Download Chapter 12 - "Chemical Formulas and Equations"

• Chemical Formulas and
Equations
• 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
• Introduction
– Empirical formula
• The simplest whole number ratio of elements in a compound
• Ionic compounds are always shown as empirical formulas.
– Molecular Formula
• The actual numbers of atoms in a molecule.
– Structural Formula
• Show the relative arrangements of atoms in a molecule
• The name, molecular formula, sketch, and structural
formula of some common molecules. Compare the kinds
and numbers of atoms making up each molecule in the
sketch to the molecular formula.
• 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.
• Molecular and Formula Weights
– The formula weight of a compound is the sum of all of
the atomic weights of the atoms in a chemical formula.
– The formula weight of an ionic compound is found by
adding up all of the atomic weights of the atoms in the
compound.
– The molecular weight is the formula weight of a
molecule
• Percent Composition of Compounds
• Mass percent is the ratio of the atomic weight of all
the atoms of a certain element to the total formula of
molecular weight of the substance.
• Chemical Equations
• Introduction
– Chemical reactions occur when bonds between the
outermost parts of atoms are formed or broken
– Chemical reactions involve changes in matter, the
making of new materials with new properties, or energy
changes.
– Chemical Reactions are described using a shorthand
called a chemical equation
• Balancing equations
– Chemical equations show the conversion of reactants (the
molecules traditionally shown on the left of the arrow)
into products (the molecules traditionally shown on the
right of the arrow).
• A + sign shows that the compounds on the left of the arrow
combine to form the product or products on the right of the
arrow.
• The arrow is read as “reacts to yield”
• Example
– C + O2  CO2
• This reads carbon plus oxygen reacts to yield carbon dioxide
• The charcoal used in a grill is basically carbon. The carbon
reacts with oxygen to yield carbon dioxide. The chemical
equation for this reaction, C + O2  CO2, contains the same
information as the English sentence but has quantitative
meaning as well.
– An equation is balanced when the total number of atoms
of each element occur the same number of times on both
sides of the equation.
– Law of conservation of mass.
• Atoms are neither created nor destroyed in a chemical reaction.
• They may, however, be combined differently or converted into
energy.
– When balancing a chemical reaction you may add
coefficients in front of the compounds to balance the
reaction, but you may not change the subscripts.
• Changing the subscripts changes the compound.
• 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.
– 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”
• 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.
– 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.
• Oxidation-reduction reactions
– An oxidation reduction reaction is one in which electrons
are transferred between atoms.
– Oxidation is the loss of electrons
– Reduction is the gain of electrons
– Oxidizing agents are substances which take electrons
away from other atoms.
• An oxidizing agent is reduced when it oxidizes another atom
– Reducing agents are substances which donate electrons to
other substances.
• A reducing agent is oxidized in the process of reducing another
atom.
• Oxidizing agents take electrons from other substances that
are being oxidized. Oxygen and chlorine are commonly
used, strong oxidizing agents.
• 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
crews 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 usuall need 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 tom 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
• Information from chemical
equations
• 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 involve 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.
– 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.
– The gram atomic weight of an element is the mass in
grams of one mole of an element that is numerically
equal to its atomic weight.
– The gram formula weight of a compound is the mass in
grams of one mole of the compound that is numerically
equal to its formula weight.
• The gram formula weight of a compound is the sum total of all
the individual atomic weight in the formula.
– The gram molecular weight is the gram formula weight
of a molecular compound.
• Quantitative uses of equations
– A balanced chemical equation can be used to interpret
the:
• Molecular ratio of reactants to products.
• Mole ratio of the reactants to products.
• Mass ratio of the reactants to 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.