ReDox Notes.ppt

Complete this sentence.

When an atom gains electrons, it will take

on a _________ charge. When it loses

electrons, it will take on a __________

charge.

Bellwork

Chemistry 1

Chapter 22 &23

Oxidation Reduction Reactions

 _{Oxidation-reduction reactions- chemical}

changes that occur when electrons are transferred between reactants.

 _Oxidation

 _{Modern definition - loss of electrons}  _Examples

 4Fe + 3O₂  2Fe₂O₃ (rusting of iron)  C + O₂  CO₂ (burning of carbon)

 C₂H₅OH + 3O₂  2CO₂ + 3H₂O (burning of

ethanol)

 _Reduction

 _{Modern definition - gain of electrons}  2Fe₂O₃ + 3C  4Fe + CO₂  _{Oxidation and reduction always occur}

simultaneously. One process cannot occur without the other.

Oxidation

Is

Loss

Reduction

Is

Gain

LEO (Lose Electrons-Oxidation) the lion goes

GER (Gain Electrons-Reduction)

Question



When K becomes K

, it ______ an

Formation of Ions

 Ex. 2Na + S  Na₂S

 _{Sodium goes from the neutral atom to the 1+}

Oxidation Numbers

 _{An easy way to determine what element is}

oxidized and which element is reduced in a reaction is to assign each element an

oxidation number.

 Oxidation numbers give the charge on each

Question



Lead loses four electrons. It take

on a charge of ___. Does this

Rules for assigning oxidation

numbers:

(Remember when there is one element left,

don’t look at the rules anymore!)

•Sum of the oxidation #’s = overall charge of compound

•Free elements = 0

•Alkali metals (Li, Na, K, Rb, Cs)= +1 •Hydrogen = +1

•Group 2, Cd, & Zn = +2 •Oxygen = -2

•Group 7A (F, Cl, Br, I) = -1

EXAMPLE:

+1 -1

(+1) + (-1) = 0

HCl

P

O

+5 -2

(2x) + (-2 x 5) = 0 X = +5

H

O

-2 +1

O

0

Free elements are always 0

MnO

Na

SO

-2 +6

x + 4(-2) + 2(+1) = 0 x = +6

NH

+1

-3 _{x + 4(+1) = 1}

x = -3

AlPO

Ca(OH)

-2

-2 +1 +3 +5

+2

+3 + x + 4(-2) = 0 x = +5

1(+2) + 2(x) + 2(+1) = 0

x = -2

TiO

-2 +4

Bellwork

Friday, May 12

 _{Write the oxidation numbers for each of the}

following.

 K₂CrO₄ LiMnO₄

 H₂SO₄ SO₃

 _{Oxidizing agent- The substance in a redox}

reaction that accepts electrons (the

substance that was reduced). Whereas only an element can be reduced, the entire

compound is called the oxidizing agent.

 _{Reducing agent- The substance in a redox}

reaction that donates electrons (the

substance that was oxidized). Whereas only an element can be oxidized, the entire

compound is called the reducing agent.

REDOX Reaction Examples

 _{Identify the element oxidized, the element}

reduced, the oxidizing agent and the reducing agent for each of the following:

MnO

+ 4HCl



MnCl

+ Cl

+ 2H

O

Mn

Cl

+4

+2

Gained 2 e

-Oxidized

If it was Reduced, then the reactant that contains Mn acts as the “Oxidizing Agent”. (MnO₂)

-1

0

Lost 1 e

-Reduced

 _{Identify the element oxidized, the element}

reduced, the oxidizing agent and the reducing agent for each of the following:

Cu + HNO0 +1 -2 ₃  Cu(NO+2 -2 ₃)₂ + NO-2₂ + H+1 -2₂O

Cu

N

0

+2

Lost 2 e

Oxidized

+5

+4 Gained 1 e

Reduced

+5 +5 +4

If it was Reduced, then the reactant that contains N acts as the “Oxidizing Agent”. (HNO₃)

Let’s Practice! Write the oxidation numbers of the elements and then indicate which substance was oxidized, which was reduced and which are the oxidizing a reducing agents.

 2Mg + O₂  2MgO AgNO₃ + Cu  CuNO₃ +

Ag

 _{Oxidized: Oxidized:}

 _{Reduced: Reduced:}

 _{Ox. Agent: Ox. Agent:}

 _{Red. Agent: Red. Agent:}

Let’s Practice! Write the oxidation numbers of the elements and then indicate which substance was oxidized, which was reduced and which are the oxidizing a reducing agents.

CH₄ + O₂  CO + 2H₂O 2Fe₂O₃ + 3C  4Fe + CO₂

 _{Oxidized: Oxidized:}

 _{Reduced: Reduced:}

 _{Ox. Agent: Ox. Agent:}

 _{Red. Agent: Red. Agent:}

C O

O₂ CH₄

C Fe

Bellwork

Monday, April 15

 _{Write the oxidation states of the following substances.}

Determine which substance is oxidized, which is reduced, and the oxidizing and reducing agents.



2HNO

+ 6HI



2NO + 3I

+ 4H

O

 _Oxidized:  _Reduced:

Electrochemistry



The interconversions of chemical

and electrical energy that involve

the transfer of electrons.



This process occurs in a device

The Activity Series and

Redox

 _{We will utilize the activity series of metals}

shown to the left to determine reactivity in

single replacement redox reactions. Electrons are lost by the more active metals and gained by the less active metals. The higher a metal is on the activity series, the easier it is

oxidized.

 _{The simplified order from most reactive to}

least reactive:

 group I&II metals plus Al,  the transition metals,

 H

 "jewelry metals” (Cu, Hg, Ag, Au).

Let’s Practice!

Circle the more active metal!



Fe / Cu



Ca / Ag



Au / H



H / Na

Batteries

 _{Electrochemical cells that}

convert chemical energy into electrical energy are called voltaic cells.

 _{The energy is produced by}

spontaneous redox reactions.

 _{Voltaic cells can be separated}

into two half cells.

 _{A half cell consists of a metal}

Batteries

 _{We write half reactions to show what happens}

in each part of the cell.

 Example Write the half reactions that

occur in the Fe2+/Ni2+ cell.

 Oxidation Fe  Fe2+ + 2e

- Reduction Ni2+ + 2e-  Ni

Diagram of voltaic cell for the

reaction of zinc and copper

.

 _{Diagram of voltaic cell for the reaction of}

zinc and copper.

 _{Oxidized: Zn  Zn}2+ + 2e

- _{Reduced: Cu}2+ + 2e-  Cu

Direction of electron flow

Solution of _{Solution of}

From anode to cathode

anode

cathode

Zinc ions

EXAMPLE:

 _{Predict what will happen if an iron nail is dipped}

into a solution of CuSO₄. Write the balanced

equation and the two half-reactions (reduction and oxidation reactions).

 _{Iron is above copper on the activity series, so the reaction}

does occur and iron is oxidized:

 Fe(s) + CuSO₄(aq)  FeSO₄(aq) + Cu(s)  _{Oxidation: Fe  Fe}2+ + 2e

Half-Cells

 _{The half cells are connected by a salt bridge. A}

salt bridge is a tube containing a solution of ions.

 _{Ions pass through the salt bridge to keep the}

charges balanced.

 _{Electrons pass through an external wire.}  _{The metal rods in voltaic cells are called}

electrodes.

 _{Oxidation occurs at the anode and reduction}

occurs at the cathode. (An Ox and Red Cat)

 _{The direction of electron flow is from the anode to}

Calculating the Charge of a Battery

 _{The potential charge of a battery can be}

calculated with a set of values from a table of reduction potentials.

 To do this, write the oxidation and reduction

half reactions.

 Look up the cell potentials from the data

table.

 Flip the sign of the cell potential for oxidation.

Example A common battery is

made with nickel and cadmium.

What is the cell potential of this

battery? (E

Cd

= -0.40V, E

=

-0.25V)

Oxidation Cd  Cd2+ + 2e- E0 = 0.40V

Reduction Ni2+ + 2e-  Ni + E0 = -0.25V

 Total = E0 =