Chapter 7
 Chemical Reactions
7.1 Describing Chemical Reactions
 What is a chemical reaction?
 Demos: Silver and Copper
 Chemical Reaction: is when a
substance undergoes a chemical
change to produce a new substance
or substances
 7 different types of chemical
reactions We will learn five types
 Chemical equations: are used to
represent a chemical change/reaction
 Ex.
2 H2 + O2  2 H2O
 Reactants: are the substances that
undergo chemical change (bonds
broken) (ingredients)
 Products: are the new substances
produced by the change (new bonds
made) (cookies)
 Reading Chemical equations
 + sign = “reacts with”
  sign = “to produce”, yields
 s = solid state
 l = liquid state
 g= gas state
Labeling the parts of a chemical equations
http://www.chem.wisc.edu/deptfiles/genchem/sstutorial/Text3/Tx3
4/tx34p1.GIF
Counting the amounts:
 Mole is an amount of a
substance that contains
6.02X1023 particles. So rather
than using atoms, molecules
or ions we use the MOLE
 Coefficient = # of moles
Mole continued
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Shoes come in pairs,
Eggs in a dozen.
Particles to chemists come in MOLES
Particles can be atoms, molecules, or
ions
 Law of conservation of mass: states
that mass is neither created nor
destroyed in a chemical reaction
 By ordinary means, we can’t make or
destroy matter
 ONLY REARRANGE IT!!!!!!!!
Balancing Equations
 When writing chemical equations, the
law of conservation of mass must be
followed.
 We must have the same amount of
each element on both sides of the
equation
 Coefficients: are used to show
proportions of reactants and
products)
 PERFECT PRACTICE,
MAKES PERFECT!!!!!!
 2H2(g)+O2(g) 2H2O(l)
 Coefficients represent the number of
units of each substance in the
reaction
 Subscripts represent the number of
atoms in the molecule. These come
from writing the correct formula.
 Symbols show the state of reactants
(s) solid, (aq) aqueous), (g) gas, (l) liquid
How to balance reactions
Chemical Equations
A chemical equation is written as an
expression similar to a mathematic
equation that can be compared to a
recipe that a chemist follows in order to
produce desired results.
Balancing a Chemical
Equation

A chemical equation is balanced when the
ions or atoms found on the reactant side of
the equation equals that found on the
product side.

The arrow can be considered the balance
point.
Balancing Equations

When balancing a chemical reaction you may
add coefficients in front of the compounds to
balance the reaction, but you may not change
the subscripts.
 Changing the subscripts changes the
compound. Subscripts are determined by the
valence electrons (charges for ionic or
sharing for covalent)
 Think
back to naming compounds/ determining
formulas. NaCl exists, because Na is + and Cl is -,
but NaCl2 does NOT exist since you would not
have a neutral compound! You can’t just add a
number to a formula to balance an equation.
Chemical Equations
4 Al(s) + 3 O2(g) ---> 2 Al2O3(s)
This equation means
4 Al atoms + 3 O2 molecules
---produces--->
2 molecules of Al2O3
AND/OR
4 moles of Al + 3 moles of O2 --produces--->
2 moles of Al2O3
Steps to Balancing Equations
There are four basic steps to balancing a chemical equation.
1. Write the correct formula for the reactants and the
products. DO NOT TRY TO BALANCE IT YET! You must
write the correct formulas first.
**And most importantly, once you write them correctly DO NOT
CHANGE THE FORMULAS!
2.
3.
4.
Find the number of atoms for each element on the left
side. Compare those against the number of the atoms of
the same element on the right side.
Determine where to place coefficients in front of formulas
so that the left side has the same number of atoms as the
right side for EACH element in order to balance the
equation.
Check your answer to see if:
 The numbers of atoms on both sides of the equation
are now balanced.
 The coefficients are in the lowest possible whole
number ratios. (reduced)
Balancing Equations
2 H2(g) + ___ O2(g) ---> ___
2 H2O(l)
___
What Happened to the Other
Oxygen Atom?????
This equation is not balanced!
Two hydrogen atoms from a hydrogen
molecule (H2) combine with one of the
oxygen atoms from an oxygen molecule
(O2) to form H2O. Then, the remaining
oxygen atom combines with two more
hydrogen atoms (from another H2 molecule)
to make a second H2O molecule.
Balance this equation!
Na + Cl2
Na- 1
Cl- 2
NaCl
Na-1
Cl-1
**note that the number of sodiums balance but
the chlorine does not. We will have to use
coefficients in order to balance this equation.
Inserting subscripts
Na + Cl2
Na- 1
Cl- 2
2 NaCl
Na- 1 2
Cl- 1 2
** Now the chlorine balances but the sodium
does not! So we go back and balance the
sodium.
Discussion question
 Why are coefficients important?
 Chemists need to know how much of
a reactant will produce a certain
amount of a product.
Finally balanced!
2Na + Cl2
Na- 1 2
Cl- 2
2 NaCl
Na-1 2
Cl-1 2
**Since the number of each element on the
reactant side and the product side of the
equation are equal, the equation is
balanced.
Balancing
Equations
2
3
___ Al(s) + ___ Br2(l) ---> ___ Al2Br6(s)
Balancing Equations
Sodium phosphate + iron (III)
oxide  sodium oxide + iron (III)
phosphate
2 Na3PO4 +
Na2O 3 +
Fe2O3 ---->
FePO4 2
Types of reactions
7.2 Types of Reactions
 7 types of reactions
 Synthesis: is a reaction in which two
or more substances react to form a
more complex single substance
 A+BC
 Ex.
2 Na + Cl2  2 NaCl
 Decomposition Reaction: is a
reaction in which a compound breaks
down into two or more simpler
substances
 ABA+B
 Ex. 2 H20  2 H2 + O2
 Compost pile, digesting food,
electrolysis (breaking down H2O with
electricity)
 Combustion Reaction: is reaction
where a substance reacts rapidly with
oxygen.
 Flammability or explosiveness
 Ex. CH4 + 2 O2  CO2 + 2 H2O
 Ex. Burning any type of fuel
 Usually produces CO2 (greenhouse gas)
 Oxidation Reaction: are reactions
where a substance reacts slowly with
the oxygen in air or water
 Happens with metals
 “rust” or “tarnish”
 Ex.
2 Ca + O2  2 CaO
 Single Replacement Reaction: is a
reaction where one element takes the
place of another element
 A+BC AC+B or D+BC BD +C
 Cu + 2 AgNO3  2 Ag + Cu(NO3)2
 Double Replacement: is a reaction
where 2 elements replace each other
 AB+CDAD+CB

CaCO3 + 2 HCl  CaCl2 + H2CO3
Discussion Question:
 What two Chemical reactions are
“opposites” of each other and why?
 Synthesis and decomposition; in
synthesis multiple substances
combine to form a new one, while
decomposition a single substance
breaks apart into multiple simpler
substances
 Exothermic Reactions: are reactions
that release energy into their
surroundings
 Give off heat (exergonic)
 Ex. Combustion Reactions
 Ex. Burning fossil fuels
 Endothermic Reaction: is a reaction
where heat energy is absorbed by its
surroundings
 Ice absorbs heat to melt
into water
 Gets colder (endergonic)
 Ex. Ice pack and decomposition of
mercury
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Increasing Reaction Rate:
1. Temperature
2. Surface area
3. Stirring
4. Concentration of Reactants
5. Catalyst: is a substance that
affects reaction rate without being
used up.
 Activation energy: What it takes to
get a reaction started
 Inhibitor: Something that stops or
inhibits a reaction from proceeding
 Rate of reaction: how fast a reaction
proceeds
Discussion Question:
 How are chemical bonds involved in
energy exchanges?
 Breaking bonds requires energy;
forming bonds releases energy.