A-level Chemistry/OCR/Atomic Structure

The Atom

The word atom comes from the Greek for "indivisible". An atom is the smallest particle of a chemical element that retains the chemical properties of the element. Atoms are composed of subatomic particles, and to understand the behaviour of an atom we must first understand its constituent particles.

Electrons

Electrons are tiny, electrically-charged particles. They have a negative charge, very little mass and they exist in the empty space surrounding the nucleus of the atom which contains all the other particles. In their elemental states atoms are not charged and will have the same number of electrons as they have protons. Electrons can behave as particles and also as waves; this is known as the wave-particle duality of matter. It is only significant for things which are of similar size to atomic particles. Electrons exist in different energy levels or orbitals, filling the lowest energy levels first.

Protons

Protons are much larger than electrons. They are 1836 times heavier than an electron and they have a positive charge equal and opposite to that of an electron.

Neutrons

Neutrons were the last of the nuclear particles to be discovered. They have no charge so they are not deflected by a magnetic field. They are almost the same weight as protons, and normally there are more neutrons than protons in a nucleus of an atom, however they can be equal/

Atomic Numbers

Atoms are normally described in terms of two key numbers, their atomic number and their mass number.

Atomic Number (Z)

The number of protons in the nucleus is the most important aspect of an atom. This number determines which element an atom belongs to. The atomic number of an atom can tell you:

• The number of protons in the nucleus of the atom
• The number of electrons in the atom when it is neutral
• The atom's position in the periodic table

Mass Number (A)

Nearly all of an atom's mass comes from the nucleus. Since we know that the mass of a proton is almost equal to that of a neutron, we can measure the mass of an atom in terms of the number of particles in its nucleus. The mass number can tell you:

• The total number of particles in the nucleus
• The number of neutrons in the nucleus (remember to subtract the atomic number)
• The relative atomic mass of an atom

Summary Table

 Particle Name Relative Mass (unified atomic mass unit) Mass (kg) Relative Charge Electron 1/1836 9.11*10-31 -1 Proton 1 1.67*10-27 +1 Neutron 1 1.67*10-27 0

Isotopes

Isotopes are Atoms with the same atomic number but different mass number

Isotopes are shown like this:

AZXY

Where A is the mass number, Z is the atomic number, Y is the charge on the atom and X is the symbol for that element.

Isotopes Of Hydrogen

For example, hydrogen has three isotopes.

 Name Protium Deuterium Tritium Symbol 11H 21H 31H Protons 1 1 1 Neutrons 0 1 2

Isotope Calculations

It is fairly simple to work out the number of each sub-atomic particle from an isotope's symbol. For a simple example, lets look at a calcium ion:

4020Ca2+

The number of protons is equal to Z.

The number of neutrons is equal to A - Z.

The number of electrons in a neutral atom is equal to Z.

The number of electrons in a positive ion is equal to Z - Y.

The number of electrons in a negative ion is equal to Z + Y.

Using these rules, you should be able to work out that a calcium-40 ion has 20 protons, 20 neutrons and 18 electrons.

Electrons in Atoms

Most of the way that atoms behave is governed by the interactions of their electrons. When you sit on a chair the force that stops you from moving through it is the repulsion between the electrons in the atoms that make up the chair and those that make up your body.

In order to understand chemical reactions, you must understand how electrons exist in atoms. Much about what we know of electrons comes from quantum theory, which states that electrons can be described by four quantum numbers. The only one you need to know about is the principle quantum number, which describes the energy level of an electron.

Energy Levels

In your GCSE chemistry you will have come across the principle quantum number in the form of electron shells. The number is given the symbol n, so that if we say 'the electron is in n = 3' what we mean is 'the electron is in the third shell, or energy level'. Each energy level can accommodate a certain number of electrons.

 n= Maxiumum Electrons 1 2 2 8 3 18

At GCSE, you would have been taught to denote the electronic configuration of that atom like this:

Examples:

Lithium (3): 2,1 Sodium (11): 2,8,1

At AS level we look deeper and divide the shells into a number of subshells. Pairs of electrons are known as atomic orbitals.
s subshells hold 2 electrons (or 1 orbital),
p subshells hold 6 electrons (or 3 orbitals),
d subshells hold 10 electrons (or 5 orbitals),
f subshells hold 14 electrons (or 7 orbitals).

The subshells are arranged in the following order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ....

note that the 4s orbital fills before the 3d!!

This sounds daunting but with practice you will learn the correct order. All you have to remember is the order the subshells come in and the number of electrons each shell holds.

To denote the electronic configuration you simply write out that order, raising the number of electrons in each subshell as a superscript.

So Hydrogen-1 (1) is simply 1s1

Lithium (3) is 1s22s1

Sodium (11) is 1s22s22p63s1

Orbitals also have a paired spin, one counterclockwise and one clockwise. You can denote this by drawing a box with an up and down arrow in it. You must remember, if asked to construct a diagram in an exam, that on each shell (say 2p) you always fill in the 3 orbitals first with arrows going in one direction (say fill in all 3 boxes with up arrows before filling in the down arrows). So if there are only meant to be 2 electrons in the 2p subshell, draw 2 up arrows in 2 separate orbitals and don't draw an up and down arrow in just one orbital.

The best evidence chemists have for the existence of the energy levels comes from ionisation energies.

Ionisation Energy

Atoms are ionised when they lose an electron. The energy required to remove the electron is known as the ionisation energy. As each electron is removed from an atom the ionisation energy required increases, so we call the energy required to remove the first electron the first ionisation energy, the energy required to remove the second electron the second ionisation energy and so on. To be more accurate, the first ionisation energy is the amount of energy needed to remove one mole of electrons from one mole of gaseous, monatomic element. The second ionisation energy is the amount of energy needed to remove one mole of electrons from one mole of gaseous ions each of which bear a single positive charge.

Electron Shells and subshells

see: http://www.bcpl.net/~kdrews/mtas/modern.html which will give plenty of info on terms and theorys