# AS-CHEM1 - Key Definitions and Equations

A selection of definitions and equations on some small, handy flashcards.

Currently incomplete and so there is more to come!!!

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• Created by: Dan Neal
• Created on: 25-11-12 13:17

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## Relative Atomic Mass

The relative atomic mass, Ar, is the average mass of an atom of an element divided by 1/12th the mass of an atom of carbon-12. (e.g. Hydrogen, H = 1 and Oxygen, O = 16).

The Ar is usually not a whole number because all the naturally occurring isotopes of an element are taken into account, relative to their percentage abundance.

The Ar can be worked out from a mass spectrum like this:

1. Read the percentage abundance of each isotope from the graph and multiply each one by the mass of the isotope.

2. Add up all the totals.

3. Divide by the total of all the percentages (which is 100). This gives the average relative atomic mass of an element sample.

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## Relative Molecular Mass

The relative molecular (or formula) mass, Mr, is the average mass of a molecule divided by 1/12th the mass of an atom of carbon-12. This means the Mr is just the total of all the Ar values added together. (e.g. H2O = (1x2) + 16 = 18)

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## Empirical Formula

The empirical formula is the lowest whole number ratio of atoms of an element in a molecule. (e.g. C4H3O2 = the lowest ratio of 4 carbons to 3 hydrogens to 2 oxygens)

The empirical formula can be calculated like this:

1. Divide the percentage (or mass) of each element by the Ar of that element.

2. Divide each of those numbers by the smallest of them.

3. This will give a rough ratio of each element in the molecule. Make sure each part of the ratio is a whole number and this is the empirical formula.

e.g. 27% Carbon and 73% Oxygen

1. Carbon = 27% ÷ 12 = 2.25 and Oxygen = 73% ÷ 16 = 4.56

2. Carbon = 2.25 ÷ 2.25 = 1 and Oxygen = 4.56 ÷ 2.25 = 2

3. 1 carbon for every 2 oxygens = CO2

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## Molecular Formula

The molecular formula is the actual number of atoms of each element in a molecule. (e.g. C8H6O4 = the actual number of carbon, hydrogen, and oxygen atoms in the molecule)

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## Electronic Configuration - Energy Levels, Sub-Shel

• Electrons are organised in energy levels - higher levels have higher energy.
• Energy levels are split up into sub-shells. These are then split into a certain number of orbitals. Each orbital can hold only two electrons.
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## Ionisation Energies

The first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous unipositive ions.

Ionisation energies generally increase across a period. The three factors that affect ionisation energy are:

• Nuclear Charge of the nucleus - The more protons in the nucleus, the stronger the attraction to the electrons (more protons = higher ionisation energy).
• Energy Level of the electron - An electron in a higher energy level is further from the nucleus and so the attraction is less (higher energy level = lower ionisation energy).
• Electron Shielding - As sub-shells are filled, they create a small amount of shielding that stops the outside electrons being attracted as much to the nucleus.
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## Moles, Mass, Mr Equation

Moles = Mass (grams) ÷ Mr

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## Concentration, Moles, Volume Equation

Concentration (mol dm-3) = Moles ÷ Volume (dm3)

Remember,

1 dm3 = 1 litre = 1000 cm3 = 0.1 m3

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## Ideal Gas Equation

PV = nRT

Where:

• P = Pressure in Pascals (Pa)
• V = Volume in metres cubed (m3)
• n = Number of moles
• R = The Gas Constant which is always 8.31 J K-1 mol-1
• T = Temperature in Kelvin (K)

Remember,

• 101000 Pascals = the usual room pressure (1 atmosphere)
• 1 m3 = 1000 dm3 = 1000000 cm3
• The temperature in Kelvin = the temperature in degrees Celsius + 273
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