Module 3: Reactive Chemistry

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:

\(\color{lightgreen}{CuCO_3}\)\(\rightarrow\)\(\color{lightblue}{CuO+CO_2}\)

\(\color{lightgreen}{\text{Green: Reactants}}\)

\(\color{lightblue}{\text{Light Blue: Products}}\)

Decomposition by light

  • Some compounds will decompose when exposed to light
  • An example is Silver Nitrate (\(AgNO_3\)):

\(\color{lightgreen}{AgNO_3}\)\(\rightarrow\)\(\color{lightblue}{2Ag + 2NO_2 +O_2 }\)

  • 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:

\(\color{lightgreen}{2C_3 H_{8(g)} +7O_{2(g)}}\)$\rightarrow$\(\color{lightblue}{2C_{(s)} + 2CO_{(g)}+ 2CO_{2(g)} +8H_2 O_{(g)}}\)

  • Some combustion reactions only have \(CO_2\) and $H_2 O$ as products
    • These are known as “complete combustion reactions”
    • An example of a complete combustion reaction is burning Methane:

\(\color{lightgreen}{CH_4 +2O_2}\)$\rightarrow$\(\color{lightblue}{CO_2 +2H_2 O}\)

Precipitation Reactions

  • When soluble ionic compounds are dissolved in water, the lattice “dissolves”, and the ions are separated
    Image of the dissolution of an ionic compound
  • 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:

\(\color{lightgreen}{Na_2 S_{(aq)}+CuSO_{4(aq)}}\)\(\rightarrow\)\(\color{lightblue}{CuS_{(s)}+Na_2SO_{4(aq)}}\)

Solubility Rules

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

IonSoluble?Exceptions
$NO_{3}^-$
$ClO_{4}^-$
$Cl^-$$Ag, Hg_2 , Pb$
$I^-$$Ag, Hg_2 ,Pb$
$SO_{4}^{2-}$$Ca, Ba, Sr, Hg, Pb, Ag$
$CO_{3}^{2-}$Alkalis and Ammonium
$PO_4 ^{3-}$Alkalis and Ammonium
$OH^-$Alkalis, $Ca, Ba, Sr$
$S^{2-}$Alkalis, Alkaline Earths, Ammonium
$Na^+$
$NH_4 ^+$
$K^+$

NAGSAG and PMS (Mnemonics)

  • NAGSAG can be used to remember common soluble ions:

N - Nitrates ($NO_{3}^-$)

A - Acetates ($C_2 H_3 O_2 ^- $)

G - Group 1 ($Li^+ , Na^+ , K^+ , etc. $)

S - Sulfates ($SO_{4}^{2-}$)

A - Ammonium ($NH_4 ^+$)

G - Group 17 ($F^- , Cl^- , Br^- , I^- , etc.$)

  • PMS can be used to remember exceptions:

P - $Pb^{2+}$ (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:

\(\color{lightgreen}{4Fe+3O_2}\)\(\rightarrow\)\(\color{lightblue}{2Fe_2 O_3}\)

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:

\(\color{lightgreen}{\text{Acid}+\text{Base}}\)\(\rightarrow\)\(\color{lightblue}{\text{Salt}+H_2 O}\)

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:

\(\color{lightgreen}{\text{Acid}+\text{Metal}}\)\(\rightarrow\)\(\color{lightblue}{\text{Salt}+H_2}\)

Acid-Carbonate Reactions

  • When an acid reacts with a carbonate compound, the products are always \(\ce{CO2}\), \(H_2 O\), and a salt
  • The general formula is:

\(\color{lightgreen}{\text{Acid}+\text{Carbonate Compound}}\)\(\rightarrow\)\(\color{lightblue}{\text{Salt}+H_2 O + CO_2 }\)

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.

$\text{Example: }\ce{FeCl2}$

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

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

  1. When Hydrogen ($\ce{H2}$) is burned in Oxygen ($\ce{O2}$), water ($\ce{H2O}$) is formed

Reactivity Series

Reactivity Series Mnemonic: Please stop calling me a careless zebra, instead try learning how copper saves gold

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

$\color{green}k=Ae^{\frac{-E_{a}}{RT}}$

$k:\text{Reaction Rate Coefficient}$

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

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

$R:\text{Universal Gas Constant}$

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

  • 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.

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References


  1. https://www.britannica.com/science/activation-energy ↩︎

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