Transcript Intro to Reactions

Ch. 8 – Chemical Reactions
Intro to Reactions
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II III IV V
Signs of a Chemical Reaction
Evolution of heat and light
 Formation of a gas
 Formation of a precipitate
 Color change

Law of Conservation of Mass

mass is neither created nor destroyed
in a chemical reaction
total mass stays the same
 atoms can only rearrange

4H
36 g
2O
4H
2O
4g
32 g
Chemical Equations
A+B  C+D
REACTANTS
PRODUCTS
Symbols used in Equations
Writing Equations
2H2(g) + O2(g)  2H2O(g)
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Identify the substances involved.
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Use symbols to show:
 How many? - coefficient
 Of what? - chemical formula
 In what state? - physical state
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Remember the diatomic elements.
Writing Equations
Two atoms of solid aluminum react
with three units of aqueous
copper(II) chloride to produce three
atoms of solid copper and two units
of aqueous aluminum chloride.
• How many?
• Of what?
• In what state?
2Al(s) + 3CuCl2(aq)  3Cu(s) + 2AlCl3(aq)
Describing Equations
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Describing Coefficients:
 individual atom = “atom”
 covalent substance = “molecule”
 ionic substance = “formula unit”
3CO2  3 molecules of carbon dioxide
2Mg
 2 atoms of magnesium
4MgO  4 formula units of magnesium
oxide
Describing Equations
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
• How many?
• Of what?
• In what state?
One atom of solid zinc reacts with
two molecules of aqueous
hydrochloric acid to produce one unit
of aqueous zinc chloride and one
molecule of hydrogen gas.
Ch. 8 – Chemical Reactions
Balancing Equations
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II III IV V
Balancing Steps
1. Write the unbalanced equation.
2. Count atoms on each side.
3. Add coefficients to make #s equal.
Coefficient  subscript = # of atoms
4. Reduce coefficients to lowest
possible ratio, if necessary.
5. Double check atom balance!!!
Helpful Tips
Balance one element at a time.
 If Hydrogen and Oxygen appear,
balance them last. Balance
Hydrogen before Oxygen.
 Update ALL atom counts after adding
a coefficient.
 If an element appears more than
once per side, balance it last.
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Balancing Example
Aluminum and copper(II) chloride react
to form copper and aluminum chloride.
2 Al + 3 CuCl2  3 Cu + 2 AlCl3
2 1
Al
1 2
3 1
Cu
1 3
6 2
Cl
3 6
Balancing Example
Aqueous nitric acid reacts with solid
magnesium hydroxide to produce
aqueous magnesium nitrate and
water
 (2,1,1,2)
 Solid Calcium metal reacts with water
to form aqueous calcium hydroxide
and hydrogen gas.
 (1,2,1,1)
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Balancing example
(hydrocarbons)
Ethane gas reacts with oxygen gas to
form carbon dioxide and water.
 Ethene gas reacts with oxygen gas to
form carbon dioxide and water.
 Butane gas reacts with oxygen gas to
form carbon dioxide and water.

Ch. 8 – Chemical Reactions
Types of Chemical
Reactions
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II III IV V
Combustion

the burning of a hydrocarbon in O2 to
produce heat
A + O2  B
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Combustion

Products:
 form CO2 + H2O
 Form carbon monoxide and/or carbon if
there is limited oxygen
C3H8(g)+ 5 O2(g)  3 CO2(g)+ 4 H2O(g)
Synthesis (Combination)

the combination of 2 or more
substances to form a compound
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only one product
A + B  AB
Synthesis (Combination)
H2(g) + Cl2(g)  2 HCl(g)
Synthesis (Combination)
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Products:
 If products are ionic - cancel charges
 If covalent - hard to tell
2 Al(s)+ 3 Cl2(g)  2 AlCl3(s)
Decomposition

a compound breaks down into 2 or
more simpler substances
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only one reactant
AB  A + B
Decomposition
2 H2O(l)  2 H2(g) + O2(g)
Decomposition
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Products:
 binary - break into elements
 others - hard to tell (more on this later)
2 KBr(l)  2 K(s) +
Br2(l)
Synthesis Reactions to know
Almost all metals react with oxygen to
form metal oxides.
 Ex/ Magnesium reacts with oxygen to
form Magnesium oxide
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Synthesis reactions to know
Nonmetals also form oxides.
 Sulfur reacts with oxygen to form
Sulfur dioxide
 Carbon reacts with oxygen to form
Carbon dioxide
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Synthesis reactions involving
water to know
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Oxides of active metals react with
water to produce metal hydroxides.
 Ex/ Calcium oxide reacts with water
to form calcium hydroxide.
Synthesis reactions to know
involving water

Many oxides of nonmetals in the
upper right portion of the periodic
table react with water to produce
oxyacids (acids with oxygen).
 Ex/ sulfur dioxide reacts with water
to form sulfurous acid
Decomposition reactions to
know
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Metal carbonates break down to
produce a metal oxide and carbon
dioxide gas.
 Ex/ calcium carbonate decomposes
to produce calcium oxide and
carbon dioxide.
Decomposition reactions to
know
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All metal hydroxides except those
containing group 1 metals
decompose when heated to yield
metal oxides and water.
 Ex/ calcium hydroxide decomposes
to produce calcium oxide and
water.
Decomposition reactions to
know
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Metal chlorates decompose to
produce a metal chloride and oxygen.
 Ex/ potassium chlorate
decomposes in the presence of the
catalyst Manganese dioxide to
produce potassium chloride and
oxygen.
Decomposition reactions to
know

Sulfites decompose into the cation
oxide and sulfur dioxide gas
 Calcium sulfite decomposes to form
calcium oxide and sulfur dioxide
gas.
Decomposition reactions to
know

Peroxides decompose into the oxide
of the cation and oxygen gas.
 Hydrogen peroxide decomposes to
form water (dihydrogen monoxide)
and oxygen gas.
 Sodium peroxide decomposes to
form . . .
Decomposition reactions to
know

Certain acids decompose into
nonmetal oxides and water.
 Ex/ Carbonic acid is unstable and
decomposes readily to produce
carbon dioxide and water.
 Sulfurous acid decomposes to
produce sulfur dioxide and water.
Decomposition reactions to
know

Ammonium salts decompose into
ammonia and the acid of the anion.
 Ex/ Ammonium acetate
decomposes to form ammonia and
acetic acid.
 Ammonium nitrate decomposes to
form ammonia and nitric acid.
Single Replacement

one element replaces another in a
compound
 metal replaces metal (+)
 nonmetal replaces nonmetal (-)
A + BC  B + AC
Single Replacement
Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
Single Replacement

Products:
 Metal atom  metal cation(+)
 Nonmetal atom  nonmetal anion(-)
 free element must be more active (check
activity series)
Fe(s)+ CuSO4(aq)  Cu(s)+ FeSO4(aq)
Br2(l)+ NaCl(aq)  N.R.
Double Replacement

Occurs between two aqueous ionic
compounds

ions in two compounds “change partners”
cation of one compound combines with
anion of the other

AB + CD  AD + CB
Double Replacement
Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2KNO3(aq)
Double Replacement

Products:
 In order for a reaction to occur on of the
following must occur
- one product must be insoluble (a precipitate)
-
(check solubility table)
A gas is produced
A covalent compound (molecular) is produced
Pb(NO3)2(aq)+ 2KI(aq)  PbI2(s)+2KNO3(aq)
NaNO3(aq)+ KI(aq)  N.R.
Examples
Solutions of Hydrochloric acid and
Sodium hydroxide are mixed
 Solutions Rubidium carbonate and
acetic acid are mixed
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Writing net ionic equations
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1. Write the skeleton equation
 Make sure charges ions are balanced.
2. Write the total ionic equation
3. Cancel out ions that appear on both
sides
4. These are called spectator ions.
Spectator ions do not take part in the
overall net reaction.
Writing net ionic equations
A solution of Lead (II) nitrate is mixed
with a solution of potassium iodide.
 A solution of silver nitrate is mixed
with a solution of sodium hydroxide.
 View clip
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Predicting products from
reactions
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If given only reactants
 Figure out what type of reaction it is
 Write the products. If one of the
products is ionic, balance charges.
 Balance the equation
 Ex/ Hydrochloric acid is mixed with
zinc metal.
Predicting (cont’d)
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Calcium metal burns in the presence of
oxygen.
Sodium carbonate is heated vigorously.
Potassium chlorate is heated vigorously.
Pentane burns in oxygen.
Aqueous silver nitrate is mixed with
aqueous sodium chloride.
Predicting products
Hydrochloric acid is mixed with
Sodium hydroxide.
 Gold metal is dropped into
hydrochloric acid.
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