Transcript NOTES CHEMICAL REACTIONS:

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CHEMICAL REACTIONS
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• Chemical reactions: Reactions that produce new
substances
• PRODUCT: substance formed during a chemical
reaction (right side of arrow)
• REACTANT: starting substance(s) in a chemical
reaction (left side of arrow)
• Law of Conservation of Mass must be satisfied!
In this unit you should
know…
• 1. How to balance chemical equations
• 2. Identify the different types of reactions
• 3. Be able to predict the products for both
single and double replacement reactions
• 4. Determine if a reaction will take place
using either the activity series of metals or
solubility rules
• 5. Understand the role of a catalyst in a
chemical reaction
Evidence of Chemical
Reactions
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Temperature change:
endothermic (colder), exothermic (hotter)
Color Change
Odor
Gas Produced (bubbles)
Precipitate: formed from 2 liquids
Balancing Equations
Steps
1) Balance atoms that appear only once on
each side.
2) Balance polyatomic ions that appear on
both sides as a single unit.
3) Balance hydrogens.
4) Balance oxygens.
5) Never change the subscripts of a
compound to balance an equation.
Types of Reactions
• Synthesis Reaction:
• 1. Two or more substances combine to
form a single compound.
• 2. Usually energy is released
(exothermic)
• 3. Basic reaction: A + B --> AB
Synthesis Reaction
Examples:
• Element + Oxygen ----> Oxide
Compound
• Magnesium + Oxygen ---> Magnesium Oxide
•
Mg
+
O2
------> 2 MgO
• Metal Oxide + Water ---> Hydroxide Compound
•
CaO + H2O ---> Ca(OH)2
(base)
Decomposition
Reactions:
• 1. Single compound is
broken down into two or
more simpler products.
• 2. Usually requires energy.
• 3. Basic reaction: AB ---> A
+B
Decomposition Reaction
Examples:
• Metal Carbonate ----> Metal oxide + carbon dioxide
•
Ca CO3 ----> CaO + CO2
• Metal Hydroxide ----> metal oxide + water
•
Ca(OH)2 ---> CaO + H2O
• Metal Chlorate ---> metal chloride + oxygen
•
2KClO3 ---> 2 KCl + O2
• Oxyacid ---> nonmetal oxide + water
•
H2SO4 ---> SO3 + H2O
SINGLE REPLACEMENT
REACTION:
• 1. One element replaces a similar element
in a compound.
• 2. A reactive metal will replace any metal
that is less reactive (see pg 288 Activity
Series of Metals)
• 3. Nonmetal will replace other nonmetals.
Activity Series: Single
Replacement Reactions Only
• One metal will
only replace
another if it is
HIGHER on the
activity series
• This is because
it is a more
reactive metal
Single Replacement
cont.
• 4. Basic Reaction:
•
A + BC ---> AC + B
•
Y + BX ---> BY + X
Single Replacement
Examples:
• Replacement of a metal in a compound by a
more reactive metal
• Use activity series to determine if one metal is
strong enough to replace the other one. If not,
then no reaction will occur
•
2Al + 3Fe(NO3)2 ---> 3Fe + 2Al(NO3)3
• Replacement of Halogens.
•
Cl2 + 2 KBr ---> 2KCl + Br2
• Metal replacing hydrogen in an acid.
•
Zn + 2HCl ---> ZnCl2 + H2
DOUBLE-REPLACEMENT
REACTIONS:
• 1.Exchange of positive ions between two
compounds.
• 2.One compound formed is usually a precipitate,
gas, or a molecular compound (often water)
• 3. Basic Equation:
•
AB + CD ---> CB + AD
• 4. Use the solubility rules to determine whether or
not a reaction will take place
Double-Replacement
Examples:
• Metal oxide + acid ---> water + salt
(metal/nonmetal)
•
MgO + 2 Hcl ---> H2O + MgCl2
• Metal carbonate + acid ---> salt + carbon dioxide +water
•
CaCO3 + 2 HCl ---> Ca Cl2 + CO2 + H2O
• Acids + metal Hydroxide ---> salt + water
•
HCl + NaOH ----> NaCl + H2O
Solubility Rules
Overview
• List of rules used to determine whether or
not a reaction will take place
• Remember! In order for a reaction to take place
you must produce a gas or a precipitate from 2
liquids.
• Solubility rules tell us whether or not a
precipitate (solid) is produced
• You will often see these letters indicating
what state of matter a substance is
•
•
•
•
Solid (s)
Liquid (l)
Gas (g)
Aqueous (aq) = soluble in water
Double Rep. Reactions:
Solubility Rules (see handoutdo not have to copy down)
• 1. Soluble: All salts containing the
ammonium or Group IA ions (Li+, Na+, K+,
Rb+, Cs+)
• 2. Soluble: All salts containing nitrate
(NO3-), acetate (C2H3O2-), and perchlorate
(ClO4-)
• 3. Soluble: All salts containing Group VIIA
ions (Cl-, Br-, I-), except those in Rule 5.
• 4. Soluble: All salts containing sulfate
(SO4-2). Exceptions are barium sulfate,
calcium sulfate, lead II sulfate, and
strontium sulfate.
Solubility Rules cont.
• 5. Insoluble: All salts containing
silver ion (Ag+), lead II ions (Pb+2),
and mercury I ions (Hg2+2)
• 6. Insoluble: All salts containing
carbonate, chromates, hydroxides,
oxides, phosphates, and sulfides
• Exceptions:
• Group IIA chromates, except barium
chromate are solulbe
• Group IIA hydroxides, except
magnesium hydroxide, are soluble
COMBUSTION
REACTIONS:
• 1. Oxygen reacting with another substance.
• 2. Usually involves hydrocarbons (contain
hydrogen & carbon)
• 3. Heat is always released.
• 4. Basic Equation: CXHY + O2 ---> H2O + CO2
•
[x & y represent a ratio of carbon &
hydrogen]
Combustion Examples:
• 4. Complete combustion:
•
C3H8 + 5O2 ---> 3CO2 + 4H2O
• 5. Incomplete combustion:
creates carbon monoxide (CO),
carbon, & water. [products
cannot be predicted]
Catalysts
• A substance that increases
the rate of a chemical reaction
by lowering activation
energies but is not itself
consumed in the reaction.
• Example: Enzymes: allow many
chemical rxns to occur at a rate that
sustains life at normal living temperatures