UNIT 6
STUDY OF
CHEMICAL REACTIONS
QUÍMICA
1 BATXILLERAT
Study of chemical reactions
Energetic criteria
Exothermic processes and reactions: release
energy to the surroundings.
Endothermic processes and reactions: absorb
energy from the surroundings.
Heat (Q)
where c is the specific heat capacity.
Enthalpy (H)
If a chemical reaction occurs at constant pressure, the energy transferred in the reaction
is called an enthalpy change (Δ
Δ H).
Kinetic criteria: rate of the reaction
Activation energy (Ea): the
minimum energy necessary
to start a reaction.
Factors that affect the rate of reactions
Increase in the concentration.
Increase in the temperature.
Increase the rate of
Increase in the contact surface
reaction.
between the reactants.
The use of catalysts modifies the rate of reactions by lowering Ea. For example: enzymes
or biocatalysts.
Direction of reactions
Complete or irreversible reaction: if the reaction takes place until the limiting reactant is
used up in a closed system.
Reversible reaction: the reaction occurs in both directions.
Forward reaction: the reactants turn into products.
Reverse reaction: the products turn into reactants.
At equilibrium
The macroscopic properties of the system remain constant with time.
At the microscopic scale, reactants and products continue reacting
(dynamic equilibrium).
Acid–base reactions
Properties of acids and bases or alkalis:
Electrolytes
Arrhenius' theory of electrolytic dissociation : there are substances called electrolytes
that, when dissolved in water, dissociate into positive ions (cations) and negative ions
(anions), and give a solution capable of conducting an electric current.
anion
cation
Strong
electrolyte:
substance that is
completely
dissociated.
undissociated
species
anion
Weak electrolyte:
substance that is
partially
dissociated.
cation
Non-electrolyte: substance that cannot dissociate.
According to Arrhenius, there are:
Strong and weak acids.
Strong and weak bases.
Limitations: only applicable to aqueous solutions and to
compounds that contain hydrogen ions or hydroxide ions.
Acids and bases according to Arrhenius
Acid is any electrolyte that, on dissolving
in water, dissociates into hydrogen ions
H+ and the respective anion.
Examples:
Strong acid
Base is any electrolyte that, on dissolving
in water, dissociates into hydroxide ions
OH- and the respective cation.
Examples:
Weak acid
Strong acids: HCl, HClO4, HBr, H2SO 4 and
HNO3.
Strong base
Weak base
Strong bases: NaOH, KOH and Ca(OH)2.
Neutralisation
Neutralisation is the reaction that occurs between the hydroxide ions of a base and the
hydrogen ions of an acid to form water molecules.
H+ + OH– H2O
Acid + base Salt + water
Examples:
Acid–base titrations
Acid–base titration: the process by which the concentration of acid or base in a solution
is determined by its reaction with a base or acid of known concentration.
Titration of a
base with an acid
Titration of an acid
with a base
Precipitation reactions
Identification of Bi3+, Al3+ and Ag+ in a solution:
Redox reactions
Oxidation-reduction (redox) reactions are transformations in which electrons are
exchanged between reactants.
Oxidation: process in which a substance loses electrons.
Reduction: process in which a substance gains electrons.
Example of a redox reaction:
Reducing agent:
Oxidation
half-reaction. Zn
Oxidising agent
Reduction
half-reaction. (oxidant): Cl
Reducing agent: chemical species
that is oxidised, that loses electrons.
Reducing agent product + n e-.
Oxidising agent: chemical species
that is reduced, that gains electrons.
Oxidising agent + n e- product.
Oxidation number
The oxidation number of an atom in a compound is defined as the charge that the atom
would have if the compound was formed of ions.
An increase in the oxidation number of an atom implies oxidation.
A drop in the oxidation number of an atom implies reduction.
To determine the oxidation number of an atom:
① For monoatomic ions, the oxidation number of the atom is equal to the charge on the ion.
② The oxidation number of an element is zero, equally for all its allotropic forms.
③ The oxidation number of oxygen in all its compounds is –2, except in peroxides, in which
it is –1.
④ The oxidation number of hydrogen is +1, except in metal hydrides, in which it is –1.
⑤ In any compound, the oxidation number of an alkali metal is always +1, and that of an
alkaline earth meatl is always +2.
⑥ The oxidation number of each halogen atom in halides is –1.
⑦ For a polyatomic compound, the sum of the oxidation numbers of all the atoms must
equal zero.
⑧ For a polyatomic ion, the sum of the oxidation numbers of all the atoms must equal the
charge on the ion.
Balancing redox reactions
Steps to follow to balance redox reactions:
① Write the chemical equation for the reaction taking into account the dissociations of the
salts, acids and bases present in the reaction medium.
② Determine the oxidation number of all the atoms.
③ Identify the species that is oxidised and the species that is reduced.
④ Write the chemical equation for each of the half-reactions of oxidation and reduction.
⑤ Balance the number of atoms in the half-reaction equations by inserting the necessary
stoichiometric coefficients and following this order:
a) Balance the number of atoms of elements that are oxidised or reduced.
b) Balance the number of oxygen atoms.
c) Balance the number of hydrogen atoms.
⑥ Balance the electric charges, placing the electrons gained or lost in each half-reaction on
the correct side.
⑦Balance the number of electrons gained and lost in each half-reaction, by multiplying each
half-reaction by the number of electrons gained or lost in the other half-reaction.
⑧Add the two half-equations and write the complete redox equation in ionic form.
⑨Write the complete reaction in molecular form, if appropriate.