I.3. Notes

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Atoms are made up of protons, neutrons and electrons. All elements are made up of atoms. Electrons have -1 charge and they go around the nucleus in orbitals. The nucleus makes up most of the mass of the atom. The diameter of the nucleus is small though compared to the whole atom. The nucleus is where you find the protons (which have a +1 charge) and neutrons (no charge. If the relative mass of protons and neutrons is 1, then an electron has negligible mass- about 1/2000.

The nuclear symbol tells you a lot about an element. The letter tells you the name of the element. The little number at the top is the mass number which tells you the total mass of the nucleus, i.e. the number of protons and neutrons. The number beneath it is the atomic number which is the number of protons in the nucleus- this number identifies the element and allows the periodic table to be ordered. All atoms of the same element have the same number of protons For neutral atoms with no overall charge, the number of electrons is the same as the number of protons told by the atomic number. The number of neutrons is the mass number minus the atomic number.

Ions have different numbers of protons and electrons. Negative ions have more electrons than protons. This is because electrons are negatively charged so if you add one to an atom it becomes negative. Positive ions have fewer electrons than protons because when you remove an electron from an atom, the charge of the protons outweighs the charge of the electrons.

Isotopes are atoms of the same element (so they have the same number of protons) with different numbers of neutrons. For example Chlorine-35 and Chlorine-37 both have an atomic number of 17 however Cl-35 has 18 neutrons whereas Cl-37 has 20. it's the number and arrangement of electrons that dcides the chemical properties of an element. Isotopes have the same configurations of electrons so they've go virtually identical chemical properties. Isotopes of an element do have slightly different physical properties though, such as different dnesities, rates of diffusion etc. This is because physical properties tend to depend more on the mass of the atom.

Relative masses are masses of atoms compared to carbon-12 because the actual mass of an atom is too small and inconvenient to use. Relative atomic mass (Ar) is the mass of an atom relative to 1/12th of the mass of a carbon-12 atom. Relative isotopic mass is the mass of an isotope of an element relative to 1/12th of the mass of a carbon-12 atom. The relative molecular mass (Mr) is the mass of a molecule relative to 1/12th of the mass of a carbon-12 atom. Relative molecular mass is similar to relative formula mass which is used for ionic or giant covalent substances: both are calculated in exactly the same way.

Amount of substance is measured using a unit called the mole which has the symbol n. One mole is roughly 6 .10^23 particles: this number is called the Avogadro constant, L. It doesn't matter what the particles are (atoms, molecules etc). Number of moles = number of particles you have / number of particles in a mole.

Molar mass is the mass of one mole. Molar mass is the same as the relative molecular mass. Avogadro's constant is the number of particles for which the weight in grams is the same as the relative molecular mass. Molar mass is measured in g mol^-1. Another important formula is number of moles = mass of substance / molar mass.

The concentration of a solution is how many moles are dissolved per 1 dm^3 of soliution. Remember that 1 dm^3 = 1 cm^3 = 1 litre.The units of concentration are mold dm^-3. The formula is number of moles = concentration . volume but make sure that the volume is in dm^3. A solution that has more moles per dm^3 than another is more concentrated and one that has few moles is more dilute. For really low concentration you end up with very small mol dm^-3 values so it's better to use parts per million (ppm.

The major gases in the atmosphere are normally given as percentages of the volume. But some gases are present in such tiny amounts that its not very convenient to write their quantities like this so they're given in parts per million. For instance, xenon makes up only 0.009 % of the atmosphere. If there's 09 parts of xenon in every one hundred parts of air, you can multiply both quantities by 100 o make the quantity large enough to work with. So you get 0.09 parts per million.

Empirical formula gives the smallest whole number ratio of atomsi na compound. The molecular formula gives the actual number of atoms in a molecule.