Download Camp 1 - drjosephryan.com Home Page

Chapter 6
Chemical Reactions
Chemical Reactions

In a chemical reaction, one or more reactants is
converted to one or more products
Reactant(s)

Product(s)
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
Chemical Equations


The following chemical equation tells us that
propane gas and oxygen gas react to form carbon
dioxide gas and water vapor
C3 H8 ( g) + O2 (g)
CO2 (g) + H2 O( g)
Propane
Carbon
dioxid e
Oxygen
Water
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
Balancing 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
C3 H8 (g) + O2 (g)
3CO2( g) + H2 O(g)
– now balance hydrogens, which occur in C3H8 and H2O
C3 H8 (g) + O2 (g)
3CO2( g) + 4H2 O(g)
– if an atom occurs as a free element, as for example Mg
or O2, balance this element last; in this case O2
C3 H8 (g) + 5O2 ( g)
3CO2 (g) + 4H2 O( g)
Balancing Equations

Practice problems: balance these equations
Ca( OH) 2 ( s) + HCl( g)
Calcium
hydroxide
CO2 ( g) + H2 O(l)
CaCl2 (s) + H2 O( l)
Calcium
chlorid e
photosynthesis
C4 H1 0 ( g) + O2 (g)
Butane
C6 H1 2 O6 (aq) + O2 (g)
Glucose
CO2 (g) + H2 O(g)
Balancing Equations

Solutions to practice problems
Ca( OH) 2 ( s) + 2 HCl(g)
Calcium
hydroxide
6CO2 (g) + 6H2 O(l)
CaCl2 (s) + H2 O( l)
Calcium
chlorid e
photosynthesis
C4 H1 0 ( l) + 13 O2 (g)
2
Bu tane
C6 H1 2 O6 (aq) + 6O2 (g)
Glucose
4 CO2 (g) + 5 H2 O( g)
– it is common practice to use only whole numbers;
therefore, multiply all coefficients by 2, which gives
2C4 H1 0 ( l) + 13O2 (g)
Butane
8CO2 ( g) + 10H2 O(g)
Formula Weight

Formula weight: the sum of the atomic
weights in atomic mass units (amu) of all
atoms in a compound’s formula
Ionic Comp ou nds
Sod ium chlorid e (N aCl)
23.0 amu N a + 35.5 amu Cl = 58.5 amu
Nickel(II) ch loride h yd rate 58.7 amu N i + 2(35.5 amu Cl) +
12(1.0) amu H) + 6(12.0 amu O) = 237.7 amu
(N iCl 2• 6H 2O)
Molecu lar Comp ou nds
Water (H 2O)
Aspirin (C9H 8 O 4)
2(1.0 amu H) + 16.0 amu O = 18.0 amu
9(12.0 amu C) + 8(1.0 amu H) +
4(16.0 amu O) = 180.0 = amu
Formula Weight

formula weight can be used for both ionic and
molecular compounds; it tells nothing about
whether a compound is ionic or molecular
 molecular 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 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 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 Mass

We can use molar mass to convert from grams to
moles, and from moles to grams
You are given one of these
and asked to find the other
Grams of A
Moles of A
Use molar mass (g/mol)
as the conversion factor
– calculate the number of moles of water in 36.0 g water
1 mol H2 O
36.0 g H 2O x
= 2.00 mol H2 O
18.0 g H 2O
Grams 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
5.63 g Na2 SO4 x
1 mol Na2SO4
= 0.0396 mol Na2SO4
142.1 g Na2 SO4
– the formula Na2SO4 tells us there are two moles of Na+
ions per mole of Na2SO4
+
0.0396 mol Na2 SO 4
2 mol Na
x
1 mol Na2SO 4
= 0.0792 mol Na+
Grams 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
0.360 g aspirin x
1 mol aspirin
180.0 g aspirin
= 0.00200 mol aspirin
– each mole of aspirin contains 6.02 x 1023 molecules
– the number of molecules of aspirin in the tablet is
0.00200 mole x 6.02 x 1023 molecules = 1.20 x 1021 molecules
mole
Stoichiometry

Stoichiometry: the study of mass
relationships in chemical reactions
– following is an overview of the the types of
calculations we study
You are given one of these
Grams of A
And asked to find one of these
Moles of A
From grams to moles,
use molar mass (g/mol)
as a conversion factor
Moles of B
Grams of B
From moles to moles, From moles to grams,
use the coefficients in use molar mass (g/mol)
the balanced equation as a conversion factor
as a conversion factor
Stoichiometry

Problem: how many grams of nitrogen, N2, are
required to produce 7.50 g of ammonia, NH3
N2 (g) + 3H2 (g)
2NH3 ( g)
– first find how many moles of NH3 are in 7.50 g of NH3
7.50 g NH 3 x
1 mol NH 3
17.0 g NH 3
= mol NH 3
– next find how many moles of N2 are required to
produce this many moles of NH3
7.50 g NH 3 x
1 mol NH 3
17.0 g NH 3
x
1 mol N2
2 mol NH 3
= mol N 2
Stoichiometry

Practice problem (cont’d)
– finally convert moles of N2 to grams of N2 and
now do the math
7.50 g NH 3 x
1 mol NH 3
17.0 g NH 3
x
1 mol N2
2 mol NH 3
x
28.0 g N 2
1 mol N2
= 6.18 g N 2
Stoichiometry

Practice problems:
– what mass of aluminum oxide is required to
prepare 27 g of aluminum?
Al2 O3 ( s) electrolysis
Al( s) + O2 ( g)
– how many grams each of CO2 and NH3 are
produced from 0.83 mol of urea?
( NH2 ) 2 CO(aq) + H2 O
Urea
urease
2NH3 (aq) + CO2 (g)
Limiting Reagent

Limiting reagent: the reagent that is used up first
in a chemical reaction
– consider this reaction of N2 and O2
N2 (g) + O2 (g)
before reaction (moles) 5.0
after reaction (moles) 4.0
1.0
0
2 NO( g)
0
2.0
– 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
Limiting Reagent

Practice Problem
– suppose 12 g of carbon is mixed with 64 g of
oxygen and the following reaction takes place
C(s) + O2 ( g)
CO2 ( g)
– complete the following table. Which is the
limiting reagent?
C
before reaction (g)
12 g
before reaction (mol)
after reaction (mol)
after reaction (g)
+
O2
CO2
64 g
0
Percent Yield

Actual yield: the mass of product formed in
a chemical reaction
 Theoretical yield: the mass of product that
should be formed according to the
stoichiometry of the balanced chemical
equation
 Percent yield: actual yield divided by
theoretical yield times 100
Percent yield =
Actual yield
x 100
Theoretical yield
Percent Yield

Practice problem:
– suppose we react 32.0 g of methanol with
excess carbon monoxide and get 58.7 g of
acetic acid
– complete this table
CH3 OH + CO
before reaction (g) 32.0
before reaction (mol)
th eoretical yield (mol)
th eoretical yield (g)
actual yield (g)
percent yield (%)
excess
CH3 COOH
0
58.7
Reactions 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
NaCl(s)

H2 O
Na+ (aq) + Cl- (aq)
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
Reactions Between Ions

Example:
– suppose we prepare these two aqueous solutions
Solution 1
Solution 2
AgNO3 (s)
NaCl(s)
H2 O
H2 O
Ag+ ( aq) + NO3 - (aq)
+
Na ( aq) + Cl ( aq)
– 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) + NO3 ( aq) + Na (aq) + Cl ( aq)
+
AgCl(s) + Na (aq) + NO3 ( aq)
Reactions Between Ions
– we can simplify the equation for the formation
of AgCl by omitting all ions that do not
participate in the reaction
Net ionic equation:
Ag+ ( aq) + Cl-(aq)
AgCl( s)
– the simplified equation is called a net ionic
equation; it shows only the ions that react
– ions that do not participate in a reaction are
called spectator ions
Reactions 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
+
HCO3 ( aq) + H3 O (aq)
CO2 ( g) + 2 H2 O( l)
Bicarbonate ion
Carbon d ioxide
– an acid neutralizes a base (Chapter 8)
– one of the ions can oxidize another (Section 4.7)
Reactions 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 (SO42-) are soluble;
exceptions are AgCl, BaSO4, and PbSO4
– most carbonates (CO32-), phosphates (PO43-), sulfides
(S2-), and hydroxides (OH-) are insoluble in water;
exceptions are LiOH, NaOH, KOH, and NH4OH which
are soluble in water
Oxidation-Reduction

Oxidation: the loss of electrons
 Reduction: the gain of electrons
 Oxidation-reduction (redox) reaction: any
reaction in which electrons are transferred
from one species to another
Oxidation-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
2+
Zn(s)
Zn (aq) + 2 e
Zn is oxidized
– copper ions gain electrons from the zinc
2+
Cu ( aq) + 2 e
-
Cu( s)
Cu
2+
is red uced
Oxidation-Reduction
– we summarize these oxidation-reduction
relationships in this way
electrons flow
from Zn to Cu2 +
2+
Zn(s)
+
Cu (aq)
loses electrons ; gains electrons ;
is red uced
is oxidized
gives electrons tak es electrons
to Cu 2+ ; is th e from Zn; is th e
oxidizin g agent
red ucing agent
2+
Zn ( aq) + Cu( s)
Oxidation-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
CH4 (g) + O2 ( g)
CO2 (g) + H2 O( g)
Methane

An alternative definition of oxidation-reduction is
– oxidation: the gain of oxygen or loss of hydrogen
– reduction: the loss of oxygen or gain of hydrogen
Oxidation-Reduction
– using these alternative definitions for the
combustion of methane
electrons are
trans ferred from
carb on to oxygen
CH4 (g)
gain s O and los es
H; is oxidized
+
O2 (g)
gain s H;
is reduced
is the reducin g is th e oxid izing
agent
agent
CO2 (g) + H2 O(g)
Oxidation-Reduction

Five important types of redox reactions
– combustion: burning in air. The products of complete
combustion of carbon compounds are CO2 and H2O.
– 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: the oxidation of iron to a mixture of iron oxides
4Fe(s) + 3O2 ( g)
2Fe2 O3 ( s)
– bleaching: the oxidation of colored compounds to
products which are colorless
– batteries: in most cases, the reaction taking place in a
battery is a redox-reaction
Heat 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
C( s) + O2 (g)

CO2 (g) + 94.0 kcal/mole C
Heat of reaction: the heat given off or absorbed in
a chemical reaction
– exothermic reaction: one that gives off heat
– endothermic reaction: one that absorbs heat
– heat of combustion: the heat given off in a combustion
reaction; all combustion reactions are exothermic
Chemical Reactions
End
Chapter 6