Transcript 2 SO 4 + H 2 O - Geneva Area City Schools

DEFINITION OF
REACTION
a chemical process in which two or
more substances act mutually on each
other and are changed into different
substances
Same as saying “combine”
Exp. Iron (Fe) + Oxygen (O2) = rust
Physical Change
Changes in physical properties
 melting
 boiling
 Condensation
No change occurs in the identity of the
substance
Exp: Ice , rain, and steam are all water
3
Chemical Change
 Atoms in the reactants are rearranged to
form one or more DIFFERENT & NEW
substances
 Old bonds are broken  new bonds form
Exp:
Fe and O2 form rust (Fe2O3)
Ag and S form tarnish (Ag2S)
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Indicators of chemical reactions
Emission of light or heat
Formation of a gas
Formation of a precipitate
Color change
Emission of odor
Learning Check E1
Classify each of the following as a
1) physical change or 2) chemical change
A. ____ a burning candle
B. ____ melting ice
C. ____ toasting a marshmallow
D. ____ cutting a pizza
E. ____ polishing silver
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Solution E1
Classify each of the following as a
1) physical change or 2) chemical change
A. __2__ a burning candle
B. __1_ melting ice
C. __2__ toasting a marshmallow
D. __1__ cutting a pizza
E. __2__ polishing silver
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Chemical Reaction
A process in which at least one new
substance is produced as a result of
chemical change.
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All chemical reactions:
have 2 parts
Reactants - the substances you start
with
 Products- the substances you end up
with

The reactants turn into the products.
Reactants  Products
A Chemical Reaction
Reactants
Products
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Describing chemical reaction
The way atoms are joined is changed
Atoms aren’t created or destroyed.
Can be described several ways:
 In a sentence


Copper reacts with chlorine to form copper (II)
chloride.
In a word equation

Copper + chlorine  copper (II) chloride
Both still mean the same thing:

Cu(s) + Cl2(g)  CuCl2(aq)
Writing a Chemical Equation
Chemical symbols give a “before-and-after”
picture of a chemical reaction
Reactants
MgO
+
Products
C
magnesium oxide
reacts with carbon
CO
to form
LecturePLUS
+
Mg
carbon monoxide
and magnesium
Timberlake
12
Matter Is Conserved
H2
+
Cl2
2 HCl
+
+
Total atoms
2 H, 2 Cl
=
Total atoms
2H, 2 Cl
Total Mass
2(1.0) + 2(35.5)
73.0 g
=
Total Mass
2(36.5)
73.0 g
=
13
Law of Conservation of Mass
In any ordinary chemical reaction,
matter is not created nor destroyed
LecturePLUS
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Symbols used in equations
(s) after the formula – solid = Cu(s)
(g) after the formula – gas = H2 (g)
(l) after the formula - liquid = H2O(l)
(aq) after the formula - dissolved in
water, an aqueous solution. = CaCl2 (aq)
Symbols used in equations
used after a product indicates a gas
(same as (g)) = O2
used after a product indicates a solid
(same as (s)) = CaCo3 
---->indicates a reaction occurred or
something is produced/yielded
Symbols used in equations
reaction.
indicates a reversible
shows that heat
is supplied to the reaction.
, or
is used to
indicate a catalyst used, in this case,
platinum.
,
indicates a
pressure other than STP
Summary of Symbols
Learning Check E3
12 oz of dough, 4 oz mushrooms, 12 slices
pepperoni, 8 oz cheese and 5 oz tomato sauce
are used to make a pizza. Write a recipe in
words for putting together a pizza.
How would you write the recipe as an
equation?
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Solution E3
Example: Combine 12 oz dough + 4 oz
mushrooms + 12 slices pepperoni + 8
oz cheese + 5 oz tomato sauce and heat
30 minutes at 350°C to produce 1 pizza
12 oz dough + 4 oz mshrm
+ 12 pep + 8 oz chse
1 pizza
+ 5 oz tom sauce
LecturePLUS
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Reading A Chemical Equation
4 NH3 + 5 O2
4 NO + 6 H2O
Four molecules of NH3 react with five
molecules O2 to produce four molecules NO
and six molecules of H2O
or
Four moles NH3 react with five moles O2 to
produce four moles NO and six moles H2O
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A Balanced Chemical Equation
Same numbers of each type of atom on each
side of the equation
Al
+
S
Al2S3
Not Balanced
2Al
+
3S
Al2S3
Balanced
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Balance Equations with
Coefficients
Coefficients in front of formulas balance
each type of atom
4NH3 +
5O2
4NO + 6H2O
4N
=
4N
12 H
=
12 H
10 O
=
10 O
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Steps in Balancing An Equation
(focus on one atom at a time)
Fe3O4 + H2
Fe + H2O
Fe: Fe3O4 + H2
3 Fe + H2O
O:
Fe3O4 + H2
3 Fe + 4 H2O
H:
Fe3O4 + 4 H2
3 Fe + 4 H2O
LecturePLUS
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Learning Check E4
Fe3O4 + 4 H2
3 Fe + 4 H2O
A. Number of H atoms in 4 H2O
1) 2
2) 4
3) 8
B. Number of O atoms in 4 H2O
1) 2
2) 4
3) 8
C. Number of Fe atoms in Fe3O4
1) 1
2) 3
LecturePLUS
Timberlake
3) 4
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Solution E4
Fe3O4 + 4 H2
3 Fe + 4 H2O
A. Number of H atoms in 4 H2O
3) 8
B. Number of O atoms in 4 H2O
2) 4
C. Number of Fe atoms in Fe3O4
2) 3
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Learning Check E5
Balance each equation. The coefficients for
each equation are read from left to right
A.
Mg
+
1) 1, 3, 2
B.
Al
+
1) 3, 3, 2
N2
Mg3N2
2) 3, 1, 2
Cl2
AlCl3
2) 1, 3, 1
LecturePLUS
3) 3, 1, 1
Timberlake
3) 2, 3, 2
27
Learning Check E5
C.
Fe2O3
+
1) 2, 3, 2,3
D.
Al
Al
+
+
Fe
2) 2, 3, 4, 3
Fe +
2) 2, 1, 1, 1
H2SO4
1) 3, 2, 1, 2
Timberlake
CO2
Al2O3
3) 3, 3, 3, 1
Al2(SO4)3
2) 2, 3, 1, 3
LecturePLUS
+
3) 1, 1, 2, 3
FeO
1) 2, 3, 3, 1
E.
C
+ H2
3) 2, 3, 2, 3
28
Solution E5
A. 3 Mg
+ N2
Mg3N2
B. 2 Al
+ 3 Cl2
2 AlCl3
C. 2 Fe2O3 + 3 C
D. 2 Al
+ 3 FeO
E. 2 Al + 3 H2SO4
LecturePLUS
4 Fe
+ 3 CO2
3 Fe +
Al2(SO4)3
Timberlake
Al2O3
+ 3 H2
29
Counting with Moles
Avogadro’s # - 6.02 x 1023
Because chemical reactions often
involve LARGE numbers of SMALL
particles, chemists use a counting unit
called  mole to measure amounts
Exp. A mole of Iron (Fe) is 6.02 x 1023
atoms of Iron
Molar Mass
The mass of one mole of a substance is
called the MOLAR MASS

For an element, the molar mass is the
same as the atomic mass expressed in
grams (g)

Exp. Atomic mass of carbon is 12.0 amu (found
on the periodic table) so the molar mass of
carbon is 12.0 grams
Molar Mass in Compounds
Calculate the molar mass by adding up
the atomic masses of its atoms and
then convert to grams
Exp. Carbon dioxide CO2 – 1 atom of
carbon & 2 atoms of oxygen

1 carbon atom (12.0 amu) + 2 oxygen
atoms (2 x 16.0 amu = 32.0 amu)

Total molar mass = 44.0 grams
Mole-Mass Conversions
Once you know the molar mass of a
substance, you can convert moles of
that substance into mass or a mass of
that substance into moles
Mole-Mass Example
Molar mass of CO2 = 44.0 grams
(12.0 + (2x16.0) = 44.0 g
 This means that 1 mole of CO2 = mass of 44.0 g
 44.0 g CO2 / 1 mol CO2 or 1 mol CO2 / 44.0 g CO2

Suppose you have 55.0 g of CO2. Calculate how
many moles of CO2 you have:
55.0 g CO2 x 1 mol CO2 / 44.0 g CO2 = 1.25 mol CO2
(Check your answer: 1.25 mol CO2 x 44.0g CO2 / 1 mol CO2 = 55.0 g CO2)
Chemical Equations
In chemical reactions the mass of a reactant or
product can be calculated by using:

Balanced chemical equation – tells you how to
relate amounts of reactants to amounts of products

Molar masses of the reactants and products –
lets you convert amounts into masses
Converting Mass to Moles
Calculate how much oxygen is required to
make 144 g of H2O:

Determine how many moles of water you are
trying to make
144 g H2O x 1 mol H2O / 18.0 g H2O = 8.00 mol H2O
 Copy down chart on page 197

Chemical Reactions
Can be determined by:

Type of reactant or # of reactants and
products
Types of Chemical Reactions
1.
2.
3.
4.
5.
Synthesis reactions
Decomposition reactions
Single displacement reactions
Double displacement reactions
Combustion reactions
1. Synthesis
Example C + O2
C + O O
General:

O C O
A + B  AB
Ex. Synthesis Reaction
Practice
Predict the products.
2 Na(s) + Cl2(g)  2 NaCl(s)
Mg(s) + F2(g) 
2 Al(s) +3 F2(g) 
• Now, balance them.
MgF2(s)
2 AlF3(s)
2. Decomposition
Example: NaCl
Cl Na
General:

Cl
+
Na
AB  A + B
Ex. Decomposition Reaction
3. Single Displacement
Example: Zn + CuCl2
Zn was oxidized
Went from neutral (0) to (+2)
Cu
Cl
+
Cl
Zn

Zn
Cl
+
Cu
Cl
Cu was reduced
Went from (+2) to Neutral (0)
General:
AB + C  AC + B
Ex. Single Replacement Reaction
Single Replacement Reactions
Write and balance the following single
replacement reaction equation:
2
Zn(s) + HCl(aq)  ZnCl2 + H2(g)
• 2 NaCl(s) + F2(g)  2NaF(s) + Cl2(g)
•2 Al(s)+ 3 Cu(NO3)2(aq) 3 Cu(s)+ 2Al(NO3)3(aq)
4. Double displacement
Example: MgO + CaS
Mg
+
O
General:
Ca
S

Mg
S
+
Ca
O
AB + CD  AD + CB
Double Replacement Reactions
Think about it like “foil”ing in algebra, first and
last ions go together + inside ions go together
Example:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
Another example:
K2SO4(aq) + Ba(NO3)2(aq) 2 KNO3(aq) + BaSO4(s)
Practice
Predict the products.
5.
HCl(aq) + AgNO3(aq) 
CaCl2(aq) + Na3PO4(aq) 
Pb(NO3)2(aq) + BaCl2(aq) 
FeCl3(aq) + NaOH(aq) 
H2SO4(aq) + NaOH(aq) 
6.
KOH(aq) + CuSO4(aq) 
1.
2.
3.
4.
5. Combustion Reactions
Combustion reactions - a
hydrocarbon reacts with
oxygen gas.
This is also called
burning!!!
In order to burn something
you need the 3 things in
the “fire triangle”:
1) Fuel (hydrocarbon)
2) Oxygen
3) Something to ignite the
reaction (spark)
Combustion Reactions
In general:
CxHy + O2  CO2 + H2O
Products are ALWAYS
carbon dioxide and water.
(although incomplete burning
does cause some by-products
like carbon monoxide)
Combustion is used to heat
homes and run automobiles
(octane, as in gasoline, is
C8H18)
Combustion
Example 8
•
6
5
C5H12 + O2  CO2 + H2O
Write the products and balance the
following combustion reaction:
•
C10H22 + O2 
Mixed Practice
State the type & predict the products.
1. BaCl2 + H2SO4 
2. C6H12 + O2 
3. Zn + CuSO4 
4. Cs + Br2 
5. FeCO3 
Synthesis Reactions
Also called combination reactions
2 elements, or compounds combine to
make one compound.
A + B

AB
Na (s) + Cl2 (g)  NaCl (s)
Ca (s) +O2 (g)  CaO (s)
SO3 (s) + H2O (l)  H2SO4 (s)
We can predict the products if they
are two elements.
Mg (s) + N2 (g)  Mg3N2 (s)
A simulation of the reaction:
2H2 + O2

2H2O
Decomposition Reactions
decompose = fall apart
one compound (reactant) falls apart into
two or more elements or compounds.
Usually requires energy
 A + B
electricity
 Na + Cl2
NaCl   

CaCO3   CaO + CO2
AB
Decomposition Reactions
Can predict the products if it is a
binary compound
Made up of only two elements
Falls apart into its elements
electricity
H2O 

 H2 (g) + O2
(g)
HgO  
Hg

(s) + O2 (g)
Decomposition Reactions
If the compound has more than two
elements you must be given one of
the products
The other product will be from the
missing pieces
NiCO3 (aq)  
 CO2 (g) + Ni (s)
H2CO3(aq)

H2 (g) + CO
2 (g)
Single Replacement
Also referred to as single displacement
One element replaces another
Reactants must be an element and a
compound.
Products will be a different element and a
different compound.
A + BC

AC + B
2Na + SrCl2  Sr + 2NaCl
F2 + LiCl  LiF + Cl2
Single Replacement
We can tell whether a reaction will happen
Some are more active than other
More active replaces less active
Double Replacement
Two things replace each other.
Reactants must be two ionic compounds
or acids.
Usually in aqueous solution
AB + CD
 AD + CB
ZnS
+ 2HCl

AgNO3 + NaCl 
ZnCl + H2S
AgCl + NaNO3
Combustion
A reaction in which a compound
(often carbon) reacts with oxygen
CH4 + O2

CO2 + H2O
C3H8 + O2

CO2 + H2O
C6H12O6 + O2 
CO2 + H2O
The charcoal used in a grill is basically
carbon. The carbon reacts with oxygen to
yield carbon dioxide. The chemical equation
for this reaction is C + O2  CO2
Acid/Base Reaction
An acid and a base react to form a salt
and water.
Always in aqueous solution
Acid (H+) + Base (OH-) → Salt + H2O
NaOH + HCl → NaCl + H2O
NH4OH + H2SO4 →
(NH4)2SO4 + H2O
How to recognize which type
Look at the reactants

Element(E), Compound(C)
E+E
C
E+C
C+C
Acid + Base
Synthesis
Decomposition
Single replacement
Double replacement
Acid/Base reaction
Look at the Products
CO2 + H2O
Combustion
Redox
Examples
H2 + O2  Synthesis
H2O  Decomposition
AgNO3 + NaCl  Double replacement
Zn + H2SO4  Single replacement
HgO  Decomposition
KBr +Cl2  Single replacement
Mg(OH)2 + H2SO3  Double replacement
Examples
Acid/Base
HNO3 + KOH 
CaPO4  Decomposition
Single replacement
AgBr + Cl2 
Zn + O2 
Synthesis
HgO + Pb Single replacement
HBr + NH4OH Acid/Base
Cu(OH)2 + KClO3  Double replacement
Summary
An equation:
Describes a reaction
Must be BALANCED because to follow Law
of Conservation of Energy
Can only be balanced by changing the
coefficients.
Has special symbols to indicate state, and if
catalyst or energy is required.
Can describe 5 different types of reactions.
Oxidation-Reduction
Reaction
A reaction in which electrons are
transferred from one reactant to
another.

A.K.A. Redox Reaction
Oxidation
Any process in which an element loses
electrons during a chemical reaction

A reactant is oxidized if it loses electrons

Exp. 2Ca + O2  2 CaO

When calcium reacts with oxygen, each neutral
calcium atom loses two electrons and becomes
a calcium ion with a charge of +2
Reduction
The process in which an element gains
electrons during a chemical reaction
A reactant is said to be reduced if it gains
electrons
 Exp. O + 2e  O2
Redox Reactions
Compare:

Contrast:
Must happen
together = when 1
atom loses eanother will gain it

Can be a partial
transfer as in the
synthesis of H2O

“Sharing”
Oxidation – loses electrons
Reduction – gains electrons
Chemical Bonds and
Energy
Energy stored in the chemical bonds of
a substance

Exp. Propane – has 10 single covalent
bonds; chemical energy of a propane
molecule is the energy stored in these
bonds
Importance of Chemical
Reactions
Chemical reactions involve the breaking
of chemical bonds in the reactants and
the formation of chemical bonds in the
products
Energy is either released or absorbed
Breaking Bonds
Requires energy
Exp. Why propane grills have an igniter,
a device that produces a spark

Spark provides enough energy to break the
bonds of reacting molecules and get the
reaction to start
Forming Bonds
Releases energy
Exp. The heat and light given off by a
propane stove result from the formation
of new chemical bonds
Exothermic and
Endothermic Reactions
Exothermic:


Releases energy
Energy released as
the products form is
greater than the
energy required to
break the bonds in
the reactants
Endothermic:


Absorbs energy
More energy is
required to break the
bonds in the
reactants than is
released by the
formation of the
products
Reaction Energy
All chemical reactions are accompanied by a change in
energy.
Exothermic - reactions that release energy to their
surroundings (usually in the form of heat)
oΔH (enthalpy) is negative – energy leaving system
Endothermic - reactions that need to absorb heat from
their surroundings to proceed.
oΔH (enthalpy) is positive – energy coming into the system
Law of Conservation of
Energy
Exothermic – chemical energy of the
reactants is converted into heat plus the
chemical energy of the products
Endothermic – heat plus the chemical
energy of the reactants is converted into
the chemical energy of the products
Both cases = total amount of energy is
the same before and after the reaction
Reaction Rate
Rate at which reactants change into
products over time
Tell you how fast a reaction is going
 Exp. Burning calories

Calories – unit of energy used in nutrition
 Pg. 212 in textbook

Factors affecting
Reaction Rates
Temperature
Surface area
Concentration
Stirring
Catalysts
Temperature
Generally an increase in temperature
will increase the reaction rate and vice
versa
Increasing the temperature causes the
particles to move faster which are thus
more likely to collide and more likely to
react
Surface Area
Smaller the particle size of a given
mass, the larger is its surface area
Increase in surface area increases the
exposure of reactants to one another
The greater the exposure, the more
collisions there are that involve reacting
particles – more collisions = more
reactions
Stirring
Stirring causes collisions between
particles of the reactants
Exp. Clothes in a washing machine
Concentration
Refers to the # of particles in a given
volume
The more reacting particles that are
present in the given volume, the more
opportunities there are for collisions
involving those particles
Catalysts
A substance that speeds up a
reaction without being used up or
changed by the reaction.
Often used to speed up a reaction or
enable a reaction to occur at a lower
temp.
Lower the energy barrier (activation
energy) required for a reaction to
take place
Reaction Energy
•Spontaneous Reactions - Reactions that proceed
immediately when two substances are mixed together. Not
all reactions proceed spontaneously.
•Activation Energy – the amount of energy that is
required to start a chemical reaction.
•Once activation energy is reached the reaction
continues until you run out of material to react.