AQA Chemistry AS Unit 1

HideShow resource information

1 Atomic Structure

1.1 Fundamental Particles

  • The subatomic particles are the protons, electrons and neutrons
  • Protons and neutrons are called nucleons as they form the nucleus
  • The electrons surround the nucleus
  • The relative mass of a neutron is 1 and of a proton is 1 whereas an electron  is 1/1836
  • The relative charge of protons is 1+, electrons is 1- and neutrons is 0
  • In a neutral atom the number of electrons is the same as protons

The Arrangement of the Sub-Atomic Particles

  • The protons and neutrons are held together due to strong nuclear force much stronger than electrostatic forces that hold electrons and protons together in the atom
  • This means it overcomes the repulsion between protons in the nucleus
  • The nucleus is surrounded by electrons which are found in a series of levels sometimes referred to as orbitals
1 of 15

1.2 The Arrangement of the Electrons

  • 1913: Niels Bohr said atoms consist of a positive central nucleus orbited by negatively charged electrons, the electrons orbit in shells of fixed size
  • 1932: James Chadwick discovered the neutron
  • Gilbert Lewis said the inertness of noble gases is due to their full outer energy level of electrons, ions were formed by atoms losing or gaining electrons to attain a full outer energy level and that atoms could bond by sharing electrons

Evolving Ideas

  • Dalton's model used to explain crystal geometries
  • Bohr's model used for simple model of ionic/covalent bonding
  • Charge cloud idea is used for sophisticated explanation of bonding and shapes of molecules
2 of 15

1.3 Mass number, atomic number and isotopes

Mass Number and Atomic Number

  • Atomic number: number of protons in one atom of an element
  • Mass number: number of protons and neutrons in the nucleus of an atom of an element
  • Relative atomic mass: Average mass of an atom of an element/ 1 twelve mass of the atom of carbon-12
  • Relative formula mass: Average mass of a formua of compound/ 1 twelve mass of an atom of carbon-12
  • Isotope: Atom with the same number of protons and electrons but different number of neutrons so different mass number
    • Isotopes have the same chemical properties as they have the same number of protons and electron arrangement but different physical properties as they are dependent on mass
3 of 15

1.4 The Mass Spectrometer

  • Determines the relative atomic mass
  • The mass spectrometer determines the mass of separate atoms

What Happens in a Mass Spectrometer?

  • The instrument is kept under high vacuum so ions don't collide with air molecules
  • A vapourised sample is always used
  • Ionisation: High energy electrons fired from an electron gun, it heads towards the atoms in the sample and knocks off an electron from the highest energy level to form positive ions
  • Acceleration: positive ions are accelerated by an electric field focusing them into a beam that passes through slits in negatively charged plates
  • Deflection: positive ions are deflected due to the magnetic field created by an electromagnet, the deflection depends on the m/z, the higher the m/z the lower the deflection. Deflection also depends on magnetic field strength as the greater it is the greater the deflection
  • Detection: Ions strike a detector, accept electrons, lose their charge and create a current which is amplified and sent to a computer to produce a mass spectrum
4 of 15

The Mass Spectrometer Continued

Mass Spectra of Element

  • Mass Spectrometer can be used to identify isotopes
  • Isotopes are separately detected as they have different masses
  • High resolution mass spectrometry can detect relative atomic mass up to 5 decimal places
  • Most work is done to nearest whole number: low resolution mass spectrometry

Low Resolution Mass Spectrometry

  • Relative Atomic Mass=
    • (mass of isotope X abundance)+(mass of isotope X % abundance)/ 100% 

Identifying Elements

  • Cl2 has ionised molecules called molecular ions:
  • Cl-35Cl-35,  Cl-35Cl-37,   Cl-37Cl-37 with m/z values of 70,72,74 respectively
5 of 15

1.5 Electron Arrangements in Atoms

Energy Levels

  • The main energy levels are divided into sub levels called s,p,d and f
  • In the first energy level the electrons spend more time close to the nucleus

Atomic Orbitals

  • Electron is considered a cloud of negative charge
  • An electron fills a volume in space called atomic orbital
  • Different sublevels have different energies
  • Sublevels are made up of orbitals that contain a maximum of two electrons of opposite spin
  • Orbitals of lower energy are filled first, and within it the sub levels of lower energy are filled
  • Orbitals of the same energy fill singly before pairing starts as electrons repel each other

Electronic Structures

  • For Cr and Cu the 3d sublevel is filled before 4s unlike other elements
6 of 15

1.6 Electron arrangements and Ionisation Energy

Ionisation Energy

  • Electrons can be removed from atoms by hitting atoms with a beam of electrons from an electron gun
  • Energy needed to remove a mole of electron from one mole of atoms in the gaseous state is the ionisation energy as the atom becomes a positive ion

Removing the Electrons one by one

  • The first electron removed needs the least energy to remove it since from a neutral atom
  • Second needs more as it is from a 1+ ion, the third needs even more as 2+
  • These are called successive ionisation energies
  • 1st Ionisation Enthalpy: enthalpy change associated with X(g) -> X+(g) + e-
  • It is in the highest energy level which is furthest from the positive nucleus
7 of 15

Electron arrangements and Ionisation Energy Cont

  • 1st Enthalpy Change is evidence for electron structure
  • Highest energy electrons more attracted to nucleus due to stronger positive nucleus due to more protons: Nuclear Charge
    • Boron's ionisation energy is lower than Berilium since its highest energy electron is removed from 2p sublevel not 2s like berylium, this requires less energy as 2p is higher in energy than 2s: this happens with magnesiu and aluminium
    • Oxygen's ionisation energy is lower than Nitrogen since its highest energy electron is in 2p4 sublevel with 2 electrons spin paired in one orbital which will repel each other so are less attracted to the nucleus. Less energy is required to remove 2p4 than 2p3: This happens with P and S
8 of 15

Electron arrangements and Ionisation Energy Cont

Trends in Ionisation Energies across a period in the periodic table

  • Ionisation generally increase across a period as nuclear charge (increasing number of protons) is increasing making it difficult to remove an electron
  • Position of highly energy electrons is the 3rd energy level for all period 3 elements
  • Shielding from positive attraction of nucleus to the highest energy electrons is constant since the same energy level is filling

Trends in Ionisation Energies down a group in the periodic table

  • Ionisation decreases down the group and the highest energy electron becomes easier to remove
  • Down the group the highest energy electron is in a higher energy level so spends less time close to the nucleus
  • Shielding from the positive nucleus to the highest energy level increases and the attraction to the nucleus of the highest energy electron is less strong
9 of 15

2 Amount of Substance

Constructing Formulae for Ionic Compounds

  • Negative ions: OH-,Cl-,NO3-,O2-,SO43-, CO32-
  • Positive ions: Na+,NH4+,Ca2+,Cu2+,Mg2+,Al3+

Relative Atomic Mass and Molecular Mass

  • Ar= average mass of one atom of an element/ 1 twelve mass of atom of carbon-12
  • Mr=average mass of one molecule/ 1 twelve mass of atom of carbon-12

The Avogadro Constant

  • Definition: the number of atoms in 12g of carbon-12

The Mole

  • It is 6.022 X 10^23 particles
  • To find the mole you have mass/relative atomic mass
10 of 15

2.2 The Ideal Gas Equation


  • P is pressure and is measured in Pa (pascals)
  • V is volume and is measured in M3 (metres cubed)
  • n is the moles
  • T is temperature and is measured in K (Kelvin): converted from degrees centigrad by +273
  • R value is 8.31
11 of 15

2.3 Empirical and Molecular Formulae

Empirical: Simplified whole number ratio of atoms of each  element in a molecule/compound

Molecular: Identifies the number of atoms of each element in a molecule

Empirical Strategy

  • Write out the masses
  • Work out the moles
  • Divide by the smallest number
  • Show whole number ratio

Molecular Strategy

  • Mr= Empirical formula mass X unknown
12 of 15

2.4 Moles in Solutions


  • Concentration= Amount of Substance/Volume of Solvent
  • Measurement is moldm-3

Mass Equation

Mass = Relative atomic mass X Mole

13 of 15

2.5 Balanced Equations

  • Reactants form products
  • Atoms are never created or destroyed so equations need to be balanced
  • Must always check using the correct formula
  • Can only change number of atoms by putting number (coefficient) in front of formulae
  • The coeffiecient shows the moles of substances reacting
  • With ionic equations the charges must be the same

Working out Amounts

  • To convert m3 to cm3 you multiply by 10^6

Using Titrations

  • Need to know the concentration of the acid and equation of the reaction
  • Steps:
    • Fill burette with acid of unknown concentration, accurately measure amount of alkali using with pipette, add alkali to conical flask with indicator. Run acid into flask till colour change then repeat
14 of 15

2.6 Atom Economies and Percentage Yields

  • Atom economy= mass of desired product/total reactants X 100
  • Factors considered: rate of reaction, percentage yield, cost of raw materials, price of produce, energy demands, pollution and health and safety
  • Yield= mole of specified product/theoretical mole of product X 100
15 of 15



Nice and concise and well divided up- thank you so much!


very helpful thank you!!!

Similar Chemistry resources:

See all Chemistry resources »See all resources »