Chapter 6Chapter 6

Chemical ReactionsChemical Reactions

Chemical ReactionsChemical Reactions

In a chemical reaction, one or more reactants is converted to one or more products

In this chapter we discuss three aspects of chemical reactions(a) mass relationships (stoichiometry)

(b) types of reactions

(c) heat gain and loss accompanying reactions

Reactant(s) Product(s)

Chemical EquationsChemical Equations

The following chemical equation tells us that propane gas and oxygen gas react to form carbon dioxide gas and water vapor

But while it tells us what the reactants and products are and the physical state of each, it is incomplete because it is not balanced

C3H8(g) + O2(g) CO2(g) +H2O(g)

Propane Oxygen Carbondioxide

Water

Balancing EquationsBalancing Equations To balance a chemical equation

– begin with atoms that appear only in one compound on the left and one on the right; in this case, begin with carbon (C) which occurs in C3H8 and CO2

– now balance hydrogens, which occur in C3H8 and H2O

– if an atom occurs as a free element, as for example Mg or O2, balance this element last; in this case O2

C3H8(g) + O2(g) 3CO2(g) + H2O(g)

C3H8(g) + O2(g) 3CO2(g) + 4H2O(g)

C3H8(g) +5O2(g) 3CO2(g) +4H2O(g)

Balancing EquationsBalancing Equations

Practice problems: balance these equationsCa(OH)2(s) + HCl(g) CaCl2(s) +H2O(l)

Calciumhydroxide

Calciumchloride

CO2(g) +H2O(l) C6H12O6(aq) + O2(g)

Glucose

photosynthesis

C4H10(g)+ O2(g) CO2(g) + H2O(g)Butane

Balancing EquationsBalancing Equations Solutions to practice problems

– it is common practice to use only whole numbers; therefore, multiply all coefficients by 2, which gives

Ca(OH)2(s) + 2HCl(g) CaCl2(s) +H2O(l)

Calciumhydroxide

Calciumchloride

6CO2(g) +6H2O(l) C6H12O6(aq) + 6O2(g)

Glucose

photosynthesis

C4H10(l) + O2(g) 4CO2(g) + 5H2O(g)Butane

132

2C4H10(l) +13O2(g) 8CO2(g) +10H2O(g)Butane

Formula WeightFormula WeightFormula weight: the sum of the atomic

weights in atomic mass units (amu) of all atoms in a compound’s formula

Ionic CompoundsSodium chloride (NaCl) 23.0 amu Na + 35.5 amu Cl = 58.5 amu

Aspirin (C9H8O4) 9(12.0 amu C) + 8(1.0 amu H) + 4(16.0 amu O) = 180.0 = amu

Water (H2O) 2(1.0 amu H) + 16.0 amu O = 18.0 amu

Nickel(II) chloride hydrate(NiCl2•6H2O)

58.7 amu Ni + 2(35.5 amu Cl) + 12(1.0) amu H) + 6(12.0 amu O) = 237.7 amu

Molecular Compounds

Formula WeightFormula Weight

formula weightformula weight can be used for both ionic and molecular compounds; it tells nothing about whether a compound is ionic or molecular

molecular weightmolecular weight should be used only for molecular compounds

in this text, we use formula weight for ionic compounds and molecular weight for molecular compounds

The MoleThe Mole

Mole (mol)– a mole of the amount of substance that contains as

many atoms, molecules, or ions as are in exactly 12 g of carbon-12

– a mole, whether it is a mole of iron atoms, a mole of methane molecules, or a mole of sodium ions, always contains the same number of formula units

– the number of formula units in a mole is known as Avogadro’s number

– Avogadro’s number has been measured experimentally– its value is 6.02214199 x 1023 formula units per mole

Molar MassMolar Mass

Molar mass:Molar mass: the formula weight of a substance expressed in grams

Glucose, C6H12O6

– molecular weight: 180 amu– molar mass: 180 g/mol– one mole of glucose has a mass of 180 g

Urea, (NH2)2CO– molecular weight 60.0 amu– molar mass: 60.0 g/mol– one mole of urea has a mass of 60.0 g

Molar MassMolar Mass We can use molar mass to convert from grams to

moles, and from moles to grams

– calculate the number of moles of water in 36.0 g water

36.0 g H2O1 mol H2O

18.0 g H2O= 2.00 mol H2Ox

Moles of AGrams of A

Use molar mass (g/mol)as the conversion factor

You are given one of theseand asked to find the other

Grams to MolesGrams to Moles Calculate the number of moles of sodium ions,

Na+, in 5.63 g of sodium sulfate, Na2SO4

– first we find the how many moles of sodium sulfate– the formula weight of Na2SO4 is 2(23.0) + 32.1 + 4(16.0) = 142.1 amu– therefore, 1 mol of Na2SO4 = 142.1 g Na2SO4

– the formula Na2SO4 tells us there are two moles of Na+ ions per mole of Na2SO4

5.63 g Na2SO4 x1 mol Na2SO4

142.1 g Na2SO4

= 0.0396 mol Na2SO4

0.0396 mol Na2SO42 mol Na+

1 mol Na2SO4

=x 0.0792 mol Na+

Grams to MoleculesGrams to Molecules

A tablet of aspirin, C9H8O4, contains 0.360 g of aspirin. How many aspirin molecules is this?– first we find how many mol of aspirin are in 0.360 g

– each mole of aspirin contains 6.02 x 1023 molecules– the number of molecules of aspirin in the tablet is

0.360 g aspirin x1 mol aspirin

180.0 g aspirin= 0.00200 mol aspirin

0.00200 mole x 6.02 x 1023 moleculesmole

= 1.20 x 1021 molecules

StoichiometryStoichiometryStoichiometry:Stoichiometry: the study of mass

relationships in chemical reactions– following is an overview of the the types of

calculations we study

Moles of AGrams of A Grams of BMoles of B

From moles to moles, use the coefficients inthe balanced equationas a conversion factor

From grams to moles,use molar mass (g/mol)as a conversion factor

From moles to grams,use molar mass (g/mol) as a conversion factor

You are given one of these And asked to find one of these

StoichiometryStoichiometry Problem: how many grams of nitrogen, N2, are

required to produce 7.50 g of ammonia, NH3

– first find how many moles of NH3 are in 7.50 g of NH3

– next find how many moles of N2 are required to produce this many moles of NH3

7.50 g NH3 x 1 mol NH3

17.0 g NH3

= mol NH3

7.50 g NH3 x 1 mol NH3

17.0 g NH3

x 1 mol N2

2 mol NH3

= mol N2

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

StoichiometryStoichiometry

Practice problem (cont’d)– finally convert moles of N2 to grams of N2 and

now do the math

x 28.0 g N2

1 mol N2

= 6.18 g N27.50 g NH3 x

1 mol NH3

17.0 g NH3

x 1 mol N2

2 mol NH3

StoichiometryStoichiometryPractice problems:

– what mass of aluminum oxide is required to prepare 27 g of aluminum?

– how many grams each of CO22 and NH33 are produced from 0.83 mol of urea?

Al2O3(s)electrolysis Al(s) + O2(g)

(NH2)2CO(aq) + 2NH3(aq) + CO2(g)H2OUrea

urease

Limiting ReagentLimiting Reagent Limiting reagentLimiting reagent: the reagent that is used up first

in a chemical reaction– consider this reaction of N2 and O2

– in this experiment, there is only enough O2 to react with 1.0 mole of N2

– O2 is used up first; it the limiting reagent– 4.0 moles of N2 remain unreacted

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

before reaction (moles) 5.0 1.0 0

after reaction (moles) 4.0 0 2.0

Limiting ReagentLimiting ReagentPractice Problem

– suppose 12 g of carbon is mixed with 64 g of oxygen and the following reaction takes place

– complete the following table. Which is the limiting reagent?

C(s) + CO2(g)O2(g)

C O2 CO2+

before reaction (g)

before reaction (mol)

after reaction (mol)

after reaction (g)

12 g 64 g 0

Percent YieldPercent YieldActual yield:Actual yield: the mass of product formed in

a chemical reactionTheoretical yield:Theoretical yield: the mass of product that

should be formed according to the stoichiometry of the balanced chemical equation

Percent yield:Percent yield: actual yield divided by theoretical yield times 100

Percent yield = Actual yieldTheoretical yield

x 100

Percent YieldPercent YieldPractice problem:

– suppose we react 32.0 g of methanol with excess carbon monoxide and get 58.7 g of acetic acid

– complete this tableCH3OH CO CH3COOH+

before reaction (g) 32.0 excess 0before reaction (mol)

theoretical yield (mol)theoretical yield (g)

percent yield (%)actual yield (g) 58.7

Reactions Between IonsReactions Between Ions Ionic compounds, also called salts, consist of both

positive and negative ions When an ionic compound dissolves in water, it

dissociates to aqueous ions

What happens when we mix aqueous solutions of two different ionic compounds?– if two of the ions combine to form a water-insoluble

compound, a precipitate will form– otherwise no physical change will be observed

NaCl(s) +Na+(aq)H2O

Cl-(aq)

Reactions Between IonsReactions Between Ions Example:

– suppose we prepare these two aqueous solutions

– if we then mix the two solutions, we have four ions present; of these, Ag+ and Cl- react to form AgCl(s) which precipitates

+Ag+(aq) Cl-(aq)

AgCl(s)

NO3-(aq) + Na+(aq) +

+ Na+(aq) + NO3-(aq)

AgNO3(s)H2O Ag+(aq) + NO3

-(aq)Solution 1

NaCl(s)H2O

Na+(aq) + Cl-(aq)Solution 2

Reactions Between IonsReactions Between Ions– we can simplify the equation for the formation

of AgCl by omitting all ions that do not participate in the reaction

– the simplified equation is called a net ionic net ionic equationequation; it shows only the ions that react

– ions that do not participate in a reaction are called spectator ionsspectator ions

Ag+(aq) Cl-(aq) AgCl(s)+Net ionic equation:

Reactions Between IonsReactions Between Ions In general, ions in solution react with each other

when one of the following can happen– two of them form a compound that is insoluble in water– two of them react to form a gas that escapes from the

reaction mixture as bubbles, as for example when we mix aqueous solutions of sodium bicarbonate and hydrochloric acid

– an acid neutralizes a base (Chapter 8)– one of the ions can oxidize another (Section 4.7)

HCO3-(aq) + H3O

+(aq) +CO2(g) 2H2O(l)Bicarbonate ion Carbon dioxide

Reactions Between IonsReactions Between Ions Following are some generalizations about which ionic

solids are soluble in water and which are insoluble– all compounds containing Na+, K+, and NH4

+ are soluble in water

– all nitrates (NO3-) and acetates (CH3COO-) are soluble in

water– most chlorides (Cl-) and sulfates (SO4

2-) are soluble; exceptions are AgCl, BaSO4, and PbSO4

– most carbonates (CO32-), phosphates (PO4

3-), sulfides (S2-), and hydroxides (OH-) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH4OH which are soluble in water

Oxidation-ReductionOxidation-Reduction

Oxidation:Oxidation: the loss of electronsReduction:Reduction: the gain of electronsOxidation-reduction (redox) reaction:Oxidation-reduction (redox) reaction: any

reaction in which electrons are transferred from one species to another

Oxidation-ReductionOxidation-Reduction Example: if we put a piece of zinc metal in a

beaker containing a solution of copper(II) sulfate– some of the zinc metal dissolves– some of the copper ions deposit on the zinc metal– the blue color of Cu2+ ions gradually disappears

In this oxidation-reduction reaction– zinc metal loses electrons to copper ions

– copper ions gain electrons from the zinc

Zn(s) +Zn2+(aq) 2e- Zn is oxidized

Cu(s)+ 2e-Cu2+(aq) Cu2+ is reduced

Oxidation-ReductionOxidation-Reduction

– we summarize these oxidation-reduction relationships in this way

electrons flowfrom Zn to Cu2+

+Zn(s) Cu2+(aq) + Cu(s)Zn2+(aq)

loses electrons;is oxidized

gains electrons;is reduced

gives electronsto Cu2+; is thereducing agent

takes electronsfrom Zn; is the oxidizing agent

Oxidation-ReductionOxidation-Reduction Although the definitions of oxidation (loss of

electrons) and reduction (gain of electrons) are easy to apply to many redox reactions, they are not easy to apply to others– for example, the combustion of methane

An alternative definition of oxidation-reduction is– oxidation:oxidation: the gain of oxygen or loss of hydrogen– reduction:reduction: the loss of oxygen or gain of hydrogen

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

Oxidation-ReductionOxidation-Reduction

– using these alternative definitions for the combustion of methane

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

gains O and losesH; is oxidized

gains H; is reduced

is the reducingagent

is the oxidizingagent

electrons are transferred from carbon to oxygen

Oxidation-ReductionOxidation-ReductionFive important types of redox reactions

– combustion:combustion: burning in air. The products of complete combustion of carbon compounds are CO2 and H2O.

– respiration:respiration: the process by which living organisms use O2 to oxidize carbon-containing compounds to produce CO2 and H2O. The importance of these reaction is not the CO2 produced, but the energy released.

– rusting:rusting: the oxidation of iron to a mixture of iron oxides

– bleaching:bleaching: the oxidation of colored compounds to products which are colorless

– batteries:batteries: in most cases, the reaction taking place in a battery is a redox-reaction

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

Heat of ReactionHeat of Reaction In almost all chemical reactions, heat is either given off or

absorbed– example: the combustion (oxidation) of carbon liberates 94.0 kcal

per mole of carbon oxidized

Heat of reaction:Heat of reaction: the heat given off or absorbed in a chemical reaction– exothermic reaction:exothermic reaction: one that gives off heat– endothermic reaction:endothermic reaction: one that absorbs heat– heat of combustion:heat of combustion: the heat given off in a combustion reaction;

all combustion reactions are exothermic

C(s) + O2(g) CO2(g) + 94.0 kcal/mole C

End End Chapter 6Chapter 6

Chemical ReactionsChemical Reactions