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Chapter 7 Chemical Reactions
7.0 Introduction
 Chemical changes always involve the formation of new substances with properties that are different from the original
substances.
 These chemical changes are what drives our whole existence
 We use chemical equations to describe these chemical changes
7.1 Chemical Reactions
 Chemical reactions in our life (digestion, photosynthesis, car battery, cooking, burning, rusting)
 In all chemical reactions the starting substances (reactants) are changed to the end substances (products)
 Reactants yield Products (an arrow  is used to represent yields)
 As the reactants are converted to products, bonds that hold atoms together are broken and new bonds are formed
(energy transfer)
 The atoms of the reactants are neither created or destroyed they are just rearranged to form new products (LCM)
 Word equations are written to describe what takes place in a chemical reaction
 These equations must show that the reactants are equal to the products
 Write word equation for the following:
1. Dihydrogen dioxide is broken down to form dihydrogen monoxide and oxygen
 2. Methane undergoes complete combustion
Combustion means oxygen in reactants – complete combustion of hydrocarbon means water and carbon dioxide are
formed
3. Iron (III) oxide is formed when Iron is oxidized.
4. pure copper and sulfur dioxide are produced by heating copper( II) sulfide in the presence of oxygen
5, Water is formed by the explosive reaction between hydrogen and oxygen
7.2 Writing Chemical Equations
 In chemistry to communicate more effectively chemical formulas are used to write chemical equations
 Example of equation: Fe + O2  Fe2O3 (all diatomics must have 2)
 The above equation is called a skeleton equation because it does not indicate the relative amounts of the reactants
and the products
 Physical states of substances in equations are indicated by: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution
(dissolved in water) –
 Some equations show physical states.
 Comes out as gas (↑) comes out as precipitate (↓)
 Fe (s) + O2 (g)  Fe2O3 (s)
 In many chemical reactions catalyst are used
 Catalyst are substances that speed up a reaction without being used up in the reaction (catalyst are written above the
arrow in an equation)
MnO2
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H2O2 (aq)  H2O(l) + O2 (g)
Write skeleton reactions for these equations
1.When calcium carbonate is heated in the absence of oxygen, calcium oxide (lime) and carbon dioxide are formed
2. Solid sulfur burns in oxygen gas to form sulfur dioxide gas
3.Oxygen gas can be made by heating potassium chlorate in the presence of the catalyst manganese (IV) oxide.
Potassium chloride is left as a solid residue.
4.When solid mercury (II) sulfide is heated in the presence of oxygen, liquid mercury metal and gaseous sulfur dioxide
are produced.
7.3 and 7.4 Balancing Equations
 All equations must be balanced which means that the reactants must be equal to the products (LCM)
 In chemical equations balance means each side of the equation has the same number of atoms of each element –
means same amount mass each side
 Coefficients are numbers that are placed in front of the reactants and the products to balance an equation (multiply
times subscript through the formula)
 The subscripts indicate how many atoms of an element are present in a formula and the coefficients indicate how
many molecules or formula units are present in the equation
 When the coefficient is one you do not write it in the equation
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Rules for Balancing equations
Write equation with correct formulas for reactants and products
Count the number of each atom on each side of the equation (polyatomic unchanged count as one unit)
Balance the elements one at a time by using coefficients, start with elements that appear only once on each side of
equation (Can not change subscripts)
 Odd – even principle you can make odds even, but can not make evens odd
4. Check to make sure all atoms and polyatomics are balanced
5. Make sure all the coefficients are in the lowest possible ratio
6. Write the ratio
 Write balanced equations for the reactions below:
1. Water is formed when hydrogen react with oxygen
2. Pure silver can be removed from silver nitrate in aqueous solution if copper reacts with the (aq) solution; copper
(II) nitrate will also be formed as a byproduct.
3. Aluminum and oxygen will combine to form aluminum oxide.
4. When sodium chloride is mixed with sulfuric acid sodium sulfate and hydrochloric acid are formed
7.5 Types of Reactions
 The only way to be certain what the products of a chemical reaction are is to carry out the reaction in the laboratory
 There are millions of compounds that will produce endless chemical reactions, therefore not all chemical reactions
can be carried out in the laboratory
 A system is used to classify chemical reactions, which allows chemist to recognize patterns and predict the products
of reactions
 One of these systems classifies chemical reactions into five general reaction types
 Some reactions fit perfectly into a group, but some fit into more than one group
7.6 Combination Reactions also known as Synthesis Reaction
 In a combination reaction two or more reactants combine to form a single product
 Reactants can be elements or compounds the product must be a compound
 Some common combination reactions
 Two nonmetals reacting or transition metals reacting with a nonmetal can produce more than one product (LMP)
 S (s) + O2 (g)  SO2(g)
or
2S (s) + 3O2 (g)  2SO3(g)
 Group A metals and nonmetals react to form metal cation and nonmetal anion

2K (s) + Cl2 (g)  2KCl (s)
 Some nonmetal oxides react with water to form acids – acids are compounds that produce hydrogen ions in aqueous
solution
 SO2(g) + H2O(l)  H2 SO3 (aq) (hydrogen sulfonate) (wetting agent detergent)
 Metallic oxides react with water to form a base – bases contain OH ions
 CaO (s) + H2O (l)  Ca (OH)2 (aq)
Aluminum combines with oxygen to
Al (s) + O2 (g) → Al2O3 (s)
4 Al (s) + 3O2 (g) → 2Al2O3 (s) combination 4:3 → 2
Complete and balance 2 possible equations that finish this reaction copper and sulfur can combine to form:
Magnesium and oxygen form:
Phosphorous and oxygen combine to form diphosphorus pentoxide
7.7 Decomposition Reactions
 In a decomposition reaction a single reactant is broken down (decomposed) into two or more products
 Most decomposition reactions require energy in the form of heat, light, or electricity
 Calcium carbonate is heated and decomposed into calcium oxide (lime) and carbon dioxide
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CaCO3(s)  CaO (s)+ CO2 (g)
∆ heated
 When Mercury (II) Oxide is heated it decomposes
2HgO(s)  2Hg (l) + O2 (g) 2  2:1
 Finish, write, and balance the following decomposition reactions:
 Hydrogen Bromide decomposes to form:
 Sodium chloride decomposes to form:
 When heated silver oxide will produce:
 When nickel (II) carbonate is heated it produces: (ratio is 11:1)
7.8 Single-Replacement Reactions AKA as Single Displacement Reactions
 In a single-replacement reaction atoms of an element replace the atoms of a second element in a compound also called
single-displacement
 Some metals will replace or displace other metals in a compound
 The activity series of metal list metals in order of decreasing reactivity (ion sheet)
 A reactive metal will replace or displace any metal below it in the activity series, but no reaction occurs if the
metal is above it in the activity series
 Mg is above Zn so: Mg (s) + Zn (NO3)2 (aq)  Mg (NO3)2 (aq) + Zn (s)
 Mg is below Li so:
Mg (s) + LiNO3 (aq)  no reaction
 Add Au after Ag –
 Metals from Li to Na will replace (H) in water and acids
 Metals from Mg to Pb will only replace H in acids
 Usually the only nonmetals that replace each other are the Halogens (F2, Cl2, Br2, and I2 (the activity of the
halogens decreases as you go down group 7A)
 Zinc is dropped into sulfuric acid
Zn (s) + H2SO4 (aq)  ZnSO4 (aq) + H2(g) 1:11:1
Finish the single replacement reactions:
Sodium is dropped into water: H2O can be written HOH
Silver is dropped into an (ag) solution of potassium nitrate
20g of Aluminum is dropped into 300 ml of sulfuric acid
7.9 Double-Replacement Reactions AKA Double Displacement
 If solutions of two ionic compounds are mixed together, either an aqueous homogeneous mixture will form or a
chemical reaction will occur
 Double-replacement reactions involve the cations of one compound trading places with the cations of another
compound
 Double-replacement reactions generally occur between two ionic compounds in aqueous solution
 One of the products usually comes out of solution either as a gas, a solid precipitate, or water
 Na2S (aq) + Cd(NO3)2 (aq)  CdS (s) + 2NaNO3 (aq)
 2NaCN (aq) + H2SO4 (aq)  2HCN (g) + Na2 SO4 (aq)
 Ca(OH)2 (aq) +2HCl (aq)  CaCl2 (aq) + 2H2 O (l)
Barium chloride is placed in potassium carbonate
BaCl2 (aq) + K2CO3 (aq)  BaCO3 (s) + 2KCl (aq) 1:11:2
Complete the following double replacement reactions:
Iron (II) sulfide is dropped into hydrochloric acid
Potassium hydroxide (aq) is combined with phosphoric acid
7.9 Combustion Reactions
 In a combustion reaction an element or a compound reacts with oxygen, often producing energy in the form of heat
and light (fuels)
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Many combustion reactions involve hydrocarbons (made of H and C)
Complete combustion of hydrocarbons produces carbon dioxide, water, heat and light
Incomplete burning produces harmful carbon monoxide and water
Combustion reactions can also be replacement and combination reactions
CH4 (q) + 2O2 (q)  CO2 (q) + 2H2O(l) + heat + light
2Mg (s) + O2 (g)  2MgO (s)
 Write a balanced equation for the complete combustion of hexacarbon
hexahydride (Benzene) C6H6
 Write a balanced equation for the incomplete combustion of Benzene
 Write a balanced equation for the complete combustion of methanol (CH3OH)
 Write a balanced equation for the incomplete combustion of octane (C8H18)
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7.11 Predicting Products of a Chemical Reaction
 Summary of the 5 types of equations (page 197)
 Another very important type of reaction is a reduction, or redox reaction in a reduction reaction Oxygen is lost from the
substance
 Page 198 problems 26 and 27
7.12 Net Ionic Equations
 Most ionic compounds dissociate, or separate, into cations and anions when they dissolve in water
 AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)
 To represent this a complete ionic equation is written with ions
 Ag+ (aq) + NO3- (aq) + Na+(aq) + Cl- (aq)→ AgCl (s) + Na+ (aq) + NO3-(aq)
 Ions that are not directly involved in the reaction appear on both sides of the equation unchanged – they are called
spectator ions - NO3-(aq) & Na+(aq)
 By crossing out the spectator ions and rewriting the equation leaving out the spectator ions a net ionic equation is
written
 A net ionic equation indicates only the substances that actually take part in the equation
 Ag+(aq) + NO3- (aq) + Na+ (aq) + Cl- (aq)→ AgCl (s) + Na+ (aq) + NO3-(aq)
 Net Ionic Equation Ag +(aq) + Cl- (aq) → AgCl (s)
 Net ionic equations are used to simplify and summarize chemical equations
 Electric charges must be balanced on each side of a net ionic equation each side equal or one side cancels out – also
atoms must be equal
 Write net ionic equation for the reaction that produces Ag and lead (II) nitrate
 Of the five types of reactions both single and double replacement reactions can be written as net ionic equations
(page 200)
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