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Chemistry - REDOX REACTIONS (8,9,10) > Chapter 8 - Redox Reactions > Flashcards

Chapter 8 - Redox Reactions Flashcards

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1
Q

What is a Redox Reaction?

A

Redox Reaction = Reduction-Oxidation Reaction

This reaction involves the TRANSFER OF ELECTRONS from one species to another.

Oxidation and Reduction occur at the same time
- SIMULTANEOUSLY

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2
Q

Reduction vs. Oxidation

A

Oxidation is defined as the LOSS OF ELECTRONS
- when one or more electrons are lost from an atom is said to be OXIDISED

Reduction is Defined as the GAIN OF ELECTRONS
- - when one or more electrons are gained from an atom is said to be REDUCED

OIL RIG**

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3
Q

Reducing agent vs Oxidising agent

A

REDUCING AGENT = REDUCTANT

  • Reductant DONATES electrons to another substance causing the SUBSTANCE TO BE REDUCED
  • the reductant itself is OXIDISED

OXIDISING AGENT = OXIDANT

  • ACCEPTS electrons from another substance causing the SUBSTANCE TO BE OXIDISED
  • the oxidant itself is REDUCED.
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4
Q

What are Half Equations?

A

Half Equations are used to represent oxidation and reduction.
- they show what is happening as the electrons are transferred in a redox reaction

Involves showing electrons as products/reactants
OXIDATION = part of products
REDUCTION = part of reactants

To Write a Half Equation

  1. Must be Balanced
  2. Charge must be balanced
  3. Indicate states
  4. Electrons are part

To Write Full Redox Equation

  1. Redox half equations are normally written
  2. multiply the half equations to the electron’s lowest common factor to have the Coefficient number
  3. Balance all species
  4. Add equations together.
  5. Cancel the electrons
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5
Q

What are Halogen Displacement reactions?

example

A

Redox reactions that do not involve oxygen

They occur when one halogen is OXIDISED
- the halide ions lose electrons to produce the elemental halogen
and the other Halogen is reduced to produce the halide ions

  • Electrons transferred from halogen (group 17 element) from the halide ions of a less reactive halogen.
  • the more reactive halogen (one from the higher up in group 17) becomes reduced as it gains electrons to form negative aqueous halide solutions.
  • the aqueous negative halide ions of the less reactive halogen become oxidised as they lose electrons and form the elemental halogen

Cl2/Br-
2Br-(aq) +Cl2(aq) —> 2Cl- (aq) +Br2(aq)
-1. 0. -1. 0

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6
Q

What is Oxidation number?

A

AKA Oxidation states

No physical meaning only:
Used to determine whether a reaction that does involve the formation of ions could be classified as oxidised or reduced.

  • Possible to determined which redox reaction occurRed by comparing the oxidation number before and after the reaction.
    INCREASE in O No. = Oxidised
    DECREASE in O No. = Reduced

The more POSITIVE an atom’s oxidation number = GREATER DEGREE OF OXIDATION

The more NEGATIVE an atom’s oxidation number = GREATER DEGREE OF REDUCTION.

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7
Q

What is Conjugate redox pair?

A

A conjugate redox pair consists of an OXIDISING AGENT (REDUCTANT) and the REDUCING AGENT ( A PRODUCT) that is formed when the OXIDISNG AGENT GAINS ELECTRONS.

In this Case the OXIDATION NUMBER OF THE OXIDISING AGENT DECREASES.

The other Conjugate redox pair in a redox reaction is MADE UP OF THE REDUCING AGENT ( A REACTANT) AND the OXIDISING AGENT (A PRODUCT) that is formed when the REDUCING AGENT LOSES ELECTRONS.

In this case the oxidation number of the reducing agent increases

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8
Q

Rules for Oxidation Numbers

A
  1. The oxidation number for FREE (UNCOMBINED) ELEMENT = 0
    - Na, C, Cl2, P4
  2. IN IONIC COMPOUNDS SIMPLE ION oxidation number = CHARGE ON THE ION
    - Na+ = +1, Cl- + -1, O^-2 + -2
  3. IN COMPOUNDS OXYGEN oxidation number = -2
    - H2O - O =-2
  4. IN COMPOUNDS HYDROGEN oxidation number = +1
    (NON METALS)
    - except in metal hydrides
    - H2O - H=+1, NaH - H = -1, CaH - h = -1
  5. The SUM of the OXIDATION numbers in a NEUTRAL COMPOUND = 0
  6. The SUM of oxidation numbers in a POLYATOMIC ION = CHARGE ON THE ION
  7. In PEROXIDES OXYGEN oxidation number = -1
  8. In COMPOUNDS WITH FLUORINE, OXYGEN has a POSITIVE OXIDATION NUMBER
    (+2)
    - OF(2) - F = -1, O = +2
    - because F is more electronegative than oxygen
  9. The MOST ELECTRONEGATIVE NUMBER is ASSIGNED THE NEGATIVE OXIDATION NUMBER
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9
Q

What happens if there is no change in the oxidation number

A

No change in oxidation number of all elements in the equation for a reaction = then the reaction is NOT A REDOX REACTION

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10
Q

Compare and Contrast Oxidising and Reducing Agents

A

REDUCING AGENT
- loses electrons
- undergoes oxidation
- reduces the oxidising agent
- is the reactant in the oxidation half equation
Reducing agent –> conjugate oxidised form + ne-

OXIDISING AGENT
- Accepts electrons
- oxidises the reducing agent
- undergoes reduction
- is the reactant in the reduction half-reaction:
Oxidising agent + ne- –> conjugate reduced form

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11
Q

Explain Transition metals and oxidation numbers

A

Transition metals and some non-metals have variable oxidation numbers that can be calculated via the rules.

  • Can form ions of different charges
  • to distinguish; indicate with a roman numeral representing the appropriate oxidation number in the name
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12
Q

How to Write and Balance redox Reactions UNDER ACIDIC CONDITIONS

A
  1. Balance all elements except H and O is the half equation
    - 2NO3^- —> N2O
  2. Balance the oxygen atoms by adding water molecules to the side with fewer oxygen atoms
    - 2NO3- —> N2O + 5H2O
  3. Balance the hydrogen atoms. by adding H+ ions (which are always present in acidic solution) to the side with fewer H+ ions
    - 2NO3-, +10H –> N2O +5H2O
  4. Balance the charge in the equation by adding electrons to the side that has the most positive charge
    - 2NO3-, +10H + 8e- –> N2O +5H2O
  5. Add the states

Note: Any H+(aq) and H2O(L) that appear on both sides of the arrow must be cancelled down when combining oxidation and reduction half-equations under acidic conditions.

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13
Q

What is the Reactivity series?

A

The reactivity series is a table that lists half-equations (REDUCTION) involving metals and their corresponding cations.

Half-equations involving STRONGER OXIDISING AGENTS (ones more easily reduced) APPEAR HIGHER in the REACTIVITY SERIES

Half equations involving STRONGER REDUCING AGENTS (ones more easily oxidised) APPEAR LOWER in the Reactivity series.

HENCE CAN BE USED TO PREDICT WHETHER A REDOX REACTION IS LIKELY TO OCCUR

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14
Q

Explain Half-equations with polyatomic ions in acidic solutions

A

Some oxidising agents will only function effectively if hydrogen ions (H+) aq are available in the reaction mixture. ie. the solution is acidic.

The H itself is neither oxidised nor reduced but it is involved in the half equation/reaction

H2SO4 is often used to acidify these reactions as it is a strong acid that can provide a high H+ concentration without the acid itself being oxidised.

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15
Q

List Some Common Redox Reactions

A
  1. Metal-metal displacement
  2. Halogen-halide ion displacement
  3. Combustion
  4. Corrosion
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16
Q

Explain metal-metal ion displacement

example

A

Metal displacement reactions involve the TRANSFER OF ELECTRONS FROM A MORE REACTIVE METAL TO THE POSITIVE IONS OF LESS REACTIVE METAL.

the reaction will occur spontaneously if the metal being oxidised is in a solution of a less reactive metal that is going to be reduced and work out

  • electron transfer from one metal element to the metal ions of a less reactive metal.
  • the more reactive metal becomes oxidised as it loses electrons to form positive aqueous metal ions
  • simultaneously, positive aqueous metal ions of the less reactive metal become reduced as they gain electrons to form a solid metallic precipitate.

Cu(s) + 2Ag+(aq) —> 2Ag(s) + Cu+2
(0). (+1). (0). (+2)

17
Q

Explain combustion reactions in redox reactions

example

A

Combustion involves the OXIDATION OF A FUEL (eg H2 gas, methane, petrol or wood) and the REDUCTION OF OXYGEN GAS.

During combustion, the various elements in the fuel, typically carbon and hydrogen are converted TO THEIR OXIDES eg. CO2 and H2O.

An incomplete combustion reaction occurs if the air (o2) supply is limited (insufficient).
Carbon monoxide (CO) and even C (soot) can be produced

examples

ETHANE (COMPLETE) COMUSTION REACTION
2C2H6(G) +7O2(G) —> 4CO2(G) +6H2O(G)
-3. +1. 0. +4. -2 +1 -2

ETHANE (INCOMPLETE) COMUSTION REACTION
2C2H6(G) +7O2(G) —> 4CO2(G) +6H2O(G)
-3. +1. 0. +2. -2 +1 -2

18
Q

Explain corrosion redox reactions

EXAMPLE

A

Corrosion involves a metal (eg. Fe, Al, Zn) being oxidised in the presence of agents such as H2O and air (O2)

  • water is important for the process, OXYGEN GAS IS THE OXIDISING AGENT

Partial corrosion of iron
2Fe (S) + O2 (g) + 2H2O(L) —> 2Fe(OH)2 (S)
0. 0. +1 -2. +2. -2. +1

Chemistry - REDOX REACTIONS (8,9,10) (3 decks)

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