Oxidation Numbers Positive oxidation number Negative oxidation number - Loses partial or total control of electrons in a bond - Gains partial or total.

Oxidation Numbers

Positive oxidation number

Negative oxidation number

- Loses partial or total control of electrons in a bond

- Gains partial or total control of electrons in a bond

Example

Mg in the +2 state (Mg+2)

Has lost partial or total control of 2 e-

- Used to tell how many electrons an atom has lost or gained in a chemical reaction

Redox - When one atom loses an electron another must receive it

Na Na+ + 1e-

Cl + 1e- Cl-

1) Oxidation - The lose of an electron by an atom

- Causes an increase in oxidation number

2) Reduction - The gain of an electron by an atom

- Causes a decrease in oxidation number

If Fe loses three electrons, we write this as

Fe Fe+3 +3e-

If the electrons appear on the products side, they are given off

oxidation

If Fe+3 gains three electrons, we write this as

Fe+3 +3e- Fe If the electrons appear on the reactants side, they are gained

reduction

Lose Electrons Oxidation

Gain Electrons ReductionLEO goes GER

L E O

G E R

Indicate if the atoms shown have lost or gained electrons

Then write down if the atom is oxidized or reduced

A. H H+ + e-

B. Cl2 + 2e- 2 Cl-

C. Fe+2 + 2e- Fe

D. Cl+5 + 6e- Cl-

E. S-2 S+4 + 6 e-

Lost or gained e- Oxidized or reduced

Lost

Gained

oxidized

reduced

Lost

Gained

Gained

oxidized

reduced

reduced

Lightening Underwater pg 207

Oxidizing Agent - Causes another substance to be oxidized- Is reduced

Ex. Sn+4 + 2 e- Sn+2

Sn+4 is the oxidizing agent

Reducing Agent- Causes another substance to be reduced

- Is oxidized

Ex. Na Na+1 + 1 e-

Na is reducing agent

Blue bottle demo

Blue is oxidized = lost electrons

Colorless is reduced = more electrons

Finding Oxidation States

The most common oxidation numbers are in the upper right corner of your periodic table

Pb

+2+4

Oxidation states

Using your periodic table, list the oxidation states of the following

Ca

Na

O

N

Examples

+2

+1 -3, -2, -1 and a bunch more!

-2

Rules for assigning oxidation numbers

1. Uncombined elements have an oxidation number of zero

Cu0 Mg0 S0

2.) The oxidation state of an ion is the same as its charge

Oxidation state of zero means the atom is not losing or gaining any electrons

Cu+ Mg+2 S-2

+1 oxid. state +2 oxid. state -2 oxid. state

Cl20

3) Group 1 metals in a compound always have a +1 oxidation state

4) Group 2 metals in a compound always have a +2 oxidation state

MgCl2 LiBr Na2O+2 +1+1

5) Hydrogen always has a +1 oxidation state in a compound

Exception - When H is attached to a group 1 or 2 metal, it has a -1 oxidation state

HCl H2O+1 +1

NaH CaH2

-1-1

6) Oxygen in a compound always has a -2 oxid. state

Exception - In a peroxide, oxygen becomes -1

Na2O2 H2O2

-1 -1

H2O HNO3

-2 -2

7) Halogens are usually -1 in compounds

SO4-2

+6 -2

-8+6 = -2

The sum of the one sulfur and the 4 oxygens should be -2

8) The sum of all oxidation numbers in a polyatomic ion must equal the charge of the ion

Examples - Assign the oxidation numbers to all elements in the following ions

NO3- CrO4

-2+6

-6+5

+5 -2-2

+6 -8= -1 = -2

Assigning Oxidation Numbers

- All of the oxidation numbers of all atoms in a neutral compound must add up to zero

For example, in H2O,

H2O+1 -2

The two hydrogens add up to +2,

+2 + -2 = 0

Since water is a neutral compound, all oxidation states must add up to zero

Neutral compounds

Practice - Assign oxidation numbers to the following

MgCl2 LiOH Na2S H2SO4 NaNO3

+2 -1 +1 -2 +1 +1 -2 +1 +6 -2 +1 +5 -2

+2 -2 +1 +1-2 +2 -2 +2 -8+6 +1 +5 -6

0 0 0 0 0

Cu(NO3)2 NaH NO2

+2 +5 -2 +1 -1 +4 -2

-12+10+2 +1 -1 +4 -4

Recognizing Redox Equations

- Not all reactions are redox

- Double replacement reactions are NOT

- Single replacement reactions are

- If a reaction is redox then the oxidation numbers of some of the elements must be different on either side of the equationIs the reaction redox?

LiOH + HCl H2O + LiCl

2H2 + O2 2 H2O

Mg + CuSO4 MgSO4 + Cu

In the reaction Mg + Cl2 --> MgCl2

There are 2 steps occurring

Mg loses 2 electrons

Each chlorine gains an electron

MgCl

Cl

Electrons are moving. If e- could move through a wire, this would be an electric current

Electrochemistry

Half reaction- Either the reduction or the oxidation portion of a redox reaction

Al+3 + 3e- --> Al

Ca --> Ca+2 + 2e-

reduction

oxidation

The reduction ½ reaction shows an atom or ion gaining e- and the oxidation number decreasing

The oxidation ½ reaction shows an atom or ion losing e- and the oxidation number increasing

Steps in Writing Half Reactions

2 Li + CaBr2 --> 2 LiBr + Ca1. Assign oxidation numbers

+2 -1-1 +10 0

2. For each atom that changes its state, write down the starting and ending oxidation states

3. Add electrons to balance the equationCa+2 + 2e- Ca0 Li0 Li+ + e-

-Notice that each set of half reactions contains oxidationand reduction

-For each half reaction, both sides have the same charge

OxidationReduction

Ca+2 --> Ca Li --> Li+

Which reaction shows conservation of charge?

Fe(s) Fe 2+ (aq) + e- Fe + 2 e- Fe 2+

Fe + 2 e- Fe 3+

Fe(s) Fe 2+ (aq) + 2 e-

Examples

Write out the half reactions

A. KBr + Na K + NaBr

B. Sr + MgO Mg + SrO

C. NiCl + CuO NiO + CuCl

K+ +e- --> K

Na --> Na+ + e-

Sr --> Sr+2 + 2e-

Mg+2 + 2e- --> Mg

Ni+1 Ni+2 + 1 e-

Cu+2 + 1e- Cu+1

Mg + CuSO4 MgSO4 + Cu

Write the ½ reaction representing oxidation.

Given the reaction

Write the reduction ½ reaction for the following

2 Al + 3 Cu+2 2 Al 3+ + 3 Cu

Mg Mg+2 + 2 e-

Cu+2 + 2e- Cu

For the following

Zn + Cr 3+ Zn 2+ + Cr

a.) Write the ½ reaction for the reduction.

b.) Write the ½ reaction for the oxidation

c.) Which species loses electrons?

d.) What happens to the number of protons in a Zn atom when it changes to Zn 2+ as the redox reaction occurs.

If one of the diatomic elements move to the zero oxidation state, we must write the reaction with two atoms

ExampleWe cannot write

Cl-1 --> Cl + e-

We must write

2 Cl- --> Cl2 + 2e-

Other examples

Br2 + 2 e- --> 2 Br-

2 H+ + 2e- --> H2

B. Electrochemical CellsElectrochemical cell - Produces electricityelectricity from a divided redox reaction

- Electrons are moved from one atom to another through a wire

Voltaic Cell -(battery)- A spontaneous chemical reaction produces a flow of e-

Anode - Is the metal in a voltaic cell that is oxidized ( It will dissolve)

Cathode - The metal in a voltaic cell where reduction will occur

- Positive ions in solution will be reduced an collect on the cathode

External conductor (Wires) -permit flow of e-

Salt Bridge - U-tube containing electrolytic solution of + and - ions

- Allows migration of ions

- Keeps compartments neutral

These two jars are called ½ cells

How to identify the anode and cathode and the direction of e- flow

1. Find out which metal is oxidized / reduced

An Ox Anode is Oxidized

Table J: Best RA will lose e- so is the oxidizing ½ reaction

Red Cat Reduction happens at the Cathode

A reaction will be spontaneous if a pure metal is on the top, and the ion is below

Examples - Determine which atom/ion pairs will react spontaneously

A. Cu+2 and Ni

B. Mn+3 and Zn

C. K and Ni+2

D. Co and Al+3

YES

NO

YES

NO

As long as the solid metal is on top, the redox will be spontaneous.

Magnesium and silver nitrate flash reaction

Thermite Reaction

2. Write ½ reactions

LEO - Loss of electrons is oxidation

GER - Gain of electrons is reduction

LEO

An Ox

GER

RED CAT

Mg0 Mg+2 + 2e- Ni+2 + 2e Ni0

3) Choose sign for anode / cathode

- ( e- flow from – to + )

- (-) lost from anode gained by cathode (+)

- e- flow from Mg to Ni

As the reaction continuesMg Mg+2 + 2e- Ni+2 + 2e- Ni

Anode Cathode

Will decrease in mass Will increase in mass

Problem - Solutions must always be electrically neutral

As soon as we add -, we must add + chargesZn

Zn+2 Cu+2

Cu

SO4-2 SO4

-2

Each solution has a total charge of zero (neutral)

If we want to add more Zn+2 to solution, we need to add some - ions

If we want to remove some Cu+2 from solution, we need to add some + charges

d. Salt Bridge Tube containing a salt solution

Na+ Cl-

Porous plugs allow ions to move out of the tube

As Zn+2 goes into solution, 2 Cl- come out of the salt bridge

As Cu+2 is removed from the solution, 2 Na+ come out of the salt bridge

Solutions are kept neutral

Zn

Zn+2

Cu+2

Cu

Cu+2

Zne-e-

2e-

2e- 2e- 2e- 2e-

2e-

SO4-2 SO4

-2

SO4-2

Na+Na+ Cl- Cl-

Zn+2

Cl-Cl-

Cl-Cl-

Cl-Cl-

Now the Zn+2 solution is neutral

Na+Na+

Na+Na+

Na+Na+

Now the Cu+2 solution is neutral

E. Electrolytic Cell

Electrolytic cell An electric current is used to force a chemical reaction to occur

Electrolysis A chemical reaction which uses electricity to break apart a compound – nonspontaneous

1. Description

Only one cell is needed

Cell contains a power source

Battery pulls electrons off of one electrode

And puts them on another electrode+e-

+-

e-

e-e-

e-

e--

The battery decides what is + and - electrodes

If we add melted NaCl (MOLTEN NaCl) to the cell

Na+ and Cl- will be free to move

Na+ move to the negative electrode and gains electrons

This is reduction The negative electrode is the cathode

Cl- move to the positive electrode and loses electrons

This is oxidation The positive electrode is the anode

+

+-

-Na+ Cl-

Na+ + e- --> Na 2 Cl- --> Cl2 + 2e-

reduction

CATHODE

oxidation

ANODE

For electrolytic cells, the charges are opposite from galvanic cells

The “red cat” and “an ox” statements still work

2. Electroplating- Electrolytic cell is used to produce pure metals (Na, Mg)

Electroplating Layering a metal onto a surface using an electrolytic cell

Cell looks the same (one cell with a power source

+

+-

-

Put the object to be plated on the – electrode ( cathode)Put the layering metal on the + electrode

AgThis is supposed to be a copper ring!

+

+-

-Ag

The battery will pull electrons off of the Ag, turning it into Ag+

Ag --> Ag+ + e-

The Ag+ goes into solution and the e- go through the wire

e-

e-

e-e-

e-

e- Ag+

The Ag+ now move over to gain its own e- on the ring

The Ag+ gets metallically bonded to the copper

Ag+ + e- --> Ag

and reduction occurs at the cathode, which is negative

Again, oxidation occurs at the anode, which is positive

Reduction Oxidation

In electroplating, the same atom is oxidized and then reduced

Multicolored Electrolysis pg 240