View Chemistry Data Sheet

Types of Reactions

Decomposition Reactions

• Decomposition reactions involve breaking down one compound into 2 or more simpler substances
• Decomposition is an ENDOTHERMIC reaction, meaning it reuqires heat input
• An example of a decomposition reaction is carbonate decomposition:

$CuC{O}_3$$\to$$CuO+C{O}_2$

$\text{Green: Reactants}$

$\text{Light Blue: Products}$

Decomposition by light

• Some compounds will decompose when exposed to light
• An example is Silver Nitrate ($AgN{O}_3$):

$AgN{O}_3$$\to$$2Ag+2N{O}_2+O_2$

• Light-based decomposition is the basis of film photography

Combustion Reactions

• Combustion reactions occur when something burns
• Combustion reactions are EXOTHERMIC (i.e. light, sound, heat are usually produced)
• Oxygen (or any oxidizer) is always a component of a combustion reaction
• An example of a combustion reaction is burning Propane:

$2C_3{H}_{8(g)}+7O_{2(g)}$$\to$$2C_{(s)}+2C{O}_{(g)}+2C{O}_{2(g)}+8H_2{O}_{(g)}$

• Some combustion reactions only have $C{O}_2$ and $H_{2}O$ as products
  • These are known as “complete combustion reactions”
  • An example of a complete combustion reaction is burning Methane:

$C{H}_4+2O_2$$\to$$C{O}_2+2H_{2}O$

Precipitation Reactions

• When soluble ionic compounds are dissolved in water, the lattice “dissolves”, and the ions are separated

  

• If two solutions are mixed together, it’s really just 4 different ions suspended in water
• However, certain combinations of ions will form an insoluble compound when mixed
• These compounds will form a PRECIPITATE, a small ionic crystal lattice
• This is known as a precipitation reaction
• An example is mixing sodium sulfide and copper sulfate solutions:

$N{a}_2{S}_{(aq)}+CuS{O}_{4(aq)}$$\to$$Cu{S}_{(s)}+N{a}_{2}S{O}_{4(aq)}$

Solubility Rules

The solubility rules are used to determine which compound is the precipitate.

NAGSAG and PMS (Mnemonics)

• NAGSAG can be used to remember common soluble ions:

N - Nitrates ($N{O}_3^{-}$)

A - Acetates ($C_2{H}_3{O}_2^{-}$)

G - Group 1 ($L{i}^{+},N{a}^{+},K^{+},etc.$)

S - Sulfates ($S{O}_4^{2-}$)

A - Ammonium ($N{H}_4^{+}$)

G - Group 17 ($F^{-},C{l}^{-},B{r}^{-},I^{-},etc.$)

• PMS can be used to remember exceptions:

P - $P{b}^{2+}$ (Lead)

M - Mercury

S - Silver

Corrosion Reactions

• Corrosion is a reaction involving a metallic element being converted into a more chemically stable form (e.g. an oxide, hydroxide, or sulfide)
• Combustion and Corrosion are both types of “oxidization reactions”
• Corrosion is EXOTHERMIC, although not as much as combustion
• An example of corrosion is iron rusting:

$4Fe+3O_2$$\to$$2F{e}_2{O}_3$

Acids and Bases

Neutralization Reactions

• When an acid and base are added together, they “neutralise” each other
• This creates water and an ionic compound known as a “salt”
• The general formula for acid-base reactions is:

$\text{Acid}+\text{Base}$$\to$$\text{Salt}+H_{2}O$

Acid-Metal Reactions (Displacement Reactions)

• Many metals will react with Acids to produce a “salt” and Hydrogen gas ($H_2$)
• The general formula for Acid-Metal reactions is:

$\text{Acid}+\text{Metal}$$\to$$\text{Salt}+H_2$

Acid-Carbonate Reactions

• When an acid reacts with a carbonate compound, the products are always $\mathrm{CO}{\phantom{A}}_2$, $H_{2}O$, and a salt
• The general formula is:

$\text{Acid}+\text{Carbonate Compound}$$\to$$\text{Salt}+H_{2}O+C{O}_2$

Redox Reactions

• Redox is short for “Reduction-Oxidization”
• Redox reactions occur between 2 substances, where electrons are LOST by one (the reductant), and GAINED by the other (the oxidant)
• An easy way to remember this is with AN OIL RIG CAT:
  • AN - at the ANode,
  • OIL - Oxidization Involves Loss of electrons
  • RIG - Reduction Involves Gain of electrons
  • CAT - at the CAThode

Rules

1. Metals are always reductants, Metal IONS are always Oxidants
2. Oxygen has an oxidation state of $2-$ (unless in a peroxide)
3. Hydrogen has an oxidation state of 1+ (except in metal hydrides)
4. Free elements have an oxidation state of 0
5. The oxidation state of an ion is equal to it’s charge
6. In compounds, the sum of all oxidation states is 0
7. The halogens (F, Cl, Br and I) typically have an oxidation state of 1- in their ionic compounds. In molecular compounds their oxidation number is typically 1- or 7-.
8. When naming ionic compounds in which variable oxidation states of metal ions are present, the oxidation state is shown in roman numerals.

$\text{Example: }\mathrm{FeCl}{\phantom{A}}_2$

$\text{Iron has an oxidation state of 2+}$

$\text{Therefore, the compound is called Iron }\text{(II)}\text{ Chloride}$

9. When Hydrogen ($\mathrm{H}{\phantom{A}}_2$) is burned in Oxygen ($\mathrm{O}{\phantom{A}}_2$), water ($\mathrm{H}{\phantom{A}}_2\mathrm{O}$) is formed

Reactivity Series

Rates of Reaction

Activation Energy $(E_a)$

• Activation energy is the minimum amount of energy required to initiate a reaction
• Activation energy is measured in $kJ/mol$ ¹

$k=A{e}^{\frac{-E_a}{RT}}$

$k:\text{Reaction Rate Coefficient}$

$A:\text{Frequency Factor of the reaction}$

$e:\text{Euler’s Number (}e\text{ on calculator, approx. 2.7182)}$

$R:\text{Universal Gas Constant}$

$T:\text{Temperature }(K)$

• According to this equation, the rate of reaction increases with temperature
• However, there are some cases where the activation energy is negative, and so higher temperatures DECREASE the rate of reaction

Catalysis

• Catalysis is the process of increasing the rate of a chemical reaction by introducing a catalyst
• A catalyst is any substance that lowers the activation energy of a reaction WITHOUT MODIFYING THE PRODUCTS
• Catalysts are not consumed by the reaction, and do not change the equilibrium constant of the reaction.
  • As a result, catalysts should be written in both the products and reactants of a chemical equation
  • Catalysts which trigger a reaction are known as Activators
• The SI unit for Catalysis is the Katal $(Kat)$

  $1Kat=1mol/s$

• There are 3 main kinds of catalysts:
  • Heterogenous Catalysts are those which exist in a different phase from the reaction being catalyzed. For example, solid catalysts the catalyze a reaction in a mixture of liquids and/or gases are heterogeneous catalysts. Surface area is critical to the functioning of this type of catalyst.
  • Homogenous Catalysts exist in the same phase as the reactants in the chemical reaction. Organometallic catalysts are one type of homogeneous catalyst.
  • Enzymes are protein-based catalysts. They are one type of biocatalyst. Soluble enzymes are homogeneous catalysts, while membrane-bound enzymes are heterogeneous catalysts.

References