1.1 Formulae and Equations

• A reducing agent loses electrons and become oxidised causing something else to be reduced.
• An oxidising agent, gains electrons from another causing it to be oxidised and itself reduced.
• The oxidation states of simple ions are equivalent to their charge.
• If the oxidation number decreases, the element is reduced.
• If the oxidation number increases, the element is oxidised.
• Electronegativity decreases as F, O, Cl, N, Br.
• If two highly electronegative elements are present in a compound, the most electronegative takes negative priority.
• The sum total of oxidation states of in a compound must balance out.
• The charge of an ion is based upon the oxidation states of the elements involved.
• Oxygen is always -2 unless in a peroxide.
• Hydrogen always has an oxidation state of +1 unless attached to a group 1 metal.
• The oxidation state of all elements is zero.
• Acids with three hydrogens are called triprotic acids.
• Acids with two hydrogens are called diprotic acids.
• Acids with one hydrogen are called monoprotic acids.
• Oxidation Number: The number of electrons that needs to be added or removed from an element in order for it to become neutral
• Oxidation:

  Gain of oxygen

  Loss of electrons

  Loss of hydrogen
• Reduction:

  Loss of oxygen

  Gain of electrons

  Gain of hydrogen
• Polyatomic: A compound or ion containing more than one type of atom
• Disproportionation: Simultaneous oxidation and reduction of the same species in a given equation
• Positive ions - cations
• Negative ions - anions
• Sodium - $Na^+$
• Potassium - $K^+$
• Silver - $Ag^+$
• Ammonium - $NH_4^+$
• Calcium - $Ca^{2+}$
• Magnesium - $Mg^{2+}$
• Lead - $Pb^{2 +}$
• Zinc - $Zn^{2+}$
• Copper - $Cu^{2+}$
• Iron (II) - $Fe^{2+}$
• Iron (III) - $Fe^{3+}$
• Aluminium - $Al^{3+}$
• Chloride - $Cl^-$
• Bromide - $Br^-$
• Iodide - $I^-$
• Fluoride - $F^-$
• Nitrate - $NO_3^-$
• Hydroxide - $OH^-$