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Chapter 3 chemical reactions

Chapter 3Chemical Reactions


Important – Read Before Using Slides in Class

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Chemical reactions
Chemical Reactions

Reactants: Zn + I2

Product: ZnI2


Chemical reactions1
Chemical Reactions

  • Evidence of a chemical reaction:

    • Gas Evolution

    • Temperature Change

    • Color Change

    • Precipitation (insoluble species forms)

  • In general, a reaction involves a rearrangement or change in oxidation state of atoms from reactants to products.


Chemical equations
Chemical Equations

Chemical Equations show:

  • thereactantsandproductsin a reaction.

  • the relative amounts in a reaction.

    Example:

    4 Al(s) + 3 O2(g)  2 Al2O3(s)

  • The numbers in the front are called

    stoichiometric coefficients

  • The letters (s), (g), (l) and (aq) are the physical states of compounds.


Reaction of phosphorus with cl 2
Reaction of Phosphorus with Cl2

Notice the stoichiometric coefficients and the physical states of the reactants and products.


Reaction of iron with cl 2
Reaction of Iron with Cl2

Notice the stoichiometric coefficients and the physical states of the reactants and products.


Chemical equations1
Chemical Equations

4 Al(s) + 3 O2(g)  2 Al2O3(s)

This equation states that:

4 Al atoms + 3 O2 molecules react to form 2 formula units of Al2O3

or...

4 moles of Al + 3 moles of O2 react to form 2 moles of Al2O3


Chemical equations2
Chemical Equations

Law of the Conservation of Matter

  • Because the same number of atoms are present in a reaction at the beginning and at the end, the amount of matter in a system does not change.

2HgO(s) 2 Hg(l) + O2(g)


Chemical equations3

Lavoisier, 1788

Chemical Equations

  • Since matter is conserved in a chemical reaction, chemical equations must be balanced for mass!

  • In other words, there must be same number of atoms of the each kind on both sides of the equatoin.


Balancing chemical reactions
Balancing Chemical Reactions

  • Steps in balancing a chemical reaction using coefficients:

  • Write the equation using the formulas of the reactants and products. Include the physical states (s, l, g, aq etc…)

  • Balance the compound with the most elements in the formula first using integers as coefficients.

  • Balance elements on their own last.

  • Check to see that the sum of each individual elements are equal on each side of the equation.

  • If the coefficients can be simplified by dividing though with a whole number, do so.


Balancing chemical equations example
Balancing Chemical Equations: Example

balance last

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s

2 O’s

1 C & 2 O’s

2 H’s & 1 O


Balancing chemical equations example1
Balancing Chemical Equations: Example

balance last

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s

2 O’s

1 C & 2 O’s

2 H’s & 1 O

balance H first

___C2H6 + O2 CO2 + ___ H2O

3

This side will always have an even # of O-atoms

This side has an odd # of O-atoms


Balancing chemical equations example2
Balancing Chemical Equations: Example

balance last

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s

2 O’s

1 C & 2 O’s

2 H’s & 1 O

balance H first

___C2H6 + O2 CO2 + ___ H2O

2

3


Balancing chemical equations example3
Balancing Chemical Equations: Example

balance last

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s

2 O’s

1 C & 2 O’s

2 H’s & 1 O

balance H first

___C2H6 + O2 CO2 + ___ H2O

2

3

balance C next

4

2C2H6 + O2 ___ CO2 + 6H2O


Balancing chemical equations example4
Balancing Chemical Equations: Example

balance last

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s

2 O’s

1 C & 2 O’s

2 H’s & 1 O

balance H first

___C2H6 + O2 CO2 + ___ H2O

2

3

balance C next

4

2C2H6 + O2 ___ CO2 + 6H2O

balance O

7

2C2H6 + ____ O2 4CO2 + 6H2O


Balancing chemical equations example5
Balancing Chemical Equations: Example

balance last

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s

2 O’s

1 C & 2 O’s

2 H’s & 1 O

balance H first

___C2H6 + O2 CO2 + ___ H2O

2

3

balance C next

4

2C2H6 + O2 ___ CO2 + 6H2O

balance O

7

2C2H6 + ____ O2 4CO2 + 6H2O

4 C’s 12 H’s 14 O’s

4 C’s 12 H’s 14 O’s


Balancing equations
Balancing Equations

___ Al(s) + ___ Br2(l)  ___ Al2Br6(s)


Balancing equations practice
Balancing Equations: Practice

___C3H8(g) + ___ O2(g) 

___ CO2(g) + _____ H2O(g)

___B4H10(g) + ___ O2(g) 

___ B2O3(g) + ___ H2O(g)


Balancing equations practice1
Balancing Equations: Practice

  • Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water.

  • Write the balanced chemical equation for this reaction.


Balancing equations practice2
Balancing Equations: Practice

  • Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water.

  • Write the balanced chemical equation for this reaction.

_ Mg(OH)2(s) + _ HCl(aq) 


Balancing equations practice3
Balancing Equations: Practice

  • Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water.

  • Write the balanced chemical equation for this reaction.

_ Mg(OH)2(s) + _ HCl(aq)  _ MgCl2(aq) + _ H2O(l)


Balancing equations practice4
Balancing Equations: Practice

  • Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water.

  • Write the balanced chemical equation for this reaction.

_ Mg(OH)2(s) + _ HCl(aq)  _ MgCl2(aq) + _ H2O(l)

Balance with a coefficient of “2” in front of both HCl and water.


Balancing equations practice5
Balancing Equations: Practice

  • Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water.

  • Write the balanced chemical equation for this reaction.

_ Mg(OH)2(s) + _ HCl(aq)  _ MgCl2(aq) + _ H2O(l)

Balance with a coefficient of “2” in front of both HCl and water.

Mg, Cl, O and H are now balanced.


Chemical equations review
Chemical Equations: Review

  • What Scientific Principles are used in the process of balancing chemical equations?

  • What symbols are used in chemical equations:

    gasses: _____

    liquids: _____

    solids: _____

    aqueous species in solution: _____

  • What is the difference between P4 and 4P in an eq.?

  • In balancing a chemical equation, why are the reactant and product subscripts not changed?


Chemical equilibrium
Chemical Equilibrium

When writing chemical reactions one starts with:

Reactants

products

2NH3(g)

N2(g) + 3H2(g)

Some reactions can also run in reverse:

2NH3(g)

N2(g) + 3H2(g)

Under these conditions, the reaction can be written:

Double arrows indicate “Equilibrium”.


Chemical equilibrium1
Chemical Equilibrium

Once equilibrium is achieved, reaction continues, but there is no net change in amounts of products or reactants.


Classifying compounds
Classifying Compounds

  • Salts (ionic compounds): Composed of a metal and non metal element(s).

  • Acids: Arrhenius definition

    Produce H+(aq) in water

    Examples: HCl, HNO3, HC2H3O2

  • Bases: Arrhenius definition

    Produce OH(aq) in water

    Examples: NaOH, Ba(OH)2, NH3


Classifying compounds1
Classifying Compounds

  • Molecular Compounds:

  • Covalently bonded atoms, not acids, bases or salts.

  • Compounds like alcohols (C2H5OH) or table sugar (C6H12O6)

  • These never break up into ions.


Classifying compounds2
Classifying Compounds

  • Classify the following as ionic, molecular, acid or base.


Classifying compounds3
Classifying Compounds

  • Classify the following as ionic, molecular, acid or base.


Reactions in aqueous solutions

That which is dissolved (lesser amount)

That which is dissolves (greater amount)

Reactions in Aqueous Solutions

Aqueous Solutions:

Water as the solvent

Solution =

solute

+solvent

There are three types of aqueous solutions:

Those with Strong Electrolytes

Those with Weak Electrolytes

& those with non-Electrolytes


Reactions in aqueous solutions1

K+(aq) + MnO4-(aq)

Reactions in Aqueous Solutions

Many reactions involve ionic compounds, especially reactions in water — aqueous solutions.

KMnO4 in water



Strong electrolyte
Strong Electrolyte

When ions are present in water, the solutions conduct electricity!

Ions in solution are called ELECTROLYTES

Examples of Strong Electrolytes:

HCl (aq), CuCl2(aq) and NaCl (aq) are strong electrolytes.

These dissociate completely (or nearly so) into ions.

Strong Electrolytes conduct electricity well.


Strong electrolytes
Strong Electrolytes

HCl(aq), CuCl2(aq) and NaCl(aq) are strong electrolytes.

These dissociate completely (or nearly so) into ions.


Weak electrolytes
Weak Electrolytes

Acetic acid ionizes only to a small extent, it is a weak electrolyte.

Weak electrolytes exist in solution under equilibrium conditions.

The small concentration of ions conducts electricity poorly.

Weak electrolytes exit primarily in their molecular form in water.


Weak electrolytes1
Weak Electrolytes

Weak electrolytic solutions are characterized by equilibrium conditions in solution:

When acetic acid dissociates, it only partially ionizes.

95%

5%

The majority species in solution is acetic acid in its molecular form.

When writing a weak electrolyte in solution, one NEVERbreaks it up into the corresponding ions!

×


Weak electrolytes2
Weak Electrolytes

Acetic acid ionizes only to a small extent, so it is a weak electrolyte.

CH3CO2H(aq)  CH3CO2-(aq) + H+(aq)


Non electrolytes
Non-Electrolytes

Some compounds dissolve in water but do not conduct electricity.

They are non-electrolytes.

Examples include:

  • sugar

  • ethanol

  • ethylene glycol

    Non-electrolytes do not dissociate into ions!


Species in Solution: Electrolytes

Strong electrolytes:

Characterized by ions only (cations & anions) in solution (water).

Conduct electricity well

Characterized by ions (cations & anions) & molecules in solution.

Weak electrolytes:

Conduct electricity poorly

Non-electrolytes:

Characterized by molecules in solution.

Do not conduct electricity



Solubility rules
Solubility Rules

  • How do we know if a compound will be soluble in water?

    • For molecular compounds, the molecule must be polar.

    • We will discuss polarity later, for now I will tell you whether or not a molecular compound is polar…

    • For ionic compounds, the compound solubility is governed by a set of SOLUBILITY RULES!

  • You must learn the basic rules on your own!!!


Water solubility of ionic compounds
Water Solubility of Ionic Compounds

If one ion from the “Soluble Compound” list is present in a compound, then the compound is water soluble.


Types of reactions in a solution
Types of Reactions in a Solution

Precipitation Reactions: A reaction where an insoluble solid (precipitate) forms and drops out of the solution.

Acid–base Neutralization: A reaction in which an acid reacts with a base to yield water plus a salt.

Gas forming Reactions: A reaction where an insoluble gas is formed.

Reduction and Oxidation Reactions (RedOx):A reaction where electrons are transferred from one reactant to another.


EXCHANGE:Precipitation Reactions

EXCHANGE

Acid-Base

Reactions

EXCHANGE

Gas-Forming

Reactions

REACTIONS

REDOX REACTIONS


Chemical reactions in water
Chemical Reactions in Water

EXCHANGE REACTIONS

The anions exchange places between cations.

A precipitate forms if one of the products in insoluble.

Pb(NO3) 2(aq) + 2 KI(aq)

 PbI2(s) + 2 KNO3 (aq)


Precipitation reactions
Precipitation Reactions

The “driving force” is the formation of an insoluble solid called a precipitate.

Pb(NO3)2(aq) + 2 KI(aq) 

2 KNO3(aq) + PbI2(s)

BaCl2(aq) + Na2SO4(aq) 

BaSO4(s) + 2 NaCl(aq)

Precipitates are determined from the solubility rules.


Precipitation reactions1
Precipitation Reactions

Which species is the precipitate?

Pb(NO3)2(?)

+ 2KI(?)

 2KNO3(?)

+ PbI2(?)


Precipitation reactions2
Precipitation reactions

Which species is the precipitate?

Pb(NO3)2(?)

+ 2KI(?)

 2KNO3(?)

+ PbI2(?)

From the solubility rules:

All nitrate salts are soluble, therefore:

Pb(NO3)2(aq)

+ 2KI(?)

 2KNO3(aq)

+ PbI2(?)


Precipitation reactions3
Precipitation Reactions

Which species is the precipitate?

Pb(NO3)2(?)

+ 2KI(?)

 2KNO3(?)

+ PbI2(?)

From the solubility rules:

All nitrate salts are soluble, therefore:

Pb(NO3)2(aq)

+ 2KI(?)

 2KNO3(aq)

+ PbI2(?)

All potassium salts are soluble, therefore:

Pb(NO3)2(aq)

+ 2KI(aq)

 2KNO3(aq)

+ PbI2(?)


Precipitation reactions4
Precipitation Reactions

Which species is the precipitate?

Pb(NO3)2(?)

+ 2KI(?)

 2KNO3(?)

+ PbI2(?)

From the solubility rules:

All nitrate salts are soluble, therefore:

Pb(NO3)2(aq)

+ 2KI(?)

 2KNO3(aq)

+ PbI2(?)

All potassium salts are soluble, therefore:

Pb(NO3)2(aq)

+ 2KI(aq)

 2KNO3(aq)

+ PbI2(?)

By the solubility rules:

PbI2 is the ppt.

Pb(NO3)2(aq)

+ 2KI(aq)

 2KNO3(aq)

+ PbI2(s)


Net ionic equations
Net Ionic Equations

Molecular Equation:all species listed as formula units or in molecular form. reactants  products

  • Note all states of each reactant or product by: (s), (l), (g) or (aq)

    Ionic Equation:All soluble (aq) species present are listed as ions.

  • Leave all (s), (l) or (g) species as is. They do not dissociate into ions

    Net Ionic Equation:

  • From the ionic equation, cancel out any species that appear on either side of the equation.

  • These are known as the “spectator ions” and they are never part of a net ionic equation!


Writing net ionic equations
Writing Net Ionic Equations

Molecular Equation:

Pb(NO3)2(aq)

+ 2KI(aq)

 2KNO3(aq)

+ PbI2(s)


Writing net ionic equations1
Writing Net Ionic Equations

Molecular Equation:

Pb(NO3)2(aq)

+ 2KI(aq)

 2KNO3(aq)

+ PbI2(s)

Total Ionic Equation:

Pb2+ (aq) + 2NO3– (aq)

+ 2K+(aq) + 2I–(aq) 

2K+(aq) + 2NO3– (aq)

+ PbI2(s)


Writing net ionic equations2
Writing Net Ionic Equations

Molecular Equation:

Pb(NO3)2(aq)

+ 2KI(aq)

 2KNO3(aq)

+ PbI2(s)

Total Ionic Equation:

Never break up any (s), (l) or (g) or molecular (aq) species!

Pb2+ (aq) + 2NO3– (aq)

+ 2K+(aq) + 2I–(aq) 

2K+(aq) + 2NO3– (aq)

+ PbI2(s)


Writing net ionic equations3
Writing Net Ionic Equations

Molecular Equation:

Pb(NO3)2(aq)

+ 2KI(aq)

 2KNO3(aq)

+ PbI2(s)

Total Ionic Equation:

Never break up any (s), (l) or (g) or molecular (aq) species!

Pb2+ (aq) + 2NO3– (aq)

+ 2K+(aq) + 2I–(aq) 

2K+(aq) + 2NO3– (aq)

+ PbI2(s)

Cancel out the spectator ions to yield the net ionic equation:

PbI2(s)

+ 2I–(aq) 

Pb2+ (aq)


Acids bases
Acids & Bases

Arrhenius Definition:

  • An acid is any substance that increases the H+(aq) concentration in an aqueous solution.

    HX(aq)  H+(aq) + X–(aq)

  • A base is any substance that increases the OH–(aq) concentration in an aqueous solution.

    MOH(aq)  M+(aq) + OH–(aq)


Acids and bases
Acids and Bases

Brönsted-Lowry:

  • An acid is any substance that donates H+(aq) to another species in an aqueous solution.

    HX(aq) + H2O(l) H3O+(aq) + X–(aq)

  • A base is any substance that accepts an H+(aq) in an aqueous solution.

    H+(aq) + NH3(aq) NH4+(aq)

H3O+(aq) = H+(aq)



Strong acids
Strong Acids

Examples:

Strong acidsare almost completely ionized in water. (strong electrolytes)

HX (aq) (X = Cl, Br & I)

hydro ___ ic acid

HNO3 (aq)

nitric acid

perchloric acid

HClO4 (aq)

sulfuric acid

H2SO4 (aq)*

* Only the 1st H is strong, sulfuric acid dissociates via:

H2SO4 (aq)  H+ (aq) + HSO4– (aq)


Acids1
Acids

An acid: H3O+ in water


Weak acids
Weak Acids

Examples:

Weak Acidsare incompletely ionized in water. (weak electrolytes) Weak acids are governed by dynamic equilibrium.

HC2H3O2 (aq)

acetic acid (vinegar)

nitrous acid

HNO2 (aq)

H2S (aq)

hydrosulfuric acid

HSO4–(aq)

hydrogen sulfate ion

Weak acids are always written in their molecular form.

See you text and home work for more examples.


Strong bases
Strong Bases

Bases:A base is a substance that produces OH– (aq)ions in water by dissociation in water:

Strong bases are almost completely ionized in aqueous solution. (Strong electrolytes)

Examples: Hydroxides of Group 1 (MOH(aq) where M = Li, Na, K ect…) and Ca, Sr, Ba.*

*Ca(OH)2, Sr(OH)2 & Ba(OH)2 are slightly soluble, but that which dissolves is present as ions only.


Bases
Bases

NaOH(aq)  Na+(aq) + OH-(aq)

Base: OH- in water

NaOH is a strong base


Weak bases

Weak Bases

Weak Bases:

NH3 acts as a base by reacting with water:

NH3(aq) + H2O(l)

NH4+(aq) + OH –(aq)

Ammonia can also accept H+ from an acid:

NH3(aq) + H+(aq)

NH4+(aq)



Reactions of Acids & Bases:

Acid-Base Neutralization

Salt + Water (usually)

Acid + Base 

HA (aq) + MOH(aq) 

MA(aq) + HOH(l)

Strong acid - Strong base neutralization: HBr(aq)/KOH(aq)

Molecular Equation:

HBr(aq) + KOH(aq) 

KBr (aq) + H2O(l)

Total Ionic Equation:

/

/

H+ (aq) + Br– (aq)

K+(aq) + Br– (aq)

+ K+(aq) + OH– (aq) 

+ H2O(l)

Br-

K+

Net Ionic equation:

H+ (aq) + OH– (aq) H2O (l)


Acid base reactions
Acid-Base Reactions

  • The “driving force” is the formation of water.

    NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(liq)

  • Net ionic equation

    OH-(aq) + H3O+(aq)  2 H2O(l)

  • This applies to ALL reactions of STRONG acids and bases.


Reactions of Acids & Bases:

Acid-Base Neutralization

Reactions of weak acids and strong bases:

Molecular Equation:

HC2H3O2(aq)

+ NaOH(aq) 

NaC2H3O2(aq)

+ H2O(l)

Total Ionic Equation:

/

/

HC2H3O2(aq)

+ Na+(aq) + OH–(aq) 

Na+(aq) + C2H3O2–(aq)

+ H2O(l)

Leave in molecular form

Net Ionic:

HC2H3O2(aq) + OH–(aq)  C2H3O2–(aq) + H2O(l)


Non metal acids
Non-Metal Acids

Nonmetal oxides can form acids in aqueous solutions:

Examples:

CO2(aq) + H2O(s)  H2CO3(aq)

SO3(aq) + H2O(s)  H2SO4(aq)

Both gases come from the burning of fossil fuels.


Bases1
Bases

Metal oxides form bases in aqueous solution

CaO(s) + H2O(l)  Ca(OH)2(aq)

CaO in water. Indicator shows solution is basic.



Gas forming reactions1
Gas-Forming Reactions

Metal carbonate salts react with acids to the corresponding metal salt, water and carbon dioxide gas.

2HCl(aq) + CaCO3(s)

 CaCl2(aq) + H2CO3(aq)

decomposes

H2O(l) + CO2(g)

Similarly:

HCl(aq) + NaHCO3(s) 

NaCl(aq) + H2O(l) + CO2(g)

acid

base

salt

water

Neutralization!!!


Gas forming reactions2
Gas-Forming Reactions

Group I metals: Na, K, Cs etc.. react vigorously with water

2K(s)

2KOH(aq)

+ H2(g)

+ 2H2O(l) 

Metals & acid:

Some metals react vigorously with acid solutions:

Zn(s)

+ 2H+(aq) 

Zn2+(aq)

+ H2(g)


Gas forming reactions3
Gas-Forming Reactions

CaCO3(s) + H2SO4(aq)  2 CaSO4(s) + H2CO3(aq)

Carbonic acid is unstable and forms CO2 & H2O

H2CO3(aq) CO2 + water

(The antacid tablet contains citric acid + NaHCO3)


Oxidation reduction reactions
Oxidation-Reduction Reactions

Thermite reaction:

Fe2O3(s) + 2Al(s) 

2Fe(s) + Al2O3(s)


Oxidation reduction reactions1
Oxidation-Reduction Reactions

REDOX = reduction & oxidation

O2(g) + 2 H2(g) 2 H2O(l)


Oxidation reduction reactions2
Oxidation-Reduction Reactions

  • Oxidationinvolves a reactant atom or compound losing electrons.

  • Reduction involves a reactant atom or substance gaining electrons.

  • Neither process can occur alone… that is, there must be an exchange of electrons in the process.

  • The substance that is oxidized is the reducing agent

  • The substance that is reduced is the oxidizing agent

oxidized

reduced

Mg(s)

+ 2H+(aq)

+ H2(g)

Mg2+(aq)

reducing

agent

oxidizing agent


Oxidation numbers
Oxidation Numbers

  • Chemists use oxidation numbersto account for the transfer of electrons in a RedOx reaction.

  • Oxidation numbersare the actual or apparent charge on atom when alone or combined in a compound.

    1. The atoms of pure elements always have an oxidation number of zero.

Examples:

Mg(s)

Hg(l)

I2(s)

O2(g)

All have an oxidation number of zero (0)


Oxidation numbers1
Oxidation Numbers

2. If an atom is charged, then the charge is the oxidation numbers.

Examples: Ion Oxidation Number

+2

1

+4

Mg2+(aq)

Cl(aq)

Sn4+(s)

+2/2 = +1 for each Hg atom


Oxidation numbers2
Oxidation Numbers

3. In a compound, fluorine always has an oxidationnumbers of 1.

4. Oxygen most often has an oxidation numberof 2.

  • *When combined with fluorine, oxygen has a positive O.N.

  • *In peroxide, the O.N. is 1.

    5. In compounds, Cl, Br & I are 1 (Except with F and O present)

    6. In compounds, H is +1, except as a hydride

    (H: 1)


Oxidation numbers3
Oxidation Numbers

Examples:

compound Oxidation Numbers

HF(g) H = +1 F = 1

H2O(l) H = +1 O = 2

OF2(g) O = +2 F = 1

Na2O2(s) Na = +1 O = 1

HCl(g) H = +1 Cl = 1

NaH(l) Na = +1 H = 1


Oxidation numbers4
Oxidation Numbers

Most common oxidation numbers:


Oxidation Numbers

7. For neutral compounds, the sum of the oxidation numbers equals zero.

For a poly atomic ion, the sum equals the charge.

Examples:

+ 2 × (−1) =

0

+2

MgCl2

+1

3

+ 4 × (+1) =


Oxidation Numbers

Determine the oxidation numberof iron in the following compound:

? +

3(1)

=

0

Fe(OH)3

Iron must have an oxidation number of +3!


Recognizing a redox reaction
Recognizing a Redox Reaction

In a RedOx reaction, the species oxidized and the species reduced are identified by the changes in oxidation numbers:

Oxidation numbers:

+1

0

0

+2

Oxidation numbers:

Since silver goes from +1 to zero, it is reduced.

Since copper goes from zero to +2, it is oxidized.

The reaction is balanced for both mass and charge.


Carbon has many oxidation states.

You have to check the electronegativity values to determine to which atoms it will give electrons and from which it will take electrons.

Unlike metals, which are almost always in a positive oxidation state, the oxidation state of carbon can vary widely, from -4 (in CH4) to +4 (such as in CO2).

Rules for carbon’s many oxidation states:

1.For every bond to something more electronegative like N,O, F Cl etc, carbon gives 1 electron and its ox # goes up +1.

2. For every bond with something less electronegative (like H, Li, Na)carbon takes an electron, and its ox # goes

down -1.


Practice finding the oxidation state of carbon

-4

CH4

CH3OH

CH2F2

  • CHF3

    CF4

-3+1 = -2

-2+2 = 0

-1+3 = +2

+4

+4

CO2

Double bond: O=C=O

Note: Carbon compounds use the rules of ORGANIC CHEMISTRY, where the carbon is always written first, whether it is acting as a metal or a non-metal.


Identify the species that is oxidized and the species that is reduced by assigning oxidation numbers in the following reaction.

Answer:

  • The carbon in methane (CH4) is oxidized (4 to 2)

see previous slide #90 for explanation of carbon oxidation

  • Chromium in dichromate is reduced (+6 to +3)


Redox reactions you do not have to memorize these
Redox Reactions you do not have to memorize these


Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)

Write the ionic equation

  • Cu(s) + 2 Ag++ NO3-Cu2+ + NO3- + 2 Ag(s)

Write the NET ionic equation, by eliminating spectator ions


Examples of redox reactions
Examples of Redox Reactions

Metal + halogen

2 Al + 3 Br2 Al2Br6


e

e

Electron Transfer in a Redox Reaction

2Ag+(aq)

+ 2Ag(s)

+ Cu(s)

Cu2+(aq)

  • Two electrons leave copper.

  • The silver ions accept them.

  • The copper metal is oxidized to copper (II) ion.

  • The silver ion is reduced to solid silver metal.


Redox reactions in our world
Redox Reactions in Our World

Batteries

Corrosion

Fuels

Manufacturing metals


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