Atomic Structure

Atoms are made up of protons, neutrons and electrons.

Proton - Relative mass 1, Relative charge +1

Nuetron - Relative mass 1, Relative charge 0

Electron -  Relative mass <1, Relative charge -1

Mass number is the total number of protons and neutrons in the nucleus of the atom.

Atomic number is the total number of protons in the nucleus and therfore the number of electrons if the atom is stable.

Ions have different numbers of protons and electrons leaving them with a charge.

Positively charged ions have fewer elcetrons than protons.

Negatively charged ions have more electrons than protons.

Isotopes are atoms of the same element with different number of neutrons.

Isotopes all have the same number of protons.

Relative masses are masses of atoms compared to Carbon-12.

Relative Atom Mass - is the average mass of an atom compared to the mass of an atom of Carbon-12.

Relative Isotopic Mass - is the mass of an atom of an isotope of an element compared to the mass of an atom of Carbon-12.

Relative Molecular Mass - is the average mass of a molecule or formular unit on a scale where an atom of Carbon-12 is 12.

Relative masses can be measured using a mass spectrometer.

Vaporisation: The sample is vapourised into a gas by heating.

Ionisation: Gas particles are bombarded with high-energy electrons from an electron gun. Electrons are knocked out of the vaporised atoms leaving behind positive ions.

Acceleration: Positive ions are accelerated using an electric field.

Deflection: The positive ions' paths are altered by chnaging the strength of the magnetic field. Lighter ions are deflected more than the heavier ions. The larger the mass/charge ration, the more deflection.

Detection: The magnetic field strength is slowly increased so ions of increasing mass reach the detector plate. The ions hit the plate, lose their charge and gain electrons. A current is produced which is proportional to the abundance. From the strength of the magnetic field the computer works out the mass to charge ratio and produces a mass spectrum.

The Y axis gives the abundance of ions, often as a percentage.

The X axis gives the mass/charge ratio.

(mass to charge ration x abundace) / total abundance

Electron shells consist of sub-shells and orbitals.

In the currently accepted model, electrons have fixed energies.

Each shell is given a number called the principal quantum number. The further the shell is from the nucleus, the higher the energy, the larger the principal number.

Not all the electrons have exactly the same energy.

Shells are divded into sub-shells that have slightly different energies. Each sub-shell can hold a maximum of two electrons.

S = 2     P = 6     D = 10     F = 14

Electron configurations can be worked out by filling the lowest energy shell up first.

EXCEPT the 4S energy shell has a lower energy level than the 3D sub-shell and therefore fills up first.

Electrons fill orbitals singley before they start to share.

Transition metals behave unusually.

Chronium and Copper donate one of their 4S electrons to the 3D sub-shell. This is because they are more stable with a full or half full D sub-shell.

Cr - 1s2 2s2 2p6 3s2 3p6 3d5 4s1     Cu - 1s2 2s2 2p6 3s2 3p6 3d10 4s1

When they become ions they lose their 4S electrons before their 3D electrons.

The outer number of electrons decides the chemical properties.

S block atoms lose their outer elctrons easily and form ions with an inert gas configuration.

The elemnets in groups 5, 6 and 7 and gain 1, 2 or 3 electrons to form negative ions with an inert gas configuration. Groups 4 to 7 can share electrons when covalent bonds are formed.

Group 0 have full outer shells and are stable.

Ionisation is the removal of one or more electrons.

The first inoisation energy if the enrgy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

ALWAYS use gas state symbols!

Lower the ionisation energy, the easier it is to form an ion.

Nuclear charge: The more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction is for the electrons.

Distance between the nucleus and the outer electrons: The closer the electron is to the nucleus, the stronger the attraction, the harder it is to form an ion.

Shielding: As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nucleus masking it easier to lose electrons.

High ionisation energy means there is a higher attraction between the electrons and the nucleus.

Ionbisation energy decreases down group two.

Each element down group two has an extra electron shell compared to the one above. The extra inner shell shields the outer electrons from the attraction of the nucleus.

The extra sheilding reduces the attraction making it easier to lose electrons.

Inonisation increases across a period.

As you move across a period, the general trend is for ionisation energies to increase.

The number of protons is increasing which means there is a stronger nuclear attraction. All the extra electrons are at roughly the same energy level, even if the outer electrons are in idfferent orbital types.

This means there's generally little extra shielding effect or extra distance to lessen the attraction of the nucleus.

However.... there are drops between groups 2 & 3 and 5 & 6.

The drop between group 2 and 3 shows the sub-shell structure.

Aluminium's outer electron is in a 3p orbital rather than a 3s. The 3p orbital has a slightly higher energy than the 3s orbital so the electron is found further from the nucleus. This also adds additional shielding.

These two factors override the eefct of the nuclear charge causing the ionisation energy to drop.

The drop between group 5 and 6 is due to electron repulsion.

The shielding is identical in the phosphorus and sulphur atoms and the electron is being removed from an identical orbital.

In phosphorus, the electron is being removed from a singly-occupied orbital.

In sulphur, the electron is being removed from an orbital containing two electrons.

The repulsion between two electrons in an orbital means that electrons are easier to remove from shared orbitals.