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iB Chemistry > Topic 9: Redox Processes > Flashcards

Topic 9: Redox Processes Flashcards

1
Q

Ways to describe Oxidation

A
  1. Addition of oxygen
  2. Loss of hydrogen
  3. Loss of electrons
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2
Q

Ways to describe Reduction

A
  1. Loss of oxygen
  2. Addition of hydrogen
  3. Gain of electrons
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3
Q

Oxidation Numbers

A
  1. The oxidation number of an atom is the charge that would exist on an individual atom if the bonding were completely ionic
  2. It is like the electronic ‘status’ of an element
  3. The oxidation number is the number of electrons which must be added or removed to become neutral
  4. The oxidation number is always written with the charge before the number
  5. In molecules or compounds, the sum of the oxidation numbers on the atoms is zero
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4
Q

Use of Oxidation Numbers

A
  1. used to tell if oxidation or reduction has taken place
  2. helps to work out what has been oxidised and/or reduced
  3. used to construct half equations and balance redox equations
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5
Q

Variable Oxidation Numbers

A
  1. Many atoms, such as S, N and Cl, can exist in a variety of oxidation states
  2. The oxidation number of these atoms can be calculated by assuming that the oxidation number of the other atom is fixed
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6
Q

Rules to deduce the oxidation number of an element

A
  1. The oxidation number of any uncombined element is 0
  2. Many atoms or ions have fixed oxidation numbers in compounds (group 1 elements are always +1, fluorine is always -1, oxygen is -2, hydrogen is +1) - there are a few exceptions
  3. The oxidation number of an element in a mono-atomic ion is always the same as the charge (Cl- is -1)
  4. The sum of the oxidation number in a compound is 0
  5. The sum of the oxidation number in an ion is equal to the charge on the ion
  6. The more electromagnetic element is always given the negative oxidation number
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7
Q

Oxidising agent

A
  1. An oxidising agent is a substance that oxidises another atom or ion by causing it to lose electrons
  2. An oxidising agent itself gets reduced – gains electrons
  3. Therefore, the oxidation number of the oxidising agent decreases
  4. Some substances can act both as oxidising and reducing agents
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8
Q

Reducing agent

A
  1. A reducing agent is a substance that reduces another atom or ion by causing it to gain electrons
  2. A reducing agent itself gets oxidised – loses/donates electrons
  3. Therefore, the oxidation number of the reducing agent increases
  4. Some substances can act both as oxidising and reducing agents
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9
Q

Redox Reaction requirements

A

For a reaction to be recognised as a redox reaction, there must be both an oxidising and reducing agent

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10
Q

Transition metals (Oxidation Number)

A
  1. Transition metals are characterized by having variable oxidation numbers.
  2. Oxidation numbers can be used in the names of compounds to indicate which oxidation number a particular element in the compound is in
  3. Where the element has a variable oxidation number, the number is written afterwards in Roman numerals.
  4. This is called the STOCK NOTATION
  5. For example, iron can be both +2 and +3 so Roman numerals are used to distinguish between them

Fe2+ in FeO can be written as iron(II) oxide
Fe3+ in Fe2O3 can be written as iron(III) oxide

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11
Q

Balancing a redox equation

A
  1. Write the unbalanced equation and identify the atoms which change in ox. no.
  2. Deduce the oxidation number changes
  3. Balance the oxidation number changes
  4. Balance the charges
  5. Balance the atoms
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12
Q

reactivity series

A
  1. Metals can be ranked in order of reactivity
  2. Metals higher in reactivity can displace less reactive metals from their compounds in solutions or from their oxides
  3. The more reactive metal acts as a reducing agent
  4. This allows metals to be ranked from most reactive (strongest reducing agents) to least reactive
  5. The more reactive a metal is the better it is at pushing electrons onto less reactive metal ions.
  6. The more reactive metal undergoes oxidation
  7. The less reactive metal ion undergoes reduction
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13
Q

Redox Titrations

A
  1. In a titration, the concentration of a solution is determined by titrating with a solution of known concentration.
  2. In redox titrations, an oxidizing agent is titrated against a reducing agent
  3. Electrons are transferred from one species to the other
  4. Indicators are sometimes used to show the endpoint of the titration. However, most transition metal ions naturally change colour when changing oxidation state
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14
Q

Types of redox titrations

A
  1. manganate(VII) titrations
  2. iodine-thiosulfate titrations
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15
Q

manganate(VII) titrations

A
  1. A redox reaction occurs between acidified manganate (VII) ions and iron (II) ions:
    MnO4– (aq) + 8H+ (aq) + 5Fe2+ (aq) → Mn2+ (aq) + 5Fe3+ (aq) + 4H2O (l)
  2. This reaction needs no indicator as the manganate (VII) is a strong purple colour which disappears at the end point, so the titration is self-indicating
  3. This reaction is often used for the analysis of iron for example in iron tablets (health supplement)
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16
Q

Iodine-Thiosulfate Titrations

A
  1. A redox reaction occurs between iodine and thiosulfate ions:
    2S2O32– (aq) + I2 (aq) → 2I–(aq) + S4O62– (aq)
  2. The light brown/yellow colour of the iodine turns paler as it is converted to colourless iodide ions
  3. When the solution is a straw colour, starch is added to clarify the end point
  4. The solution turns blue/black until all the iodine reacts, at which point the colour disappears.
  5. This titration can be used to determine the concentration of an oxidizing agent, which oxidizes iodide ions to iodine molecules
  6. The amount of iodine is determined from titration against a known quantity of sodium thiosulfate solution
  7. This reaction can be used for the analysis of chlorine in bleach
17
Q

The Winkler Method

A
  1. The Winkler Method is a technique used to measure dissolved oxygen in freshwater systems
  2. Dissolved oxygen is used as an indicator of the health of a water body, where higher dissolved oxygen concentrations correlate with high productivity and little pollution
18
Q

biological oxygen demand

A
  1. The biological oxygen demand (BOD) is the amount of oxygen used to decompose the organic matter in a sample of water over a specified time period, usually 5 days, at a specified temperature
  2. A high BOD indicates a greater quantity of organic waste in the water, which means a lower level of dissolved oxygen
19
Q

Voltaic Cells

A
  1. Voltaic (or Galvanic) cells generate electricity from spontaneous redox reactions
  2. Instead of electrons being transferred directly from the zinc to the copper ions a cell is built which separates the two redox processes
  3. Each part of the cell is called a half cell
  4. If a rod of metal is dipped into a solution of its own ions, an equilibrium is set up
  5. This is a half cell and the strip of metal is an electrode
  6. The position of the equilibrium determines the potential difference between the metal strip and the solution of metal. This is known as an electrode potential.
  7. Note that a chemical reaction is not taking place - there is simply a potential difference between the rod and the solution
20
Q

Potential Difference (Voltaic Cells) Dependent on

A
  1. the nature of the ions in solution
  2. the concentration of the ions in solution
  3. the type of electrode used
  4. the temperature
21
Q

Creating an emf

A
  1. If two different electrodes are connected, the potential difference between the two electrodes will cause a current to flow between them. Thus an electromotive force (emf) is established and the system can generate electrical energy
  2. A typical electrochemical cell can be made by combining a zinc electrode in a solution of zinc sulphate with a copper electrode in a solution of copper sulphate
  3. The circuit must be completed by allowing ions to flow from one solution to the other
  4. This is achieved by means of a salt bridge - often a piece of filter paper saturated with a solution of an inert electrolyte such as KNO3(aq)
  5. The e.m.f can be measured using a voltmeter
  6. Voltmeters have a high resistance so that they do not divert much current from the main circuit
  7. The combination of two electrodes in this way is known as an electrochemical or voltaic cell, and can be used to generate electricity
22
Q

Cathode

A
  1. positive electrode
  2. most favours reduction
  3. Reduction takes place
  4. cations flow to the positive half cell since there is a surplus positive charge from the metals becoming cations in the negative half cell
23
Q

Anode

A
  1. negative electrode
  2. oxidation takes place
  3. anions flow to the negative half cell to replace the negative charge of the electrons
24
Q

Electrolytic Cells

A
  1. An ionic compound conducts electricity when it is molten or in solution
  2. The current causes the ionic compound to split up and form new substances.
  3. This process is called electrolysis, a word which comes from Greek and means “splitting by electricity”
25
Q

Electrolysis Uses

A
  1. purifying copper
  2. plating metals with silver and gold
    extracting reactive metals, such as aluminium
  3. making chlorine, hydrogen and sodium hydroxide
26
Q

Electrolysis cells

A
  1. Electrolysis cells can be constructed using a beaker or crucible as the cell depending whether the ionic compound is in solution or molten
  2. In electrolysis, the substance that the current passes through and splits up is called the electrolyte
  3. The electrolyte contains positive and negative ions
  4. Metals will always be formed at the cathode and non-metals at the anode
27
Q

Electrolysis movement of ions

A
  1. Negative ions move to the anode and lose electrons - this is oxidation
  2. Positive ions move to the cathode and gain electrons - this is reduction
  3. Electrically neutral atoms or molecules are released

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