Week 1: Reactions in Aqueous Solution – Oxidation and Reduction (Redox Reactions)
Aqueous Solutions (aq)
Any solution in which H2O is the solvent. (acids must be in aq. solutions)
aq > in an aqueous solution
Solute
- Can be solid, liquid or gas
- Can be ionic or covalent
Solvent
- Can be any liquid o Water (aq); ethanol; hydrocarbons (e.g. hexane)
• Ionic compounds – often form strong electrolytes (100% dissociated) Example:
• Molecular compounds – non-, weak, and strong electrolytes (~3% dissociated, weak electrolyte)
Example:
Aqueous Solutions: Electrical Conductivity
Properties in aqueous solution: non/electrolyte
(aq) – i.e. surrounded by water molecules
• Non-electrolyte – solution does not conduct electricity
• Strong-electrolyte – solution conducts electricity well (ions are free to move – movement of +ve & -
ve charges)
• Weak-electrolyte – solution conducts electricity, but not that well
Types:
Electrolyte: substances that dissociate into ions in water
, lOMoAR cPSD| 23050439
Chemistry 1102: Lecture Notes | Semester 2 – 2018
Non-electrolyte: substances that dissolve in water but DO NOT dissociate into ions
Solute + Solvent
Electrolyte
Strong Weak Non-
electrolyte electrolyte electrolyte
100% Few% 0% dissociated
dissociated dissociated into ions
into ions into ions
Salts dissociate => ions (ionisation)
• Ionic compounds: strong electrolytes
• Molecular compounds: non/weak/strong electrolytes (depending)
Electrical conductivity
• Depends on # of ions in solution o Non-electrolyte: does not conduct electricity o Strong
electrolyte: conducts electricity very well o Weak electrolyte: conducts electricity, but not that well
Exchange Reactions (metathesis/double-displacement reactions)
Reactions in which anions and cations exchange partners
• A+ B- + C+D- A+D- + C+B-
, • If a combination of anion-cation is NOT soluble, it could form (net reaction occurs):
o Precipitate (s) o Weak or
non-electrolyte (aq)
o Gas (g)
Redox Reactions
Reactions where electrons are transferred between species, NOT just an exchange of ions OIL
RIG
Oil – oxidation is loss (loss of electrons)
o Increase in oxidation number
Rig – reduction is gain (of electrons) o
Decrease in oxidation number
, lOMoAR cPSD| 23050439
Chemistry 1102: Lecture Notes | Semester 2 – 2018
Cu atom loses 2 electrons to become a 2x charged cation (Oxidation is loss, so Cu is oxidized)
2 Ag+ cations gain an electron each to become two Ag atoms (Reduction is gain, so Ag + is reduced)
Half reactions/equations
Oxidation step (loss of e-): Cu(s) Cu 2+(aq) + 2e -
Reduction step (gain of e-): 2 Ag + (aq) + 2e - 2 Ag(s)
1. Each half reaction contains same charge on either side of reaction
2. The same # of e- that are lost in oxidation are gained in reduction
*The species that is being oxidised is the reducing agent or reductant because it is reducing the other
species.
*The species that is being reduced is the oxidising agent or oxidant because it is oxidising the other
species.
Why redox reactions?
Essential for energy storage and conversion in biological organisms. (example: respiration)
**example: refer to pg. 1 black book
Determining when a reaction is a redox reaction:
Oxidation Numbers (oxidation states) o Method for assigning electrons to
individual atoms within a molecule o A change in oxidation number =
redox reaction
o Oxidation # is determined by electronegativities (if atoms lusts for electrons it will take
them from something that doesn’t lust for them as much **Learn electronegativity periodic table Rules
for assigning oxidation numbers:
1. Any atom in its elemental state has an oxidation number of zero (no electronegativity difference
between bonded atoms) – e.g. C ; H2; Mg; Cl2 o H2, He2, N2, O2, F2, Cl2 – Br2, I2 0
2. Monatomic ions: the oxidation number is equal to the net charge on the species (it has already
gained or lost electrons) – e.g. O2- oxidation # is -2; Cu+ oxidation # is +1; F3+ oxidation # is +3
3. Fluorine (F): always have oxidation # of -1. **it is the most electronegative element – e.g. NaF, HF
o F, O Group 7
4. Halogens: Clorine (Cl), Bromine (Br) & Iodine (I): oxidation numbers of -1, EXCEPT in compounds
with Oxygen (O) or Fluorine (F) – e.g. CIO4- -- because Cl, Br & I have high electronegativities, but not
as high as O or F | HCl oxidation # of Cl is -1; KBr oxidation number of Br is -1
5. Hydrogen (H) atoms are usually +1 in most compounds, EXCEPT in metal hydrides (NaH, CaH2)
where oxidation number is -1 (because H has low-ish electronegativity, but not as low as metals) o
6. Oxygen (O): -2 in most compounds. Except peroxides (which have O=O bond) where oxidation # is
-1 (because O bond w/ other O involves NO electronegativity difference)
o Peroxide (H2O2)-1
7. The sum of oxidation #s for all atoms in a polyatomic species is equal to the species’ charge
Balancing redox reactions:
SOHC
Species oxygen Hydrogen Charge
Aqueous Solutions (aq)
Any solution in which H2O is the solvent. (acids must be in aq. solutions)
aq > in an aqueous solution
Solute
- Can be solid, liquid or gas
- Can be ionic or covalent
Solvent
- Can be any liquid o Water (aq); ethanol; hydrocarbons (e.g. hexane)
• Ionic compounds – often form strong electrolytes (100% dissociated) Example:
• Molecular compounds – non-, weak, and strong electrolytes (~3% dissociated, weak electrolyte)
Example:
Aqueous Solutions: Electrical Conductivity
Properties in aqueous solution: non/electrolyte
(aq) – i.e. surrounded by water molecules
• Non-electrolyte – solution does not conduct electricity
• Strong-electrolyte – solution conducts electricity well (ions are free to move – movement of +ve & -
ve charges)
• Weak-electrolyte – solution conducts electricity, but not that well
Types:
Electrolyte: substances that dissociate into ions in water
, lOMoAR cPSD| 23050439
Chemistry 1102: Lecture Notes | Semester 2 – 2018
Non-electrolyte: substances that dissolve in water but DO NOT dissociate into ions
Solute + Solvent
Electrolyte
Strong Weak Non-
electrolyte electrolyte electrolyte
100% Few% 0% dissociated
dissociated dissociated into ions
into ions into ions
Salts dissociate => ions (ionisation)
• Ionic compounds: strong electrolytes
• Molecular compounds: non/weak/strong electrolytes (depending)
Electrical conductivity
• Depends on # of ions in solution o Non-electrolyte: does not conduct electricity o Strong
electrolyte: conducts electricity very well o Weak electrolyte: conducts electricity, but not that well
Exchange Reactions (metathesis/double-displacement reactions)
Reactions in which anions and cations exchange partners
• A+ B- + C+D- A+D- + C+B-
, • If a combination of anion-cation is NOT soluble, it could form (net reaction occurs):
o Precipitate (s) o Weak or
non-electrolyte (aq)
o Gas (g)
Redox Reactions
Reactions where electrons are transferred between species, NOT just an exchange of ions OIL
RIG
Oil – oxidation is loss (loss of electrons)
o Increase in oxidation number
Rig – reduction is gain (of electrons) o
Decrease in oxidation number
, lOMoAR cPSD| 23050439
Chemistry 1102: Lecture Notes | Semester 2 – 2018
Cu atom loses 2 electrons to become a 2x charged cation (Oxidation is loss, so Cu is oxidized)
2 Ag+ cations gain an electron each to become two Ag atoms (Reduction is gain, so Ag + is reduced)
Half reactions/equations
Oxidation step (loss of e-): Cu(s) Cu 2+(aq) + 2e -
Reduction step (gain of e-): 2 Ag + (aq) + 2e - 2 Ag(s)
1. Each half reaction contains same charge on either side of reaction
2. The same # of e- that are lost in oxidation are gained in reduction
*The species that is being oxidised is the reducing agent or reductant because it is reducing the other
species.
*The species that is being reduced is the oxidising agent or oxidant because it is oxidising the other
species.
Why redox reactions?
Essential for energy storage and conversion in biological organisms. (example: respiration)
**example: refer to pg. 1 black book
Determining when a reaction is a redox reaction:
Oxidation Numbers (oxidation states) o Method for assigning electrons to
individual atoms within a molecule o A change in oxidation number =
redox reaction
o Oxidation # is determined by electronegativities (if atoms lusts for electrons it will take
them from something that doesn’t lust for them as much **Learn electronegativity periodic table Rules
for assigning oxidation numbers:
1. Any atom in its elemental state has an oxidation number of zero (no electronegativity difference
between bonded atoms) – e.g. C ; H2; Mg; Cl2 o H2, He2, N2, O2, F2, Cl2 – Br2, I2 0
2. Monatomic ions: the oxidation number is equal to the net charge on the species (it has already
gained or lost electrons) – e.g. O2- oxidation # is -2; Cu+ oxidation # is +1; F3+ oxidation # is +3
3. Fluorine (F): always have oxidation # of -1. **it is the most electronegative element – e.g. NaF, HF
o F, O Group 7
4. Halogens: Clorine (Cl), Bromine (Br) & Iodine (I): oxidation numbers of -1, EXCEPT in compounds
with Oxygen (O) or Fluorine (F) – e.g. CIO4- -- because Cl, Br & I have high electronegativities, but not
as high as O or F | HCl oxidation # of Cl is -1; KBr oxidation number of Br is -1
5. Hydrogen (H) atoms are usually +1 in most compounds, EXCEPT in metal hydrides (NaH, CaH2)
where oxidation number is -1 (because H has low-ish electronegativity, but not as low as metals) o
6. Oxygen (O): -2 in most compounds. Except peroxides (which have O=O bond) where oxidation # is
-1 (because O bond w/ other O involves NO electronegativity difference)
o Peroxide (H2O2)-1
7. The sum of oxidation #s for all atoms in a polyatomic species is equal to the species’ charge
Balancing redox reactions:
SOHC
Species oxygen Hydrogen Charge
