Chemistry F321 1.1 Hemal Gangani

1. The Atom
1. Made up of subatomic particles electrons, neutrons and protons. Relative mass and charges of 1,1,1/2000, 1+, 0, 1-, respectively.
2. Mass Number (on top) no. of protons and neutrons in the nucleus of an atom. Atomic Number (on the bottom) numbers of protons.
3. Neutral atoms number of protons and electrons same. For negative ions electrons more than protons and other way round for positive ions.
4. Isotopes: same number of protons but a different number of neutrons. Same chemical properties. Slightly different physical properties i.e. density or rate of diffusion.
2. Atomic Models
1. 1. Thomson's Model: Experiments made him conclude that atoms weren't solid and indivisible. Came up with Plum pudding model- positively charged sphere with negative electrons or "corpuscles" embedded within.
2. 2. Rutherford Model: Fired alpha particles at an extremely thin gold sheet. Most passed straight through and some deflected backward - disproving the plum pudding model. New model, the nuclear model positively charged nucleus where most mass is concentrated and cloud of electrons. Most of the atom found to be empty space.
3. 3. Changes of Rutherfords Model: Discovered protons, explained different charges in the nuclei. Rutherford predicted they'd also be other particles to make the nuclei heavier, these were discovered by James Chadwick - the neutron.
4. 4. Bohr Model: Four main changes: Electrons existed in fixed orbits. Each shell has a fixed energy. When electron more between shells electromagnetic is emitted or absorbed. Radiation has fixed frequency because energy levels are fixed. This explained by Noble Gases were inert and other trends in reactivity.
3. Relative Mass
1. Relative Atomic Mass: the weighted mean mass of an atom of an element compared to the mass of 1/12 of the mass of an atom in carbon 12
2. Relative Isotopic Mass: the mass of an atom of an isotope compared to the mass of 1/12 of the mass of an atom in carbon-12
3. When calculating relative atomic mass: Add: percentage of abundance x mass number, together and divide by 100. On mass spectrum? Multiply each relative isotopic mass by it's abundance and the divide by sum of abundances.
4. Relative molecular mass and formula mass is the same: add up all Ar's in compound. Only difference is formula is for ionic compounds as they aren't made of molecules.
4. The Mole
1. Number of particles in a mole is the Avogadro's Constant (Na) which is 6 x 10 to power of 23.
2. Molar mass is Mr but when a g mol -1 stuck on the end.
3. Number of Moles = mass of substance / molar mass
5. Concentration Calculations
1. For Liquids: Number of Moles = Concentration x Volume (in dm3)
2. For Gases: Number of Mole = Volume (in dm3) / 24
6. Formulas
1. Empirical Formula: Smallest Whole Number Ratio of Atoms in a Compound.
  1. M - Mass M - Moles R- Ratio (MMR)
2. Molecular Formula: Actual Numbers of the atoms in a molecule.
  1. 1. Find the Empirical Mass by adding all Mrs in Empirical Formula. 2. Divide the Molecular Mass by the Empirical Mass. 3. Multiply the empirical formula by the number found from previous step.
3. 7. Calculating Formulas
  1. Empirical Formulas from experimental data: 1. Moles = mass / Mr to find out the amount of moles per product made. 2. Use this to find out the moles of each atom you began with. 3. Write down the ratio of moles. 4. Divide to get a whole number ratio. 5. YOU HAVE AN EMPIRICAL FORMULA.
  2. Empirical Formulas from percentage compositions: Just convert % to grams.
  3. Molecular Formulas from experimental data: 1. Moles = mass / Mr to find out the amount of moles per product made. 2. Use this to find out the moles of each atom you began with. 3. Write down the ratio of moles. 4. Divide to get a whole number ratio. 5. Work out the mass of the empirical formula. 6. Divide the molecular mass by empirical mass. 7. Multiply the number from previous steps by the empirical formula.
8. Balancing Equations
1. You're an A-level student you should be able to do this, nevertheless, you can only add numbers to whole compounds on full equations. In ionic equations balance the charges by adding electrons and make sure their is the same number of atoms on each side.
9. Equations and Calculations
1. Calculating Masses: 1. Write out a balanced question. 2. Work out how many moles of the reactants you have. 3. Use the molar ratio from the balanced equation to work out the number of moles of product that will be formed from this much reactant. 4. Calculate the mass of that many moles of product.
2. Calculating Gas Volumes: 1. Write out a balanced question. 2. Work out how many moles of the reactants you have. 3. Use the molar ratio from the balanced equation to work out the number of moles of product that will be formed from this much reactant. 4. Put these numbers in a gas volume equation.
3. State Symbols: a = solid, aq = aqueous, l=liquid and g = gas.
10. Acids, Bases and Salts
1. Acids are proton donors and Bases are proton acceptors.
2. H+ (aq) + H2O (l) ---> H3O+ (aq) this a hydroxonium ion that has formed.
3. Release of H+ ions equations: H2SO4 (l) + water ---> 2H+ (aq) SO4 2- (aq) & HCl (l) + water ---> H+ + Cl-
4. Acceptance of H+ ions equations: NH3 (aq) + H+(aq) ---> NH4+ (aq) & OH- + H+ ---> H2O(l)
5. General Formulas for different types of reactions: Acid + Base = Salt + Water, Metal Oxide + Acid = Salt + Water, Metal Hydroxide + Acid = Salt + Water, Metal + Acid ---> Metal Salt + Hydrogen & Metal Carbonate + Acid = Metal Salt + Carbon Dioxide + Water
11. Anhydrous and Hydrated Salts
1. Water in lattice is called water of crystallisation. A solid containing water of crystallisation is hydrated. A salt is anhydrous if it doesn't contain any water of crystallisation
2. First you find the mass of water lost by taking the mass of the anhydrous salt away from the mass of the hydrated salt. Then you find the number of moles of water lost. Next find the number of moles of anhydrous salt that's produced. Now you look at the amount of moles of anhydrous salt to moles of water. you scale down to 1:n and round off.
12. Titrations
1. Allows you to find out exactly how much acid is need to neutralise a quantity of alkali. Method: Measure you alkali using a pipette and put it in a flask along with some indicator. Add acid to the alkali using a burette. Add a bit at a time. Swirl to make sure it's all mixed properly.
2. Indications: methyl orange - turns yellow to red when adding acid to alkali. phenolphthalein - turns pink to colourless when adding acid to alkali.
3. When calculating concentrations or volume rearrange the following equation: Number of Moles = Concentration x Volume (in dm3)
13. Oxidation States
1. 14. Redox Reactions
  1. OIL RIG
  2. Oxidising agent gets reduce and a reducing agent gets oxidised.
2. Oxidation states of an element tell you how many electrons it has accepted or donated. Here are the rules: Uncombined elements: 0, Monoatomic ion: same as it's charge. Neutral compound it's 0, Combined Oxygen is -2 but in peroxides it's -1, Combined hydrogen +1 expect hydrides where it's -1 and H2 = 0.
3. Systematic names: oxidation number is the value of attached roman numeral.

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Chemistry F321 1.1 Hemal Gangani

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