Transcript types of reactions

CHAPTER 6
CHEMICAL REACTIONS AND EQUATIONS
Chemical reactions are responsible for just about
everything that occurs around us all the time.
How do you know a chemical reaction has taken place
when a cake is baked?
What are some other examples of chemical reactions
in real life?
MACROSCOPIC CHANES INDICATE A MICROSCOPIC CHANGE
OR CHEMICAL REACTION HAS TAKEN PLACE.
INDICATIONS OF A CHEMICAL REACTION
1. Gas is released (bubbles)
2. Light is released or absorbed
3. Odor is formed
4. Color change
5. Heat is released or absorbed
6. Precipitate is formed (solid from mixed liquids)
• During a chemical reaction energy is either released or
absorbed.
• Endothermic reaction: energy is absorbed
• Exothermic reaction: energy is released (explosions)
- seen as heat and or light
CHEMICAL REACTION (same as chemical change)
change of one substance into another substance
• original substance can not be recovered
• proper amounts of starting materials are needed,
and all must be used to make finished product
• all amounts going in = amounts coming out
ex: building a bicycle
BICYCLE MATERIALS
Frame, bolts, screws, nuts, wheels, handlebars, seat,
gears, etc…..
Would the bicycle work properly if there were nuts and
bolts left over when it was built?
In chemical reactions the
LAW OF CONSERVATION OF MASS
holds true
C + O2  CO2
Same numbers of atoms must be
on each side of reaction
MASS IS CONSERVED
activity
CHEMICAL EQUATION
• method of using chemical symbols to represent a chemical
reaction on paper
• describes what goes on in a chemical reaction
• easier than writing out words
Ex:
2 mols. of
hydrogen
2 H2
+
one mol. of
oxygen
+
O2
becomes

2 mols. of
water
2 H2O
PARTS OF A CHEMICAL EQUATION
reactants
“goes to”
product
2 H (g) + O (g) -----------------
2 H2O
“becomes”
coefficient
physical state
subscript
Reactant- chemicals that will react , left side of the arrow
Product- chemicals that are produced, right side of arrow
Physical state- solid, liquid, gas
Arrow- shows direction of reaction and what reactants become
Coefficient- large numbers in front of chemical formulas
represent number of atoms of each element in molecule
Subscript- small number at lower right of chemical symbols
represent number of atoms of that element in molecule
COUNTING ATOMS
Rules
1. Multiply subscript x coefficient
2. If parenthesis is present multiply
subscript of each
x
subscript
each element
outside
inside parenthesis
parenthesis
3. Add total atoms
x
coefficient
COUNTING ATOMS
2 H2O
3 H2SO4
5 C6H12O6
H 2x2=4
O 1x2= 2
total
6
H 2x3=6
S 1x3=3
O 4 x 3 = 12
total
21
C 6 x 5 = 30
H 12 x 5 = 60
O 6 x 5 = 30
total
120
Mg(OH)2
Mg 1 x 1 =
O 1x2x1=
H 1x2x1=
total
1
2
2
5
3 Na(PO4)3
Na 1 x 3 =
3
P 1x3x3= 9
O 4 x 3 x 3 = 36
total
48
BALANCING EQUATIONS
When a chemical reaction occurs between reactants there
must always be equal number of atoms of reactants and
products.
WHAT LAW DOES THIS DEMONSTRATE?
REACTANT S
NUMBER
OF ATOMS
PRODUCTS
=
NUMBER
OF ATOMS
BALANCING EQUATIONS
Rules
1. Write all the reactants on the left side of arrow, products on
right side of arrow
2. NEVER change subscripts (this will change the reactant
substance), ONLY change coefficients
Ex: H2O water
H2O2 hydrogen peroxide
3. Find numbers of atoms of each element
Ex: 2 H2O
O=2
H=4
BALANCING EQUATIONS
Rules, cont.
4. Make sure to add up all sources of same element on each side of arrow
ex:
2 H2O + C2H4O6  do same with products
C 2
= C2
H 4+4= H8
O 2+6= O8
5. Balance equation so that numbers of elements are equal on each side of
the arrow (conservation of mass)
6. Simplify (reduce) equation to lowest terms (find common denominator)
ex: 4 NH3 + 2 O2  4 NO3 + 6 H2O (all divisible by 2, so simplify)
2 NH3 + O2  2 NO3 + 3 H2O
STEPS FOR BALANCING EQUATIONS
H2 + O2  H2O
1. List out the atoms on each side of equation
H
H
O
O
2. Count number of each type of atom in reactants
H
O
2
2
3. Count number of each type of atom in products
H
O
2
1
4. Note which atoms are not balanced
H2 + O2  H2O
H 2
2
O 1
H
O 2
5. Select one atom to balance
- easiest to start with atom which is by itself on one
side
- order: metal, nonmetal, save H & O last
H2O + O2  2 H2O
6. update atom counts on each side of equation
H 2
H 4
O 2
O
1
7. continue to update atom counts until both sides have
equal numbers of atoms of each element
2H2O + O2  2 H2O
BALANCED EQUATION!!!
8. Reduce coefficients if possible
4 NH3 + 12 O2 
4 NO3 + 6 H2O
ALL COEFFICIENTS ARE DIVISIBLE BY 2 SO SIMPLIFY
2 NH3 + 6 O2  2 NO3 + 3 H2O
Al + Fe2O3 
Al
1
Fe
2
O
3
Al2O3 + Fe
Al
2 unbalanced
Fe 1 unbalanced
Fe 1
2 Al + Fe2O3  Al2O3 + Fe
Al
Fe
O
2
2
3
Al
Fe
O
2
1 unbalanced
3
2 Al + Fe2O3  Al2O3 + 2 Fe
ALL ATOMS EQUAL ON BOTH SIDES
EQUATION IS BALANCED !
Mg Cl2 + K2S  MgS + KCl
Mg 1
Mg 1
Cl
2
Cl
1 unbalanced
K
2
Cl
1 unbalanced
S
1
S
1
Mg Cl2 + K2S  MgS + 2 KCl
Mg 1
Mg 1
Cl
2
Cl
2
K
2
K
2
S
1
S
1
ATOMS ARE EQUAL ON BOTH SIDES
EQUATION IS BALANCED !
P4O10 + H2O 
P
4
H
2
O
11
P4O10 + H2O 
P
4
H
2
O
11
P4O10 +
P
H
O
6 H2O 
4
12
16
H2PO4
P
1
H
3
O
4
4 H2PO4
P
H
O
4
12
16
4 H2PO4
P
H
O
4
12
16
PbO2
PbO2
PbO2
+
Pb
Cl
H
O
+
Pb
Cl
H
O
HCl
1
1
1
2

PbCl2
+
Pb
Cl
H
O
Cl2 +
1
4
2
1
H2O
4 HCl  PbCl2 + Cl2 + H2O
1
4
4
2
+ 4 HCl
Pb
1
Cl
4
H
4
O
2

PbCl2
Pb
Cl
H
O
1
4
2
1
+
Pb
Cl
H
O
Cl2 +
1
4
4
2
2 H2O
Example of Parenthesis
(NH4)2 Cr2 O7
N
2
Cr
2
H
8
O
7
(NH4)2 Cr2 O7
N
2
Cr
2
H
8
O
7
+
+


Cr2O3
N
Cr
H
O
+
N2
+
H2O
2
2
2
4
Cr2O3 + N2
N
2
Cr
2
H
8
O
7
EQUATION IS BALANCED !
+
4 H2O
Practice Problems
TYPES OF REACTIONS
Reactions are classified into types because it makes it easier to predict
what will happen during and at the results of that reaction.
Types of Reactions
Synthesis
two or more substances or combine
to form a new single product
- most exothermic
a
+ b

ab
2H2 (g) + O2 (g)  2H2O (g)
element + element = new cpd
4 Fe (s) + 3 O (g)  2 FeO (rust)
cpd
+ cpd = new cpd
CO2 (g) + H2 (g)  H2CO3
element + cpd = new cpd
O2 (g) + CO (g) 
CO2 (g)
Types of Reactions
Decomposition:
compound breaks down into two or
more simpler substances
ab 
a +
b
2H2O (l)  2H2 (g) + O2
NH4NO3 (s)  N2O (g) + 2 H2O
- most endothermic
Types of Reactions
Single displacement:
-
one element takes the place of another in
a compound
element can replace first or second part of
a compound
more reactive element will replace less
reactive element
-
a + bc  ac + b
or
d + bc  bd + c
Zn
+ Cu(NO3)2  Zn(NO3)2 + Cu
ele
Fe
cpd
+
ele
Cl2
CuSO4 
cpd
+
2NaBr
new cpd
FeSO4
new cpd
 2NaCl
new ele
+
Cu
ele
+ Br2
Real Life Example of Single Displacement Reaction
Salicylic acid (C7H6O3)

Aspirin (C9H8O4)
Types of Reactions
Double displacement
positive portions of 2 ionic
compounds are interchanged
ab + cd  ad + cb
AgNo2 + NaCl  AgCl + NaNO3
cpd
cpd
cpd
cpd
Pb(NO3)2 + 2 KCL  PbI2 + 2 KNO2
Types of Reactions
Combustion:
- Compound usually containing
carbon rapidly combines with
oxygen to
form one or more oxides
(compound containing oxygen)
- called burning
- heat is produced
ab
CH4
+
O2  oxide of a + oxide of b
+ 2O2
 CO2 + 2H2O
2C2H2 + 5O2  4CO2
+ 2H2O + energy
Most reactions fall into these five categories,
and some may fall into more than one .
NH4NO3 (s)  N2O (g) + 2 H2O
Combustion- two new products contain oxygen (oxides)
Decomposition- one reactant is broken down into two
simpler
substances
NATURE OF REACTIONS
Reversible reaction
• reaction that changes direction based on energy flow
Moves right: more product is made
Moves left: more reactants are formed
• shown with double arrow each way or one arrow pointing both directions

•
•


reactants do not become used up because as new product is formed
other new products break down and supply new reactants
reactions will try to reach a state of equilibrium
equilibrium – state where no net change in amount of reactants
or products (a stable state)
Equilibrium Example
equilibrium does not mean there are equal numbers of products and
reactants, amount of material started with = amount of material in end
Equilibrium Example
Rechargeable Batteries
Dynamic Equilibrium:
state where 2 exactly opposite chemical reactions are
occurring at the same place, same time, and same rate
(speed) where reactions continuously occur
•Two opposing forces are being exerted but they are in a
state of balance
•Amounts of all chemical entities are constant but do not
have to be the same
ex: tug of war where both teams are pulling eqally
against each other
Dynamic equilibrium example:
• person rowing boat upstream at exactly same speed as current
- if current speeds up, he has to speed up
- if current slows, he has to slow down to stay in same
place
• NET RESULT: boat has NO net movement even though person
is still rowing and stream is still flowing
Equilibrium does not mean there are equal
numbers of reactants and products
It means:
amount of material = amount of material
started with
in end
CaCO3  CaO + CO2
Double arrow shows that reaction can go in either direction
Process of Dynamic Equilibrium
1. at first reactions goes to right more quickly until more of the
reactants are used up
2. as more product is formed , the reverse reaction starts to
occur more frequently
3. eventually reaction moves at same speed in either direction
4. EQUILIBRIUM has been achieved
• DYNAMIC because molecules are always moving back
and forth
• reactions try to stay at equilibrium because this is a
stable state
What happens if you add stress to a reaction?
• If one side pulls harder there is more stress
• Other side has to react and pull harder in order to keep
stability (equilibrium)
** same is true of chemical reactions **
Le Chateliers Principle
If a change of condition (stress) is applied to a chemical
system, the reaction will return to equilibrium by shifting
to counteract the stress
3 stresses
1. change concentration of reactants or products
2. change temperature
3. change pressure in system containing gases
FACTORS THAT AFFECT DIRECTION OF REACTION
1. addition or removal of reactants or products
• addition of reactants/ removal of products:
 pushes reaction to right
•
addition of products/ removal of reactants:
 pushes reaction to left
FACTORS THAT AFFECT DIRECTION OF REACTION
2. addition or removal of energy
- addition: endothermic reactions absorb energy so more
product made (reaction moves to right)
- removal: exothermic reactions give off energy to make
more reactant (reaction moves to left)
In order for any reaction to occur energy is needed.
Activation energy: amount of energy needed to make atoms
collide
ex: H and O
both highly reactive elements
can coexist in the same container for years if just sitting quietly
add spark (energy) and they will cause an explosion
Hindenberg- spark ignited H with O and exploded
Space shuttle- liquid O is fine to go into space, but
exploded when heat energy was added
due to loss of heat shield shingles
Reaction rate: how fast a reaction occurs
Rate of reaction = amt substance changed/ amount of time
or
amt of substance made in an amt of time
Importance of Reaction Rate
• Industry
• Kitchen
• Body processes
• Nature
FACTORS THAT AFFECT REACTION RATE
1. Temperature
2. Concentration of reactants
3. Catalysts
4. Inhibitors
FACTORS THAT AFFECT REACTION RATE
1. Temperature
temperature
temperature
rate
rate
2. Concentration
concentration
concentration
rate
rate
** Higher amounts of substance and temperature
will increase numbers of collisions **
limiting reactant: reactant which is depleated first
(not the one with the lesser amount)
• concentration of reactants is important because if
we run out of one of the reactants it can limit and stop
the whole reaction
ex: How many smores can be made with the following?
6 graham crackers
3 marshmallows
2 pieces of chocolate
Which ingredient is the limiting reactant?
How many bicycles can be made if the following is needed?
1 bicycle =
1 frame +
2 wheels
We have:
100 frames
120 Wheels
Which component is the limiting reactant?
How many bikes can be made?
FACTORS THAT AFFECT REACTION RATE
3. Catalysts:
- substance that speeds up a reaction without being itself
changed or used up
- lowers activation energy needed to start reaction
- most powerful catalysts found in nature
ex: enzymes
4. Inhibitors:
Substance that slows down a reaction, doesn’t completely
stop reaction
Ex: stabalizers in products, retardents
Study for the test !