1. Smallest unit of a pure element Compound is two or more different types of atom bonded together in fixed proportion Are bonded together either by sharing two electrons (Covalent) or by electrostatic attraction of ions (Ionic) Atoms have a nucleus containing protons and neutrons surrounding this are orbits containing electrons Atoms Atoms have equal numbers of protons and electrons. When electrons are added or removed atoms turn into ions

2. Ernest Rutherford’s gold foil experiment showed that an atom is mostly empty space with a small, dense, positively-charged nucleus. J.J. Thompson discovered the electron and developed the “plum-pudding” model of the atom. + - + - Positive & negative + - + - + particles spread throughout - + - + entire atom. - Dalton’s model of the atom was a solid sphere of matter that was uniform throughout. The Bohr Model of the atom placed electrons in “planet-like” orbits around the nucleus of an atom. The current, wave-mechanical model of the atom has electrons in “clouds” (orbitals) around the nucleus Atomic History

3. If an atom has more electrons than protons it is called an anion and is negative Ions are atoms with an unequal number of + protons and – electrons. They always have a charge When an ionic compound dissolves two oppositely charged ions are formed When atoms or ions become more negative (gain electrons) it is called reduction. All pure elements have a charge of zero ions Metals always form cations because they give away electrons. When an atom loses electrons it is called oxidation

4. Group number = Valence electrons Group 13 through 18 valence electrons = group # - 10. Eg Group 14 has 4 Valence electrons Period number equals number of shells with electrons in. Shell # e - to fill shell 1 2 2 8 3 8 4 18 Metals always try to give e- away. Non metals try to collect them Periodic Table 7 3 Li Mass Number Atomic Number Protons + Neutrons Protons Lowest energy Metal (M) Non metal (N) M-M = Metallic M-N = Ionic N-N = Covalent Exceptions polyatomic ions eg NH 4 CL

5. Alkali MetalsAlkali Earth MetalsHalogensNoble gases Transition Metals = Most reactive in group = Metalloids Non-metals Metals = Only two liquids at STP = Most electronegative = Least electronegative Giving electrons is easier when you have lots Collecting electrons is easier when you have only a few

6. = First ionization Energy (increases when traveling right and up) = Atomic Radius (increases when traveling left and down) = Electronegativity (increases when traveling right and up) but doesn’t include noble gases Arrows point in the direction of increase

7. = First ionization Energy (increases when traveling right and up) = Atomic Radius (increases when traveling left and down) = Electronegativity (increases when traveling right and up) but doesn’t include noble gases

8. Polar bonds are a type of covalent where there is a big difference in electronegativities Covalent bonding: Each bond contains two electrons shared between two non metal atoms Ionic bonds occur when one atom gives it’s valence electron(s) to another atom making two oppositely charged ions Polar molecules must be unsymmetrical. Examples include NH 3 and H 2 O Bonding Metallic bonds are metal ions floating in a sea of electrons

9. substanceNaCliodinediamondwater bondingioniccovalent polar structureLatticesimple covalentgiant molecularsimple molecular physical properties Melting point Solubility Conduction High melting and boiling point. Hard and brittle. Soluble in water. Conducts electricity when molten or in solution. Low melting point. Soft, shiny black solid which forms a purple vapour. Low solubility in water. Does not conduct electricity. Very high melting and boiling point. Hard and brittle. Insoluble in water. Does not conduct electricity. Medium / low melting and boiling point. usually liquids. Medium solubility in water. Does not conduct electricity. Covalent don’t normally conduct substanceNaCliodinediamondwater Bonding

10. Every mole contains 6.022 x 10 23 Particles. They can be Atoms Ions, Molecules or even electrons Solutions Moles = Concentration x volume (Moles/Liter) (Liters) Moles = Mass in grams GFM Moles are used to make sure that the right number of particles react with each other and so a molar ration is used in balanced equations Moles One Mole of any gas at Standard temperature 273 K and pressure 1 Atmosphere 101.3kPa occupies 22.4 Liters

11. Brønsted-Lowry A base is a proton acceptor. An acid is a proton donor. Arrhenius Definition An acid is a substance which forms H+ ions as the only positive ion in aqueous solution Neutralization needs equal amounts of H + & OH - Always makes water and a salt. Using Titration allows you to work out the Concentration of a an acid or alkali Providing you know 3 of the items below M 1 V 1 = M 2 V 2 pH = -log [H 3 O + ] pH + pOH = 14 [H 3 O + ] x [OH - ]=1 x 10 - 14 Acids

12. Alpha and Beta have charges so are most dangerous inside the body, but Gamma rays can travel straight through you so are most dangerous outside the body Gamma Particles are very very powerful and have no mass since they are a type of light ray. Alpha particles are helium nuclei    Sheet of paper Few mm of aluminium Few cm of lead + + 4 2 He - - Fission reactions split heavy nuclei into smaller ones. Fusion reactions occur when light nuclei combine to form a heavy nucleus and a lot of energy. Beta Particles are high energy electrons and B- particles are positrons Positrons are positive electrons (anti matter) Nuclear

13. Alkane only single bonds between carbon atoms. C n H 2n+2 Every Carbon must have 4 bonds Hydrocarbons with the same molecular formula but different structures are called isomers Alkene has at least one carbon carbon double bond C n H 2n Organic Alkyne has at least one carbon carbon triple bond C n H 2n - 2

14. Organic substitution reactions occur when an alkane and a halogen (Group 17) reacts so that one or more hydrogen atoms on the alkane are replaced with oxygen. Organic addition reactions occur when an alkene or alkyne combine with a halogen to make one product (halide). Esterification occurs when an organic acid and an alcohol react to make water and an ester. Saponification occurs when an ester reacts with a base to make alcohol and a soap. Fermentation reactions occur when yeast catalyze a sugar (C 6 H 12 O 6 ) to make carbon dioxide and ethanol. Polymers are long chains of repeating units called monomers. Polymers form by polymerization reactions. Addition polymerization occurs when unsaturated monomers join in a long polymer chain. nC 2 H 2  (C 2 H 2 )n Condensation polymerization occurs when monomers join to form a polymer by removing water. Water is a product! Natural polymers include starch, cellulose, and proteins. Synthetic polymers include plastics such as nylon, rayon, and polyester.

15. Reactants are on the left side of the reaction arrow and products are on the right. Endothermic reactions absorb heat. The energy value is on the left side of the reaction arrow in a forward reaction. Bond breaking Exothermic reactions release energy and the energy is a product in the reaction. Bond Making Synthesis reactions occur when two or more reactants combine to form a single product. Example: 2H 2 (g) + O 2 (g)  2H 2 O(g) Decomposition reactions occur when a single reactant forms two or more products. Example: CaCO 3 (s)  CaO(s) + CO 2 (g) Single replacement reactions occur when one element replaces another element in a compound. Example: Mg + 2HCl  MgCl 2 + H 2 Double replacement reactions occur when two compounds react to form two new compounds. Example: AgNO 3 + KCl  AgCl + KNO 3 Physical changes do not form new substances. They merely change the appearance of the original material. (The melting of ice) Chemical changes result in the formation of new substances. (The burning of hydrogen gas to produce water vapor) Only coefficients can be changed when balancing chemical equations!

16. Remember many reactions are reversible so both of these diagrams refer to the same reaction. They show you the different paths forward and backward Products Reactants Exothermic Endothermic Notice that the activation energies for the two routes are different ΔG = ΔH -TΔS T must be in Kelvin A reaction is spontaneous if G is less than zero Entropy (chaos)

17. Electrolytic Cells Reactions are non spontaneous Redox reactions require electrical energy to occur Voltaic Cells Reactions are spontaneous Redox reactions produce electrical energy Electrochem All metals want to give away electrons and become Cations. The one highest up Table J gets its way Anions in the electrolyte are attracted towards the anode where they undergo oxidation.

18. -3 -2 -1 0 1 2 3 4 5 Oxidation Reduction Charges ANode OXidation REDuction CAThode OIL RIG xsoesa isdi dsun ac tt ii oo n n of electrons

19. The solubility of dissolved gases reduces as the Temperature increases. And increases as the pressure increases - - Questions often use different masses of water instead of 100g So be careful to check Solubility Solubility is usually measured as the amount of solute that will dissolve in 100g of water at a specific temperature Ionic solids Gases solubility Temperature solubility Temperature unsaturated Super saturated On the line = saturated Dissolved particles increase the boiling point and decrease the freezing point of the solvent. The more particles the more the solvent is effected