Ch. 8: Chemical Reactions
and Chemical Quantities
Dr. Namphol Sinkaset
Chem 200: General Chemistry I
I. Chapter Outline
I.
II.
III.
IV.
V.
VI.
Introduction
Chemical Change
Writing and Balancing Equations
Reaction Stoichiometry
Limiting Reactants and % Yield
Three Types of Reactions
I. Reaction Chemistry
•
•
The stereotypical chemist sits in a lab
and runs reactions w/ glassware and
chemicals.
In reaction chemistry, we are concerned
with two things.
1) What will happen if we mix this and that?
2) How much reactant do we need or how
much product will form?
II. Chemical Change
• Chemical change occurs when atoms
rearrange, such that the original
compound becomes a different one.
• Contrast with physical change, in which
the composition remains the same.
• These are closely related to chemical
and physical properties.
II. Chemical/Physical Properties
• A chemical property is only displayed
when a substance changes its
composition.
 e.g. flammability – when a substance
exhibits flammability, it becomes one or
more different substances
• A physical property is displayed without
changing composition.
 e.g. smell or color
II. Chemical/Physical Change
III. Chemical Reactions
• Chemical changes occur via chemical
reactions, of which there are several
types.
• We represent chemical reactions with
chemical equations.
• Chemical equations must be balanced.
Why?
III. Balancing Equations
• Steps for balancing:
1) Translate into formulas, if necessary. Reactants
on left, products on right.
2) Balance atoms one at a time. Start w/ most
complicated compound or element that’s NOT O
or H. Single elements are easy to balance last.
3) Can ONLY change coefficients, not subscripts!!
Adjust to smallest whole numbers, if necessary.
4) Check work!
5) Specify states.
III. Balancing Practice
•
e.g. Balance the following.
a) Solid calcium carbonate reacts with nitric
acid to form carbon dioxide gas, liquid
water, and aqueous calcium nitrate.
b) PCl3 + HF  PF3 + HCl
c) C4H10(g) + O2(g)  CO2(g) + H2O(g)
III. Last Thoughts on Balancing
• Coefficients distribute over all atoms in
a formula.
• Never change a formula when trying to
balance. e.g. MgO cannot be changed
to MgO2 to balance O atoms.
• Never add other reactants/products to
try to achieve balance.
• A balanced equation stays balanced
when multiplied by a constant.
IV. Calculations w/ Equations
• A balanced equation allows calculations
of amounts of reactants or products.
• If you know the # of moles of one
substance in a balanced equation, you
know the # of moles of any of the other
substances.
• Numerical relationships between
chemical amounts in a balanced
chemical equation are called reaction
stoichiometry.
IV. Mole Ratios
• Many stoichiometric relationships exist
in any one balanced equation.
• e.g. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l)
 1 mole C3H8 reacts w/ 5 moles O2
 1 mole C3H8 produces 3 moles CO2
 5 moles O2 leads to 4 moles H2O
• These can be converted to mole ratios.
1 mole C3H8
5 moles O2
1 mole C3H8
3 moles CO2
5 moles O2
4 moles H2O
IV. Mass-to-Mass Conversions
• How much CO2 produced from burning gasoline?
• 2 C8H18 + 25 O2  16 CO2 + 18 H2O
IV. Sample Problem
• How many grams of iron form when 135
g of aluminum reacts according to the
reaction below?
Fe2O3(s) + 2Al(s)  Al2O3(s) + 2Fe(l)
IV. Sample Problem
• How many kg of HNO3 forms, assuming
adequate O2 and H2O, if 16 kg of NO2
reacts according to the equation below?
NO2(s) + O2(s) + H2O(l)  HNO3(aq)
V. Limiting Reactants
• Consider the following reaction.
 2CO(g) + O2(g)  2CO2(g)
• Submicroscopically, 2 molecules CO
react with 1 molecule O2 to give 2
molecules CO2.
• How much CO2 forms if we have 2
molecules CO and 10 molecules of O2?
• We call CO the limiting reactant.
V. Sample Problem
• How many grams of Al2S3 can be
formed according to the reaction below
if 10.0 g Al is reacted with 15.0 g S8?
16Al(s) + 3S8(s)  8Al2S3(s)
V. Sample Problem
• What is the maximum amount of silicon
dioxide that can be produced when 168
g of CaSiO3 reacts with 125 g of CO2
with adequate water according to the
reaction below?
2CO2(g) + CaSiO3(s) + H2O(l)  SiO2(s) + Ca(HCO3)2(aq)
V. Real-life Reactions
• In reality, we cannot have complete
conversion to products.
• Even if there was complete conversion,
difficult to actually collect all of the product.
V. Reaction Yields
• Stoichiometry gives us theoretical yield.
• What is collected in lab is the actual yield.
• The efficiency of a reaction is commonly
expressed as percent yield.
V. Sample Problem
• Elemental iron can be obtained using
the reaction below. If 167 g of Fe2O3
reacts with 85.8 g of CO to produce
72.3 g of Fe, what is the percent yield?
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
VI. Three Chemical Reactions
• We will now look briefly at 3 examples
of chemical reactions:
 Combustion reactions
 Alkali metal reactions
 Halogen reactions
VI. Combustion Reactions
• A combustion reaction involves the
reaction of a substance with O2 to form
one or more oxygen-containing
compounds.
 One of the products is often H2O.
 Combustion reactions also emit heat.
VI. Typical Combustion Reactions
• In the combustion reactions you will
commonly see, the reactant will only
have C, H, and O.
 All of the C becomes CO2 and all of the H
becomes H2O.
 CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
 If there are other elements in the reactant,
they would form an oxide.
VI. Sample Problem
• Write a balanced reaction for the
complete combustion of liquid C2H5SH.
VI. Alkali Metal Reactions
• Alkali metals have ns1 valence econfigs, so they are quite reactive.
• They react with nonmetals.
 2M + X2  2MX
• They react with water.
 2M(s) + 2H2O(l)  2M+(aq) + 2OH-(aq) + H2(g)
VI. Na/Cl2 and Alkali Metals/H2O
VI. Halogen Reactions
• Halogens have ns2np5 outer e- configs, so
they are quite reactive.
• React w/ metals to form metal halides.
 2Fe(s) + 3Cl2(g)  2FeCl3(s)
• React w/ H2 to form hydrogen halides.
 H2(g) + I2(g)  2HI(g)
• React w/ each other to form interhalogens.
 Br2(l) + F2(g)  2BrF(g)