1. Equation Writing
1.1. Ionic and Compound Formulae
1.1.1. Cations and Anions
1.1.1.1. Cations are ions that have lost a certain amount of valence electrons in order to have a positive overall charge. Anions are ions that have gained a certain amount of valence electrons in order to have a negative overall charge. Cations and anions are drawn together by electrostatic forces of attraction present between ions of opposite charges, thus forming ionic compounds.
1.1.1.1.1. List of cations to be learned include
1.1.1.1.2. List of anions to be learned include
1.1.2. Ionic Compounds
1.1.2.1. Formed when ions (atoms with positive or negative charges aka electrons gained or lost respectively) attract each other mutually via electrostatic forces of attraction present between ions of differing charges in a process known as ionic bonding.
1.1.2.1.1. Examples of ionic compounds include
1.1.3. Molecules
1.1.3.1. Differs from ionic bonding (involving electron transfer). Formed when atoms share their valence (outermost shell) electrons to form covalent bonds.
1.1.3.1.1. Examples of molecular GASES include
1.1.4. Polyatomic Ions
1.1.4.1. Mixture of ionic compounds and molecules. They are bound together by covalent bonds, and yet adopt a positive or negative overall charge due to electron loss or gain.
1.1.4.1.1. List of polyatomic cations to be learned include
1.1.4.1.2. List of polyatomic anions to be learned include
1.2. Logic to remember ionic formulae
1.2.1. Periodic Table Groups
1.2.1.1. Group I / II / III ions have 1 / 2 / 3 positive charge(s) as they all lose 1 / 2 / 3 valence electron(s) to attain a stable and low-energy noble gas electronic configuration
1.2.1.2. Group V / VI / VII ions have 3 / 2 / 1 negative charge(s) as they all gain 3 / 2 / 1 valence electron(s) to attain a stable and low-energy noble gas electronic configuration
1.2.1.3. Transition metals (located between Groups II and III in the periodic table) can adopt varying ionic charges. Ion charge is usually specified as Copper (II) and Iron (III) having 2 and 3 positive charges respectively. Lead is not a transition metal but still has ionic charge listed after its name (as in Lead (II) = 2 positive charges). Silver and Zinc, while being transition metals, do not have any charge specified as they usually only form ions of 1 particular charge. (Zinc would be Zn 2+ while Silver would be Ag +)
1.2.2. Electronegativity and Oxidisation number
1.2.2.1. For polyatomic ions as there is no clear logic, this is a higher level method, therefore use only if you have confidence, otherwise use plain memory.
1.2.2.2. Any ion more electronegative than the oxygen ion that it is attached to is given an oxidisation number (for example, Nitrogen, being in Group V of the Periodic Table, is given an oxidisation number of +5 rather than -3 as it is less electronegative than oxygen). Since the oxygen ions total up 6- in loxidisation number (3x 2- charges of oxygen), the overall charge for Nitrate ions is - .
1.3. Logic to remember compound formulae
1.3.1. Charge totals
1.3.1.1. The Gases are simple molecular structures. As such, they form covalent bonds and share electrons
1.3.1.2. The ionic compounds are all binary (meaning that they combine one cation with one anion). The sum of the charges always totals to 0, making it electrically-neuteral. Conventionally, they are written cation first followed by anion. If more than 1 polyatomic ion is required to balance the equation, it must be bracketed followed by a subscripted number stating the amount of ions used.
1.3.1.2.1. For example, in magnesium chloride, Mg has a charge of 2+ while chloride has a charge of -, therefore 2 chloride ions are required to balance the compound of magnesium chloride, making the final equation MgCl2. ((2+) + (2*(1-)) = 0)
1.3.1.2.2. For example, in ammonium sulfate, NH4 has a charge of + while SO4 has a charge of 2-, meaning 2 ammonium ions are required to balance the compound of ammonium sulfate, making the final equation (NH4)2SO4. ((2*(1+) + (2-) = 0)
1.4. Writing Chemical Equations
1.4.1. Chemical compounds and standalone elements often react with one another in chemical reactions to form new chemical compounds. This process is represented by chemical equations.
1.4.1.1. An example of a chemical equation: A + BC → AB + C. Element A reacts with ionic compound BC to form ionic compound AB and element C. A + BC is known as the reactants while AB + C is known as the products. Use a Right Arrow (→) to denote changes.
1.4.1.2. All valid chemical equations must have correct chemical formulae for compounds and must be balanced by using the lowest possible whole numbers.
1.4.1.3. Another example of a chemical equation A2B reacts with C to give A and B2C. The equation is as follows: 2(A2B) + C → 4A + B2C. (Brackets are for clarity, in the exam the 2 before A2B can be written as normal whole number while the 2 between A and B must be in subscript. Likewise, the 4 before the A on the products can be whole while the 2 in B2C must be in subscript.) 2 molecules of A2B are required to balance the equation, which explains the addition of a 2 in front of A2B to form 4 A atoms and 2 B atoms to balance the equation.
2. Acids and Bases
2.1. Acids
2.1.1. An acid is a molecule that dissociates in water to form hydrogen ions as the only cations, and other anions. Examples include Hydrochloric Acid (HCl, dissociates into H+ and Cl- ions) and Nitric Acid (HNO3 dissociates into H+ and NO3- ions). Only hydrogen atoms attached to electronegative atoms can dissociate in water.
2.1.2. Stuff with acid in it
2.1.2.1. Citrus fruits (Citric Acid)
2.1.2.2. Stomach (Hydrochloric Acid)
2.1.2.3. Vinegar (Ethanoic Acid)
2.1.2.4. Fizzy drinks (Carbonic Acid)
2.1.2.5. Coca-Cola (Phosphoric Acid)
2.1.3. General properties of Acids
2.1.3.1. Sour
2.1.3.2. Changes moist blue litmus paper red
2.1.3.3. Irritates at low concentrations
2.1.3.4. Burns at high concentrations
2.2. Hydrogen Ion Concentration and its determining factors
2.2.1. Acid Strength
2.2.1.1. Strong acids refers to acids which have all their molecules dissociate in water for the maximum possible amount of hydrogen ions. An example of this is hydrochloric acid.
2.2.1.2. Weak acids refers to acids which only have a fraction of their molecules dissocated in water to form hydrogen ions. An example of this is ethanoic acid.