Atomic Number: 3, Atomic Symbol: Li, Mass Number: 7
Atomic Number (Z)
Number of protons in the nucleus
Mass Number (A)
Total number of protons and neutrons in the atom
Number of neutrons
A - Z
Isotopes
Atoms with the same number of protons, but different numbers of neutrons
Relative isotopic mass
Mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12
Relative atomic mass
Average mass of one atom compared to one twelfth of the mass of one atom of carbon-12
Relative molecular mass
Average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12
Isotopes have similar chemical properties because they have the same electronic structure. They may have slightly varying physical properties because they have different masses.
Mass spectrometer
Can be used to determine all the isotopes present in a sample of an element and to therefore identify elements
Can measure m/z (mass/charge ratio) and abundance for each isotope
Species for a peak in a mass spectrum
24Mg+
25Mg+
26Mg+
Calculating relative atomic mass
1. R.A.M = (isotopic mass x % abundance) / 100
2. R.A.M = (isotopic mass x relative abundance) / total relative abundance
Mass spectra for Cl2 and Br2
Cl has two isotopes Cl35 (75%) and Cl37(25%)
Br has two isotopes Br79 (50%) and Br81(50%)
Peaks in mass spectra for Cl2 and Br2
Cl35Cl35 +
Cl35Cl37 +
Cl37Cl37 +
Br79Br79 +
Br79Br81 +
Br81Br79 +
Br81Br81 +
Sometimes two electrons may be removed from a particle forming a 2+ ion. 24Mg2+ with a 2+ charge would have a m/z of 12
Molecular ion
The peak with the largest m/z, due to the complete molecule and equal to the Mr of the molecule
Mass spectrum for butane
Molecular ion C4H10+
Fragments at 43, 29
First ionisation energy
The energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge
Second ionisation energy
The energy required when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge
Factors affecting ionisation energy
Attraction of the nucleus
Distance of the electrons from the nucleus
Shielding of the attraction of the nucleus
Successive ionisation energies are always larger
Removing the first electron
Positive ion is formed, increasing the attraction on the remaining electrons
The pattern in the first ionisation energy gives us useful information about electronic structure
Helium has the largest first ionisation energy
Going down a group
First ionisation energies decrease
Going across a period
First ionisation energies generally increase
Sodium has a much lower first ionisation energy than Neon
There is a small drop in first ionisation energy from Mg to Al
n
In the first shell closest to the nucleus and has no shielding effects from inner shells
Two electrons of opposite spin in the same orbital
An early model of the atom was the Bohr model (GCSE model) with electrons in spherical orbits. Early models of atomic structure predicted that atoms and ions with noble gas electron arrangements should be stable.
The A-level model
Electrons are arranged on principle energy levels numbered 1,2,3,4.. with sub energy levels labelled s, p, d and f
s holds up to 2 electrons, p holds up to 6 electrons, d holds up to 10 electrons, f holds up to 14 electrons
Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus
Each orbital has its own approximate, three dimensional shape. It is not possible to draw the shape of orbitals precisely.
An atom fills up the sub shells in order of increasing energy (note 3d is higher in energy than 4s and so gets filled after the 4s)
For oxygen 1s2 2s2 2p4
Spin diagrams
An arrow is one electron, the arrows going in the opposite direction represents the different spins of the electrons in the orbital, a box represents one orbital
When filling up sub levels with several orbitals, fill each orbital singly before starting to pair up the electrons