C2 revision Revision PowerPoint for EDEXCEL Chemistry Unit 2

Periodic Table Mendeleev created the periodic table as we know it in He arranged elements in order of their properties leaving gaps where he thought there should be other elements. This was found to be correct when other elements that were discovered filled these blanks. He arranged them by the following properties: Atomic mass Density Melting Point Formula of the oxide Patterns emerged down the groups and across the periods which confirmed his predictions.

Structure of the atom Atoms are made up of PROTONS, NEUTRONS and ELECTRONS. They are found in the positions shown on the diagram. Electron shells Nucleus containing protons and neutrons ParticleRelative Mass Relative Charge Proton1+1 Neutron10 Electron1/1840 Electrons fill shells from the middle in the order of 2, 8, 8, 18 (how many elements are in each period)

Data from Symbols Top number (usually; always biggest) – MASS NUMBER The number of protons and neutrons Bottom number (usually; always smallest) – ATOMIC NUMBER The number of protons (also the same as electrons) Al Protons = 13 Electrons = 13 Neutrons = = 14 Protons and neutrons are packed together tightly in the nucleus (high density) Electrons are spread out in shells (low density)

Electronic Configuration The shells fill up from the middle. The shells can only fit a certain number of electrons – this is the number of elements in the period ShellPeriodNumber of Electrons 1 st 12 2 nd 28 3rd38

Ionic Bonding Ionic bonds form between METALS and NON-METALS. Ionic bonding involves the transfer of ELECTRONS. Metallic Ions are POSITIVELY charged (ANIONS). Non-metallic elements are NEGATIVELY charged (CATIONS). Loose electron +

Common Ions

Properties of Ionic compounds Conduct electricity when MOLTEN (melted) and in an AQUEOUS SOLUTION (dissolved in water) DO NOT conduct electricity as a SOLID Have high MELTING and BOILING points Usually SOLID at ROOM TEMPERATURE

Solubility and Precipitates Precipitates are SALTS that are formed in chemical reactions that DO NOT DISSOLVE in the solvent used in the reaction.

Ion Tests Ion tests for metals (ANIONS) are usually done by FLAME TESTS. Each ION produces a certain flame colour. Testing for CATIONS is done through chemical testing. 1.Chloride – add nitric acid and silver nitrate, if a white precipitate forms then chloride ions are present. 2.Sulphate – add hydrochloric acid and barium chloride, if a white precipitate forms then sulphate ions are present. 3.Carbonate – add an acid and if the gas produced turns limewater MILKY the carbonate ions were present. (carbonate release carbon dioxide)

Covalent Bonding Covalent bonds are usually between 2 non-metals. They are strong bonds. They involve the sharing of electrons (shown by dot cross diagrams) The electrons in the bond are called a ‘bonding pair’, the electrons not in a bond are called a ‘lone pair’

Properties of Covalent Compounds Simple Covalent Small molecules – oxygen gas, carbon dioxide WEAK Low melting and boiling points due to WEAK bonds between molecules (intermolecular). Giant Covalent Giant molecules STRONG High melting and boiling points due to lots of STRONG bonds in the molecule (intramolecular)

Diamond and Graphite HardSoft Giant StructureLayers

Miscible or Immiscible Miscible substances mix together easily. They are separated by (FRACTIONAL) DISTILLATION or CHROMATOGRAPHY. Immiscible substances do not mix together. They are like oil (low density) and water (high density). They are separated suing a SEPARATING FUNNEL.

Metallic Bonding DELOCALISED Conduct electricity due to the sea of free electrons, we say the electrons are DELOCALISED. Malleable (hammer into shape) due to layers that can slip over each other

Alkali Metals React with water to form a HYDROXIDE (OH - ) More reactive down the group (easier to lose the outside electron) 1 electron in outer shell (+1 ion) Increase in density down the group The atomic radius increases The shielding affect increases

Halogens Fluorine – pale green Chlorine – pale yellow Bromine – brown liquid, orange vapour Iodine – Grey solid, purple vapour 7 electrons in outer shell (-1 ion) Less reactive down the group More reactive Halogens displace lesser reactive halogens

Displacement Reactions In a displacement reaction the more reactive element takes the place of the lesser reactive element. For example Potassium Iodide + Chlorine → Potassium Chloride + Iodine This happens because CHLORINE is more reactive than IODINE

Noble Gases UNREACTIVE. Noble gases are UNREACTIVE. They have a full outer shell of electrons. LOW DENSITY They have a very LOW DENSITY NOBLE GASES Uses of NOBLE GASES: Helium – balloons Neon – Bulbs and lamps Argon – Welding to stop reactions with OXYGEN

Exothermic and Endothermic Exothermic reactions give out heat energy (feel hot) HEAT EXITS THE REACTION Endothermic reactions take in energy from the surroundings (feel cold) HEAT ENTERS THE REACTION Energy is created from making and breaking bonds. Energy is released from making bonds. Energy is used to break bonds.

Rates of Reaction Temperature GAIN ENERGY Hotter = faster; they GAIN ENERGY Concentration High concentration (more particles) = faster Surface Area Big surface area (lots of small pieces) = faster Catalysts Adding a catalyst speeds up the reaction but is never used up (used and then reformed)

Collision Theory For a reaction to occur then molecules must collide together with enough energy (ACTIVATION ENERGY) for the bonds to break. This causes a chemical reaction and new bonds will form between the atoms. If the energy to break the bonds is more than that of the bonds made then the reaction is ENDOTHERMIC. If the energy to break the bonds is less than that to make the bonds then the reaction is EXOTHERMIC.

Relative Atomic Mass and Relative Molecular Mass The relative atomic mass of an atom is the same as the TOP (larger) number (mass number) E.g. Carbon = 12Oxygen = 16 The relative formula mass is the total mass of all atoms in the molecule E.g.Carbon Dioxide (CO 2 ) (1 x Carbon) + (2 x Oxygen) (12) + (2 x 16) = 44

Empirical Formula From the masses of reactants, it is possible to calculate an empirical formula. The empirical formula is the simplest ratio of atoms of each substance in the formula.

Theoretical Yield (what you should make) We can use masses in a reaction to help us calculate the amount of reactant and product. Step 1 - Write out the equation for the reaction. Make sure it is balanced. Step 2 - Work out the relative masses of the substances needed in the calculation. Remember to multiply by the number of molecules that are present. Step 3 - Convert the relative masses into the units in the question. Step 4 - Find the ratio by dividing both numbers by the smallest relative mass. Step 5 - Find the mass of the unknown by multiplying the mass of the known by the ratio of the unknown.

Percentage Yield Percentage Yield = Actual Yield Theoretical Yield The Actual Yield is how much you have made from the reaction (from the question). The Theoretical Yield is how much you should make if you have no loss what so ever (100% efficient). X 100