Transcript 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
dioxide
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
ph otosynthes is
C4 H1 0 ( g) + O2 (g)
Bu tane
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
6 CO2 (g) + 6 H2 O(l)
CaCl2 (s) + H2 O( l)
Calcium
chlorid e
ph otosynthes is
C4 H1 0 ( l) + 13 O2 (g)
2
Butane
C6 H1 2 O6 (aq) + 6 O2 (g)
Glucose
4CO2 (g) + 5H2 O( g)
– it is common practice to use only whole numbers;
therefore, multiply all coefficients by 2, which gives
2 C4 H1 0 ( l) + 1 3 O2 (g)
Butane
8 CO2 ( g) + 1 0 H2 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 SO 4 x
1 mol Na2SO 4
= 0.0396 mol Na2SO 4
142.1 g Na2 SO 4
– the formula Na2SO4 tells us there are two moles of Na+
ions per mole of Na2SO4
+
0.0396 mol Na2 SO4
2 mol Na
x
1 mol Na2SO4
= 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 N H 3 x
1 mol N H 3
17.0 g N H 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
2NO( 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
Solu tion 1
Solu tion 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
Zn(s)
Zn2 + (aq) + 2e-
Zn is oxidized
– copper ions gain electrons from the zinc
Cu2 + ( aq) + 2e-
Cu( s)
Cu2+ is reduced
Oxidation-Reduction
– we summarize these oxidation-reduction
relationships in this way
electrons flow
from Zn to Cu2 +
Zn(s)
+
loses electrons ;
is oxidized
gives electrons
to Cu 2+ ; is th e
red ucing agent
2+
Cu (aq)
gains electrons ;
is red uced
tak es electrons
from Zn; is th e
oxidizin g 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)
Methan e

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