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CHEMICAL REACTIONS

• 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

Temperature change: endothermic (colder), exothermic (hotter)

Color Change Odor Gas Produced (bubbles) Precipitate: formed from 2 liquids

Balancing Equations Steps1) 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 (base)

• CaO + H2O ---> Ca(OH)2

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 handout- do 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