Transcript Document

Unit 8
Chemical Reactions
PAGE 123
Describing Chemical
Reactions
• A chemical reaction is the
process by which one or
more substances are
changed into one or more
different substances.
• They are described by
chemical equations
Describing Chemical
Reactions
• In a chemical reaction:
bonds are broken in the
reactants
– Bonds are formed in the
products
–
Parts of a Chemical
Equation
Reactants
• original substances in a
chemical reaction
• written on the left hand
side of a chemical equation
Products
• The resulting substances
produced in a chemical
reaction.
• written on the right hand
side of the chemical equation
yields
Subscripts in (parenthesis)
• represents the physical states
of the compounds (elements)
• Ex: (s)= solid, (l)=liquid, (g)= gas,
(aq)= aqueous (dissolved in water)
• Words or symbols are
placed over/under the
arrow(s) to indicate certain
conditions under which the
reaction is carried out.
– : heat is applied
– catalyst
• substance that speeds up a
chemical reaction without
becoming part of the
reaction.
Law of conservation of
mass
According to the law of
conservation of mass,
• the total mass of reactants must
equal the total mass of products
for any given chemical reaction.
Law of conservation of
mass
Balance atoms
2 H2 (g) + O2 (g)  2 H2O
Mass must be equal
50 g + 45 g  95g
(g)
Translate Chemical
Equations into Words
Chemical Equations
• 1) include all of the
symbols appearing in the
equation (with the
exception of the
coefficients)
Chemical Equations
• 3. Diatomic Elements:
Elements that combine
with each other when
found alone in nature
Diatomic Elements
• 7 diatomic elements (know)
H2 O2 F2 Br2 I2 N2 Cl2
Example #2
Mg(s) + 2HCl(aq)
MgCl2(aq) + H2(g)
Solid magnesium reacts with
aqueous hydrochloric acid
to produce aqueous magnesium
chloride and hydrogen gas
Check for Understanding
CaCO3(s)
CaO(s)
+
CO2(g)
Solid calcium carbonate is
heated and produces solid
calcium oxide and gaseous
carbon dioxide
Check for Understanding
#2
3KOH(aq) + Fe(NO3)3(aq)
Fe(OH)3(s) + 3KNO3(aq)
Aqueous potassium hydroxide
reacts with aqueous iron (III)
nitrate to produce solid iron (III)
hydroxide and aqueous potassium
nitrate
Writing Chemical Equations
from a Written Description
Example #1
Aluminum metal reacts with
oxygen in the air to form
solid aluminum oxide.
4 Al
(s)
+ 3 O2 (g)  2 Al2O3
(s)
Example #2
When solid mercury(II) sulfide is
heated with oxygen gas, liquid
mercury metal and gaseous sulfur
dioxide are produced.
HgS
(s)
+ O2 (g)
Hg
(l)
+ SO2
(g)
Check for Understanding
Oxygen gas can be produced by heating
solid potassium chlorate in the
presence of the catalyst manganese
dioxide. Potassium chloride is a solid
residue.
2 KClO3
(s)
MnO2
2 KCl
(s)
+ 3O2 (g)
Check for Understanding
Aqueous potassium nitrate
and a precipitate of barium
chromate are formed when
aqueous solutions of barium
nitrate and potassium
chromate are mixed.
Check for Understanding
Aqueous potassium nitrate and a
precipitate of barium chromate are
formed when aqueous solutions of
barium nitrate and potassium
chromate are mixed.
Ba(NO3)2 (aq) + K2CrO4 (aq) 
2 KNO3
(aq)
+
BaCrO4 (s)
Types of Chemical
Reactions
Synthesis (Combination)
Reactions
Two or more substances
combine to form a new
compound.
.
Formation of only ONE product.
Synthesis (Combination)
Reactions
Examples:
2H2
+
CaO +
O2
2H2O
H2O
Ca(OH)2
Decomposition Reactions
• A single compound
undergoes a reaction that
produces two or more
simpler substances.
Decomposition Reactions
Decomposition Reactions
Decomposition reactions are
the opposite of synthesis
reactions.
There is only one reactant.
Examples
• CaCO3  CaO
+
CO2
• Na2CO3  Na2O + CO2
Single-Replacement
Reactions
• one element replaces a similar
element in a compound.
• a single element reacts with a
compound.
Single-Replacement
Reactions
Distinguishing
Characteristics
• Examples:
Pb
+
2HNO3
F2
+
2NaI 
 Pb(NO3)2
2NaF
+
+
I2
H2
Predict the Products
Zn + CuCl2  ZnCl2
+
Cu
Double Replacement
Reactions
The ions of two compounds
exchange places in an
aqueous solution to form
two new compounds.
Two ionic compounds “swap”
ions.
Double Replacement
Reactions
Usually forms a
precipitate
•
precipitate: when two aqueous
solutions are mixed and a solid
product that is formed
AgNO3(aq) + KCl(aq) AgCl(s) + KNO3(aq)
Precipitate
Examples
• PbSO4 + 2KCl  K2SO4 + PbCl2
• AgNO3(aq) + KCl(aq) ----> AgCl(s) + KNO3(aq)
Predict the Products
2 HOH
2 KOH + H2SO4  K2SO4 + H2O
Combustion Reactions
A hydrocarbon (CxHy)
combines with oxygen,
releasing a large amount of
energy in the form of light
and heat.
General Form
CxHy +
O2  CO2
+
H 2O
Example
• C3H8
+ 5O2  3 CO2
• CH4 + O2  CO2 + H2O
+
4 H 2O
Predict the Products
1
2
6
5
2 C2H6 + 10 O2  4 CO2 + 12 H2O
1)
2 Li + 2 H2O  2 LiOH + H2
single replacement / displacement
2)
NH4Cl  NH3
+
decomposition
HCl
3)
AgNO3 + NaCl  AgCl + NaNO3
double replacement / displacement
4) 2 C6H14 + 19 O2
12 CO2 + 14 H2O
combustion
5) N2
+
3 H2
2 NH3
combination / synthesis
Oxidation – Reduction
Processes
Redox Reaction
Any chemical reaction that
involves the transfer of one
or more electrons between
atoms. One reactant gains
electrons and the other
reactant loses electrons.
Examples of Redox Reactions
• Combustion of gasoline
• Burning of wood
• Energy from food
• Bleaching stains
• Iron rusting
Oxidation and Reduction
•
Opposing reactions
• Must occur together
– Redox (Reduction – Oxidation)
Oxidation
• Electrons are fully or partially
lost
Mg0 (s)  Mg
2+
+ 2e-
Electrons lost:
written as a product
Reduction
•
Electrons are fully or partially
gained
S0(s) + 2 e-  S2electrons gained:
written as a reactant
Mnemonic Device
• LEO the lion goes GER
LEO: Loss of Electrons is Oxidation
GER: Gain of Electrons is Reduction
Mnemonic Device
• OIL RIG
– Oxidation Is Loss of electrons
– Reduction Is Gain of electrons
Complete Redox Reaction
• Add together the reduction halfreaction with the oxidation halfreaction to get the complete redox
reaction.
Complete Redox Reaction
+
Mg0 (s)  Mg
2+
+ 2e-
2+
+ S2-
S0(s) + 2 e-  S2-
Mg0 (s) + S0(s)  Mg
Mg +
S
 Mg+2
+
S
-2
Assigning Oxidation
Numbers
Oxidation Numbers
• Numbers assigned to all of
the elements involved in the
reaction to determine if
electrons have been
transferred between atoms in
a reaction.
Oxidation Numbers
• The oxidation number is usually equal
to the charge on the ion if it was
formed.
Rule #1
• Free elements are assigned an
oxidation state of 0.
– Al = 0
– Na = 0
– H2 = 0
Rule #2
• The oxidation state for any simple
one-atom ion is equal to its
charge.
– Na+ = +1
– Be2+= +2
– F- = -1
Rule #3
• The alkali metals (Li, Na, K,
Rb, Cs and Fr) in compounds
are always assigned an
oxidation state of +1.
– LiOH:
Li = +1
– Na2SO4: Na= +1
Rule #4
• Fluorine in compounds is
always assigned an oxidation
state of -1.
– HF: F= -1
– MgF2: F= -1
Rule #5
• The alkaline earth metals (Be, Mg,
Ca, Sr, Ba, and Ra) and also Zn and
Cd in compounds are always
assigned an oxidation state of +2.
Similarly, Al & Ga are always +3.
• MgF2: Mg = +2
• CaO: Ca= +2
• Al2O3: Al=+3
Rule #6
• Hydrogen in compounds is assigned an
oxidation state of +1.
Exception - Hydrides, ex. LiH (H=-1).
• H2SO4: H=+1
• HCl: H=+1
Rule #7
• Oxygen in compounds is assigned an
oxidation state of -2.
Exception - Peroxide, ex. H2O2 (O = -1).
– H3PO4: O=-2
– H2O: O=-2
Rule #8
• The sum of the oxidation
numbers of all atoms in a
neutral compound is 0.
– H2SO4: add up to 0
– CO2: add up to 0
Rule #9
• The sum of the oxidation
numbers of all atoms in a
polyatomic ion equals the charge
on the ion.
–SO42-: add up to -2
–NH4+: add up to +1
Examples
• Al(s):
Rule #1
0
– Al = ____
Examples
• CaCl2:
+2
Ca
– Ca = ____
-1
– Cl=____
Rule #5
Rule #8
1 (+2) + 2 (Cl) = 0
Cl = -1
Examples
• HNO3:
+1
– H= ____
+5
– N= ____
-2
– O= ____
Rule # 6
Rule # 7
Rule # 8
1(1) + 1(N) + 3(-2) = 0
N = +5
Examples
• SO42-:
+6
– S= ____
-2
– O= ____
Rule #7
Rule #9
1(S) + 4(-2) = -2
S= +6
Examples
• H2O:
Rule #6
+1
– H= ____
-2
– O= ____
Rule #7
Examples
• (NH4)2CO3
NH4+
CO32-
Rule 6 Rule 7
Rule 9
1(N)+4(+1)=+1
N= -3
1(C)+3(-2)=-2
C=+4
-3
N=____
+1
H=____
+4
C=___
-2
O=___
Check for Understanding
• S8:
0
– S= ____
• AsO43-:
+5
– As= ____
-2
– O= ____
Check for Understanding
OH• Cr(OH)3:
+3
– Cr= ____
-2
– O= ____
+1
– H= ____
Check for Understanding
• (NH4)3PO4:
-3
– N= ____
+1
– H=_____
+5
– P= ____
-2
– O= ____
NH4+
PO4-3
Oxidized Element
• Atoms that lose electrons in
a chemical reaction.
– Elements that lose electrons
are undergoing oxidation and
are said to be oxidized.
Oxidized Element
• The substance that loses electrons
is the oxidized element.
• Atoms that are oxidized will have
an increase in their oxidation
number
Oxidized Element
Increase in oxidation number
Mg0 (s) + S0(s)  Mg
2+
Mg is oxidized
+ S2-
Reduced Element
• Atoms that gain electrons in
a chemical reaction.
– Elements that gain electrons
are undergoing reduction and
are said to be reduced.
Reduced Element
• The substance that gains electrons
is the reduced element.
• Atoms that are reduced will have
an decrease in their oxidation
number
Reduced Element
Decrease in oxidation number
Mg0 (s) + S0(s)  Mg
2+
S is reduced
+ S2-
Oxidizing Agent
• The element or compound that
is reduced.
• It is called the oxidizing
agent because it oxidizes the
other element or compound.
Oxidizing Agent
Decrease in oxidation number
Mg0 (s) + S0(s)  Mg
2+
+ S2-
S is reduced
S is the oxidizing agent
(it is oxidizing Mg)
Reducing Agent
• The element or compound that
is oxidized.
• It is called the reducing agent
because it reduces the other
element or compound.
Reducing Agent
Increase in oxidation number
Mg0 (s) + S0(s)  Mg
2+
+ S2-
Mg is oxidized
Mg is the reducing agent
(it is reducing S)
Examples
• Are the following redox reactions?
• If yes:
–
–
–
–
What
What
What
What
element is oxidized?
element is reduced?
is the oxidizing agent?
is the reducing agent?
Examples
0
Mg(s) +
S(s)
0
+2 -2
 MgS
Mg
What
element
is
oxidized?
Step
#
2:
Do
the
oxidation
numbers
Step
Oxidation Numbers
Yes#1:
– ItAssign
is a redox
change?
What element is reduced?S
- if is
no the
 itoxidizing
is not a redox
What
agent?
S
reaction
What is the reducing agent?Mg
- if yes  it is a redox reaction
Examples
0
0
+1
+2
2AgNO3(aq) + Cu(s)  Cu(NO3)2(aq) + 2Ag(s)
NO3-
NO3-
Step #
2: Do the
oxidation
What
element
is oxidized?
Cunumbers
change?
What
element is reduced?Ag
Yes
–
It
is
a
redox
Step
#1:
Assign
Oxidation
Numbers
if
no

it
is
not
a redox
What is the oxidizing agent?
AgNO3(aq)
reaction
What is the reducing
agent?Cu(s)
- if yes  it is a redox reaction
Check for Understanding
•
Are the following redox
reactions?
• If yes:
– What element is oxidized?
– What element is reduced?
– What is the oxidizing agent?
– What is the reducing agent?
Check for Understanding
0
0
+3 -2
4Fe(s) + 3O2(g)  2Fe2O3(s)
What element
is
oxidized?
Fe
Yes
What element is reduced?O
What is the oxidizing agent?
O2
What is the reducing agent?
Fe
Check for Understanding
+1 -2
+2 -2+1
0
0
Ca(s) + H2O(l)  Ca(OH)2(aq) + H2(g)
OHYesis oxidized?Ca
What element
What element is reduced?H
What is the oxidizing agent?
H2O
What is the reducing agent?
Ca
Check for Understanding
+1 -1
+1 -2 +1
+1 -1
+1 -2
HCl + NaOH  NaCl + H2O
No – it is not a redox
Check for Understanding
0
+3 -2
+2 -2
+4 -2
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
What element is oxidized?C
Yesis reduced?Fe
What element
What is the oxidizing agent?
Fe2O3
What is the reducing agent?
CO
Check for Understanding
+1
0
+2
+0
2AgNO3(aq) + Cu(s)  Cu(NO3)2(aq) + 2Ag
NO3-
NO3-
(s)
What element
is
oxidized?
Cu
Yes
What element is reduced?Ag
What is the oxidizing agent?
AgNO3
What is the reducing agent?
Cu
Check for Understanding
+1-1 +1
+1
+1 -1
NaCl + AgNO3  NaNO3 + AgCl
NO3-
NO3-
No – it is not a redox
Check for Understanding
0
0
+1-1
2 H2(g) + O2(g)  2 H2O(l)
What element is oxidized?H
Yes is reduced?O
What element
What is the oxidizing agent?
O2
What is the reducing agent?
H
2
Check for Understanding
-4+1
0
+4 -2
+1 -2
CH4 + 2 O2  CO2 + 2 H2O
What element is oxidized?C
Yesis reduced?O
What element
What is the oxidizing agent?O2
What is the reducing agent?
CH4