Module 1: Properties and Structure of Matter

1.1 Properties of Matter

• Categories of matter: pure substance and mixtures

• Pure substance: made of a single chemical species

• Chemical species: can be an element, a molecule or a compound

• Single particles of a pure substance are held together by chemical bonds called intramolecular bonds (including covalent, ionic and metallic bonds)

• All pure substances are homogenous (uniform)

• Mixtures are any combination of more than one pure substance

• Mixtures can be separated by a physical process as they aren’t chemically bonded

• Categories of mixtures: homogenous and heterogenous

• Homogenous mixtures are uniformly mixed

• Any solution of a dissolved solute in a solvent is a homogenous mixture

• Heterogenous mixtures are not uniformly mixed

• Sieving: particle size

• Filtration: solid substance, liquid substance, therefore particle size

• Vaporisation: liquid has a much lower boiling point

• Distillation: big difference in boiling points

• Fractional distillation: significant but small difference in boiling points

• Sedimentation and decantation: density

• Separating funnel: immiscible liquids; different densities

• Adding a solvent then filtering: one substance is soluble in the chosen solvent, while the others are insoluble

• %weight = mass of component/total mass of mixture x 100

• IUPAC naming conventions for ionic compounds: positive ion (cation) first, negative ion (anion) second, anions ending changed to ‘ide’

• IUPAC naming conventions for covalent compounds (simple inorganic): least electromagnetic element first, the most electromagnetic element second, second element ending is changed to ‘ide’ and both elements need a Greek prefix

• Periodicity: relationship between elements’ physical and chemical properties and their position in the periodic table

• Thermal physical properties: boiling point, melting point, heat capacity, heat conductivity

• Mechanical physical properties: hardness, viscosity, density, ductility, brittleness

• Optical physical properties: colour, transparency, reflectivity, refractivity, absorption, lustre

• Electrical physical properties: electrical conductivity, resistance, electric charge

• Chemical properties: acidity and basicity, combustibility, ability to oxidise or reduce

• Periodic Law: the properties of the elements are periodic functions of their atomic number

• Properties of metals: shiny, malleable, good conductors of heat and electricity

• Properties of non-metals: appear dull, poor conductors of heat and electricity

• Properties of metalloids: conduct heat and electricity moderately well, possess some properties of metals and non-metals

1.2 Atomic Structure and Atomic Mass

• Number of protons defines an element

• Mass number = protons + neutrons

• Isotope defined by = number of neutrons

• Radioactive decay is spontaneous and doesn’t require an input of energy to occur

• In nuclear reactions, the nucleus gains stability by undergoing a change of some kind

• Radioisotope = an isotope of an element that is unstable and undergoes radioactive decay

• Isotope notation = mass number → 12 C

                              atomic number → 6  

• All elements with an atomic number greater than 83 are unstable because the nucleus is too large and will decay regardless of the neutron to proton ratio

• Quantum mechanics = all matter can be described as both a wave and a particle

• Electrons are so small the effects of quantum mechanics are significant

• The smaller the object the more its wave-like characteristics come into play

• Two qualities of waves that affect the arrangement of electrons: constructive and deconstructive interference which causes electrons to affect the other electrons around them and that waves are stretched out over space from their peak to their trough, which means the exact position of the electron is unknown

• Highest probability of an electron = at the peak of the wave

• Erwin Schrodinger = developed mathematical equation that uses the wave-like characteristics of electrons to determine where an atom’s electrons are most likely to be

• Energy levels = orbitals

• Groups of orbitals = shells

• Orbitals come in many shapes

• Energy levels of electrons can be directed observed by measuring the wavelength of the photons they emit when excited

• Electron shells subdivided into electron orbitals

• Each orbital only contains 2 electrons

• First shell = 1 s orbital

• Second shell = 1 s orbital, 3 p orbitals

• Third shell = 1 s orbital, 3 p orbitals, 5 d orbitals

• Fourth shell = 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals

• Shells referred to as the principle energy level (n)

• Hund’s rule: states that for an orbital group also called sublevel (like 2p) electrons will fill all orbitals with a single electron before a second electron can be added

• Writing spdf notation: -

  • ordered left to right

  • shell number first, orbital type second

  • number of electrons are represented in superscript after the orbital group (sublevel) eg. 2p^2

• 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 7s, 5f, 6d, 7p

• Using isotopic composition to calculate relative atomic mass: multiply each percentage by the mass number, add them together, divide by 100

• Alpha particles: relatively heavy, positively charged, can be stopped by a piece of paper

• Beta particles: much lighter, negatively charged, can pass through paper and 0.5mm sheet of aluminium, stopped by 0.5mm sheet of lead

• Gamma rays: no charge, extremely penetrating, stopped by 5cm of lead or15cm of concrete

• Alpha particles: are helium nuclei (2 protons and 2 neutrons stuck together)

• Beta particles: are electrons

• Gamma rays: are a type of electromagnetic radiation

1.3 Periodicity

• d

1.4 Bonding

Module 2: Introduction to Quantitative Chemistry

2.1 Chemical Reactions and Stoichiometry

2.2 Mole Concept

• 1mol = 6.022x10^23 particles

• 6.022x10^23 = Avogadro’s number

• Number of particles = mols x Avogadro’s number

• mol=m/Mm

• Percent composition% = mass of element/mass of compound x 100

• Empirical formula

• Limiting reagent

2.3 Concentration and Molarity

• j

2.4 Gas Laws

Module 3: Reactive Chemistry

3.1 Chemical Reactions

3.2 Predicting Reactions of Metals

3.3 Rates of Reactions

Module 4: Drivers of Reactions

4.1 Energy Changes in Chemical Reactions

4.2 Enthalpy and Hess’s Law

• Enthalpy =

4.3 Entropy and Gibbs Free Energy

• Entropy =

• Gibbs free energy =

Module 7: Organic Chemistry

7.1 Nomenclature

7.2 Hydrocarbons

7.3 Products of Reactions Involving Hydrocarbons

7.4 Alcohols

7.5 Reactions of Organic Acids and Bases

7.6 Polymers