Electron Transfer Reactions Electron transfer reactions are oxidation-
reduction or redox reactions. Results in the generation of an electric current
(electricity) or be caused by imposing an electric current.
Therefore, this field of chemistry is often called ELECTROCHEMISTRY.
Terminology for Redox Reactions OXIDATION—loss of electron(s) by a
species; increase in oxidation number; increase in oxygen.
REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.
OXIDIZING AGENT—electron acceptor; species is reduced. (an agent facilitates something; ex. Travel agents don’t travel, they facilitate travel)
REDUCING AGENT—electron donor; species is oxidized.
You can’t have one… without the other! Reduction (gaining electrons) can’t happen without
an oxidation to provide the electrons. You can’t have 2 oxidations or 2 reductions in the
same equation. Reduction has to occur at the cost of oxidation
LEO the lion says GER!oselectrons
xidation
ainlectrons
eduction
GER!
Another way to remember
OIL RIGxidation
s ose
eduction
s ain
Oxidation Numbers
OBJECTIVES Determine the oxidation number of an atom of any element in a pure substance.
Oxidation Numbers
OBJECTIVES Define oxidation and reduction in terms of a change in oxidation number, and identify atoms being oxidized or reduced in redox reactions.
Assigning Oxidation Numbers
• An “oxidation number” is a positive or negative number assigned to an atom to indicate its degree of oxidation or reduction.
• Generally, a bonded atom’s oxidation number is the charge it would have if the electrons in the bond were assigned to the atom of the more electronegative element
Rules for Assigning Oxidation Numbers
1)The oxidation number of any uncombined element is zero.
2)The oxidation number of a monatomic ion equals its charge.
11
2
0022
ClNaClNa
Rules for Assigning Oxidation Numbers
3)The oxidation number of oxygen in compounds is -2, except in peroxides, such as H2O2 where it is -1.
4)The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1.
2
2
1
OH
Rules for Assigning Oxidation Numbers
5) The sum of the oxidation numbers of the atoms in the compound must equal 0.
2
2
1
OH2(+1) + (-2) = 0 H O
2
122)(
HOCa(+2) + 2(-2) + 2(+1) = 0 Ca O H
Rules for Assigning Oxidation Numbers
6) The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge.
3
2?ON
X + 3(-2) = -1N O
24
2?OS
thus X = +5 thus X = +6
X + 4(-2) = -2S O
Review of Oxidation numbers
The charge the atom would have in a molecule (or anionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 02. In monatomic ions, the oxidation number is equal to the
charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -23. The oxidation number of oxygen is usually –2. In H2O2 and O2
2- it is –1.
4.4
4. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.
6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1.
HCO3-
O = -2 H = +13x(-2) + 1 + ? = -1
C = +4
Oxidation numbers of all the atoms in HCO3
- ?
4.4
Oxidation and Reduction (Redox)
A process called “reduction” is the opposite of oxidation, and originally meant the loss of oxygen from a compound Oxidation and reduction always occur simultaneously The substance gaining oxygen (or losing electrons) is oxidized, while the substance losing oxygen (or gaining electrons) is reduced.
Oxidation and Reduction (Redox)
Today, many of these reactions may not even involve oxygen Redox currently says that electrons are transferred between reactants Mg + S → Mg2+ + S2-
• The magnesium atom (which has zero charge) changes to a magnesium ion by losing 2 electrons, and is oxidized to Mg2+
• The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons, and is reduced to S2-
(MgS)
Oxidation and Reduction (Redox)11
2
0022
ClNaClNa
Each sodium atom loses one electron:
Each chlorine atom gains one electron:
eNaNa10
10 CleCl
LEO says GER :
eNaNa10
Lose Electrons = Oxidation
Sodium is oxidized
Gain Electrons = Reduction10
CleCl Chlorine is reduced
LEO says GER :
- Losing electrons is oxidation, and the substance that loses the electrons is called the reducing agent. - Gaining electrons is reduction, and the substance that gains the electrons is called the oxidizing agent.
Mg(s) + S(s) → MgS(s)
Mg is oxidized: loses e-, becomes a Mg2+ ion
S is reduced: gains e- = S2- ion
Mg is the reducing agent
S is the oxidizing agent
Oxidation and Reduction (Redox)
It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds? In water, we learned that oxygen is highly electronegative, so: the oxygen gains electrons (is reduced and is the oxidizing agent), and the hydrogen loses electrons (is oxidized and is the reducing agent)
Not All Reactions are Redox Reactions
- Reactions in which there has been no change in oxidation number are NOT redox reactions.
Examples:
)()()()( 3
2511111
3
251aqONNasClAgaqClNaaqONAg
)()()()(22
2
1
4
26
2
1
4
26
2
1121lOHaqOSNaaqOSHaqHONa
Corrosion•Damage done to metal is costly to prevent and repair
• Iron, a common construction metal often used in forming steel alloys, corrodes by being oxidized to ions of iron by oxygen.•This corrosion is even faster in the presence of salts and acids, because these materials make electrically conductive solutions that make electron transfer easy
Corrosion•Luckily, not all metals corrode easily•Gold and platinum are called noble metals because they are resistant to losing their electrons by corrosion
•Other metals may lose their electrons easily, but are protected from corrosion by the oxide coating on their surface, such as aluminum
• Iron has an oxide coating, but it is not tightly packed, so water and air can penetrate it easily
Corrosion•Serious problems can result if bridges, storage tanks, or hulls of ships corrode•Can be prevented by a coating of oil, paint, plastic, or another metal
• If this surface is scratched or worn away, the protection is lost
•Other methods of prevention involve the “sacrifice” of one metal to save the second•Magnesium, chromium, or even zinc (called galvanized) coatings can be applied
Reducing Agents and Oxidizing Agents
• • An increase in oxidation number = oxidation• A decrease in oxidation number = reduction
Na0
Na1
e
Cl0
e Cl 1
Sodium is oxidized – it is the reducing agent
Chlorine is reduced – it is the oxidizing agent
Trends in Oxidation and ReductionActive metals:
Lose electrons easily Are easily oxidized Are strong reducing agents
Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents
Balancing Redox Equations
OBJECTIVES Describe how oxidation numbers are used to identify redox reactions.
Balancing Redox Equations
OBJECTIVES Balance a redox equation using the oxidation-number-change method.
Balancing Redox Equations
OBJECTIVES Balance a redox equation by breaking the equation into oxidation and reduction half-reactions, and then using the half-reaction method.
Identifying Redox Equations In general, all chemical reactions can be assigned to one of two classes:1)oxidation-reduction, in which
electrons are transferred:• Single-replacement, combination,
decomposition, and combustion2) this second class has no electron
transfer, and includes all others:• Double-replacement and acid-
base reactions
Identifying Redox Equations In an electrical storm, nitrogen and oxygen react to form nitrogen monoxide:
N2(g) + O2(g) → 2NO(g)
• Is this a redox reaction?• If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox.
YES!
Balancing Redox Equations It is essential to write a correctly balanced equation that represents what happens in a chemical reaction• Fortunately, two systematic methods
are available, and are based on the fact that the total electrons gained in reduction equals the total lost in oxidation. The two methods:
1)Use oxidation-number changes2)Use half-reactions
Using Oxidation-Number Changes
Sort of like chemical bookkeeping, you compare the increases and decreases in oxidation numbers.• start with the skeleton equation• Step 1: assign oxidation numbers to all atoms; write above their
symbols• Step 2: identify which are oxidized/reduced• Step 3: use bracket lines to connect them• Step 4: use coefficients to equalize• Step 5: make sure they are balanced for both atoms and charge
Using half-reactions A half-reaction is an equation showing just the oxidation or just the reduction that takes place they are then balanced separately, and finally combined
Step 1: write unbalanced equation in ionic formStep 2: write separate half-reaction equations for oxidation and reductionStep 3: balance the atoms in the half-reactions
Using half-reactions continued• Step 4: add enough electrons to one side of each half-reaction to balance the charges
• Step 5: multiply each half-reaction by a number to make the electrons equal in both
• Step 6: add the balanced half-reactions to show an overall equation
• Step 7: add the spectator ions and balance the equation
Choosing a Balancing Method1) The oxidation number change
method works well if the oxidized and reduced species appear only once on each side of the equation, and there are no acids or bases.
2) The half-reaction method works best for reactions taking place in acidic or alkaline solution.
Balancing Redox Equations
19.1
1. Write the unbalanced equation for the reaction ion ionic form.
The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?
Fe2+ + Cr2O72- Fe3+ + Cr3+
2. Separate the equation into two half-reactions.
Oxidation:
Cr2O72- Cr3+
+6 +3Reduction:
Fe2+ Fe3++2 +3
3. Balance the atoms other than O and H in each half-reaction.
Cr2O72- 2Cr3+
Balancing Redox Equations4. For reactions in acid, add H2O to balance O atoms and H+ to balance H
atoms.
Cr2O72- 2Cr3+ + 7H2O
14H+ + Cr2O72- 2Cr3+ + 7H2O
5. Add electrons to one side of each half-reaction to balance the charges on the half-reaction.
Fe2+ Fe3+ + 1e-
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
6. If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients.
6Fe2+ 6Fe3+ + 6e-
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
19.1
Balancing Redox Equations7. Add the two half-reactions together and balance the final equation by
inspection. The number of electrons on both sides must cancel. You should also cancel like species.
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
6Fe2+ 6Fe3+ + 6e-Oxidation:
Reduction:
14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O
8. Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6x2 = 24 = 6x3 + 2x3
19.1
9. For reactions in basic solutions, add OH- to both sides of the equation for every H+ that appears in the final equation. You should combine H+ and OH- to make H2O.
CHEMICAL CHANGE --->ELECTRIC CURRENT
To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire.
Zn
Zn2+ ions
Cu
Cu2+ ions
wire
saltbridge
electrons
This is accomplished in a GALVANIC or VOLTAIC cell.
A group of such cells is called a battery.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf
Galvanic Cells
19.2
spontaneousredox reaction
anodeoxidation
cathodereduction
- +
