1. Atomic Structure

An atom is the samllest particle which makes us the chemical structure of elements

An atom consists of a central nucleus

This is made of protons and neutrons

It is surrounded by electrons

In an atom - equal number of protons and electrons

Atomic number - the number of protons in the nucleus

Mass number - sum of protons and neutrons in the nucleus

Mass number with a decimal place, relative atomic mass

Proton; relative mass, one, relative charge +1

Neutron; relative mass, one, relative charge 0

Electron; relative mass, one, relative charge -1

Isotopes - atoms of the same element which have different masses

They have the same number of protons and electrons

They differ in the number of neutrons

Have the same chemical properities

Have the same electron configuration

Have slightly different physical properties

Relative abundance - the ratio of each isoptope occuring naturally on earth

Relative atomic mass = average mass of one atom of an element

                                         1/12 mass of one atom of carbon-12

Calculating A  = ∑(m/z) x (relative intenisty for each isotope)

                                                   total relative intensity

Relative molecular mass = average mass of a molecule

                                       1/12 mass of one atom of carbon

Separates atoms and molecules - according to their mass

Measures the relative masses of isotopes

Finds the relative abundance of isotopes in a sample of an element

Records the mass spectrum

The sample is in gaseous form

Bombared with a stream of electrons from an electron gun

These electrons knock an electron off the atom

This forms a positivie ion

M_(g)              M⁺_(g) + e^-

Positively charged ions are accelerated towards the electromagnet

This is done using a negative electric field

Slit present narrow the beam

The accelerated ions are deflected in the magnetic field

The amount of deflection repends on the mass-to-charge ratio

The smaller the ratio, the greater the deflection

Deflection depends on the mass and the charge of the ion

When the ion reaches the detector, it produces a tiny current

Recorded as one ion, with an m/z value

The current produced is proportional to the abundance of the ions in the sample

The sample is then removed using a spectrum

Description of electron arrangement within an atom

Indicated population and location of electrons

Electrons are arranged in electron shells, and sub-shells

Orbital - region where the electron is most likely to be found

Can hold up to two electrons

Four orbitals, s, p, d, f

Aufbau principle - electrons enter the lowest energy orbital available

They fill the lowest energy level first

Hunds rule - electrons prefer to occupy on their own

Ony pair up when no empty orbitals of the same energy are available

Fill up: 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶4d¹⁰

Chromimum promotes an electron from 4s to 3d

Therefore both orbitals are half filled

1s²2s²2p⁶3s²3p⁶4s¹3d⁵

Copper promotes an electron from 4s to 3d

Therefore its orbital is fully filled

1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰

Atoms and ions that contain the same number of electrons

For example:

Ne - 1s²2s²2p⁶

O2+ - 1s²2s²2p⁶

Na+ - 1s²2s²2p⁶

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms, to form one mole of gaseous ions with a 1+ charge

Measured in kJmol-1

All ionisation energies endothermic - energy required to remove an electron

1. The charge on the nucleus

1. The more positively charged the nuclues, the more strongly electrons are attracted 1. to it

2. The distance of the electron from the nucleus

2.  An electron closer to the nucleus, more attraced than one further away

3. The number of electrons between outer electrons and nucleus

3. Electrons in the inner energy shells shield the outmost electrons from the nucleus' 3. pull

4. Wheter the electron is on its own in an orbital

4. Two electrons in the same orbital will repel each other

Ionisation energy increases across a period

Nuclear charge increased across a period

Magnesium and aluminium are exceptions to this

Aluminium's ionisation energy is lower than magnesium's

The outer electron in aluminium, p sub-level

The outer electron in magnesium, s sub-level

Therefore less energy required to remove it

Phosphorus and sulfur are exceptions to this

Sulfur's ionisation energy is lower than phosphorus'

In phosphorus, electrons in 3p sub-level, unpaired

In sulfur, two electrons in 3p sub-level, paired

Some repulsion between paired electrons

Reduced their attraction to the nucleus

Less energy required to remove one paired electron

Second ionisation energy is the energy required to remove a second electron from each ion in mole mole of gaseous +1 ions, to give gaseous +2 ions

E⁺_(g)                        E²⁺_(g) + e^-