Unit 9
eactivity 3.2: Electron transfer reactions
R
tructure 3.1.6
S
- The oxidation state is a number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge
that an atom would have if the compound were composed of ions.
● Deduce the oxidation states of an atom in an ion or a compound.
- Rule Number 1:The oxidation state of any uncombinedelement is zero
- Rule Number 2:The oxidation state of each of theatoms in a compound counts separately and all of these oxidation states
add up to zero.
- Rule Number 3:The oxidation state of a chemical in a simple monatomic ion is equal in size and sign to the charge on the ion.
- Rule number 4:In a polyatomic ion (an ion with more than one atom in it like the hydroxide ion), the algebraic sum of the
oxidation state of the atoms is equal to the charge on the ion.
- Rule number 5:Many elements have fixed oxidationstate in their common compounds, the rules below have to applied in the
order in which they are given
- Gp I metals in all compounds +1
- Gp II metals in all compounds +2
- Aluminium in all compounds +3
- Hydrogen in most compounds +1 except in hydrides (e.g. MgH2)
- Fluorine in all compounds -1
- Chlorine, bromine and iodine in most compounds -1 except with oxygen and fluorine
- Oxygen in most compounds -2 except in peroxides and with fluorine
- Rule number 6:Assign oxidation numbers to all other chemicals in a compound using rules 2-5
- The oxidation state of an element is zero because they have neither been reduced or oxidized
- Oxidation states are shown with a + or – sign followed by the number, e.g. +2, –1
- Examples should include hydrogen in metal hydrides (–1) and oxygen in peroxides (–1).
- The terms “oxidation number” and “oxidation state” are often used interchangeably, and either term is acceptable in assessment.
- Naming conventions for oxyanions use oxidation numbers shown with Roman numerals, but generic names persist and are acceptable.
- Examples include NO3- nitrate, NO2- nitrite, SO42– sulfate and SO32 – sulfite.
3.2.1: Oxidation and reduction
- Oxidation and reduction can be described in terms of electron transfer, change in oxidation state, oxygen gain/loss or hydrogen
loss/gain.
- Definition 1:oxidation is gain of oxygen, reductionis loss of oxygen
- Definition 2:oxidation is loss of hydrogen, reductionis gain of hydrogen
- Definition 3:oxidation is loss of electrons, reductionis gain of electrons
- Definition 4:oxidation is increase in oxidation state, reduction is decrease in oxidation state
● Identify the oxidized and reduced species and the oxidizing and reducing agents in a chemical reaction.
- The oxidised species is the reducing agent and the reduced species is the oxidizing agent
- Include examples to illustrate the variable oxidation states of transition element ions and of most main group nonmetals
- Include the use of oxidation numbers in the naming of compounds.
3.2.2: Half equations
- Half-equations separate the processes of oxidation and reduction, showing the loss or gain of electrons.
● Deduce redox half-equations and equations in acidic or neutral solutions.
, - Step 1:Write the skeletons of the oxidation and reduction half-reactions separately. (The skeleton reactions contain the
formulas of the compounds oxidized and reduced, but the atoms and electrons have not yet been balanced.)
- Step 2:Balance all elements other than H and O.
- Step 3:Balance the oxygen atoms byadding H2O
moleculeswhere needed.
- Step 4:Balance the hydrogen atoms byadding H+ionswhere neededin acidic solutions and OH- in basicsolutions
- Step 5:Balance the charge byadding electrons, e-
- Step 6:If the number of electrons lost in the oxidationhalf-reaction is not equal to the number of electrons gained in the
reduction half-reaction, multiply one or both of the half-reactions by a number that willmake the numberof electrons
gained equal to the number of electrons lost.
- Step 7:Add the 2 half-reactions as if they were mathematicalequations and cancel out. The electrons will always cancel. If
the same formulas are found on opposite sides of the half-reactions, you can cancel them. If the same formulas are found on
the same side of both half-reactions, combine them.
- Step 8:Check to make sure that the atoms and the charges balance.
- Why are some redox titrations described as “self-indicating”
- Indicators are sometimes used to show the endpoint of the titration. However, most transition metal ions naturally change
colour when changing oxidation state. Therefore titrations including coloured transition metals are known as self
indicating
- Examples of these titrations:
- iron and potassium permanganate that change from colourless to pink
- iodine and sodium thiosulphate form brown to clear
- The winkler method
- REMOVE SPECTATOR IONS FROM HALF EQUATIONS BUT ADD THEM BACK IN FOR THE FULL EQUATION
- Example: 2PbSO4 (s)+ 2H2O(l) ⇌ Pb(s)+ PbO2(s) + 2H2S O4(aq)
- Pb2+ + 2e- → Pb
- Pb2+ + 2H2O→ PbO2 + 4H+ +2e-
3.2.3: Ease of oxidation/reduction: the reactivity series
- The relative ease of oxidation and reduction of an element in a group can bepredicted from its position in the periodic table.
- The reactions between metals and aqueous metal ions demonstrate the relative ease of oxidation of different metals.
● Predict the relative ease of oxidation of metals.
T he alkali metals are most likely to be oxidised and are therefore great reducing agents
-
- When a more reactive metal is placed in a solution of a less reactive metal ions, a displacement reaction takes place
- Observations such as colour and temperature changes as well as formation of solid compounds
- Common reactivity Mg>Zn>Fe>Cu
- Metals react with acid to produce hydrogen gas and salt
- Effervescence and gas production can be observed and tested for with a lit splint that goespop
● Predict the relative ease of reduction of halogens.
H alogens are the most likely to be reduced and are good oxidising agents
-
- Halogens readily gain electrons and have a high electron affinity, so are therefore easily reduced to their ions
- Using the periodic table the reactivity series of the halogens isF >Cl >Br >I
2 2 2 2
- Reactivity increases down group 1
- Reactivity increases up group 17
- Observations can include the production of coloured gases corresponding to the reactivity of the halogens
dead>clear>yellow>brown
eactivity 3.2: Electron transfer reactions
R
tructure 3.1.6
S
- The oxidation state is a number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge
that an atom would have if the compound were composed of ions.
● Deduce the oxidation states of an atom in an ion or a compound.
- Rule Number 1:The oxidation state of any uncombinedelement is zero
- Rule Number 2:The oxidation state of each of theatoms in a compound counts separately and all of these oxidation states
add up to zero.
- Rule Number 3:The oxidation state of a chemical in a simple monatomic ion is equal in size and sign to the charge on the ion.
- Rule number 4:In a polyatomic ion (an ion with more than one atom in it like the hydroxide ion), the algebraic sum of the
oxidation state of the atoms is equal to the charge on the ion.
- Rule number 5:Many elements have fixed oxidationstate in their common compounds, the rules below have to applied in the
order in which they are given
- Gp I metals in all compounds +1
- Gp II metals in all compounds +2
- Aluminium in all compounds +3
- Hydrogen in most compounds +1 except in hydrides (e.g. MgH2)
- Fluorine in all compounds -1
- Chlorine, bromine and iodine in most compounds -1 except with oxygen and fluorine
- Oxygen in most compounds -2 except in peroxides and with fluorine
- Rule number 6:Assign oxidation numbers to all other chemicals in a compound using rules 2-5
- The oxidation state of an element is zero because they have neither been reduced or oxidized
- Oxidation states are shown with a + or – sign followed by the number, e.g. +2, –1
- Examples should include hydrogen in metal hydrides (–1) and oxygen in peroxides (–1).
- The terms “oxidation number” and “oxidation state” are often used interchangeably, and either term is acceptable in assessment.
- Naming conventions for oxyanions use oxidation numbers shown with Roman numerals, but generic names persist and are acceptable.
- Examples include NO3- nitrate, NO2- nitrite, SO42– sulfate and SO32 – sulfite.
3.2.1: Oxidation and reduction
- Oxidation and reduction can be described in terms of electron transfer, change in oxidation state, oxygen gain/loss or hydrogen
loss/gain.
- Definition 1:oxidation is gain of oxygen, reductionis loss of oxygen
- Definition 2:oxidation is loss of hydrogen, reductionis gain of hydrogen
- Definition 3:oxidation is loss of electrons, reductionis gain of electrons
- Definition 4:oxidation is increase in oxidation state, reduction is decrease in oxidation state
● Identify the oxidized and reduced species and the oxidizing and reducing agents in a chemical reaction.
- The oxidised species is the reducing agent and the reduced species is the oxidizing agent
- Include examples to illustrate the variable oxidation states of transition element ions and of most main group nonmetals
- Include the use of oxidation numbers in the naming of compounds.
3.2.2: Half equations
- Half-equations separate the processes of oxidation and reduction, showing the loss or gain of electrons.
● Deduce redox half-equations and equations in acidic or neutral solutions.
, - Step 1:Write the skeletons of the oxidation and reduction half-reactions separately. (The skeleton reactions contain the
formulas of the compounds oxidized and reduced, but the atoms and electrons have not yet been balanced.)
- Step 2:Balance all elements other than H and O.
- Step 3:Balance the oxygen atoms byadding H2O
moleculeswhere needed.
- Step 4:Balance the hydrogen atoms byadding H+ionswhere neededin acidic solutions and OH- in basicsolutions
- Step 5:Balance the charge byadding electrons, e-
- Step 6:If the number of electrons lost in the oxidationhalf-reaction is not equal to the number of electrons gained in the
reduction half-reaction, multiply one or both of the half-reactions by a number that willmake the numberof electrons
gained equal to the number of electrons lost.
- Step 7:Add the 2 half-reactions as if they were mathematicalequations and cancel out. The electrons will always cancel. If
the same formulas are found on opposite sides of the half-reactions, you can cancel them. If the same formulas are found on
the same side of both half-reactions, combine them.
- Step 8:Check to make sure that the atoms and the charges balance.
- Why are some redox titrations described as “self-indicating”
- Indicators are sometimes used to show the endpoint of the titration. However, most transition metal ions naturally change
colour when changing oxidation state. Therefore titrations including coloured transition metals are known as self
indicating
- Examples of these titrations:
- iron and potassium permanganate that change from colourless to pink
- iodine and sodium thiosulphate form brown to clear
- The winkler method
- REMOVE SPECTATOR IONS FROM HALF EQUATIONS BUT ADD THEM BACK IN FOR THE FULL EQUATION
- Example: 2PbSO4 (s)+ 2H2O(l) ⇌ Pb(s)+ PbO2(s) + 2H2S O4(aq)
- Pb2+ + 2e- → Pb
- Pb2+ + 2H2O→ PbO2 + 4H+ +2e-
3.2.3: Ease of oxidation/reduction: the reactivity series
- The relative ease of oxidation and reduction of an element in a group can bepredicted from its position in the periodic table.
- The reactions between metals and aqueous metal ions demonstrate the relative ease of oxidation of different metals.
● Predict the relative ease of oxidation of metals.
T he alkali metals are most likely to be oxidised and are therefore great reducing agents
-
- When a more reactive metal is placed in a solution of a less reactive metal ions, a displacement reaction takes place
- Observations such as colour and temperature changes as well as formation of solid compounds
- Common reactivity Mg>Zn>Fe>Cu
- Metals react with acid to produce hydrogen gas and salt
- Effervescence and gas production can be observed and tested for with a lit splint that goespop
● Predict the relative ease of reduction of halogens.
H alogens are the most likely to be reduced and are good oxidising agents
-
- Halogens readily gain electrons and have a high electron affinity, so are therefore easily reduced to their ions
- Using the periodic table the reactivity series of the halogens isF >Cl >Br >I
2 2 2 2
- Reactivity increases down group 1
- Reactivity increases up group 17
- Observations can include the production of coloured gases corresponding to the reactivity of the halogens
dead>clear>yellow>brown
