Chemical changes

• Metal reactivity - The more reactive a metal is the more vigorous its reactions are. This means it can lose electrons in reactions more easily to form positive ions (cations).
• Metal reactivity series - The comparison of the relative reactivity of different metals with either an acid or water put in order of most reactive to least reactive.
• Metal reactivity series - 1. Potassium (K) 2. Sodium (Na) 3. Lithium (Li) 4. Calcium (Ca) 5. Magnesium (Mg) 6. Carbon (C) 7. Zinc (Zn) 8. Iron (Fe) 9. Hydrogen (H) 10. Copper (Cu) - Metals below carbon on the reactivity series can be extracted from their ores by reduction with carbon, while elements above cannot. Metals above hydrogen will react with acids but metals below will not.
• Acid + metal —> salt + hydrogen e.g. magnesium + hydrochloric acid —> magnsium chloride + hydrogen
• Metal + water —> metal hydroxide + hydrogen e.g. sodium + water --> sodium hydroxide + hydrogen
• A metal’s reactivity when reacting with an acid can be determined in a test tube by the amount of bubbles of hydrogen gas given off.
• Oxidation - Gain of oxygen
• Metal oxides are found in the ground, often as ores that metals need to be extracted from.
• Oxidation - Loss of oxygen
• Some metals can be extracted from their ores chemically by the reduction of carbon e.g. iron oxide + carbon —> iron + carbon dioxide. This can only happen to metals below carbon (less reactive than itself) in the reactivity series (Zinc, Iron, copper). Metals highe than carbon have to be extracted from their ores by electrolysis. Some really unrewctive metals are found in the earth itself in its pure form.
• Oxidation - Loss of electrons
• Reduction - gain of electrons
• A redox reactions - When reduction and oxidation happens at the same time in redox reactions
• Redox reactions - Metals + acids - Iron + sulfuric acid makes the i atoms lose 2 electrons to become positive ions (oxidised) and the hydrogen in the sulfruic acid gain 2 electrons to become just hydrogen atoms
• Redox reactions - Halogen displacement reactions - Chlorine + Potassium bromide. Chlorine will displace bromine so --> Potassium Chloride + Bromine is the product. The chlorine gains electrons (reduction) + Bromine loses electrons (oxidation)
• Half equations - shows oxidation and reduction. E.g. Ca(s)→Ca2+(aq)+2e− (oxidation) and Fe2+(aq)+2e−→Fe(s) (reduction)
• Acids - Form acidic solutions in water, Produce hydrogen H+ ions in aqueous solutions, pH less than 7
• Alkali - water-soluble bases that produce hydroxide ions in solutions OH-, pH greater than 7
• Base - Any substance that reacts with an acid to form salt + water
• All alkalis are bases, but not all bases are alkalis.
• Neutralisation - Acid + Base --> Salt + Water, pH of 7, H+ + OH- --> H20
• Titrations -Allows you to find the exact volume of acid (or alkali) needed to neutralise a measured volume of alkali (or acid)
• Titrations - 1) Use a pipette and a pipette filler and add a set volume of alkali to a conical flask (2) Add two or three drops of indicator (3) Use a funnel to fill a burette with acid of a known concentrations and record the initial volume of acid (3) Use the burette to add acid drop-by-drop to the alkali and regularly swirl the flask (4) The indicator will permanently change colour once neutralised (5) Record the final volume of acid in the burette and use the initial volume to calculate the volume of acid needed to neutralise the solution (6) Repeat steps until you get concordant titres.
• Titration equipment - Pipette, burette, acid, conical flask, alkali
• Titration indicator - Single indicators need to be used as they have a sudden colour change e.g. Phenolphthalein - Colourless in acid --> Pink in Alkalis e.g. Litmus solution - Red in acid --> Blue in alkali e.g. Methyl orange --> red in acid --> yellow in alkalis
• Ionization


  The process of breaking an atom or molecule into ions
• Strong acids - Fully ionise in aqueous equations e.g Hydrochloric acid - HCl(aq)→H+(aq)+Cl−(aq) e.g. Sulfuric acid - H2​SO4​(aq)→H+(aq)+HSO4−​(aq) e.g Nitric acid - HNO3​(aq)→H+(aq)+NO3−​(aq)
• Weak acids - Partially ionise in aqueous solutions e.g. Carbonic acid - H2​CO3​(aq)⇌H+(aq)+HCO3−​(aq) e.g. ethanoic acid, citric acid
• pH scale - A measure of the concentration of H+ ions in the solution. Strong acids have a lower pH than weak acids for a given concentration because strong acids full ionise producing a greater concentration of hydrogen ions. As the pH scale decreases by one unit the concentration of hydrogen ions times by 10 (or on order of magnitude)
• Hydrochloric acids --> ___ chlorides
• Sulfuric acids --> ___ sulfates~
• Nitric acid --> ___ nitrates
• Ethanoic acids --> ___ ethanoate
• Metal carbonates - When acids react with metal carbonate they make a salt + water + carbon dioxide e.g. Hydrochloric acid + sodium carbonate --> sodium chloride + water + carbon dioxide
• Making soluble salts -

  A) dilute sulfuric acid

  B) boiling

  C) copper oxide

  D) glass rod

  E) blue

  F) dissolving

  G) reacted

  H) unreacted

  I) basin

  J) boiling water

  K) 24 hours

  L) gently pat dry

  
• Electrolysis - Solid ionic compounds cannot conduct electricity because the ions are locked in place and not free to move. When ionic compounds are melted (molten) or dissolved in water, the ions are free to move. These liquids can now conduct electricity. Known as electrolytes.
• Electrolysis - Electrodes - Made of an inert, good conducting material e.g. copper, graphite
• Electrolysis set up

  A) Negative electrode , Cathode, Negative terminal

  B) Positive electrode, Anode, Positive terminal

  C) DC power supply

  
• Metals which are more reactive than carbon need to be extracted from compounds by electrolysis
• Electrolysis of aluminium oxide - 1) Aluminium needs to extracted from its ore (bauxite) by electrolysis (2) Aluminium oxide is mixed with cryolite (lowers its melting point reducing energy and cost) (3) Apply an electrical current to the molten aluminium oxide (4) Al3+ ions are attracted to the cathode where they pick up 3 electrons to form aluminium atoms = Al3+ + 3e- --> Al (reduction) (5) O2- ions are attracted to the anode where they lose 2 electrons to form oxygen atoms which will combine to form O2 molecules = 2O2- --> O2 + 4e- (oxidation)