Chemistry 3

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Created by pxkemon over 5 years ago
GCSE AQA Chemistry - Unit 1
James Jolliffe
GCSE - AQA: C1.1 The Fundamental Ideas in Chemistry
Olly Okeniyi
GCSE AQA Chemistry 1 Fuels & The Environment
Lilac Potato
AS Chemistry - Enthalpy Changes
Sarah H-V
OP doplnovaci otazky II.
Helen Phamova
chemistry: c2
kristy baker
Chemistry unit 2
36jessieh
Flame tests
Joshua Rees
GCSE AQA Chemistry - Unit 3
James Jolliffe
GCSE AQA Chemistry - Unit 2
James Jolliffe
Chemistry 3
1 Water
1.1 Hard Water
1.1.1 Contains Calcium (Ca^2+) or Magnesium (Mg^2+) Ions
1.1.2 Advantages
1.1.2.1 Dissolved Compounds good for health
1.1.2.2 Calcium helps build strong bones and teeth, as well as helping to reduce the risk of heart disease
1.1.3 Disadvantages
1.1.3.1 More soap needed to form a lather (more expensive)
1.1.3.2 Leads to deposits (scale) on heating systems like kettles, reducing efficiency
1.1.4 Removing hardness
1.1.4.1 To soften hard water, the ions need to be removed
1.1.4.1.1 Method one: Add sodium carbonate solution (washing soda.) The carbonate ions react with the calcium carbonate and magnesium carbonate, which precipitate out of the solution as they are both insoluble
1.1.4.1.2 Method two:Pass the hard water through an ion exchange column which contains a resin which supplies hydrogen or sodium ions. As the hard water passes through the resin, the calcium and magnesium ions in it are replaced by the hydrogen or sodium ions from the resin.
1.1.5 Permanent hard water remains hard upon boiling, whereas temporary hard water is softened upon boiling
1.2 Soft Water
1.2.1 Doesn't contain many dissolved compounds
1.2.2 Readily forms a lather with soap
1.3 Drinking water
1.3.1 Production
1.3.1.1 Passed through filter bed to remove solid particles
1.3.1.1.1 Sedimentation allows small particles to sink to bottom
1.3.1.1.1.1 Chlorine gas added to kill harmful micro-organisms
1.3.1.1.1.1.1 Fluoride added to reduce tooth decay (excessive fluoride can cause tooth discolouration)
2 Periodic table
2.1 Newlands
2.1.1 1864 - Arranged known elements (63) in order of atomic weight
2.1.1.1 Strictly following atomic weight created problems because some elements were placed in the wrong group
2.1.2 Noticed similar properties amongst every eighth element in the series (law of octaves). Makes sense since noble gases (group 8) weren't discovered until 1894
2.1.2.1 Noticed periodicity (repeated patterns) but the missing elements caused problems
2.2 Dmitri Mendeleev
2.2.1 Noticed that not all elements had yet been discovered
2.2.1.1 Left gaps when he created his table in (1869.) Used this gaps to predict existence of other elements
2.3 Modern periodic table
2.3.1 Discovery of sub atomic particles (protons, neutrons, and electrons) and electronic structure in the early 20th century helped build a more accurate table
2.3.1.1 Arranged in terms of electronic structure.
2.3.1.2 Arranged in periods (rows) according to amount of electrons in outer energy level (shell)
2.3.1.2.1 From each period (left to right) an energy shell is gradually filled. In the next period, the next energy level is filled
2.3.1.2.1.1 This arrangement means that electrons with same amount of electrons in their outer shell are in the same group (column
2.4 Groups
2.4.1 Group 1 - Alkali metals
2.4.1.1 Six metals
2.4.1.2 Low melting and boiling points which decreases as you go down the group
2.4.1.3 Low density (lithium, sodium, and potassium are less dense than water
2.4.1.4 More reactive as you go down the group
2.4.1.5 Reactions
2.4.1.5.1 Stored in oil because react vigorously with oxygen and water.
2.4.1.5.2 Forms metal hydroxide ( & hydrogen) when reacted with water
2.4.2 Group 7 - The Halogens
2.4.2.1 Five elements
2.4.2.2 Non-metals
2.4.2.3 melting and boiling points increase as you go down the group
2.4.2.3.1 At room temp chlorine and fluorine are gases, bromine is a liquid
2.4.2.4 Have coloured vapours (chlorine's and bromine's vapours smell particularly strong
2.4.2.5 reactivity decreases as you go down the group
2.4.2.6 Exist as molecules made up of pairs of atoms - diatomic
2.4.2.7 Reactions
2.4.2.7.1 React with metals to produce ionic salts
2.4.2.7.2 Halogen atom gains one electron to produce a halide ion that carries a negative charge

Annotations:

  • i.e. Chloride, bromide, or iodide ion e.g. lithium 2Li(s) + chlorine Cl2(g) = lithium chloride 2LiCl(s)
2.4.2.7.3 React with other non-metallic compounds to produce molecular compounds

Annotations:

  • e.g.  hydrogen H2(g) + chlorine Cl2(g) = hydrogen chloride 2HCl(g)
2.4.3 Transition metals
2.4.3.1 Metallic group of elements between groups 2 and 3
2.4.3.2 Form coloured compounds
2.4.3.3 Have ions with different charges, e.g. Fe^2+ & Fe^3+
2.4.3.4 Can be used as catalysts to speed up chemical reactions
2.4.3.5 good conductors of heat and electricity
2.4.3.6 Easily bent or hammered into shape
2.4.3.7 In comparison to group 1
2.4.3.7.1 Higher densities and melting points, except mercury which is liquid at room temp
2.4.3.7.1.1 Harder and stronger (except mercury)
2.4.3.7.1.1.1 Much less reactive and don't react as vigorously with oxygen or water
3 Energy
3.1 Measure in Joules (J) Takes 4.2 joules of energy to heat up 1g of water by 1ºc
3.2 1 calorie = 4.2 joules
3.3 Information about energy in food products is given in kilocalories (kcal)
3.4 Whenever chemical change happens, so does energy change. Energy is either taken in or given out. Measured using calorimetry
3.4.1 Method
3.4.1.1 Place 100g of water in a calorimeter (a metal or glass container) and measure the temperature of the water
3.4.1.1.1 Find the mass (in grams) of fuel to be burned
3.4.1.1.1.1 Burn the fuel under the water in the calorimeter for a few minutes
3.4.1.1.1.1.1 Record the new temperature and calculate the temperature change
3.4.1.1.1.1.1.1 Weigh the fuel and calculate how much fuel has been used
3.4.1.1.1.1.1.1.1 Energy released (Q) can be calculated by Q = mc ∆T mass of water x 4.2 (constant) x temperature change
3.5 Amount of energy produced in a chemical reaction can be calculated by mixing the reactants in an insulated container.
3.5.1 This enables the temperature change to be measured before heat is lost to the surroundings
3.5.1.1 Suitable for neutralisation reactions and reactions involving solids e.g. zinc and acid
3.6 Energy changes in a chemical reaction can be drawn on an energy level diagram
3.6.1 Energy is lost in exothermic reactions so the products have less energy than the reactants
3.6.2 Energy is taken in for endothermic reactions, meaning that energy is being gained to the products have more energy than the reactants
3.7 Activation energy is the amount of energy needed to start a reaction
3.7.1 Catalysts reduce the activation energy needed for a reaction
3.8 Hydrogen can be used as a fuel in combustion engines and in fuel cells, which produce electricity which can be used to power vehicles
3.8.1 Hydrogen + oxygen = water + heat
4 Analysis and Quantitative Chemistry
4.1 Flame tests can be used to identify metal ions
4.1.1 Lithium, sodium, potassium, calcium, and barium compounds can be recognised by the colours they produce in a flame test
4.1.1.1 Method
4.1.1.1.1 Heat, and then dip a piece of nichrome, (a nickel-chromium alloy) wire in hydrochloric acid to clean it
4.1.1.1.1.1 Dip the wire into the compound
4.1.1.1.1.1.1 Put into Bunsen flame.
4.1.1.1.1.1.1.1 Green = Barium
4.1.1.1.1.1.1.2 Brick red = Calcium
4.1.1.1.1.1.1.3 Crimson red = Lithium
4.1.1.1.1.1.1.4 Lilac = Potassium
4.1.1.1.1.1.1.5 Yellow = Sodium
4.2 Carbonates react with dilute acids to form carbon dioxide gas (and a salt and water)

Annotations:

  • Calcium carbonate CaCo3(s) + hydrochloric acid 2HCl(aq) = calcium chloride CaCl2(aq) + CO2(g) + Water H2O(l)
4.2.1 Carbon dioxide turns limewater milky
4.3 Some metal ions form precipitates i.e. insoluble solids that come out of solution when sodium hydroxide solution is added to them
4.3.1 E.g. when sodium hydroxide is added to calcium chloride solution, a white precipitate of calcium hydroxide is formed (as well as sodium chloride solution)
4.4 If dilute hydrochloric acid and barium chloride solution are added to a solution containing sulfate, a white precipitate of barium sulfate is formed
4.4.1 Titration
4.4.1.1 Used to find out how much of an acid is needed to neutralise an acid
4.4.1.2 When neutralisation takes place, the hydrogen ions H+) from the acid join with the hydroxide ions (OH-) from the alkali to form water (neutral pH)

Annotations:

  • hydrogen ion + hydroxide ion = water molecue
4.4.1.3 Method
4.4.1.3.1 Wash and rinse a pipette with the alkali that you will use
4.4.1.3.1.1 Use the pipette to measure out a known and accurate volume of the alkali
4.4.1.3.1.1.1 Place the alkali in a clean, dry conical flask. Add a suitable indicator e.g. phenolphthalein
4.4.1.3.1.1.1.1 Place the acid in a burette that has been carefully washed and rinsed with the acid.
4.4.1.3.1.1.1.1.1 Take an initial reading of the acid in the burette
4.4.1.3.1.1.1.1.1.1 Add the acid to the alkali until the indicator changes to show neutrality This is called the end point
4.4.1.3.1.1.1.1.1.1.1 Take a reading of the volume of acid in the burette (final reading)
4.4.1.3.1.1.1.1.1.1.1.1 Calculate the volume of acid added by subtracting the final reading from the initial reading
4.5 Different strength acids and alkalis can reacts to form a neutral solution. A suitable indicator must be used in titrations. For example, if you have a strong acid and strong alkali you should use any suitable acid base indicator e.g. litmus
5 Ammonia
5.1 Haber process used to produce ammonia
5.2 Raw materials are nitrogen (from the fractional distillation of liquid air) and hydrogen (from natural gas and steam)