Chemical reactions
Chapter 10
• The products of chemical research have substantially
increased food supplies but have also increased the
possibilities of pollution. Balancing the benefits and
hazards of the use of chemicals requires a knowledge
of chemistry and a knowledge of the alternatives.
Chemical formulas
1. Empirical formula
– Identifies elements
present in terms of
simplest whole
number ratios
– Examples: table
salt, NaCl; water,
H2O
2. Molecular formula
– Identifies actual
number of atoms in
a molecular
compound
– Example: water,
H2O; not table salt,
NaCl (ionic
compound)
Chemical formulas, cont.
3. Structural formula
– Represents
arrangement of
atoms within a
molecule
– Related to 3-D
structure of molecule
Empirical or molecular
formula?
Empirical
• Ionic - lacking discrete
unit, or molecule
• Composed of both
metallic and nonmetallic
elements
• Electronegativity
difference > 1.7
Molecular
• Covalent compounds
• Usually nonmetals bound
to nonmetals
• Molecular and empirical
formulas can be different
– Glucose – molecular
C6H12O6 versus empirical
CH2O
Molecular and formula weights
• Formula weight
– Sum of atomic
weights of all atoms
in chemical formula
• Molecular weight
– Formula weight of a
molecular substance
– Term often used for
nonmolecular
substances, as well
Percent composition of
compounds
• Finding the mass percentage of an individual
element from the formula weight
Percentage composition
example
• If you know the name of an ingredient, you can write a chemical
formula, and the percent composition of a particular substance
can be calculated from the formula. This can be useful
information for consumer decisions.
Chemical reactions
• Occur through formation and breaking
of chemical bonds between atoms
• Involve changes in matter, creation of
new materials and energy exchange
• Chemical equations - concise
representation of chemical reactions
Chemical equations
•
•
•
•
Reactants - substances existing before reaction
Products - substances existing after reaction
Word representation not sufficiently precise
Chemical symbols and formulas needed for
quantitative purposes
Balancing equations
• Law of conservation
of mass - atoms are
neither created nor
destroyed in
chemical reactions
• Change coefficients
in front of chemical
formulas, not
coefficients within
formulas, to balance
mass of reactants = mass of products
• The meaning of subscripts and coefficients used with
a chemical formula. The subscripts tell you how many
atoms of a particular element are in a compound. The
coefficient tells you about the quantity, or number, of
molecules of the compound.
• Tanks like these grow larger as they are filled with natural gas,
then collapse back to the ground as the gas is removed. Why
do you suppose the tanks are designed to inflate and collapse?
One reason is to keep the gas under a constant pressure. The
height of each tank varies with the amount of gas inside, so
more gas means a greater volume rather than a greater
pressure. A rigid gas tank with a constant volume would be
under very high pressure when full and very low pressure when
nearly empty, which would make it difficult to pump gas into or
out of the tank.
– There are four basic steps to balancing a chemical
equation.
• Write the correct formula for the reactants and
the products in an unbalanced equation.
• Inventory the number of each kind of atom on
both sides of the unbalanced equation.
• Determine where to place coefficients in front of
formulas to balance the equation.
• Take another inventory to determine if:
–The numbers of atoms on both sides of the
equation are now balanced.
–The coefficients are in the lowest possible
whole number ratios.
– When balancing equations remember the following:
• Atoms are neither lost nor gained during a
chemical reaction.
• A correct formula of a compound cannot be
changed by altering the number or placement of
subscripts.
• A coefficient in front of a formula multiplies
everything in the formula by that number.
• Compare the numbers of each kind of atom in the
balanced equation with the numbers of each kind of
atom in the sketched representation. Both the
equation and the sketch have the same number of
atoms in the reactants and in the products.
Stepwise balancing procedure
– A general approach to balancing reactions is:
• Look first to formulas of compounds with the
most atoms and try to balance the atoms or
compounds they were formed from or
decomposed to.
• Polyatomic ions that appear on both sides of the
equation should be balanced as independent
units with their charge.
• Try both the “Crossover technique” and the use
of “fractional coefficients”
– Conventions
• gas (g)
• Liquid (l)
• Aqueous solution (aq)
• Escaping gas ()
• Solid formation ()
• Change of temperature ()
• One of two burners is
operating at the
moment as this hot air
balloon ascends. The
burners are fueled by
propane (C3H8), a
liquified petroleum gas
(LPG). Like other forms
of petroleum, propane
releases large amounts
of heat during the
chemical reaction of
burning.
• Hydrocarbons are composed of the elements hydrogen and
carbon. Propane (C3H8) and gasoline, which contain octane
(C8H18) are examples of hydrocarbons. Carbohydrates are
composed of the elements of hydrogen, carbon, and oxygen.
Table sugar, for example, is the carbohydrate C12H22O11.
Generalizing, all hydrocarbons and carbohydrates react
completely with oxygen to yield CO2 and H2O.
Types of chemical reactions
• Oxidation-reduction (redox) reactions
– Oxidation - loss of electrons by an atom
– Reduction - gain of electrons by an atom
– Oxidizing agents - substances taking electrons from other
substances (oxygen, chlorine)
– Reducing agents - supply electrons to oxidizing agents (hydrogen,
carbon)
• Oxidizing agents take electrons from other substances
that are being oxidized. Oxygen and chlorine are
commonly used, strong oxidizing agents.
Alternative classification
1. Combination reactions
2. Decomposition reactions
3. Replacement reactions

(1-3 = redox reaction subclasses)
4. Ion exchange reactions
• Types of chemical reactions
– Combination reactions
• This is a synthesis reaction where several atoms
or molecules combine to form one or more new
compounds.
• The combining substances can be elements,
compounds, or combinations of these two.
• Example
–2 Mg (s) + O2 (g)  2 MgO (s)
• Magnesium and oxygen in this reaction combine
to form magnesium oxide.
• Rusting iron is a
common example of
a combination
reaction, where two
or more substances
combine to form a
new compound.
Rust is iron (III)
oxide formed on
these screws from
the combination of
iron and oxygen
under moist
conditions.
– Decomposition reactions
• A decomposition reaction is one in which a
compound is broken down
–Into elements.
–Into Simpler compounds
–Into both elements and simpler compounds
• In a decomposition reaction there usually needs
to be some sort of an input of energy to cause
the decomposition to proceed.
»Example
»

2HgO (s)  2 Hg (s) + O2 
• Mercury (II) oxide is decomposed by heat,
leaving the silver-colored element mercury
behind as oxygen is driven off. This is an
example of a decomposition reaction, 2 HgO 
2 Hg + O2 . Compare this equation to the
general form of a decomposition reaction.
– Replacement reactions
• A replacement reaction is one where an atom or a
polyatomic ion is exchanged for another atom or
polyatomic ion.
• These types of reactions occur as some elements have a
greater ability to hold or attract electrons to themselves.
• Elements that have the least ability to hold electrons are
the most reactive.
• A metal will replace any element that occurs above it in
the activity series.
• Metal ions above hydrogen in the activity series will
replace hydrogen as hydrogen ionizes from acids in
solution.
• Example
– Zn (s) + H2SO4 (aq)  ZnSO4 + H2 
• The activity series for common metals, together with some
generalizations about the chemical activities of the metals. The
series is used to predict which replacement reactions will take
place and which reactions will not occur. (Note that hydrogen is
not a metal and is placed in the series for reference to acid
reactions.)
• This shows a reaction between metallic aluminum and
the blue solution of copper (II) chloride. Aluminum is
above copper in the activity series, and aluminum
replaces the copper ions from the solution as copper
is deposited as a metal. The aluminum loses electrons
to the copper and forms aluminum ions in solution.
– Ion exchange reactions
• An ion exchange reaction is a reaction that takes
place when the ion of one compound interacts
with the ions of another compound forming.
–A solid that comes out of the solution
(precipitates)
–A gas
–Water
• Example
• 3Ca(OH)2 (aq) + Al2(SO4)3 (aq)  3CaSO4 (aq) + 2Al(OH)3
• Must refer to a solubility table, like the one in appendix B of
your
textbook, to determine if an ion exchange has taken place.
No ion exchange has taken place if the new products are both
soluble.
Information - chemical
equations
• Atoms are
conserved
• Mass is conserved
• Law of combining
volumes (gases)
– Gases at the same
temperature and
pressure contain
equal numbers of
molecules
• Introduction.
– Information from a balanced chemical equation tells
us information about:
• Atoms
• Molecules
• Atomic weights.
– The coefficients in the balanced reaction is the
number of atoms or molecules involved in the
reaction.
– In 1808 Gay-Lussac determined that gases
combine in small, whole number volumes when the
temperature and pressure were held constant.
• This is the Law of Combining Volumes.
– Avogadro proposed an explanation for the law of
combining volumes in 1811.
• It was proposed that gases at the same temperature
contained the same number of molecules.
– This had two implications for the coefficients in a
balanced equation
• The coefficients represent the number of molecules of
each substance
• It also represents the ratios of the combining volumes.
• Reacting gases combine in ratios of small, whole-number
volumes when the temperature and pressure are the same for
each volume. (A) One volume of hydrogen gas combines with
one volume of chlorine gas to yield two volumes of hydrogen
chloride gas. (B) Two volumes of hydrogen gas combine with
one volume of oxygen gas to yield two volumes of water vapor.
• Avogadro's
hypothesis of
equal volumes of
gas having equal
numbers of
molecules offered
an explanation for
the law of
combining
volumes.
Units of measurement used
with equations
• Atomic mass unit = 1/12
mass of carbon-12
• One mole of a
substance contains
Avogadro’s number
(6.02x1023) of the basic
chemical unit of that
substance (atoms,
molecules, ions, …)
• A mole of carbon-12
atoms is defined as
having a mass of
12.00g
• Units of measurement used with equations.
– We use a mole concept to bring together the
concepts of counting numbers and atomic weights
of elements.
– The mole is derived from the following information.
• Atomic weights are an average of the relative masses of
all of the isotopes of the given element.
• The number of C-12 atoms in exactly 12.00 g of C12 is
6.02 X 1023.
– This called Avogadro’s number.
• An amount of a substance that contains Avogadro’s
number of atoms, ions, molecules, or any other chemical
unit is called a mole.
• A mole of C-12 atoms is defined as having a mass of
exactly 12.00 g, a mass that is equal to its atomic weight.
• The mole concept for
(A) elements, (B)
compounds, and (C)
molecular
substances. A mole
contains 6.02 X 1023
particles. Since every
mole contains the
same number of
particles, the ratio of
the mass of any two
moles is the same as
the ratio of the
masses of individual
particles making up
the two moles.
Molar weights
• Gram atomic weight mass in grams equal to
atomic weight
• Gram formula weight mass in grams equal to
formula weight
• Gram molecular weight
- mass in grams equal
to molecular weight
Quantitative use of equations
Possible interpretations:
1. Molecular ratios of
reactants and
products
2. Mole ratios of
reactants and
products
3. Mass ratios of
reactants and
products
– The molecular ratio leads to the concept of the mole
ratio since any number of molecules can react as
long as they are in the correct ratio
– Since 6.02 X 1023 molecules is the number of
particles in a mole, the coefficients therefore
represent the number of moles involved in a
chemical reaction.
– The gram formula weight of a compound is the
mass in grams of one mole that is numerically equal
to its formula weight.
• The equation also describes the mass ratios of the
reactants to products.