Redox Reactions
HSC Chemistry Syllabus
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Apply the definitions of oxidation and reduction in terms of electron transfer and oxidation numbers to a range of reduction and oxidation (redox) reactions
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Construct relevant half-equations and balanced overall equations to represent a range of redox reactions
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Predict the reaction of metals in solutions using the table of standard reduction potentials
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Predict the spontaneity of redox reactions using the value of cell potentials
Redox Reactions
This video explores what reaction reactions are? It explains the acronym OILRIG which refers to oxidation being the loss of electrons, while reduction is the gain of electrons.
What is Redox Reaction?
Redox chemistry is all about the movement of electrons between different species in a reaction. To break it down:
- Reduction refers to the gain of electrons. For example, the transformation of iron ions into solid metallic iron as iron ion accepts electrons to become atomic iron.
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Oxidation refers to the loss of electrons. Consider magnesium turning into magnesium ions and releasing electrons; this process exemplifies oxidation.
These two processes always occur together – where one element loses electrons (oxidation) another gains them (reduction).
An easy mnemonic to help you remember the differences between oxidation and reduction is "OILRIG" which stands for "oxidation is loss, reduction is gain". This will help you identify whether a reaction involves oxidation or reduction, such as when magnesium loses electrons, it's undergoing oxidation.
Metal Displacement Reactions
Metal displacement reactions, also known as single replacement reactions, are a fascinating type of chemical reaction where a more reactive metal displaces a less reactive metal ion from its ionic compound. These reactions are a specific category of redox reactions, which involve the transfer of electrons between substances.
Let's consider the reaction between copper (Cu) metal and a silver nitrate (AgNO₃) solution, which contains silver ions (Ag⁺). When copper metal is placed in a solution of silver nitrate, you will observe the gradual deposition of silver metal on the copper surface. Simultaneously, the solution, initially colourless, will turn blue due to the formation of copper ions (Cu²⁺), which are blue in aqueous solution.
The overall reaction can be written as:
$$Cu(s) + 2AgNO_3(aq) \rightarrow Cu(NO_3)_2(aq) + 2Ag(s)$$
The reactivity series of metals explains why this reaction occurs. The reactivity series ranks metals based on their ability to lose electrons and form positive ions. Copper is more reactive than silver, meaning copper is more willing to lose electrons and be oxidised compared to silver. Thus, copper displaces silver from its ionic compound. This is discussed in more detail under the section Reduction Potentials below.
Half-Equations in Redox Reactions
Redox reactions can be split into half-equations, which separately describe the oxidation and reduction occurring in a reaction.
The above reaction between copper metal and silver ion can be first simplified into a net ionic equation by removing spectator ions.
$$Cu(s) + 2Ag^+(aq) \rightarrow Cu^{2+}(aq) + 2Ag(s)$$
In this metal displacement reaction, copper has lost 2 electrons to form the `Cu^{2+}` ion while each `Ag^+` ion has gained one electron to form silver metal.
This transfer of electrons can be demonstrated by the following equations:
$$\text{oxidation half-equation:} \hspace{2cm} Cu \rightarrow Cu^{2+} + 2e^–$$
$$\text{reduction half-equation:} \hspace{2cm} 2Ag^+ + 2e^– \rightarrow 2Ag$$
The top half of the equation is the oxidation half equation since we know that the copper has lost electrons to ionise, and the bottom half is the reduction half equation since silver ion has gained electrons to become hydrogen gas.
It is important to notice that the number of electrons lost in oxidation should always equal to the number of electrons gained in reduction. This is why there are 2 silver ions gaining 2 electrons to become 2 silver atoms in the reduction half equation.
When the two half equations are added, the two electrons on both sides of the equation can be cancelled to reform the net ionic equation:
$$Cu(s) + 2Ag^+(aq) + 2e^- \rightarrow Cu^{2+}(aq) + 2Ag(s) + 2e^- $$
$$Cu(s) + 2Ag^+(aq) \rightarrow Cu^{2+}(aq) + 2Ag(s)$$
Oxidising and Reducing Agents
An oxidising agent, or oxidant, is a substance that causes another substance to undergo oxidation. In the process, the oxidising agent itself gets reduced or undergoes reduction. In the reaction between copper and silver ion, silver ion is the oxidising agent because it causes copper to lose electrons (oxidation).
A reducing agent, or reductant, is a substance that causes another substance to undergo reduction. In the process, the reducing agent itself gets oxidised or undergoes oxidation. In the reaction above, copper is the reducing agent because it causes silver ions to gain electrons (reduction).
Reduction Potentials
A reduction potential table, also known as an electrochemical series, is a tabulated list of reduction potentials of various half-reactions. These potentials indicate the tendency of a species to gain electrons and be reduced. The table is crucial in predicting the spontaneity of redox reactions and determining the relative strengths of oxidising and reducing agents.
The reactivity of metals is related to their position in the reduction potential table. More reactive metals, which are easily oxidised to form their respect metal ions are found at the top of the table, like potassium (K).
The metals found at the bottom of the reduction potential table are less reactive because their metal ions have greater tendency to undergo reduction by gaining electrons. An example of a relatively inert metal is silver (Ag).
Key Features of the Reduction Potential Table
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Standard Conditions:
The values listed are standard reduction potentials, measured under standard conditions: 25°C (298 K), 1 mol/L concentration for species in aqueous states, and 1 atm pressure for gases.
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Reference Electrode:
The standard hydrogen electrode (SHE) is used as the reference electrode, with a defined potential of 0.00 V.
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Arrangement:
Species with higher (more positive) reduction potentials are stronger oxidising agents and more readily gain electrons. Species with lower (more negative) reduction potentials are stronger reducing agents and more readily lose electrons.
Spontaneity of Redox Reactions
The concept of spontaneity in reaction tells us whether a reaction will proceed with or without external energy. A reaction is spontaneous if it occurs without external energy input.
Reduction potentials are used help us predict this. When the total potential of a reduction half-equation and an oxidation half-equation is greater than zero, the redox reaction occurs spontaneously.
Now let's consider the reaction between copper and silver ions.
$$Cu(s) + 2Ag^+(aq) \rightarrow Cu^{2+}(aq) + 2Ag(s)$$
The half equations of this equation are:
$$Cu(s) \rightarrow Cu^{2+}(aq) + 2e^–$$
$$2Ag^+(aq) + 2e^– \rightarrow 2Ag(s)$$
Using the reduction potential table, the reduction of silver ion is 0.77 V while the reduction of copper ion (`Cu^{2+}`), is 0.34 V. Since the reduction potential of silver ion is greater (more positive), it preferentially undergoes reduction. This means the copper atoms in copper metal undergoes the opposite reaction of reduction, that is oxidation.
The oxidation potential of copper to form `Cu^{2+}` ions is -0.34 V.
The total potential of the reaction between silver ion and copper atom is the sum of the potential of these two half equations:
0.77 + (–0.34) = 0.43 V
The positive value of this electrical potential indicates that this reaction occurs spontaneously.
Examples
Using the reduction potential table, determine whether a chemical reaction occurs between the following species. If there is a reaction, write the redox half-equations and identify the oxidising and reducing agents.
- Copper metal in nickel(II) chloride solution
- Iron metal in aluminium chloride solution
- Aluminium metal and copper(II) nitrate
- Silver nitrate and iron metal
- Zinc metal and magnesium chloride
- Copper(II) chloride and zinc metal
Watch the video above for answers.