Introduction_to_Chemical_Reactions_2011-2012.ppt

Introduction to Chemical
Reactions

What is a Chemical Reaction?
It is a chemical change in which one or
more substances are destroyed and one
or more new substances are created.
BEFORE
H2 gas
and
O2 gas
AFTER
H2O liquid

Parts of a Chemical Reaction
Reactants  Products
Reactants: Substances that are destroyed by the
chemical change (bonds break
break).
Products: Substances created by the chemical
change (new bonds form
form).
The arrow () is read as “yields”.

Other symbols in chemical
reactions
• (s) = solid
• (l) = liquid
• (g) = gas
• (aq) = aqueous solution (the substance is
dissolved in H2O)
• “+” separates two or more reactants or
products
• “” yield sign separates reactants from
products

Evidence for a Chemical Reaction
1) Evolution of light or heat.

2) Temperature change (increase or
decrease) to the surroundings.

3) Formation of a gas (bubbling or an odor)
other than boiling.

4) Color change (due to the formation of a
new substance).

5) Formation of a precipitate (a new solid
forms) from the reaction of two aqueous
solutions.

Word Equations
• Statements that indicate the reactants and
products in a chemical reaction.
• Ex. Iron (s) + chlorine (g)  iron (III) chloride (s)
• This is read as:
“Solid iron and chlorine gas react (combine) to produce
solid iron (III) chloride”

Translating Word Equations to
Skeleton Equations
• A skeleton equation uses chemical formulas
rather than words to identify the reactants and
products of a chemical reaction.
• The word equation
Iron (s) + chlorine (g)  iron (III) chloride (s)
• The skeleton equation
Fe(s) + Cl2(g)  FeCl3 (s)
A skeleton equation is not yet “balanced” by coefficients!

One more example…
• 6 Na (s) + Fe2O3 (s)  3 Na2O (s) + 2 Fe (s)
– The numbers preceding the chemical formulae are
coefficients. They are used to balance the reaction.
– The numbers within the chemical formulae are
subscripts.
– You can read the above balanced reaction as:
• “6 atoms of solid sodium plus 1 formula unit of solid
iron (III) oxide yields 3 formula units of solid sodium
oxide and 2 atoms of solid iron” or…
• “6 moles of solid sodium plus 1 mole of solid iron (III)
oxide yields 3 moles of solid sodium oxide plus 2
moles of solid iron”
• Chemical reactions can never be read in terms of
grams, only in terms of particles or groups of particles
(moles).

Conservation of Mass
During a chemical reaction, atoms are neither
created nor destroyed (Conservation of
Mass).
Hydrogen and oxygen gas react to form
water:
H2 (g) + O2 (g)  H2O (l)

Conservation of Mass
H2 (g) + O2 (g)  H2O (l)
What is wrong with this equation above? Doesn’t
it appear that one oxygen atom “went missing”?
According to conservation of mass, the proper way
to write this reaction is:
2H2 (g) + 1O2 (g)  2H2O (l)
The red coefficients represent the # of molecules
(or the # of moles) of each reactant or product.

Not All Properties are Conserved
During Chemical Reactions!
CONSERVED NOT CONSERVED
Mass
Types of atoms
Number of each atom
Color
Physical state (solid,
liquid, gas)
Volume
Number of moles of
reactants/products

There are 5 basic types….
• Single Replacement (Displacement)
(Redox)
• Double Replacement (Displacement)
(Metathesis)
• Synthesis (Combination)
• Decomposition
• Combustion

A single uncombined
element replaces
another element in
an ionic compound.
There are two
reactants and two
products.
1) SINGLE REPLACEMENT
REACTION
Ex: Zn + CuSO4  ZnSO4 + Cu

Single Replacement Reactions
Single replacement reactions have the
general form, A + BC  AC + B.
Question: Do all single replacement
reactions actually occur?
Answer: Not necessarily…

Examine the reaction:
Zn + CuSO4  ZnSO4 + Cu
This reaction does occur!’
Now let’s try:
Cu + ZnSO4  No Reaction
Conclusion: Zn will replace Cu in
solution, but not vice versa!

How do we know which reactions will occur
and which ones will not?
We look at the “activity series”.
Elements with higher activities replace
elements with lower activities during a
single-replacement reaction, but not vice-
versa.

HIGHEST ACTIVITY
Li
Rb
K
Ba
Ca
Na
Mg
Al
Mn
Zn
Cr
Fe
Ni
Sn
Pb
H
Cu
Hg
Ag
Pt
Au
LOWEST ACTIVITY
Activity Series for
Metals

Activity Series for Nonmetals
Highest Activity
F
Cl
Br
I
Lowest Activity

Predicting the Products of Single
Replacement Reactions
1) Write the reactants.
2) Identify the cation and anion of the reactant
that is a compound.
3) Use the activity series to see if the single
element will replace one of the elements in
the compound. If no reaction will occur,
just write “NR” for the products and you
are done.
4) Identify the reactant that is the element.
Determine its charge when it becomes an
ion.
5) Perform criss-cross to predict the new
compound on the products side of the
reaction.
6) Write both new products.
7) Balance the reaction.

Single Replacement Between
Metals and Water
• Some metals have a higher activity than hydrogen
and can replace it in a single replacement
reaction. In these reactions, you may think of
water (H2O) as H(OH).
• Ex: Na + H2O  ?
Na + HOH  ?
Na + H+
OH-
 Na+
OH-
+ H
2Na + 2H2O  2NaOH + H2

Parts of two
aqueous ionic
compounds switch
places to form two
new compounds.
There are two
reactants and two
products.
2) DOUBLE REPLACEMENT
REACTION
Example:
AgNO3 + NaCl 
AgCl + NaNO3

Double Replacement Reactions
The general form of a double replacement reaction is:
AB + CD  AD + CB
Just like single replacement reactions, not all double
replacement reactions actually occur.
We can experimentally attempt a D.R. reaction. The
reaction occurs if:
1) A solid precipitate is produced, or
2) A gas is produced, or
3) Water is produced.
If none of the above are produced and both products are
(aq), then there is no reaction (NR)!

Examples of Double Replacement
Reactions:
Pb(NO3)2 (aq) + 2NaI (aq)  PbI2 (s) + 2NaNO3 (aq)
(precipitate forming)
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
(water-forming, acid-base, neutralization)
CaCO3 (s) + 2HCl (aq)  CaCl2 (aq) + H2CO3
(gas-forming)
H(OH)
H2O (l) + CO2 (g)

How do you determine if one of the products
of a double replacement reaction will be a
precipitate?
• Use the solubility rules….
Soluble compounds
These compounds break down when put in water.
Example: In water, NaCl  Na1+
and Cl1-
.
We say that NaCl…
 has dissolved.
 is soluble.
 forms an aqueous solution (aq).

The Solubility Rules
Insoluble compounds
These compounds do NOT
break down when put in
water.
Example: In water, CaCO3
does NOT break down
into Ca2+
and CO3
2-
ions.
The CaCO3 stays as a solid,
(s) or (ppt).
This is fortunate for many
sea-creatures!
Seashells are made of CaCO3!

The Solubility Rules
You do not have to memorize these rules,
but you do have to know how to use them
to determine if a product is a precipitate.
See the chart on the next slide…..
Let’s check NaCl and CaCO3… Are these
compounds soluble or insoluble in
aqueous solution?

Predicting the Products of Double
Replacement Reactions…
Step Example
1) Write the two reactants (both are ionic
compounds)
2) Identify the cations and anions in both of the
compound reactants
3) Pair up each cation with the anion from the
other compound
(i.e. – switch the cations)
4) Write the formula for each product using the
criss-cross method
5) Write the complete equation for the double
replacement reaction
6) Balance the equation.
7) Use the solubility rules chart to figure out which
product is a precipitate (s) and which product
is an aqueous solution (aq). If both products
are (aq) it is really not a reaction.

Two or more simple substances
(the reactants) combine to form
a more complex substance (the
product).
3) SYNTHESIS REACTION
Ex: 2Mg + O2
 2MgO

SYNTHESIS REACTION
Types of synthesis:
a)Element A + Element B Compound
Na(s) + Cl2 (g)  2NaCl(s)
a)Element + Compound A  Compound B
O2(g) + 2SO2(g)  2SO3(g)
a)Compound A + Compound B  Compound C
CaO(s) + H2O(l)  Ca(OH)2 (s)

Synthesis Reactions (cont’d)
• Metallic and nonmetallic elements react to form ionic
compounds. The resultant compound should be charge
balanced by the criss-cross method.
Ex. 4Li + O2  2Li2O
• Nonmetals react with each other to form covalent (molecular)
compounds. You should be able to draw a valid Lewis
Structure for the product.
2H2 + O2  2H2O
or
H2 + O2  H2O2
But NOT
H2 + O2  2OH

A more complex substance (the
reactant) breaks down into two
or more simple parts (products).
Synthesis and decomposition
reactions are opposites.
4) DECOMPOSITION REACTION
Ex: 2H2O  2H2 + O2
Electrolysis of
Water

DECOMPOSITION REACTIONS
(Cont’d)
Decomposition of a compound produces two or
more elements and/or compounds
The products are always simpler than the
reactant.
Gases are often produced (H2, N2, O2, CO2, etc.)
in the decomposition of covalent compounds.
Ionic compounds may be decomposed into pure
elements by using electricity (electrolysis). This is
how pure metals are obtained from salts.

The Decomposition of Water by
Electrolysis
2H2O  2H2 + O2
An electrical
current can be
used to chemically
separate water into
oxygen gas and
hydrogen gas.
Notice that twice
as much hydrogen
is produced
compared to
oxygen!

Electrolysis of Molten Sodium
Chloride Many pure metals are
obtained by using
electrolysis to separate
metallic salts (ex. NaCl
is used to obtain pure
Na).

5) COMBUSTION REACTIONS
a) All involve oxygen (O2) as a reactant,
combining with another substance
b) All combustion reactions are are
exothermic
c) Complete combustion of a
hydrocarbon always produces CO2
and H2O
d) Incomplete combustion of a
hydrocarbon will produce CO and
possibly C (black carbon soot) as well
Ex: CH4 + 2O2 => CO2 + 2H2O (complete combustion – blue flame)
Ex: CH4 + 1.5O2 => CO + 2H2O (incomplete combustion – yellow flame)
Ex: CH4 + O2 => C + 2H2O (incomplete combustion – yellow flame, soot)

Combustion (cont’d)
• Any synthesis reaction which involves O2 as a
reactant is also considered to be a combustion
reaction!
Ex. 2Mg + O2  2MgO
(metal oxide)
This is called the combustion of magnesium or
the synthesis of magnesium oxide. The
combustion of a metal always produces a metal
oxide (in this case, magnesium oxide). Make
sure the metal product is criss-crossed
correctly!

TRY TO CLASSIFY THESE:
1) C4H8 + 6O2  4CO2 + 4H2O
2) HCl + NaOH  H2O + NaCl
3) 2KNO3(s)  2KNO2(s) + O2(g)

TRY TO CLASSIFY THESE:
4) 2Ag + S  Ag2S
5) MgCO3(s)  MgO(s) + CO2(g)
6) Cl2 + 2KBr  2KCl + Br2

Check Your Answers…
1) Combustion (of a hydrocarbon)
2) Double replacement (water forming)
3) Decomposition
4) Synthesis
5) Decomposition
6) Single Replacement

Counting Atoms
SnO2 + 2H2 → Sn + 2H2O
SUBSCRIPT COEFFICIENT

Rules for Counting Atoms
1)Coefficients propagate to the right through the
entire compound, whether or not parentheses
are present.
2) Subscripts affect only the element to the left of
the subscript, unless…
3) If a subscript occurs to the right of a
parentheses, the subscript propagates to the left
through the parentheses.
4) When a coefficient and subscript “meet”, you
must multiply the two.

Examples of Counting Atoms
SnO2 + 2H2 → Sn + 2H2O
2 C4H10 + 13 O2 → 8 CO2 + 10 H2O
Cu + 2AgNO3 → Cu(NO3)2 + 2Ag
3Pb(NO3)2 + 2AlCl3 → 3PbCl2 + 2Al(NO3)3

Classwork
Complete “The Count” worksheet
on counting atoms in chemical
reactions.

Warm-Up
2Ca3(PO4)2 + 6 SiO2 + 10C 
6 CaSiO3 + P4 +10CO
Atom # Atoms on
Left Side
# Atoms on
Right Side
Ca
P
O
Si
C

Rules for Balancing
Chemical Reactions
__H2 + __ O2  __H2O
Balancing is about finding the
right coefficients!

Rules for Balancing
Chemical Reactions
1) You can change the coefficients, but
NEVER the subscripts!
__H2 + __ O2  __H2O
Off Limits!

Rules for Balancing
Chemical Reactions
2) The coefficients must reduced to
represent the lowest possible numbers.
4H2 + 2 O2  4H2O

Rules for Balancing
Chemical Reactions
3) It is OK to use fraction coefficients, but
you must get rid of them in the end
(multiply through by denominator).
H2 + ½ O2  H2O

Rules for Balancing
Chemical Reactions
4) Often, it is helpful to save the following
elements until the end (do other
elements first):
H, C, O

Rules for Balancing
Chemical Reactions
5) Do a final balance check for each
element!
2H2 + O2  2H2O

Practice
1) K + Br  KBr
2) HgO  Hg + O2
3) Na + H2O  NaOH + H2

Practice
4) CaO + H2O  Ca(OH)2
5) Al + HCl  AlCl3 + H2

Energy Changes Accompanying
Chemical Reactions
All chemical reactions involve a net release or absorption of
energy. Therefore, heat energy moves between the
chemical system and the surroundings. This exchange of
heat can be monitored by keeping track of changes in
temperature of the surroundings (calorimetry).
Remember, q = mcpT
where  q = change in heat (in Joules)
m = mass of H2O (in grams)
cp = specific heat capacity of
H2O (J/g ◦C )
T = change in temperature
of H2O (in ◦C)

Where does the energy come from
during a chemical reaction?
• During chemical reactions, bonds are broken and new bonds are
formed.
• The heat energy that moves between the system and
surroundings during chemical reactions is basically the energy that
is used to break bonds and the energy that is released when
bonds form. (i.e. bond energy)
• The energy change that accompanies any chemical reaction is
called the enthalpy (heat) of reaction or H0
rxn.
H0
rxn = Hfinal – Hinitial
• H0
simply means that the energy changes during chemical
reactions are generally measured at “standard state” conditions of
298 K (25◦
C) and 1 atm pressure.
• It is important to note that absolute amounts of energy within
a chemical system cannot be measured. We can only
measure changes in energy within a chemical system. Hence
we use the “” sign.

Exothermic Reactions
A chemical reaction is exothermic if energy is given off by the system to
the surroundings (the energy exits):
Reactants  Products + Energy Released
The temperature of the surroundings (including the temperature probe)
increases during exothermic reactions because the system releases
energy. The H0
rxn is negative because Hfinal is less than Hinitial. In other
words, the system lost energy. (sign goes with the system)
The majority of chemical reactions are exothermic because nature favors a
low chemical potential energy.
System Surroundings
reactants
products
Chemical
Potential
Energy
(H)
Reaction progress
Hrxn is (-)

Example: An Exothermic Reaction
The “Smashing” Thermite Reaction:
2Al(s) + Fe2O3 (s)  2Fe (s) + Al2O3 (s)
Reaction Progress
Chemical
Potential
Energy
(H)

Endothermic Reactions
A chemical reaction is endothermic if energy is absorbed by the system
from the surroundings (the energy enters):
Reactants + Energy Absorbed  Products
The temperature of the surroundings (including the temperature probe)
decreases during endothermic reactions because the system absorbs
energy. The H0
rxn is positive because Hfinal is more than Hinitial. In other
words, the system gained energy. (sign goes with the system)
Endothermic chemical reactions are generally unfavorable but may occur
only if they are accompanied by an increase in entropy or disorder of the
system (due to more particles formed, liquids/gases formed, mixtures
formed, volume of gas increases).
System Surroundings
reactants
products
Chemical
Potential
Energy
(H)
Reaction progress
Hrxn is (+)

Example: An Endothermic Reaction
Ba(OH)28H2O (s) + 2NH4(NO3) (s) 
Ba(NO3)2 (aq) + 2NH3 (g) + 10 H2O (l)
Reaction Progress
Chemical
Potential
Energy
(H)

Do you have to actually perform and
observe a chemical reaction to know if it is
exothermic or endothermic?
• No – you can calculate H0
rxn from data that has
already been measured and tabulated by
thermo-chemists (see handout).
• H0
f = standard heat of formation for a compound
(in kJ/mol). It is determined by forming the
compound from its elements in their stable forms
at conditions of 298K and 1 atm of pressure
inside of a calorimeter.
• For most compounds, H0
f is negative because
bond formation is exothermic!
• H0
f of an element is always 0 kJ/mol by def.

H0
rxn = nH0
f (products) - nH0
f (reactants)
• Not as hard as it looks 
• Basically, you just
1) multiply the coefficient of each product times its
standard heat of formation and add together for all
products
2) multiply the coefficient of each reactant times its
standard heat of formation and add together for all
reactants
3) take the difference of 1 and 2
(always products - reactants)
4) If the difference is (-) the reaction is exothermic;
if the difference is (+) the reaction is endothermic.

Try this…
• Calculate the H0
rxn for the thermite reaction
using tabulated data (see handout):
2Al (s) + Fe2O3 (s)  2Fe (s) + Al2O3 (s)
H0
rxn = nH0
f (reactants)

Try this…
• Calculate the H0
rxn for this reaction based on tabulated
data:
Ba(OH)28H2O (s) + 2NH4(NO3) (s) 
Ba(NO3)2 (aq) + 2NH3 (g) + 10 H2O (l)
H0
rxn = nH0
f (reactants)
Compound H0
f
(kcal/mol)
NH4(NO3) (s) -87.73
Ba(OH)28H2O (s) -798.8
Ba(NO3)2 (aq) -227.62
NH3 (g) -11.02
H2O (l) -68.32 1 kcal = 4.184 kJ

Summarizing H0
rxn
• If H0
rxn is (-) the reaction is exothermic and the
bonds formed are stronger and more stable
than the bonds broken.
• If H0
rxn is (+) the reaction is endothermic and the
bonds formed are weaker and less stable than
the bonds broken. However, if the entropy of
the system has increased to sufficiently to
counteract this increase in enthalpy, then the
reaction can still occur.

Bond Enthalpies
Bond Enthalpies
• Another way to determining an enthalpy
Another way to determining an enthalpy
change
change (H0
rxn) for a chemical reaction is to
for a chemical reaction is to
compute the difference in bond enthalpies
compute the difference in bond enthalpies
between reactants and products
between reactants and products
• The energy to required to break a covalent
The energy to required to break a covalent
bond in the gaseous phase is called a
bond in the gaseous phase is called a bond
bond
enthalpy (bond dissociation energy).
enthalpy (bond dissociation energy).
• Bond enthalpy tables give the average
Bond enthalpy tables give the average
energy to break a chemical bond. Actually
energy to break a chemical bond. Actually
there are slight variations depending on the
there are slight variations depending on the
environment in which the chemical bond is
environment in which the chemical bond is
located
located

Bond Enthalpy Table
Bond Enthalpy Table
The average bond enthalpies for several types of
The average bond enthalpies for several types of
chemical bonds are shown in the table below:
chemical bonds are shown in the table below:

Bond Enthalpies
Bond Enthalpies
• Bond enthalpies can be used to calculate the
Bond enthalpies can be used to calculate the
enthalpy change (
enthalpy change (
H
H0
0
rxn) for a chemical
) for a chemical
reaction.
reaction.
• Energy is required to break chemical bonds
Energy is required to break chemical bonds.
.
Therefore when a chemical bond is broken its
Therefore when a chemical bond is broken its
enthalpy change
enthalpy change carries a
carries a positive sign
positive sign.
.
• Energy is released when chemical bonds
Energy is released when chemical bonds
form
form. When a chemical bond is formed its
. When a chemical bond is formed its
enthalpy change
enthalpy change is expressed as a
is expressed as a negative
negative
value.
value.
• By combining the enthalpy required and the
By combining the enthalpy required and the
enthalpy released for the breaking and
enthalpy released for the breaking and
forming chemical bonds, one can calculate
forming chemical bonds, one can calculate
the overall enthalpy change for a chemical
the overall enthalpy change for a chemical
reaction.
reaction.

Bond Enthalpy Calculations
Bond Enthalpy Calculations
Example : Calculate the enthalpy change
Example : Calculate the enthalpy change (H0
rxn)
for the reaction N
for the reaction N2 + 3 H
+ 3 H2 
 2 NH
2 NH3
Bonds broken (energy in)
1 N≡N: = 945
3 H-H: 3(435) = 1305
Total = 2250 kJ/mol
Bonds formed (energy out)
2x3 = 6 N-H: 6 (390) = - 2340 kJ/mol
H0
rxn = [energy used for breaking bonds] + [energy released in forming bonds]
Net enthalpy change (H0
rxn)
= + 2250 + (-2340) = - 90 kJ/mol (exothermic reaction)
H - H
H - H
H - H
You may have
to draw a
Lewis
Structure to
know what
type of bonds
are present!

Another Way to Think About It
Chemical
Potential
Energy
(H)
of
System
Start
+ 2250 kJ/mol
(energy in
when bonds
break)
-2340 kJ/mol
(energy out
when bonds
form)
H0
rxn = -90 kJ/mol (net)
released by the system to the
surroundings

Introduction_to_Chemical_Reactions_2011-2012.ppt

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