Chemistry - EDEXCEL IGCSE - Physical Chemistry

PHYSICAL CHEMISTRY

Acids, Alkalis and Salts:

Universal indicator colours:

Red - acidic

Orange - weak acid

Green - neutral 

Blue - weak alkaline

Purple - alkaline

Acid - contains an excess of H⁺

Alkalis - contains an excess of OH^- ions

Metal + acid —> salt + hydrogen

Metal oxide + acid —> salt + water

Metal carbonate + acid —> salt + carbon dioxide + water

General rules for the solubility of salts in water:

• All common sodium, potassium and ammonium salts are soluble. 
• All nitrates are soluble
• Common chlorides except silver chloride and lead chloride are soluble
• Common sulphates except barium, lead and calcium sulphates are soluble
• Common carbonates except sodium, potassium and ammonium carbonates are insoluble

Preparing soluble salts:

1. Ready an acid with an excess of a solid metal (until there is some left at the bottom). Filter and head unless there are magnesium or carbonates. Place in an evaporating basin to form crystals.
2. Titration - only use this method if the metal is soluble in acid. Place 25cm³ of metal hydroxide in a conical flask using a pipette and add some phenolphthalein. Add the appropriate acid using a burette until the solution turns colourless. Repeat the experiment without phenolphthalein, heat gently to make the solution more concentrated and leave to cool to form crystals.

Preparing an insoluble salt:

Precipitate reactions - mix two solutions, one with the correct positive ion and the other with the correct negative ion. Filter out the spectator ions and water. Wash the residue with pure distilled water to remove any remaining spectator ions. After obtaining the purer residue of the salt, leave the salt to dry in a warm place.

Energetics:

Exothermic - a reaction where heat energy is given out

Endothermic - a reaction where heat energy is taken in

Measure the temperature at the beginning of the experiment and at the end. The amount of heat energy that has been released or taken in is the calorimetry of the reaction.

∆H = the energy released or taken per mole of the substance

∆H = temperature change * mass * heat capacity (J/g) / number of moles

Activation energy - energy required to start the reaction 

Breaking bonds require energy, this takes in heat and so is endothermic.

Making of bonds releases energy, this releases heat and so is exothermic

Calculating enthalpy change:

C-H = 413kJ/mol O=O = 495kJ/mol C=O = 803kJ/mol H-O - 463kJ/mol

CH₄ + 2O₂ —> CO₂ + 2H₂O

4*C-H + 2*O=O —> 2*C=O + 4*H-O

4*413 + 2*497 - 2*803 - 4*463 = -812kJ

This means that the reaction is exothermic. If the result was positive then it would be endothermic.

This also means that 812kJ energy is released for every mole of methane

Rates of reaction:

Factors affecting rate of reaction in terms of particle collision theory

1. Surface area - higher surface area means particles collide more frequently as there is more contact between the reactants.
2. Concentration / pressure - a higher concentration or pressure will mean there is a higher chance of particles colliding as there are more particles
3. Temperature - particles receive more kinetic energy meaning they move faster and so collide more frequently. 
4. Catalyst - provides an alternative pathway which requires a lower activation energy.

Equilibria:

Some reactions happen both ways with reactants becoming products and products becoming reactants. The symbol ⇌ is used to show this.

e.g. CuSO₄.5H₂O ⇌ CuSO₄ + 5H₂O

       NH₄Cl ⇌ HCl + NH₃

Dynamic Equilibrium is when in a reversible reaction, the forward and reverse reactions are at the same rate 

Le Chatelier’s Principle states that a reaction will move in such a way as to minimise the effect of a change to the system in order to return to its dynamic equilibrium

Because of this, for a reversible endothermic reaction, increasing the temperature will cause more products to be produced as a forward reaction occurs. However, for an exothermic reversible reaction when the temperature is increased fewer products will be made. This is because if fewer products are produced in an exothermic reaction then less heat is produced, this balances out the increased temperature. 

e.g. H₂ + I₂ ⇌ 2HI     ∆H = -10kJ (so exothermic)

Increasing the temperature here favours the endothermic reaction so more HI decomposes while when the temperature is decreased more HI will be formed.

The same principle applies to changing the pressure.

e.g. N₂ + 3H₂ ⇌ 2NH₃

      4 moles of gas - 2 moles of gas

So under high pressure, more ammonia will form and less nitrogen and hydrogen. The opposite happens when the pressure is decreased.