chapter 7 chemical reactions

Chapter 7 Chemical Reactions

2006, Prentice Hall

H2

O2

H2O

react to form?

+ heat

2

CHAPTER OUTLINE

Chemical Reactions Chemical Equation Balancing Equations Types of Chemical Reactions Activity Series of Metals Aqueous Reactions Precipitation Reactions Neutralization and Other Reactions Heat in Chemical Reactions

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CHEMICALREACTIONS

A chemical reaction is a rearrangement of atoms in which some of the original bonds are broken and new bonds are formed to give different chemical structures.

In a chemical reaction, atoms are neither created, nor destroyed.

A chemical reaction, as described above, is supported by Dalton’s postulates.

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CHEMICALREACTIONS

6 oxygen atoms 6 oxygen atoms=In a chemical reaction, atoms are neither created, nor destroyed

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CHEMICALREACTIONS

A chemical reaction can be detected by one of the following evidences:

1. Change of color

2. Formation of a solid

3. Formation of a gas

4. Exchange of heat with surroundings

Evidence of Chemical Change

Color ChangeFormation of Solid PrecipitateFormation of a Gas

Emission of LightRelease or Absorption of Heat

Why use Chemical Equations?

1. Shorthand way of describing a reaction

2. Provides information about the reaction– Formulas of reactants and products– States of reactants and products– Relative numbers of reactant and product

molecules that are required– Can be used to determine amounts of the

reactants and products

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CHEMICALEQUATIONS

A chemical equation is a shorthand expression for a chemical reaction.

Word equation:

Aluminum combines with ferric oxide to form iron and aluminum oxide.

Chemical equation:

Al + Fe2O3 Fe + Al2O3

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CHEMICALEQUATIONS

Reactants are separated from products by an arrow.

Al + Fe2O3 Fe + Al2O3

Coefficients are placed in front of substances to balance the equation.

2 Al + Fe2O3 2 Fe + Al2O3

Subscripts

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CHEMICALEQUATIONS

Reaction conditions are placed over the arrow.

Al + Fe2O3 Fe + Al2O3

The physical state of the substances are indicated by the symbols (s), (l), (g), (aq).

2 Al (s) + Fe2O3 (s) 2 Fe (l) + Al2O3 (s)

heat

solid liquid

Symbols Used in Equations

1. energy symbols used above the arrow for decomposition reactions

– = heat– h = light– shock = mechanical– elec = electrical

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BALANCINGEQUATIONS

A balanced equation contains the same number of atoms on each side of the equation, and therefore obeys the law of conservation of mass.

Many equations are balanced by trial and error; but it must be remembered that coefficients can be changed in order to balance an equation, but not subscripts of a correct formula.

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BALANCINGEQUATIONS

The general procedure for balancing equations is:

Write the unbalanced equation:

CH4 + O2 CO2 + H2O

Make sure the formula for each

substance is correct

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BALANCINGEQUATIONS

The general procedure for balancing equations is:

Balance by inspection:

CH4 + O2 CO2 + H2O

Count and compare each

element on both sides of the

equation

1 C = 1 C

4 H 2 H

2 O 3 O

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BALANCINGEQUATIONS

Balance elements that appear only in one substance first.

Balance H

CH4 + O2 CO2 + H2O

1 CH4 + O2 CO2 + 2 H2O

4 H present on each side

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BALANCINGEQUATIONS

Balance O

1 CH4 + O2 CO2 + 2 H2O

1 CH4 + 2 O2 CO2 + 2 H2O

4 O present on each side

When finally done, check for the smallest coefficients possible

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Examples:

AgNO3 + H2S Ag2S + HNO3

Al(OH)3 + H2SO4 Al2(SO4)3 + H2O

Fe3O4 + H2 Fe + H2O

C4H10 + O2 CO2 + H2O

2 AgNO3 + H2S Ag2S + 2 HNO3

2 Al(OH)3 + 3 H2SO4 Al2(SO4)3 + 6 H2O

Fe3O4 + 4 H2 3 Fe + 4 H2O

2 C4H10 + 13 O2 8 CO2 + 10 H2O

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TYPES OFCHEMICAL REACTIONS

Chemical reactions are can be classified into five types: Based on what the atoms do

1. Synthesis or combination

2. Decomposition

3. Single replacement

4. Double replacement

5. Combustion

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SYNTHESIS orCOMBINATION

In these reactions, 2 elements or compounds combine to form another compound.

A + B AB

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DECOMPOSITION

In these reactions, a compound breaks up to form 2 elements or simpler compound.

AB A +B

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SINGLEREPLACEMENT

In these reactions, a more reactive element replaces a less reactive element in a compound.

A + BC B + AC

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DOUBLEREPLACEMENT

In these reactions, two compounds combine to form two new compounds.

AB + CD AD + CB

The cation from one compound replaces the cation in another compound.

+ +

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COMBUSTION

A reaction that involves oxygen as a reactant and produces large amounts of heat is classified as a combustion reaction.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

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Examples:

Classify each of the reactions below:

1. Mg + CuCl2 MgCl2 + Cu

2. CaCO3 CaO + CO2

3. 2 HCl + Ca(OH)2 CaCl2 + 2 H2O

4. 4 Fe + 3 O2 2 Fe2O3

Single replacement

Mg is more reactive than Cu

Decomposition

Double replacement

Synthesis

Single Displacement

22 ZnCl Cu(s) (aq)CuCl Zn(s)

The Zinc Replaces the Copper

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ACTIVITY SERIESOF METALS

Activity series is a listing of metallic elements in descending order of reactivity.

Hydrogen is also included in the series since it behaves similar to metals.

Activity series tables are available in textbooks and other sources.

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ACTIVITY SERIESOF METALS

Elements listed higher will displace any elements listed below them.

For example Na will displace any elements listed below it from one of its compounds.

2 Na (s) + MgCl2 (aq) 2 NaCl (aq) + Mg (s)

2 Na (s) + AgCl (aq) NaCl (aq) + Ag (s)

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ACTIVITY SERIESOF METALS

Elements listed lower will not displace any elements listed above them.

For example Ag cannot displace any elements listed above it from one of its compounds.

Ag (s) + CuCl2 (aq) No Reaction

Ag (s) + HCl (aq) No Reaction

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Example 1:

Use activity series to complete each reaction below. If no reaction occurs, write “No Reaction”.

Pb (s) + HCl (aq) Metals FeNiSnPbH

CuAg

Pb is more reactive than H

PbCl2 (aq) + H2 (g)2

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Example 2:

Use activity series to complete each reaction below. If no reaction occurs, write “No Reaction”.

Ni (s) + CuCl2 (aq) Metals FeNiSnPbH

CuAg

Ni is more reactive than Cu

NiCl2 (aq) + Cu (s)

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AQUEOUSREACTIONS

Many ionic solids dissolve in water to form ions.

2H O(s) (aq

2-4 4 aq)

+(2 )K 2 CrO CrO K +

These substances are called electrolytes.

2H O(s) (aq) (

-3 2 3 q

2+a )Ba B(NO ) 2 N + Oa

NaCl(s) Na+(aq) + Cl(aq)

H2O

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AQUEOUSREACTIONS

When ionic substances dissolve in water they separate into ions.

K2CrO4 Ba(NO3)2

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electrolytes are substances whose water solution is a conductor of electricity

electrolytes are ions dissolved in water (Na+ + Cl-)

AQUEOUSREACTIONS

Types of Electrolytes

• salts = water soluble ionic compounds

• acids = form H+1 ions in water solution– react and dissolve many metals

– strong acid = strong electrolyte, weak acid = weak electrolyte

• bases = water soluble metal hydroxides– increases the OH-1 concentration

When will a Salt Dissolve?

• a compound is soluble in a liquid if it dissolves in that liquid– NaCl is soluble in water

• a compound is insoluble if a significant amount does not dissolve in that liquid– AgCl is insoluble in water

• though there is a very small amount dissolved, but not enough to be significant

AgCl remains solid = precipitate

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AQUEOUSREACTIONS

Aqueous reactions occur only when one of the following conditions is present:

1. Formation of a solid: Precipitation

2. Formation of water: Neutralization

3. Formation of a gas: Unstable product

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PRECIPITATIONREACTIONS

An aqueous chemical reaction that produces a solid as one of its products is called a precipitation reaction.

The insoluble solid formed in these reactions is called a precipitate.

2 4 (aq) 3 2 (aq) 4 (s) 3 (aq)K CrO + Ba(NO ) BaCrO + 2 KNO

Precipitate

Example of a Precipitation Reaction

Pb(NO3)2(aq) + 2 KI(aq) 2 KNO3(aq) + PbI2(s)

Let’s Look Closer at PbI2 Formation

Pb(NO3)2(aq) + 2 KI(aq) 2 KNO3(aq) + PbI2(s)

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SOLUBILITYRULES

S

O

L

U

B

L

E

NO3 No exceptions

Na+, K+

NH4+

No exceptions

Cl, Br, I Except those containing Ag+ , Pb2+

SO42

Except those containing Ba2+ , Pb2+, Ca2+

Chemists use a set of solubility rules to predict whether a product is soluble or insoluble.

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SOLUBILITYRULES

I

N

S

O

L

S2, CO32

PO43

Except those containing Na+ , K+, NH4

+

OHExcept those containing

Na+ , K+, Ca2+, NH4+

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Write balanced equations for each reactions shown below. Indicate if no reaction occurs.

Example 1:

NaCl (aq) + AgNO3 (aq) NaNO3 (?) + AgCl (?)

NaCl (aq) + AgNO3 (aq) NaNO3 (aq) + AgCl (s)

soluble

precipitate

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Write balanced equations for each reactions shown below. Indicate if no reaction occurs.

Example 2:

NH4Cl (aq) + KNO3 (aq) NH4NO3 (?) + KCl (?)

NH4Cl (aq) + KNO3 (aq) NH4NO3 (aq) + KCl (aq)

soluble

No Reaction

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Write balanced equations for each reactions shown below. Indicate if no reaction occurs.

Example 3:

PbCl2 (aq) + NaI (aq) PbI2 (?) + 2 NaCl (?)

precipitate

2

PbCl2 (aq) + 2 NaI (aq) PbI2 (s) + 2 NaCl (aq)

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Molecular, Complete Ionic, and Net Ionic Equations

A molecular equation is a chemical equation showing the complete, neutral formulas for every compound in a reaction.

A complete ionic equation is a chemical equation showing all of the species as they are actually present in solution.

A net ionic equation is an equation showing only the species that actually participate in the reaction.

© 2012 Pearson Education, Inc.

Writing Chemical Equations for Reactions in Solution: Molecular and Complete Ionic Equations

• A molecular equation is an equation showing the complete neutral formulas for every compound in the reaction.

• Complete ionic equations show aqueous ionic compounds that normally dissociate in solution as they are actually present in solution.

• When writing complete ionic equations, separate only aqueous ionic compounds into their constituent ions.

• Do NOT separate solid, liquid, or gaseous compounds.

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Writing Chemical Equations for Reactions in Solution: Net Ionic Equations

• In the complete ionic equation, some of the ions in solution appear unchanged on both sides of the equation.

• These ions are called spectator ions because they do not participate in the reaction.

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Writing Chemical Equations for Reactions in Solution: Proper Net Ionic Equations

• To simplify the equation, and to more clearly show what is happening, spectator ions can be omitted.  

• Equations such as this one, which show only the species that actually participate in the reaction, are called net ionic equations.

Ag+(aq) + Cl− (aq) AgCl(s)

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NEUTRALIZATIONREACTIONS

The most important reaction of acids and bases is called neutralization.

In these reactions an acid combines with a base to form a salt and water.

2HCl (aq) + NaOH (aq) NaCl (aq) + H O (l)¾¾®

Acid Base Salt

Salts are ionic substances with the cation donated from the base and the anion donated from the acid.

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Write balanced equations for each of the neutral-ization reactions shown below:

Examples:

HNO3 + Ba(OH)2 Ba(NO3)2 + H2O2 2

H2SO4 + NaOH Na2SO4 + H2O2 2

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GAS FORMINGREACTIONS

Some chemical reactions produce gas because one of the products formed in the reaction is unstable.

Three such products are:

H2CO3 (aq) CO2 (g) + H2O (l)Carbonic acid:

Sulfurous acid: H2SO3 (aq) SO2 (g) + H2O (l)

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

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GAS FORMINGREACTIONS

When either of these products appears in a chemical reaction, they should be replaced with their decomposition products.

2 HCl + Na2CO3 2 NaCl + H2CO3

2 HCl + Na2CO3 2 NaCl + CO2 (g) + H2O (l)

2 HNO3 + K2SO3 2 KNO3 + H2SO3

2 HNO3 + Na2SO3 2 KNO3 + SO2 (g)+ H2O (l)

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Enthalpy: A Measure of the Heat Evolved or Absorbed in a Reaction

• Chemical reactions can be exothermic (they emit thermal energy when they occur).

• Chemical reactions can be endothermic (they absorb thermal energy when they occur).

• The amount of thermal energy emitted or absorbed by a chemical reaction, under conditions of constant pressure (which are common for most everyday reactions), can be quantified with a function called enthalpy.

© 2012 Pearson Education, Inc.

Enthalpy: A Measure of the Heat Evolved or Absorbed in a Reaction

• We define the enthalpy of reaction, ΔHrxn, as the amount of thermal energy (or heat) that flows when a reaction occurs at constant pressure.

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Sign of ΔHrxn

• The sign of ΔHrxn (positive or negative) depends on the direction in which thermal energy flows when the reaction occurs.

• Energy flowing out of the chemical system is like a withdrawal and carries a negative sign.

• Energy flowing into the system is like a deposit and carries a positive sign.

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Exothermic and Endothermic reactions  

• (a) In an exothermic reaction, energy is released into the surroundings. (b) In an endothermic reaction, energy is absorbed from the surroundings.

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HEAT IN CHEMICALREACTIONS

Reactions that release heat are classified as exothermic.

Reactions that absorb heat are classified as endothermic.

H2 (g) + Cl2 (g) 2 HCl (g) + 185 kJ

In exothermic reaction, heat is produced and can be written as a product.

In endothermic reaction, heat is required and can be written as a reactant.

N2 (g) + O2 (g) + 181 kJ 2 NO (g)

Endothermic

Exothermic

© 2012 Pearson Education, Inc.

Sign of ΔHrxn

• When thermal energy flows out of the reaction and into the surroundings it is a ??? reaction and has a + or – enthalpy?

• The enthalpy of reaction for the combustion of CH4, the main component in natural gas:

• The magnitude of ΔHrxn tells us that 802.3 kJ of heat are emitted when 1 mol CH4 reacts with 2 mol O2.

© 2012 Pearson Education, Inc.

Stoichiometry of ΔHrxn

• The amount of heat emitted or absorbed when a chemical reaction occurs depends on the amounts of reactants that actually react.

• We usually specify ΔHrxn in combination with the balanced chemical equation for the reaction.

• The magnitude of ΔHrxn is for the stoichiometric amounts of reactants and products for the reaction as written.

© 2012 Pearson Education, Inc.

Stoichiometry of ΔHrxn

• The balanced equation and ΔHrxn for the combustion of propane is:

• When 1 mole of C3H8 reacts with 5 moles of O2 to form 3 moles of CO2 and 4 moles of H2O, 2044 kJ of heat are emitted.

© 2012 Pearson Education, Inc.

  Example 1:

• An gas tank in a home barbecue contains 11.8 x 103 g of propane (C3H8).

• Calculate the heat (in kJ) associated with the complete combustion of all of the propane in the tank.

© 2012 Pearson Education, Inc.

  Example 1:

Classifying Reactions• Also we can classify reactions by what happens:• Redox reactions are the exchange of e-

• Redox are all reactions except?

Note

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OXIDATION-REDUCTIONREACTIONS

Reactions known as oxidation and reduction (redox) have many important applications in our everyday lives.

Rusting of a nail or the reaction within your car batteries are two examples of redox reactions.

In an oxidation-reduction reaction, electrons are transferred from one substance to another.

If one substance loses electrons, another substance must gain electrons.

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OXIDATION-REDUCTIONREACTIONS

Oxidation is defined as loss of electrons, and reduction is defined as gain of electrons.

One way to remember these definitions is to use the following mnemonic:

Combination, decomposition, single replacement and combustion reactions are all examples of redox reactions.

Oxidation Is Loss of electrons

Reduction Is Gain of electrons

OIL

RIG

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OXIDATION-REDUCTIONREACTIONS

In general, atoms of metals lose electrons to form cations, and are therefore oxidized, while atoms of non-metals gain electrons to form anions, and are therefore reduced.

Ca + S CaS

Ca Ca2+ + 2 e-

S + 2 e- S2-

Oxidation

Reduction

For example, in the formation of calcium sulfide from calcium and sulfur

Therefore, the formation of calcium sulfide involves two half-reactions that occur simultaneously, one an oxidation and the other a reduction.

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COMBUSTION

A reaction that involves oxygen as a reactant and produces large amounts of heat is classified as a combustion reaction.

Combustion reactions are a subclass of Oxidation-Reduction reactions

occurs inthe cylindersof the engine

2 C8H18(g) + 25 O2(g) 16 CO2(g) + 18 H2O(g)

Combustion Products• predicting the products of a combustion

reaction; simply combine each element in the other reactant with oxygen

Reactant Combustion Product

contains C CO2(g)

contains H H2O(g)

contains S SO2(g)

contains N NO(g) or NO2(g)

contains metal M2On(s)

Combustion Reactions• combustion reactions are always exothermic

• in combustion reactions, O2 combines with the elements in another reactant to make the products

4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + energy

CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) + energy The flame on a gas stove results from the oxidation of carbon in natural gas.

Reverse of Combustion Reactions• since combustion reactions are exothermic,

their reverse reactions are endothermic

• the reverse of a combustion reaction involves the production of O2

energy + 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g)

energy + CO2(g) + 2 H2O(g) → CH4(g) + 2 O2(g)

• reactions in which O2 is gained or lost are redox reactions

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REDOX INBIOLOGICAL SYSTEMS

Many important biological reactions involve oxidation and reduction.

In these reactions, oxidation involves addition of oxygen or loss of hydrogen, and reduction involves loss of oxygen or gain of hydrogen.

For example, poisonous methyl alcohol is metabolized by the body by the following reaction:

CH3OH H2CO + 2H•

methyl alcohol formaldehyde

Oxidation (loss of

hydrogen)

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REDOX INBIOLOGICAL SYSTEMS

The formaldehyde is further oxidized to formic acid and finally carbon dioxide and water by the following reactions:

2 H2CO + O2 2H2CO2

formic acidformaldehyde

Oxidation (gain of oxygen)

formic acid

2 H2CO2 + O2 CO2 + H2O

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REDOX INBIOLOGICAL SYSTEMS

In many biochemical oxidation-reduction reactions, the transfer of hydrogen atoms produces energy in the cells.

The oxidation of a typical biochemical molecule can involve the transfer of two hydrogen atoms to a proton acceptor such as coenzyme FAD to produce its reduced form FADH2.

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REDOX INBIOLOGICAL SYSTEMS

In summary, the particular definition of oxidation-reduction depends on the process that occurs in the reaction.

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THE END