves as well as those of particles

1932 - James Chadwick - discovered the neutron

Gilbert Lewis

• inertness of noble gases was relate to their having full outer shells
• ions were formed by atoms losing/gaining electrons to become stable
• atoms could bond by sharing electrons

Mass number

number of neutrons + number of protons

Atomic number

number of protons

Isotopes

• same number of protons and electrons, different number of neutrons
• different isotopes of same leectrons react, chemically, in the same way
• different isotopes of the same element vary in mass number due to the differing number of neutrons

Ionisation - beam of electrons from 'electron fu' knock out electrons from atoms/molecules so they form positive ions (most end with a +1 charge)

Acceleration - positive ions are attracted towards negatively charged plates and accelerated to a high speed (dependent on mass)

• some ions pass through the slits in the plates, forming them into a beam

Deflection - beam of ions moves into a magnetic field at right angles to its direction of travel

• magnetic field defelcts the beam into an arc of a ircle
• deflection depeds ont he ratio of mass/charge, so heavier ions are deflected less than lighter ones
• stronger magnetic field = greater deflection

Detection - computer works out the mass/charge ratio by detecting the strength of the magnetic field when a particular ion hits the detector

Energy levels - electrons in different shells = different amounts of energy, so can be put on to an energy level diagram

• shells ---> main energy levels, labelled 1, 2, 3, etc.
• main energy levels divided into sub-levels, labelled s, p, d, and f

Quantum mechanics - finds the probability of finding an electron in a given volume of space called an atomic orbital

Atomic orbitals

• electron is a cloud of negative charge, which fills a volume of space called an atomic orbital
• any single atomic orbital can hold up to 2 electrons, but p-orbitals always come in groups of 3, for example, so they can hold up to 6 electrons in total

Spin

• two electrons in same orbital must have opposite spins
• electrons usually represented by arrows pointing up or down to show different direction of spin

Ionisation energy

• the amount of energy required to remove a mole of electrons from a mole of atoms, in a gaseous state, measured in kJ/mol
• first electron requires least amount of energy - being removed from neutral atom
• second electron requires more energy - being removed from 1+ ion
• third is more energy, etc.

Trends

• generally increase across a period due to higher nuclear charge
• Magnesium to aluminium - drops due to outer electron in aluminium being in a higher energy orbital than magnesiums outer electron
• Phosphorus to sulfur - drops due to one of the 3p orbitals in sulfur containing 2 electrons which repulse each other makes it easier to remove one of them
• decreases down a group due to increase in shielding and atomic radii