Oxidation - Reduction Reactions

Oxidation - Reduction Chemistry

Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances -- this can be a complete transfer to form ionic bonds or a partial transfer to form covalent bonds In all redox reactions:

• one substance loses electrons -- this substance is oxidized • one substance gains electrons -- this substance is reduced There are lots of processes in the natural world (and in the laboratory) that involve redox reactions • e.g., corrosion, batteries, photosynthesis/respiration, etc.

Reaction between zinc and sulfuric acid Sulfuric acid (solution of Zn strip

H+

Reaction between zinc and sulfuric acid

2-

and SO4 ions)

Overall reaction

H2 bubbles

Zn(s) + H2SO4(aq)

ZnSO4(aq) + H2(g)

Zinc loses electrons Overall ionic reaction

• zinc is oxidized

Zn(s)

Zn2+(aq) + 2e-

-- all dissolved ions are explicitly shown

Zn(s) + 2 H+(aq) + SO42–(aq)

Zn2+(aq) + SO42–(aq) + H2(g)

Hydrogen gains electrons Net ionic reaction

• hydrogen is reduced

2 H+(aq) + 2eElectrons are transferred from zinc to hydrogen

H2(g)

-- includes only substances that undergo change -- ions that are present but do not react (spectator ions) are not shown

Zn(s) + 2 H+(aq)

Zn2+(aq) + H2(g)

Oxidation - Reduction Reactions Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances

• Oxidation occurs when a substance loses electrons • Reduction occurs when a substance gains electrons In a redox reaction, oxidation and reduction occur simultaneously -- one cannot occur in the absence of the other

Reaction between Cu and AgNO3 Cu(s) Ag+(aq)

Ag(s) initial

final

Cu2+(aq) NO3–(aq)

NO3–(aq) Oxidation of Cu: Cu(s) ! Cu2+(aq) + 2eReduction of Ag+: 2 Ag+(aq) + 2e- ! 2 Ag(s) Electrons are transferred from Cu atoms to Ag+ ions in solution

Cu loses electrons (oxidation)

initial

Cu(s)

Ag+(aq)

Cu(s) + 2 Ag+(aq)

final

Ag(s)

Cu2+(aq)

Cu2+(aq) + 2 Ag(s)

Ag+ gains electrons (reduction)

Voltaic cells

Overall Reaction: Cu(s) + 2 AgNO3(aq) ! 2 Ag(s) + Cu(NO3)2(aq) Overall Ionic Equation: Cu(s) + 2 Ag+(aq) + 2 NO3–(aq) ! 2 Ag(s) + Cu2+(aq) + 2 NO3–(aq) Net Ionic Equation: Cu(s) + 2 Ag+(aq) ! 2 Ag(s) + Cu2+(aq)

Zinc-copper voltaic cell

A voltaic cell is a device that produces an electric current from a spontaneous redox reaction

• the

oxidation reaction and the reduction reaction are physically separated and connected with a wire

• electrons transferred during the redox reaction must pass through the wire, producing an electric current

• the electric current is used to perform work - e.g., lighting a bulb, running an electric motor, etc. Chemical potential energy (energy stored in chemical bonds) is converted to electricity that is used to perform work

Oxidation of Zn: Zn(s) ! Zn2+(aq) + 2eReduction of Cu2+: Cu2+(aq) + 2e- ! Cu(s) Net Ionic: Zn(s) + Cu2+(aq) ! Cu(s) + Zn2+(aq) Electrons are transferred from Zn atoms to Cu2+ ions in solution

Dry cell batteries

Oxidation number

A dry cell battery is a small, efficient voltaic cell that contains a non-liquid electrolyte

metal cap (+)

The oxidation number of an atom is an integer value that represents the number of electrons gained, lost, or unequally shared by that atom

Alkaline-type dry cell battery

carbon rod (positive electrode)

Electrolyte is NaOH or KOH

zinc case (negative electrode)

Electrolyte is source of OHand H2O for redox reactions

manganese (IV) oxide moist paste of NaOH or KOH (electrolyte)

metal bottom (–)

• an oxidation number of zero indicates that the atom has the same number of electrons assigned to it as there are in the free, neutral atom • a positive oxidation number indicates that the atom has fewer electrons assigned to it than in the neutral atom • a negative oxidation number indicates that the atom has more electrons assigned to in than in the neutral atom

Oxidation of Zn: Zn(s) + 2 OH-(aq) ! ZnO(s) + H2O(l) + 2eReduction of Mn4+: 2 MnO2(s) + H2O(l) + 2e- ! Mn2O3(s) + 2 OH-(aq)

Oxidation number An element in its free state (uncombined with other elements) has an oxidation number of zero

•• Ba barium (Ba)

•• •• ••Cl •• Cl •• •• ••

chlorine (Cl2)

The oxidation number of an element that has gained or lost electrons to form an ion is that same as its positive or negative charge +2

Ba2+ barium ion (oxidation number: +2)

Clchloride ion (oxidation number: –1)

-1

BaCl2 In an ionic compound, the ions retain their oxidation number

Oxidation numbers of common monoatomic ions

Oxidation number In covalent compounds (shared electrons), oxidation numbers are assigned by an arbitrary system based on relative electronegativities of the atoms

H •• H hydrogen (H2)

•• •• •• Cl •• Cl •• •• •• chlorine (Cl2)

In diatomic molecules containing only one element (nonpolar covalent bonding), the bonding pair of electrons is shared equally between the atoms (" electronegativity = 0) • each atom is assigned an oxidation number of zero

Oxidation number In covalent compounds containing different elements, the bonding atoms are shared unequally between the atoms -- the higher the electronegativity of the atom, the greater its affinity for the shared electrons In these types of compounds, oxidation numbers are determined by assigning both bonding electrons to the most electronegative atom Example: Water +1

Electronegativity Oxygen: 3.5 Hydrogen: 2.1

–2

•• H •• O •• H ••

both bonding electrons assigned to oxygen

+1

both bonding electrons assigned to oxygen

Many elements have multiple oxidation numbers

Many elements have multiple oxidation numbers

It depends on the types of compounds they form

It depends on the types of compounds they form

N oxidation number

N2

N2O

NO

N2O3

NO2

N2O5

NO3–

0

+1

+2

+3

+4

+5

+5

Note: The oxidation number of oxygen is –2 in all of these compounds Elemental copper (Cu)

Copper (II) sulfate (CuSO4)

Cu oxidation state: 0

Cu oxidation state: +2

Rules for assigning oxidation numbers 1. All elements in their free state (uncombined with other elements) have an oxidation number of zero (e.g., Na, Cu, Mg, H2, O2, Cl2, etc.) 2. H is +1, except in metal hydrides, where it is -1 (e.g., NaH, CaH2) 3. O is -2, except in peroxides, where it is -1, and in OF2, where it is +2 4. The metallic element in an ionic compound has a positive oxidation number 5. In covalent compounds, the most electronegative element is assigned a negative oxidation number 6. The sum of the oxidation numbers of the elements in a neutral compound is zero

Rules for determining the oxidation number of an element within a compound Step 1: Write the oxidation number of each known atom below the atom in the formula Step 2: Multiply each oxidation number by the number of atoms of that element in the compound Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero

7. The sum of the oxidation numbers of the elements in a polyatomic ion is equal to the charge of the ion

Example: Determine the oxidation number of carbon in carbon dioxide

CO2 -2 2(-2) = -4 4 + C + (-4) = 0 + 4

Example: Determine the oxidation number of sulfur in sulfuric acid

H2SO4 +1 2(+1) = +2

-2 4(-2) = -8

(–2) + 8 +2 + S + (-8) = 0 (–2) + 8

C = +4

S = +6

(oxidation number for carbon)

(oxidation number for sulfur)

Step 1: Write the oxidation number of each known atom below the atom in the formula

Step 1: Write the oxidation number of each known atom below the atom in the formula

Step 2: Multiply each oxidation number by the number of atoms of that element in the compound

Step 2: Multiply each oxidation number by the number of atoms of that element in the compound

Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero

Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero

Example: Determine the oxidation number of chromium in Cr2O72-

Example: Determine the oxidation number of potassium and nitrogen in KNO3

Cr2O72-

KNO3

-2

K+

NO3–

Recognize that KNO3 is an ionic compound between K+ and NO3-

7(-2) = -14 2Cr + (-14) = -2 (the charge on the ion) Cr = +6 (oxidation number for chromium) Step 1: Write the oxidation number of each known atom below the atom in the formula

The oxidation number of potassium in K+ is +1 (the charge on the ion) For nitrogen:

NO3– -2 3(-2) = -6

Step 2: Multiply each oxidation number by the number of atoms of that element in the compound

N + (-6) = -1 (the charge on the ion)

Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero

N = +5 (oxidation number for nitrogen)

Oxidation - Reduction Reactions

Oxidation - Reduction Reactions

Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances

Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances

-- this can be a complete transfer to form ionic bonds or a partial transfer to form covalent bonds

-- this can be a complete transfer to form ionic bonds or a partial transfer to form covalent bonds

L ose E lectrons O xidation

G ain E lectrons R eduction

In redox reactions, the oxidation numbers of the elements involved in the reaction change

• oxidation of an element (loss of electrons) results in an increase in its oxidation number

• reduction of an element (gain of electrons) results in a decrease in its oxidation number

Reaction can be rewritten to emphasize electron transfer

Reaction between zinc and sulfuric acid

Zn(s) + H2SO4(aq)

ZnSO4(aq) + H2(g)

Zn + 2 H+ + SO42-

Zn2+ + SO42- + H2

0

+2

0

+2

+1

0

+1

0

•zinc loses electrons

•zinc loses electrons

•the oxidation number of Zn increases •zinc is oxidized

•the oxidation number of Zn increases •zinc is oxidized

•hydrogen gains electrons

•hydrogen gains electrons

•the oxidation number of H decreases •hydrogen is reduced

•the oxidation number of H decreases •hydrogen is reduced

Electrons are transferred from zinc to hydrogen

Electrons are transferred from zinc to hydrogen

Reaction between zinc and sulfuric acid

Oxidizing and reducing agents Oxidizing agent: The reactant that causes another substance to be oxidized – i.e., the reactant that causes an increase in the oxidation state of another substance The oxidizing agent is reduced in a redox reaction Reducing agent: The reactant that causes another substance to be reduced – i.e., the reactant that causes a decrease in the oxidation state of another substance The reducing agent is oxidized in a redox reaction

Zn(s) + H2SO4(aq)

ZnSO4(aq) + H2(g)

0

+2

+1

0

• • • •

zinc loses electrons the oxidation number of Zn increases zinc is oxidized zinc is the reducing agent (it causes hydrogen to be reduced)

• • • •

hydrogen gains electrons the oxidation number of H decreases hydrogen is reduced sulfuric acid is the oxidizing agent (it causes zinc to be oxidized)

Example: Is the following a redox reaction?

Example: Is the following a redox reaction?

Neutralization reaction between hydrochloric acid and potassium hydroxide

Thermite reaction

HCl (aq) + KOH (aq) +1

-1

Element

+1

H2O (l) + KCl (aq)

-2 +1

+1

Oxidation number Oxidation number before reaction after reaction

H

+1

+1

O

–2

–2

K

+1

+1

Cl

–1

–1

-2

+1

-1

All oxidation numbers unchanged No redox reactions occurred

Homework assignment Chapter 6 Problems: 6.76, 6.77, 6.78, 6.79, 6.84, 6.87, 6.88, 6.89, 6.90, 6.91, 6.92

2 Al (s) + Fe2O3 (s) 0

Element

+3

Al2O3 (l) + 2 Fe (l)

-2

+3

-2

Oxidation number Oxidation number before reaction after reaction

Al

0

+3

Fe

+3

0

O

–2

–2

0

Which element is oxidized?

Al

Which element is reduced?

Fe