CPC 022
DAY 4
Chemical Reactions
Dr. Olivieri
Joseph Black was an Irish
born Chemist who taught
the world that the
combination of gas and
solid was possible.
Although it sounds
elementary, Black
discovered that there were
different types of air. He
changed what the world
knew of gasses.
Copyright Dr. James DeHaven 2006
1/Day 4
•
Chemical Reactions, an introduction
Chemical Reactions
When one pure substance is changed into another pure substance, we
say that a chemical reaction has taken place.
During a chemical reaction, existing chemical bonds are broken
and (usually) new chemical bonds are formed.
A chemical reaction can involve changes in several pure substances at
once...it is not limited to one pure substance at a time.
Almost all biochemical processes, the mining and refining of metals,
the combustion of hydrocarbons, the fabrication of plastics...all involve
chemical reactions.
•
Chemical Equations
A chemical equation describes, in symbols, what is taking place
during a chemical reaction.
When potassium chlorate is heated in the presence of manganese
dioxide, oxygen gas is released and solid potassium chloride is left
over. The manganese dioxide is unchanged. It is a catalyst, a
substance that speeds up the reaction but does not participate
in the reaction.
This is a verbal description of a chemical reaction. The symbolic
description the chemical equation, is shown below:
2/Day 4
o What they mean
Chemical Equations 2
The catalytic decomposition of KClO3 is represented...as seen on the
previous page..by the equation:
Several features of this equation are worth noting:
A - The physical states of the participants are labeled by
(s) for solid and (g) for gas.
B - The triangle under the reaction arrow means that the
reaction must be heated.
C - The MnO2 above the reaction arrow indicates its status
as a catalyst.
D - The stoichiometric coefficients tell us that, for every 2
moles of KClO3 we get 3 moles of O2 and 2 moles of KCl.
o Symbols
Chemical Equations 3
Every chemical equation is of the form:
It may contain any or all of the following symbols:
•
Stoichiometric coefficient
3/Day 4
If there is no number in front of a reactant or product, the number 1 is implied.
description of state
o (s) = solid
o (g) = gas
o (l) = liquid
(aq) - aqueous solution (dissolved in water)
•
o
Description of product state
Sometimes arrows are used to denote the formation of a solid (precipitate) or of a gas.
o These are only used on the product side.
o An arrow pointing up indicates that a gas is evolved.
o An arrow pointing down indicates the formation of a precipitate.
•
o
•
A triangle above or below the reaction arrow means that the reactants are heated.
Any chemical formula found above or below the reaction arrow means a catalyst is present.
•
o
(Pronounced delta H) = some number at
the end of the equation.

heat is given off if (delta h) is
negative (exothermic).

heat is absorbed if (delta h) is
positive (endothermic).
4/Day 4
These are alternate "reaction
arrows" implying that the final
mixture contains significant
amounts of reactant in
equilibrium with product.
o Interpreting chemical equations
Chemical Equations 4
Let's look at a few chemical equations and interpret them.
This says that 1 mole of aqueous hydrochloric acid reacts with 1 mole
of sodium bicarbonate to give a mole of sodium chloride plus a mole of
gaseous carbon dioxide and a mole of liquid water.
2 moles of potassium iodide in aqueous solution react with a mole of
aqueous lead nitrate to give two moles of aqueous potassium nitrate
plus a mole of lead (II) iodide precipitate.
Now you try it. Write verbal description of these:
5/Day 4
Here are the solutions
Chemical Equation 5 Answers
An aqueous solution of carbonic acid (H3CO3) upon heating evolves
gaseous CO2 and liquid water. For every mole of H2CO3 decomposed,
one mole of CO2 and 1 mole of H2O are formed.
When a mole of aqueous HCl reacts with a mole of aqueous AgNO3,
one mole of aqueous nitric acid is formed, one mole of solid silver
chloride is precipitated, and 15 kilojoules of heat are given off.
Now try going the other way:
In the presence of MnO2 , liquid hydrogen peroxide decomposes into
liquid water and oxygen gas. For every 2 moles of peroxide, 2 moles of
water are formed and one mole of oxygen is formed.
Here is the solution:
Chemical Equation Answer
In the presence of MnO2 , liquid hydrogen peroxide decomposes into
liquid water and oxygen gas. For every 2 moles of peroxide, 2 moles of
water are formed and one mole of oxygen is formed.
6/Day 4
o Balancing them
Balancing
One obvious question at this point is: Where do the stoichiometric
coefficients come from?
Answer: Conservation of mass. If mass is conserved and if atoms don't
change their identity (and they don't), then you must have the same
number of atoms on each side of the equation.
Remember this one?
Notice that there are
6 O's on the left
hand side
6 O's on the right hand side
2 K's on the left
hand side
2 K's on the right hand side
2 Cl's on the left
hand side
2 Cl's on the right hand side
This is what is crucial in balancing an equation...that there are equal
numbers of each atom on each side.
Chemical Equation Answer
In the presence of MnO2 , liquid hydrogen peroxide decomposes into
liquid water and oxygen gas. For every 2 moles of peroxide, 2 moles of
water are formed and one mole of oxygen is formed.
7/Day 4
o Some examples of balancing equations
Balancing Examples
Here are some more examples:
Step 1 balance the carbons.
Step 2 balance the hydrogens.
We have a problem here... we have 7 oxygens on the right hand side,
but O2 is diatomic, meaning that we must always have an even
number of oxygens on the left hand side. The way out is to multiply
through by 2.
Now the right hand side has 8+6=14 oxygens and so we do the final
step.
Now this works.
More Examples
8/Day 4
•
This reaction is already balanced, that happens
sometimes.
You try these:
Balancing Answers
Remember...you never change the subscripts to balance an equation.
9/Day 4
You only change the stoichiometric coefficients.

Moles to moles calculations
When we write a balanced equation...for example:
We are saying that for every 2 moles of H2O2 that decompose, we get
2 moles of H2O and 1 mole of O2.
But suppose we didn't start out with 2 moles of H2O2. Suppose we
started out with 1 mole of H2O2.
I hope you see that we would get 1 mole of H2O and 1/2 mole of O2.
Halving the starting material gives half the products.
Calcs 2
Let's look more deeply at the equation from the previous page:
This equation actually defines a series of relationships as shown below:
10/Day 4
R1 says that for every mole of H2O2 that decomposes, 1 mole of
H2O(l) is formed.
R2 says that, for every mole of H2O2 that decomposes, 1/2 mole of O2
is formed.
R3 says that, for every mole of H2O that is formed, 1/2 mole of O2 is
formed.

Example Problems
So let's look at a problem
If we decompose H2O2 according to the equation:
and we consume 5 moles of H2O2 in the process, how many moles of
O2 and H2O are formed?
11/Day 4
So the general procedure is to use the stoichiometric coefficients to
make your "stoichiometric ratios" you then use the ratio that makes
sense dimensionally.
Let's try another
If I use 7 moles of BaS...
(1) How many moles of H3PO4 are required?
(2) How many moles of Ba3PO4 do I get?
(3) How many moles of H2S do I get?
1.
2.
3.
Now you try one-
If I start with 5 moles of CuFeS2, how many moles of copper, iron (II)
oxide, and sulfur dioxide are formed?
12/Day 4
Answer

Mass to mass calculations
Mass to Mass
When you are given the mass of a reactant and asked to calculate the
mass of product produced you simply build on the mole to mole
technique you just learned.
(A) Make sure you have a balanced equation.
(B) Calculate the number of moles of the compound whose
mass you know.
(C) Use the stoichiometric ratios to calculate the number of
moles of the "unknown" compound.
(D) Convert "the moles" of your answer in letter "C" to
grams.
13/Day 4
Let's try an example!
Mass to Mass Example
How many grams of water and carbon dioxide are produced from the
combustion of 56 grams of methane.
•
•
Balance
the
equatio
n
Convert
grams
of CH4
to
moles
of CH4
Use the
correct
stoichoimetric
ratios
Day 4, Page 14
14/Day 4
•
Moles
product
to
grams
product

Example problems
Try This Mass to Mass Example -
If we start with 109.5g of HCl, how much MgCl2 do we get? How much
H2O do we get?
Mass to Mass Answer
15/Day 4

Meaning of limiting reagents
Theoretical Yield
The quantities we have been calculating so far assume that everything
goes perfectly...that there are no spills, side reactions or other
unexpected problems...that, in effect, we get the maximum possible
product from our reactants. We have therefore been calculating what
is known as theoretical yield, the best yield from a reaction that is
possible.
16/Day 4
o Electrolytes
Reactions in Solution: The Idea of an Electrolyte
Electrolytes
•
•
•
•
An electrolyte is something which, when dissolved in
water, causes the resulting solution to conduct electricity.
Since the carriers of electricity in solution are ions, then an
electrolyte must produce some ions in water.
Water itself is so weakly ionized that almost no current can flow
through it. It is generally considered a non-electrolyte. Sugar
(sucrose) being a molecular compound does not ionize in water
and is also a non-electrolyte.
NaCl, BaCl2, NaOH, AlCl3...almost any combination of metal and
non metal forming an ionic compound will ionize completely on
entering solution in water and is a strong electrolyte.
Strong acids and strong bases are always strong
electrolytes.
•
Reactions in Solution: The Idea of an Electrolyte Part Le
Deux
Weak Electrolytes
•
•
Some compounds only partially ionize. Only a few percent (or
less) of the molecules actually form ions at a given time.
Solutions of such compounds conduct electricity, but not as
effectively as solutions of strong electrolytes.
These weakly ionized compounds are referred to as weak
electrolytes. Many weak acids such as acetic acid, citric acid
and lactic acid are weak electrolytes.

Solubilities
Reactions in Solution: Solubility
Compounds containing the following anions are soluble:
•
•
•
NO3 - nitrate
CH3COO - acetate
ClO3 - chlorate
17/Day 4
All halides are soluble except:
•
•
•
AgX, Hg2X2, PbX2, Hgx2
(note HgCl is soluble however)
"X" often is used to represent a "generic" halide ion, i.e., X = F-,
Cl-, Br-, or I-
All sulfates are soluble except sulfates of:
•
•
Ca2+, Sr2+, Ba2+, Ag+, Hg22+, Pb2+
Thus Na2SO4 would be soluble but CaSO4 would not.
Sulfides are insoluble unless they are sulfides of:
•
•
•
alkali metals
alkaline earth metals
ammonium ions
Carbonates are insoluble except for:
•
•
(NH4)2 CO3
alkali metal carbonates
Ditto for phosphates
Hydroxides are insoluble except for:
•
•
alkali metal hydroxides
hydroxides of Ba2+, Sr2+, Ca2+
o Kinds of reactions
Many reactions can be classified as one of the following 5 types (with
examples).
(1) Combination
18/Day 4
(2) Decomposition
(3) Single Displacement (or replacement)
(4) Double Displacement (or metathesis)
19/Day 4
(5) Combustion
o Driving forces behind reactions
Driving Forces
The circumstances that "drive reactions to completion".
(1) Formation of a precipitate
•
Frequently seen in aqueous ionic metathesis reactions...by
removing product from the solution it depletes the mixture of its
constituent ions.
(2) Formation of a strong electrolyte
•
Not a driving force - If the ions are still freely floating around
in the solution, no identifiable product has been formed.
•
Each side of the equation represents an equimola solutions of
Na+, Cl-, K+, and NO3- ions. So where is the reaction?
(3) Formation of a gas
•
A gaseous substance can leave the site of the reaction simply by
diffusing into the surrounding atmosphere.
20/Day 4
(4) Formation of a weak or non-electrolyte
(5) Making and breaking covalent bonds
Any combustion reaction will also be driven by this force.
(6) Oxidation/Reduction
•
•
Oxidation is an increase in positive charge; Reduction is a
decrease in positive charge. They always occur together.
e.g. in 2nd reaction
2+
o Fe goes to Fe
- oxidation
2+
o Cu
goes to Cu - reduction
o
Aqueous metathesis and driving forces
Ionic Equations
Metathesis and displacement reactions, especially those occurring in
aqueous solutions (which is most of them), can be better understood
using the ideas of spectator ions and net ionic equations.
Consider a typical acid-base reaction:
21/Day 4
Recognizing that this is really a reaction involving dissolved ions, we
can rewrite this as a full ionic equation.
Note that Na+ and Cl- don't really do anything. They are the same on
the right hand side as they are on the left hand side. They watch but
do not participate and we therefore call them spectator ions.
In a net ionic equation we eliminate the spectator ions:
This net ionic reaction is special. It represents the net ionic equation
for any acid-base neutralization reaction, and vividly illustrates the role
of the formation of a nonelectrolyte as a driving force in these kind of
reactions. Note how the nonionic water molecule has, in effect,
scavenged the hydrogen and hydroxide ions that formed it, rendering
them unavailable for any further participation in ionic reactions.

Standard Chemical Equation
Ionic Equations: Part II
The idea of the net ionic equation also helps in understanding
displacement reactions in aqueous solution.
Chemical Equation:
Full ionic equation:
Net ionic equation:
22/Day 4

Full Ionic Equation
Ionic Equations: Practice Problems
Write the full Ionic equations and net ionic equations for the following.
Identify any spectator ions.
Ionic 4 Answers
1.) Full:
•
All ions are spectator ions...no driving force…reaction does not
go to completion and therefore we cannot write a net ionic
equation--or, more precisely, the net ionic equation is identical
to the full ionic equation.
2.) Full:
•
K+ and NO3- are spectator ions.
Net Ionic:
23/Day 4
3.) Full:
•
K+ and PO43- are spectator ions.
Net:
4.) Full:
•
NO3- is a spectator ion.
Net:
o Oxidation reduction
Oxidation Reduction
Basically, oxidation and reduction occur when one substance
transfers electrons to another. The formal definition of oxidation
and reduction requires us to study something called "oxidation
numbers" which we will cover in DAY 6. For now, the working
definitions shown below will suffice.
An atom is said to be
oxidized, during a reaction,
when it
•
•
•
gains positive charge
bonds to additional oxygen
atoms
detaches from hydrogen
atoms.
An atom is said to be
reduced, during a reaction,
when it
•
•
•
loses positive charge
bonds to additional
hydrogen atoms
detaches from oxygen
atoms.
24/Day 4
Note that the definitions pertaining to loss or gain of hydrogen and
oxygen apply to covalently bonded compounds only.
One thing to note: oxidation and reduction are processes and not
states. When we say that a substance is oxidized, we mean that it is
undergoing, or has undergone, a process in which it gained positive
charge, lost hydrogen, or gained oxygen.
A couple of examples:
Carbon is oxidized because it loses H atoms and acquires 0 atoms.
Here is an example of an ionic reaction:
Iron is oxidized because its charge increases from 0 to +2.
Note that copper is reduced (from +2 to 0).
This is a good illustration of the fact that oxidation and reduction
always occur together.
Oxidizing Agents/Reducing Agents
A couple more definitions/explanations, and we are done:
•
•
•
When a substance causes oxidation it is said to be an oxidizing
agent.
The oxidizing agent is, itself reduced.
The reducing agent is, itself, oxidized.
Remember the reactions we examined on the previous page?
25/Day 4
For these examples:
•
•
•
Iron is oxidized and is the reducing agent because it causes
copper to be reduced.
Cu2+ is reduced and is the oxidizing agent because it causes iron
to be oxidized.
Likewise, methane is the reducing agent, and oxygen the
oxidizing agent in the first reaction.

Idea of activity
Activity Series
The more metallic an element, the more readily it becomes positively
charged, i.e. the more easily it is oxidized.
The Activity Series rates metals on the basis of their ease of
oxidation - the most easily oxidized (the most active) at the top (see
column below)
Note how the very reactive alkali and alkaline earth metals are at the
top of the series.
On the other hand the metals used in coinage, the ones that resist
corrosion (i.e. oxidation) the best are at the bottom.
Li
Fe
K
Co
Ba
Ni
Ca
Sn
Na
Pb
Mg
Cu
Al
Ag
Mn
Hg
Zn
Pt
Cr
Au
26/Day 4