Topic 6 (REDOX) Flashcards
(234 cards)
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Oxidation & Reduction
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Redox reactions
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Oxidation and reduction take place together at the same time in the same reaction
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These are called redox reactions
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Oxidation can be defined as:
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The gain of oxygen
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The loss of electrons
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An increase in oxidation number
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Reduction can be defined as:
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The loss of oxygen
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The gain of electrons
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A decrease in oxidation number
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Names using oxidation numbers
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Transition elements can bond in different ways by forming ions with different charges
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When naming
the charge on the ion is shown by using a Roman numeral after the element's name
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e.g. iron can form ions with a 2+ charge
called iron(II) ions or a 3+ charge
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The Roman numeral is the oxidation number of the element
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When iron reacts with oxygen to form iron oxide
the formula depends on the oxidation state of the iron ions
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The compound where iron has a 2+ charge has the formula FeO and is called iron(II) oxide
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The compound where iron has a 3+ charge has the formula Fe2O3 and is called iron(III) oxide
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Examiner Tips and Tricks
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You may see the term oxidation state used instead of oxidation number. Although there is a subtle difference between the two terms (this is beyond the scope of this course)
they are often used interchangeably. Usually oxidation number is used to refer to the Roman numerals found within the name.
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Oxidation and reduction in terms of oxygen
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Oxidation is a reaction in which oxygen is added to an element or a compound
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Reduction is a reaction in which oxygen is removed from an element or compound
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Example: Identifying the loss and gain of oxygen in an equation
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zinc oxide + carbon → zinc + carbon monoxide
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ZnO + C → Zn + CO
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In this reaction
the zinc oxide has been reduced since it has lost oxygen
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The carbon atom has been oxidised since it has gained oxygen
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Oxidation and reduction in terms of electrons and oxidation number
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Redox reactions can also be defined in terms of electron transfer
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Oxidation is a reaction in which an element
ion or compound loses electrons
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The oxidation number of the element is increased
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This can be shown in a half equation
e.g. when silver reacts with chlorine
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Ag → Ag+ + e-
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Reduction is a reaction in which an element
ion or compound gains electrons
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The oxidation number of the element is decreased
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This can be shown in a half equation
e.g. when oxygen reacts with magnesium
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O2 + 4e- → 2O2-
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Example: Identifying Redox Reactions
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zinc + copper sulphate → zinc sulphate + copper
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Zn + CuSO4 → ZnSO4 + Cu
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The ions present (with state symbols) in the equation are:
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Zn(s) + Cu2+(aq) + SO42-(aq) →Zn2+(aq) + SO42-(aq) + Cu(s)
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The spectator ions (those that do not change) are SO42-(aq)
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These can be removed and the ionic equation written as:
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Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
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By analysing the ionic equation
we can split the reaction into two half equations by adding in the electrons to show how the changes in charge have occurred:
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Zn(s) → Zn2+(aq) + 2e-
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Cu2+(aq) +2e- → Cu(s)
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It then becomes clear that zinc has been oxidised as it has lost electrons
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Copper ions have been reduced as they have gained electrons
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Examiner Tips and Tricks
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Use the mnemonic OIL-RIG to remember oxidation and reduction in terms of the movement of electrons: Oxidation Is Loss – Reduction Is Gain.
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Did this video help you?
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Identifying Redox Reactions
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Oxidation Number
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The oxidation number (also called oxidation state) is a number assigned to an atom or ion in a compound which indicates the degree of oxidation (or reduction)
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It shows the number of electrons that an atom has lost
gained or shared in forming a compound
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The oxidation number helps you to keep track of the movement of electrons in a redox process
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It is written as a +/- sign followed by a number (not to be confused with charge which is written by a number followed by a +/- sign)
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E.g. aluminium in a compound usually has the oxidation state +3
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A few simple rules help guide you through the process of determining the oxidation number of any element
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Table of Rules for Assigning Oxidation Numbers
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6-3-2-oxidation-number-rules-table
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Redox reactions can be identified by the changes in the oxidation number when a reactant goes to a product
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Worked Example
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The equation for the reaction between chlorine and potassium iodide is shown below.
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Cl2 + 2KI → 2KCl + I2
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Identify which species has been:
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a) Oxidised
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b) Reduced
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Answer:
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The species that has been oxidised is iodine
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The oxidation number of I- is -1
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The oxidation number of iodine in I2 is 0
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The oxidation number has increased so the iodine has been oxidised (lost electrons)
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2I-(aq) → I2(s) +2e-
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The species that has been reduced is chloride ions
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The oxidation number of chlorine as Cl2 is 0.
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The oxidation number of Cl- is -1
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The oxidation number has decreased so the Cl- has been reduced (gained electrons)
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Cl2(g) + 2e- → 2Cl-(aq)
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Identifying Redox Reactions by Colour Changes
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The tests for redox reactions involve the observation of a colour change in the solution being analysed
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Two common examples are acidified potassium manganate(VII)
and potassium iodide
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Potassium manganate(VII)
KMnO4
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When acidified potassium manganate(VII) is added to a reducing agent its colour changes from purple to colourless
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6-3-2-test-for-reducing-agent-using-potassium-manganatevii
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Diagram to show the colour change when potassium manganate(VII) is added to a reducing agent
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Potassium iodide
KI
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When added to an acidified solution of an oxidising agent such as aqueous chlorine or hydrogen peroxide (H2O2)
the solution turns a red-brown colour due to the formation of iodine
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2KI (aq) + H2SO4 (aq) + H2O2 (aq) → I2 (aq) + K2SO4 (aq) + 2H20 (l)
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The potassium iodide is oxidised as it loses electrons and hydrogen peroxide is reduced
therefore potassium iodide is acting as a reducing agent as it will itself be oxidised:
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2I- → I2 + 2e-
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6-3-2-test-for-oxidising-agent-using-potassium-iodide6-3-2-test-for-oxidising-agent-using-potassium-iodide
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Diagram to show the colour change when potassium iodide is added to an oxidising agent
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Oxidising & Reducing Agents
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Oxidising agent
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A substance that oxidises another substance
and becomes reduced in the process
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An oxidising agent gains electrons as another substance loses electrons
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Common examples include hydrogen peroxide
fluorine and chlorine
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Reducing agent
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A substance that reduces another substance
and becomes oxidised in the process
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A reducing agent loses electrons as another substance gains electrons
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Common examples include carbon and hydrogen
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The process of reduction is very important in the chemical industry as a means of extracting metals from their ores
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Example
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CuO + H2 → Cu + H2O
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In the above reaction
hydrogen is reducing the CuO and is itself oxidised as it has lost electrons
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H2 → 2H+ + 2e-
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The CuO is reduced to Cu by gaining electrons and has oxidised the hydrogen
so the oxidising agent is therefore copper oxide
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Cu2+ +2e- → Cu
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Worked Example
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When iron reacts with bromine to form iron(II) bromide
a redox reaction reaction occurs:
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Fe + Br2 → FeBr2
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What is acting as the reducing agent in this reaction?
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Answer
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Step 1 - Write half equations to work out what has gained/lost electrons
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Fe → Fe2+ + 2e-
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Br2 + 2e- → 2Br-
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Fe loses electrons; Br2 gains electrons
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Step 2 - Deduce what has been oxidised/reduced (remember OIL RIG)
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Fe has been oxidised as it has lost electrons
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Br2 has been reduced as it has gained electrons
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Step 3 - Identify the reducing agent
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Fe is the reducing agent as it has been oxidised by losing electrons and caused Br2 to be reduced as it gained electrons
