nature of chemical reactions types of chemical...

Nature of Chemical Reactions

Types of Chemical Reactions

Synthesis Reactions

Decomposition Reactions

Single-Replacement Reactions

Double-Replacement Reactions

Combustion Reactions

Identifying Oxidation-Reduction Reactions

Nature of Chemical Reactions Objectives

1. Define chemical Reactions.

2. Describe four indications that a chemical reaction

has occurred.

3. State the law of conservation of mass and describe

how it relates to a chemical reaction.

4. Identify and use the common symbols used in

writing chemical equations.

5. Translate chemical equations into word equations.

6. Translate word equations into chemical equations.

7. Balance chemical equations.

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Characteristics of Chemical Reactions

A chemical reaction is a process

in which the chemical and

physical properties of the

original substance changes as

new substances with different

physical and chemical

properties are formed.

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Characteristics of Chemical Reactions

In any chemical reaction, there are

always two kinds of substances; the

substances that are present before the

change, called the reactants, and the

substances that are formed by the

change, called the products.

Reactants → Products

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Identifying Reactants and Products

Identify the reactants and products

in the following equation:

2Al(s) + 3Cl2(g) → 2AlCl3(s)

reactant(s)

products(s)

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Al, Cl2

AlCl3

Indications of a Chemical Reaction

What are the four observations that generally

indicate that a chemical reaction has occurred?

1. Production of a gas.

2. Formation of a precipitate. (an insoluble solid

that forms in an aqueous reaction)

3. Change in energy.

Endothermic reaction – energy is absorbed

Exothermic Reaction – energy is released

4. Change in color or odor.

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Law of Conservation of Mass

Chemical equations are balanced to satisfy

the Law of Conservation of Mass

The Law of Conservation of Mass states that

mass cannot be created or destroyed by

ordinary physical or chemical means.

This means that the total mass of the

reactions must equal the total mass of the

products.

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Law of Conservation of Mass

Remember that atoms don’t

change in a chemical reaction; they

just rearrange. For a chemical

equation to accurately represent a

reaction, the same number of each

kind of atom must be on the left

side of the arrow as are on the right

side.

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Symbols Used in Chemical Equations

+ Used to separate two reactants or two

products

→ “yields”, separates reactants from products

⇄ Used in place of a → for reversible reactions

(s) Designates a reactant or product in the solid

state

(l) Designates a reactant or product in the liquid

state

(aq) Designates an aqueous solution, the

substance is dissolved in water.

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Symbols Used in Chemical Equations

(g) Designates a reactant or product in the

gaseous state

Indicates that heat is supplied to the

reaction.

A formula written above or below the yield

sign indicates its use as a catalyst (in this

example, platinum).

A catalyst is a substance that speeds up a reaction

without being used up itself or permanently

changed.

In biology, a catalyst is commonly called an enzyme.

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Writing Sentences to Describe Balanced Equations

Sentences can be written to describe

(interpret) what is happening in a balanced

equation.

Example 1.

2NaOH(aq) + CO2(g) → Na2CO3(s) + H2O(l)

Aqueous sodium hydroxide reacts with

carbon dioxide gas to produce solid sodium

carbonate and water.

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Writing Sentences to Describe Balanced Equations

Example 2.

2Al(s) + 3Cl2(g) → 2AlCl3(s)

Aluminum metal reacts with chlorine gas to

produce solid aluminum chloride.

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You Try It

Write sentences to describe each of the

following equations. (Do not worry about the

coefficients at this time.)

i. 2H2(g) + O2(g) → 2H2O(l)

Hydrogen gas reacts with oxygen gas to

produce water.

ii. Zn(s) + CuCl2(aq) → ZnCl2(aq) + Cu(s)

Zinc metal reacts with aqueous copper(II)

chloride to produce aqueous zinc

chloride and copper metal.

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Counting Atoms in Chemical Compounds

Before we can balance an equation, we must first make

sure that everyone can count the atoms present in a

compound. Here are some examples.

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KClO3 ___K, ___Cl, ___O

Mg(OH)2 ___Mg, ___O, ___H

Al2(SO4)3 ___Al, ___S, ___O

2Al(NO3)3 ___Al, ___N, ___O

1 1 3

1 2 2

2 3 12

2 6 18

You Try It

Determine the number of each atom present

in each of the following.

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AlPO3 ___Al, ___P, ___O

Ba(NO3)2 ___Ba, ___N, ___O

HC2H3O2 ___H, ___C, ___O

3Fe3(PO4)2 ___Fe, ___P, ___O

1 1 3

1 2 6

4 2 2

9 6 24

Identifying Balanced Chemical Equations

An equation in which the number of atoms of

each element is the same on both sides of

the equation is called a balanced chemical

equation.

Pb(NO3)2 + 2KI → PbI2 + 2KNO3

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Reactant Side Element Product Side

Pb

N

O

K

I

1 1 2 2 6 6

2 2 2 2

You Try It

Classify each of the following equations as

balanced or unbalanced.

a. 2H2O2(aq) → 2H2O(l) + O2(g)

b. Mg(s) + HCl(aq) → MgCl2(aq) + H2(g)

c. Al4C3(s) + H2O(l) → CH4(g) + Al(OH)3(s)

d. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

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balanced

unbalanced

unbalanced

balanced

Balancing Chemical Equations

The easiest equations to balance are the

ones in which the skeleton equation is

already written.

Example 1. Cl2 + Na → NaCl

Add coefficients to balance the equation.

NEVER change subscripts.

What coefficient must be added in front of

NaCl in order to balance the chlorine?


Na in order to balance the sodium?

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2 2


Example 2.

Mg + HCl → MgCl2 + H2


HCl in order to balance the chlorine and the

hydrogen?

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2


Example 3.

Fe + O2 → Fe2O3

Although your first instinct might be to balance the

Fe first, you should actually balance the O first.

We have 2 O’s on the reactant side and 3 O’s on the

product side. What is the least common multiple of 2

and 3?

What number can we put in front of O2 to obtain 6?

What number can we put in front of Fe2O3 to obtain

6?

What coefficient must be added in front of Fe to

balance the Fe? Back to main menu

2 3 4

6

You Try It

Balance the following chemical equations.

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2 i. Cl2 + KI KCl + I2

ii. Li + O2 Li2O

iii. HC2H3O2 + CaCO3 → Ca(C2H3O2)2 + CO2 + H2O

iv. C2H6O(l) + O2(g) → CO2(g) + H2O(g)

v. Al(s) + HCl(aq) → AlCl3(aq) + H2(g)

vi. Pb(NO3)2 + K2CrO4 PbCrO4 + KNO3

vii C3H6 + O2 CO2 + H2O

viii. Al(OH)3 Al2O3 + H2O

ix. HCl + Fe2O3 FeCl3 + H2O

2

4 2

2

2 3 3

2 6 2 3

2

2 6 6 9

2 3

2 6 3


If the skeleton equation is not written for you,

you must write your own.

Example 1. When an electric current is

passed through water, the water molecules

break down to produce hydrogen and oxygen.

Bubbles of each gas are evidence of the

reaction.

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→ H2(g) H2O(aq) + O2(g) 2 2


Example 2. Solutions of aluminum sulfate

and barium chloride react to produce solid

barium sulfate and aqueous aluminum

chloride.

aluminum sulfate = Al3+, SO42- = Al2(SO4)3

barium chloride = Ba2+, Cl- = BaCl2

barium sulfate = Ba2+, SO42- = BaSO4

aluminum chloride = Al3+, Cl- = AlCl3

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main menu

→ BaSO4(s) Al2(SO4)3(aq) + BaCl2(aq)

2

3

+ AlCl3(aq)

3

You Try It

Write a balanced equation for each of the

following reactions.

a. When magnesium and oxygen react, the

product is solid magnesium oxide.

b. When nitrogen and hydrogen react, the

product is ammonia.

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N2(g) + H2(g) → NH3(g)

2 Mg(s) + O2(g) → MgO(s)

3 2

2

You Try It

Write a balanced equation for each of the

following reactions.

c. Aluminum reacts with oxygen to produce

aluminum oxide (rust).

d. Solutions of calcium chloride and sodium

sulfate react to produce aqueous sodium

chloride and solid calcium sulfate.

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CaCl2(aq) + Na2SO4(aq)→ NaCl(aq) + CaSO4(s)

2 Al(s) + O2(g) → Al2O3(s)

2

4 3

Synthesis Reactions

In a synthesis reaction, two

substances – either elements or

compounds – combine to form a

single compound.

The general equation for a

synthesis reaction is

A + B → C

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Synthesis Reactions – General Rules

a. Metal + Nonmetal → *Binary Ionic Compound

*You must remember to balance the charges

when writing the formula for the binary ionic

compound.

Na + O2 →

What product is going to be formed?

Sodium oxide

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Na2O 2 4

Na+, O2- Na2O


b. Nonmetal + Nonmetal → Binary Covalent

Compound

S + O2 → SO2

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c. Metal oxide + Water → *Base (metal hydroxide)

*You must remember to balance the charges

when writing the formula for the base formed.

CaO + H2O →

What product is going to be formed?

Calcium hydroxide

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Ca(OH)2

Ca2+, OH- Ca(OH)2


d. Nonmetal oxide + Water → Acid

CO2 + H2O →

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H2CO3


e. Metallic oxide + Nonmetallic oxide → Ternary Ionic

Compound (salt)

Na2O + CO2 →

Na2CO3


f. Here are two synthesis reactions that must be

memorized.

N2(g) + 3H2(g) 2NH3(g)

NH3(g) + H2O(l) NH4OH(aq)

You Try It Balance the following equations. You may need to

complete the equation before balancing it.

a. Solid lithium metal reacts with oxygen gas.

b. Magnesium metal burns in oxygen gas.

Li + O2 Li2O 2 4

Mg + O2 MgO 2 2

You Try It c. Magnesium metal reacts with chlorine gas

d. Solid calcium oxide is heated in the presence

of sulfur trioxide.

e. Sulfur dioxide gas is bubbled into distilled water.

Mg + Cl2 MgCl2

CaO + SO3 CaSO4

SO2 + H2O H2SO3

You Try It f. Solid sodium oxide is added to water.

g. Calcium metal is heated strongly in nitrogen

gas.

h. Solid tetraphosphorus decoxide reacts with

water to produce phosphoric acid.

Na2O + H2O NaOH

Ca + N2 Ca3N2

P4O10 + H2O H3PO4

2

3

4 6

You Try It i. Iodine crystals react with chlorine gas to form

solid iodine trichloride.

j. Solid strontium oxide reacts with water.

I2 + Cl2 ICl3

SrO + H2O H2SrO2

2 3

Decomposition Reactions In a decomposition reaction, a compound breaks down

into two or more simpler substances.

The compound may break down into individual

elements, a compound and an element, or into simpler

compounds.

Many decomposition reaction take place only when

energy is added in the form of heat or light.

Electrolysis is the decomposition of a substance by an

electric current.

The general equation for a decomposition reaction is

C → A + B

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Decomposition Reactions – General Rules

a. Binary compounds break into their *elements.

*You must remember to watch out for the diatomic

elements (H2, Br2, O2, N2, Cl2, I2, F2).

K2O →

What products are going to be formed?

Potassium

Oxygen

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K 2 4 + O2


b. Metal carbonates generally break down to

produce a metal oxide and carbon dioxide.

K2CO3 →


Potassium oxide

Carbon dioxide

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K2O + CO2

K+, O2- K2O


c. Metal hydroxides generally break down to

produce a metal oxide and water.

Sr(OH)2 →


Strontium oxide

Water

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SrO + H2O

Sr2+, O2- SrO


d. Metal chlorates generally break down to produce

a metal chloride and oxygen.

KClO3 →


Potassium chloride

Oxygen

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KCl + O2

K+, Cl- KCl

2 3 2


e. Many acids decompose to produce water and a

nonmetal oxide.

H2CO3 →


Water

Nonmetal oxide

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+ CO2 H2O


f. Ternary Ionic compounds (salts) decompose to

produce a metal oxide and a nonmetal oxide.

CaSO3 →


Metal oxide

Nonmetal oxide

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+ SO2 CaO


g. Here are four decomposition reactions that must

be memorized.

2H2O2 2H2O + O2

NH4OH NH3 + H2O

(NH4)2CO3 2NH3 + H2O + CO2

NH4NO3 N2O + 2H2O

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Balance the following decomposition reactions.

Use the examples above to help you predict the

products when necessary. You do not have to

indicate the states of the reactants and products.

a. Solid barium carbonate is heated.

BaCO3 →

b. Solid silver oxide is heated.

Ag2O

You Try It!

BaO + CO2

Ag + O2 2 4

c. Solid sodium chlorate is heated.

NaClO3 →

d. An aqueous solution of sulfuric acid is

heated.

H2SO4 →

e. Aluminum can be obtained from

aluminum oxide with the addition of a

large amount of electrical energy.

Al2O3 →

You Try It!

NaCl + O2

H2O(l) + SO3(g)

2

4 Al + O2

2 3

2 3

f. Heating tin(IV) hydroxide gives tin(IV)

oxide and steam.

Sn(OH)4 →

g. Solid sodium carbonate is heated.

Na2CO3 →

h. Solid magnesium sulfite is heated.

MgSO3 →

You Try It!

SnO2 + 2H2O

Na2O + CO2

MgO + SO2

i. Powdered magnesium carbonate is

heated strongly.

MgCO3 →

j. Solid calcium hydroxide decomposes.

Ca(OH)2(s) →

You Try It!

MgO(s) + CO2(g)

CaO(s) + H2O(g)


In a single-replacement reaction, an uncombined

element replaces an element in a compound.

There are three general equations for a single

replacement reaction.

A + BC → AC + B (If A is a metal)

A + BC → BA + C (If A is a halogen)

A + BC → Base + H2

(If A is an active metal and BC is H2O)

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Examples

Zn + 2HCl → ZnCl2 + H2

Cl2 + 2LiBr → 2LiCl + Br2

2K + 2H2O → 2KOH + H2

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Whether one metal will

replace another from a

compound can be

determined by using the

activity series of metals.

The activity series of

metals lists metals in order

of chemical reactivity.

A reactive metal will

replace any metal found

below it in the activity

series.


The halogens can also take place

in single-replacement reactions.

The order of reactivity for the

halogens from increasing to

decreasing reactivity is fluorine,

chlorine, bromine, iodine.

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Examples

a. Li + KCl →

b. Mg + K2SO4 →

c. Cl2 + KBr →

d. Li + H2O →

Back to main menu

+ K LiCl

no reaction

KCl + Br2

LiOH + H2

2 2

2 2 2

You Try It

Complete and balance the following

equations.

a. Zn + Pb(NO3)2 →

b. K + Ba(C2H3O2)2 →

c. Al + NiSO4 →

d. Na + H2O →

e. Zn + BaCl2 →

f. F2 + AlCl3 →

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+ Pb Zn(NO3)2

KC2H3O2 + Ba

Al2(SO4)3

no reaction

+ Ni

2 2

2 3

2

3

NaOH + H2 2 2

AlF3 + Cl2 3 2 2 3


In a double-replacement reaction, the

negative ions of two compounds

exchange places in an aqueous

solution to form two new compounds.

The general equation for a double

replacement reaction is

AB + CD → AD + CB

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Double-Replacement Reactions For a double-replacement reaction to occur, one of

the following statements is usually true concerning

at least one of the products of the reaction.

a. It is a gas that bubbles out of the mixture.

FeS(s) + HCl(aq) →

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2 FeCl2(aq) + H2S(g)

Other gases commonly formed include:

H2S

CO2 (formed from the decomposition of H2CO3)

SO2 (formed from the decomposition of H2SO3

NH3 (formed from the decomposition of NH4OH)


b. It is a molecular compound such as water.

This is common in acid-base neutralization

reactions.

HCl(aq) + NaOH(aq) →

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H2O(l) + NaCl(aq)

Double-Replacement Reactions c. It is only slightly soluble and precipitates from

solution.

KI(aq) + Pb(NO3)2(aq) →

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2 KNO3(aq) + PbI2(s) 2

Identifying Precipitates A solubility chart can be used to identify insoluble

solids (precipitates).

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Identifying Precipitates

Example. Label the precipitate formed in the

following reaction.

Na2CO3(aq) + Ca(NO3)2(aq) → 2NaNO3( ) + CaCO3( )

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aq s

You Try It

Label the precipitate formed in each of the following

reactions.

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aq s i. NH4Cl(aq) + AgNO3(aq) → NH4NO3( ) + AgCl( )

ii. BaCl2(aq) + Na2SO4(aq) → BaSO4( ) + 2NaCl( ) aq s

a. AgNO3 + NaCl →

b. Mg(NO3)2 + KOH →

c. LiOH + Fe(NO3)3 →

d. Pb(NO3)2 + NaOH →

e. NH4Cl(aq) + NaOH(aq) →

f. K2SO3(aq) + HBr(aq) →

You Try It

Complete and balance the following

equations.

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+ NaNO3 AgCl

Mg(OH)2 + KNO3

LiNO3 + Fe(OH)3

2 2

3 3

2 2 Pb(OH)2 + NaNO3

NH3 + H2O + NaCl KBr + SO2 + H2O 2 2

NH4OH

+ H2SO3


In a combustion reaction, a substance combines

with oxygen, releasing a large amount of energy in

the form of heat and light.

For our purpose, we will talk about hydrocarbons

burning in the presence of oxygen.

The general equation for a complete combustion

reaction is

CxHy + O2 → CO2 + H2O

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Types of Combustion Reactions

Complete Combustion – the products

are CO2 and H2O.

CH4 + 2O2 → CO2 + 2H2O

Incomplete Combustion – the products

are C and/or CO and H2O.

2C3H8 + 7O2 → 6CO + 8H2O

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Example Problem. Write a

balanced equation for the complete

combustion of ethane, C2H6.

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C2H6 + O2 → CO2 + H2O 2 4 6 7

You Try It

Write balanced equations for the

complete combustion of the

following substances.

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+ O2 + H2O 3 4 5

2

a. C3H8

b. C5H10

c. C6H14

+ O2

+ O2

→ CO2

→ CO2

→ CO2

+ H2O

+ H2O

10 10 15

2 12 14 19

Oxidation-Reduction Reactions (Redox)

Oxidation-Reduction reactions are

the chemical changes that occur

when electrons are transferred

between reactions.

Examples include the burning of

gasoline and the rusting of a nail.

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Oxidation Oxidation originally meant the

combination of an element with

oxygen to give oxides.

However, today it is also defined as

the loss of electrons. (Oxygen

does not have to be present for

oxidation to occur.)

Example: 4Fe + 3O2 → 2Fe2O3

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Reduction Reduction originally meant the loss

of oxygen from a compound.

Today it is also defined as the gain

of electrons.

Example: 2Fe2O3 + 3O2 → 4Fe + 3CO2

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Oxidizing Agent The oxidizing agent in a redox

reaction gains electrons.

It is the substance being reduced.

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Reducing Agent The oxidizing agent loses

electrons.

It is the substance being oxidized.

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Helpful Mnemonics Leo the Lion says Ger

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L – Loss

e – of electrons

o – is oxidation

G – Gain

e – of electrons

r – is reduction

Oil Rig

O – Oxidation

i – is

l – loss of electrons

R – Reduction

i – is

g – gain of electrons

Rules for Assigning Oxidation Numbers

1. The oxidation number of a

monatomic ion is equal to its

charge.

Examples:

Br- equals

Fe3+ equals

Back to main menu

-1

+3


2. For a polyatomic ion, the sum

of the oxidation numbers must

equal the ionic charge of the

ion.

Examples:

SO42- equals

NO3- equals

Back to main menu

-2

-1


3. The oxidation number of a

metal cation is the same as its

ionic charge.

Examples:

sodium is

calcium is

Back to main menu

+1

+2


4. The oxidation number of

hydrogen in a compound is +1

except in metal hydrides, for

example, NaH, where it is -1.

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5. The oxidation number of

oxygen in a compound is -2

except in peroxides, for

example, H2O2, where it is -1.

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6. The oxidation number of an

uncombined element is 0.

For example, the oxidation

number of the potassium atoms

in potassium metal, K, and of

the nitrogen atoms in nitrogen

gas, N2, is zero.

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7. For any neutral compound, the

sum of the oxidation numbers

of the atoms in the compound

must equal 0.

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8. Assign the oxidation numbers in

the following order.

a. Metals (determined by looking at their group number)

b. Hydrogen (usually +1)

c. Oxygen (usually -2)

d. Transition Metals and Everything Else

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What is the oxidation number of each element in the following?

SO2

Back to main menu

+4 -2

-4 +4 Start with oxygen. According to rule 5, the oxidation

number of oxygen in a compound is usually -2.

What number does sulfur need to be in order for the

overall compound to have a sum of zero?

=0

S = +4

O= -2


KClO3

Back to main menu

+1 -2

-6 +1 Start with potassium. According to rule 3, the oxidation

number of potassium is the same as its ionic charge,

which is +1.

Next do oxygen. According to rule 5, the oxidation


What number does chlorine need to be in order for the


=0

K = +1

O= -2

Cl = +5 +5

+5


KClO2

Back to main menu

+1 -2

-4 +1 Start with potassium. According to rule 3, the oxidation

number of potassium is the same as its ionic charge,

which is +1.

Next do oxygen. According to rule 5, the oxidation


What number does chlorine need to be in order for the


=0

K = +1

O= -2

Cl = +3 +3

+3


CO32-

Back to main menu

+4 -2

-6 +4 Start with oxygen. According to rule 5, the oxidation


What number does carbon need to be in order for the

polyatomic ion to have an overall charge of -2?

= -2

C = +4

O= -2

You Try It. Determine the oxidation number of each element in each of the following.

1. Na2Cr2O7 Na = O = Cr =

2. BaH2 Ba = H =

(Hint: BaH2 is a metal hydride.)

3. Li2O2 Li = O =

(Hint: Li2O2 is a peroxide.)

4. ClO3- Cl = O =

Back to main menu

+2

-2 +6 +1

-1

+1 -1

+5 -2

Oxidation Number Changes

The changes in oxidation number can be

used to determine which elements are

oxidized and which elements are reduced.

Remember – an increase in the oxidation

number of an atom signifies oxidation and a

decrease in the oxidation number signifies

reduction.

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Use the change in oxidation number to

identify which elements are oxidized and

reduced in each of these reactions. Also

identify the oxidizing (OA) and reducing (RA)

agents.

a. F2(g) + 2HBr(aq) → 2HF(aq) + Br2(l)

Back to main menu

0 +1

-1

+1

-1

0

Fluorine: 0 to -1; reduced, OA

Bromine: -1 to 0; oxidized, RA

+1 =0

-1

+1 =0

-1






agents.

b. 2KClO3(s) → 2KCl(s) + 3O2(g)

Back to main menu

+1 -2 +5 +1 -1 0

Chlorine: +5 to -1; reduced, OA

Oxygen: -2 to 0; oxidized, RA

-6 +1 +5 =0 +1 =0 -1






agents.

c. SO3(g) + H2O(l) → H2SO4(aq)

Back to main menu

+6 -2 +1 -2 +1

Since there is no change in oxidation

number, this is not a redox reaction.

-6 +6 +2

-2

-8 +6

+6

+2 -2 =0 =0 =0

You Try It





agents.

a. 2H2(g) + O2(g) → 2H2O(g)

Back to main menu

0 0 +1 -2

+2

hydrogen: 0 to +1; oxidized, RA

oxygen: 0 to -2; reduced, OA

-2 =0

You Try It





agents.

b. 2Al(s) + 6HCl(aq) → 2AlCl3(s) + 3H2(g)

Back to main menu

0 0 +1 +3

+3

aluminum: 0 to +3; oxidized, RA

hydrogen: +1 to 0; reduced, OA

-3

-1 -1

=0 =0 +1 -1

You Try It





agents.

c. 2NaCl(aq) + Li2S(aq) → Na2S(s) + LiCl(aq)

Back to main menu

+1 +1 +1 -2

+2

Since there is no change in oxidation

number, this is not a redox reaction.

-2

-1

+2

-2

-2

+1 -1

=0 =0 =0 =0 +1 -1 +1 -1

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