Chapter 11
Chemical Reactions
Chemistry 2
Describing Chemical Reactions 11.1
Writing Chemical Equations 11.1
• Reactants  products
▫  = yields, gives, or reacts to produce
• Iron + Oxygen  iron(III) oxide
• Hydrogen peroxide  water + oxygen
▫ Hydrogen peroxide decomposes to form water and oxygen gas
• Skeleton equation – chemical equation that does not indicate
the relative amount of the reactant and products
• Fe(s) + O2(g)  Fe2O3(s)
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(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous (dissolved in water)
• Catalyst – substance that speeds up the reaction but is not used
up in the reaction
▫ Formula is written above the arrow
• Practice Problems Page 324 # 1-2
Balancing Chemical Equations 11.1
• Coefficients – small whole #s placed in front of
formulas to balance
• Each side of equation should have same # of each type
of atom
• Law of Conservatio nof mass: mass is neither created
or destroyed in chemical reaction
▫ Atoms rearranged
▫ Bonds broken and bond formed
Here are some practice problems.
1. __NaCl + __BeF2  __NaF + __BeCl2
2. __FeCl3 + __Be3(PO4)2  __BeCl2 + __FePO4
3. __AgNO3 + __LiOH  __AgOH + __LiNO3
4. __CH4 + __O2  __CO2 + __H2O
5. __Mg + __Mn2O3  __MgO + __Mn
Practice Problems: Page 327 # 3 – 4, Page 328 #5-6
Types of Chemical Reactions 11.2
Classifying Reactions 11.2
• 5 types of reactions
1. Combination or Synthesis – 2 or mores substances
react to form a single new substance
 Group A metal + nonmetal = 2K + Cl2  2KCl
 2 nonmetals can have more than 1 product
 S + O2  SO2
 2S + 3O2  2SO3
 Transition metal and nonmetal can have more than 1
product
 Fe + S  FeS
 2Fe + 3S  Fe2S3
 Practice Problems page 331 # 13 - 14
2. Decomposition Reactions – chemical change in which a
single compound breaks down into two or more simpler
products
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Opposite of synthesis
2HgO  2Hg + O2
Difficult to predict product
Most required energy input (endothermic)
Practice Problems page 332 # 15 – 16
3. Single-Replacement Reaction – chemical change in
which one element replaces a second element in a
compound
▫ 2K + 2H2O  2KOH + H2
▫ Both reactant and product contain of an element and a
compound
▫ Activity series – list metals in order of decreasing reactivity =
table 11.2 page 333
 Halogens = reactivity decreases down column
 Br2 + NaI  NaBr + I2
 Br2 + NaCl  no reaction
• Metals from Li to Na will replace H from acids and
water
• Metals from Mg to Pb will replace H from acids only
• Practice Problems Page 334 # 17
4. Double Replacement Reactions – a chemical change
involving an exchange of positive ions bvetween 2
compounds
▫ To occur, one of the following is usually true:
 1 of produces is slightly soluble and precipitates from
solution
 Na2S(aq) + CD(NO3)(aq)  CdS(s) + 2NaNO3(aq)
▫ CdS precipitated out
 One of products is a gas
 One product is a molecular compound
▫ Practice Problems Page 335 # 18 - 19
5. Combustion Reaction – chemical change in which an
element or a compound reacts with oxygen often
producing energy in the form of heat and light
▫ ALWAYS involves oxygen
▫ Often uses hydrocarbons
 Complete combustion form CO2 and water (and energy)
 Incomplete combustion forms CO
 Supply of oxygen is limited
▫ Practice Problems Page 337 # 20 - 21
Predicting the Products of a Chemical
Reaction 11.2
• # of elements/compounds reacting is a good indicator
of possible reaction type
▫ Combination = 2 or more reactants  single product
▫ Decomposition = single compound  2 or more
substances
▫ Single-replacement = element + compound  element +
compound
▫ Double-Replacement = 2 ionic compounds  2 new
compounds
▫ Combustion = Oxygen + hydrocarbon (usually)  water
and carbon dioxide
Reactions in Aqueous Solution 11.3
Net Ionic Equations 11.3
• Many of chemical reactions take place in water (aqueous
solution)
▫ 70% Earth surface covered by water
▫ 66% of human body is water
• Complete Ionic equation – an equation that shows
dissolved ionic compounds as dissociated free ions
▫ 2Na1+(aq) + SO42-(aq) + Ba2+(aq) + 2Cl1-(aq)  2Na1+(aq) +
2Cl1-(aq) + BaSO4(s)
 Cross out ions that appear unchanged on both sides =
spectator ions
▫ Na1+(aq) + SO42-(aq) + Ba2+(aq) + 2Cl1-(aq)  2Na1+(aq) +
2Cl1-(aq) + BaSO4(s)
 Write the net ionic equation
▫ Ba2+(aq) + SO42-(aq)  BaSO4(s)
 Then balance
Predicting the Formation of a
Precipitate 11.3
• Mixing ionic compounds can sometimes form a
precipitate (insoluble salt)
• Solubility Rules for Ionic Compounds
▫ Salts of alkali metals (1A) and ammonia (NH4)+= Soluble
▫ Nitrate (NO3)- salts and chlorate (ClO3)- salts = Soluble
▫ Sulfate (SO4)2- slats except compounds with Pb2+, Ag+,
Hg22+, Ba2+, Sr2+, and Ca2+ = Soluble
▫ Chloride salts, except compound with Ag+, Pb2+, and
Hg22+ = Soluble
▫ Carbonates (CO3)2-, Phosphates (PO4)3-,
Chromates(CrO4)2-, Sulfides, and Hydroxides (OH)-
• Precipitate example:
▫ Na2Co3 (aq) + Ba(NO3)2(aq)  ?? Precipitate formed
 Separate = 2Na+(aq) + CO32- (aq) + Ba2+ (aq) + 2NO3-(aq)
 Would form NaNO3 and BaCO3
▫ Na = Alkali = soluble
▫ Nitrate salts = soluble Carbonates generally insoluble = BaCO3 will
precipitate out
▫ Ba2+ (aq) + CO32- (aq)  BaCO3 (s)
• Practice Problems Page 343 # 28 - 29