Chapter 4 Chemical Quantities and Aqueous Reactions

1
Chapter 4Chemical Quantities and Aqueous
Reactions
Chemistry A Molecular Approach, 1st Ed.Nivaldo
Tro
Roy Kennedy Massachusetts Bay Community
College Wellesley Hills, MA
2008, Prentice Hall
2
Reaction Stoichiometry

• the numerical relationships between chemical
  amounts in a reaction is called stoichiometry
• the coefficients in a balanced chemical equation
  specify the relative amounts in moles of each of
  the substances involved in the reaction

2 C8H18(l) 25 O2(g) ? 16 CO2(g) 18 H2O(g) 2
molecules of C8H18 react with 25 molecules of
O2 to form 16 molecules of CO2 and 18 molecules
of H2O 2 moles of C8H18 react with 25 moles of
O2 to form 16 moles of CO2 and 18 moles of H2O 2
mol C8H18 25 mol O2 16 mol CO2 18 mol H2O
3
Predicting Amounts from Stoichiometry

• the amounts of any other substance in a chemical
  reaction can be determined from the amount of
  just one substance
• How much CO2 can be made from 22.0 moles of C8H18
  in the combustion of C8H18?
• 2 C8H18(l) 25 O2(g) ? 16 CO2(g) 18 H2O(g)
• 2 moles C8H18 16 moles CO2

4
Example Estimate the mass of CO2 produced in
2004 by the combustion of 3.4 x 1015 g gasoline

• assuming that gasoline is octane, C8H18, the
  equation for the reaction is
• 2 C8H18(l) 25 O2(g) ? 16 CO2(g) 18 H2O(g)
• the equation for the reaction gives the mole
  relationship between amount of C8H18 and CO2, but
  we need to know the mass relationship, so the
  Concept Plan will be

5
Practice

• According to the following equation, how many
  milliliters of water are made in the combustion
  of 9.0 g of glucose?
• C6H12O6(s) 6 O2(g) 6 CO2(g) 6 H2O(l)
• convert 9.0 g of glucose into moles (MM 180)
• convert moles of glucose into moles of water
• convert moles of water into grams (MM 18.02)
• convert grams of water into mL
• How? what is the relationship between mass and
  volume?

density of water 1.00 g/mL
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Limiting Reactant

• for reactions with multiple reactants, it is
  likely that one of the reactants will be
  completely used before the others
• when this reactant is used up, the reaction stops
  and no more product is made
• the reactant that limits the amount of product is
  called the limiting reactant
• sometimes called the limiting reagent
• the limiting reactant gets completely consumed
• reactants not completely consumed are called
  excess reactants
• the amount of product that can be made from the
  limiting reactant is called the theoretical yield

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Things Dont Always Go as Planned!

• many things can happen during the course of an
  experiment that cause the loss of product
• the amount of product that is made in a reaction
  is called the actual yield
• generally less than the theoretical yield, never
  more!
• the efficiency of product recovery is generally
  given as the percent yield

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Limiting and Excess Reactants in the Combustion
of Methane

• CH4(g) 2 O2(g) CO2(g) 2 H2O(g)
• Our balanced equation for the combustion of
  methane implies that every 1 molecule of CH4
  reacts with 2 molecules of O2

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Limiting and Excess Reactants in the Combustion
of Methane
CH4(g) 2 O2(g) ? CO2(g) 2 H2O(g)

• If we have 5 molecules of CH4 and 8 molecules of
  O2, which is the limiting reactant?

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Example 4.4Finding Limiting Reactant,
Theoretical Yield, and Percent Yield
11

• Example
• When 28.6 kg of C are allowed to react with 88.2
  kg of TiO2 in the reaction below, 42.8 kg of Ti
  are obtained. Find the Limiting Reactant,
  Theoretical Yield, and Percent Yield.

12
Practice How many grams of N2(g) can be made
from 9.05 g of NH3 reacting with 45.2 g of CuO?2
NH3(g) 3 CuO(s) ? N2(g) 3 Cu(s) 3 H2O(l)
13
Solutions

• when table salt is mixed with water, it seems to
  disappear, or become a liquid the mixture is
  homogeneous
• the salt is still there, as you can tell from the
  taste, or simply boiling away the water
• homogeneous mixtures are called solutions
• the component of the solution that changes state
  is called the solute
• the component that keeps its state is called the
  solvent
• if both components start in the same state, the
  major component is the solvent

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Describing Solutions

• since solutions are mixtures, the composition can
  vary from one sample to another
• pure substances have constant composition
• salt water samples from different seas or lakes
  have different amounts of salt
• so to describe solutions accurately, we must
  describe how much of each component is present
• we saw that with pure substances, we can describe
  them with a single name because all samples
  identical

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Solution Concentration

• qualitatively, solutions are often described as
  dilute or concentrated
• dilute solutions have a small amount of solute
  compared to solvent
• concentrated solutions have a large amount of
  solute compared to solvent
• quantitatively, the relative amount of solute in
  the solution is called the concentration

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Solution ConcentrationMolarity

• moles of solute per 1 liter of solution
• used because it describes how many molecules of
  solute in each liter of solution

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Preparing 1 L of a 1.00 M NaCl Solution
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Example 4.5 Find the molarity of a solution
that has 25.5 g KBr dissolved in 1.75 L of
solution
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Using Molarity in Calculations

• molarity shows the relationship between the moles
  of solute and liters of solution
• If a sugar solution concentration is 2.0 M, then
  1 liter of solution contains 2.0 moles of sugar
• 2 liters 4.0 moles sugar
• 0.5 liters 1.0 mole sugar
• 1 L solution 2 moles sugar

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Example 4.6 How many liters of 0.125 M NaOH
contains 0.255 mol NaOH?
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Dilution

• often, solutions are stored as concentrated stock
  solutions
• to make solutions of lower concentrations from
  these stock solutions, more solvent is added
• the amount of solute doesnt change, just the
  volume of solution
• moles solute in solution 1 moles solute in
  solution 2
• the concentrations and volumes of the stock and
  new solutions are inversely proportional
• M1V1 M2V2

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Example 4.7 To what volume should you dilute
0.200 L of 15.0 M NaOH to make 3.00 M NaOH?
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Solution Stoichiometry

• since molarity relates the moles of solute to the
  liters of solution, it can be used to convert
  between amount of reactants and/or products in a
  chemical reaction

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Example 4.8 What volume of 0.150 M KCl is
required to completely react with 0.150 L of
0.175 M Pb(NO3)2 in the reaction 2 KCl(aq)
Pb(NO3)2(aq) ? PbCl2(s) 2 KNO3(aq)
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What Happens When a Solute Dissolves?

• there are attractive forces between the solute
  particles holding them together
• there are also attractive forces between the
  solvent molecules
• when we mix the solute with the solvent, there
  are attractive forces between the solute
  particles and the solvent molecules
• if the attractions between solute and solvent are
  strong enough, the solute will dissolve

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Table Salt Dissolving in Water
Each ion is attracted to the surrounding water
molecules and pulled off and away from the crystal
When it enters the solution, the ion is
surrounded by water molecules, insulating it from
other ions
The result is a solution with free moving charged
particles able to conduct electricity
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Electrolytes and Nonelectrolytes

• materials that dissolve in water to form a
  solution that will conduct electricity are called
  electrolytes
• materials that dissolve in water to form a
  solution that will not conduct electricity are
  called nonelectrolytes

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Molecular View of Electrolytes and
Nonelectrolytes

• in order to conduct electricity, a material must
  have charged particles that are able to flow
• electrolyte solutions all contain ions dissolved
  in the water
• ionic compounds are electrolytes because they all
  dissociate into their ions when they dissolve
• nonelectrolyte solutions contain whole molecules
  dissolved in the water
• generally, molecular compounds do not ionize when
  they dissolve in water
• the notable exception being molecular acids

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Salt vs. Sugar Dissolved in Water
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Acids

• acids are molecular compounds that ionize when
  they dissolve in water
• the molecules are pulled apart by their
  attraction for the water
• when acids ionize, they form H cations and
  anions
• the percentage of molecules that ionize varies
  from one acid to another
• acids that ionize virtually 100 are called
  strong acids
• HCl(aq) ? H(aq) Cl-(aq)
• acids that only ionize a small percentage are
  called weak acids
• HF(aq) ? H(aq) F-(aq)

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Strong and Weak Electrolytes

• strong electrolytes are materials that dissolve
  completely as ions
• ionic compounds and strong acids
• their solutions conduct electricity well
• weak electrolytes are materials that dissolve
  mostly as molecules, but partially as ions
• weak acids
• their solutions conduct electricity, but not well
• when compounds containing a polyatomic ion
  dissolve, the polyatomic ion stays together
• Na2SO4(aq) ? 2 Na(aq) SO42-(aq)
• HC2H3O2(aq) ? H(aq) C2H3O2-(aq)

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Classes of Dissolved Materials
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Solubility of Ionic Compounds

• some ionic compounds, like NaCl, dissolve very
  well in water at room temperature
• other ionic compounds, like AgCl, dissolve hardly
  at all in water at room temperature
• compounds that dissolve in a solvent are said to
  be soluble, while those that do not are said to
  be insoluble
• NaCl is soluble in water, AgCl is insoluble in
  water
• the degree of solubility depends on the
  temperature
• even insoluble compounds dissolve, just not
  enough to be meaningful

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When Will a Salt Dissolve?

• Predicting whether a compound will dissolve in
  water is not easy
• The best way to do it is to do some experiments
  to test whether a compound will dissolve in
  water, then develop some rules based on those
  experimental results
• we call this method the empirical method

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Solubility RulesCompounds that Are Generally
Soluble in Water
Compounds Containing the Following Ions are Generally Soluble Exceptions (when combined with ions on the left the compound is insoluble)
Li, Na, K, NH4 none
NO3, C2H3O2 none
Cl, Br, I Ag, Hg22, Pb2
SO42 Ag, Ca2, Sr2, Ba2, Pb2
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Solubility RulesCompounds that Are Generally
Insoluble
Compounds Containing the Following Ions are Generally Insoluble Exceptions (when combined with ions on the left the compound is soluble or slightly soluble)
OH Li, Na, K, NH4, Ca2, Sr2, Ba2
S2 Li, Na, K, NH4, Ca2, Sr2, Ba2
CO32, PO43 Li, Na, K, NH4
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Precipitation Reactions

• reactions between aqueous solutions of ionic
  compounds that produce an ionic compound that is
  insoluble in water are called precipitation
  reactions and the insoluble product is called a
  precipitate

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2 KI(aq) Pb(NO3)2(aq) ? PbI2(s) 2 KNO3(aq)
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No Precipitate Formation No Reaction
KI(aq) NaCl(aq) ? KCl(aq) NaI(aq) all ions
still present, ? no reaction
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Process for Predicting the Products ofa
Precipitation Reaction

• Determine what ions each aqueous reactant has
• Determine formulas of possible products
• Exchange ions
• () ion from one reactant with (-) ion from other
• Balance charges of combined ions to get formula
  of each product
• Determine Solubility of Each Product in Water
• Use the solubility rules
• If product is insoluble or slightly soluble, it
  will precipitate
• If neither product will precipitate, write no
  reaction after the arrow

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Process for Predicting the Products ofa
Precipitation Reaction

1. If either product is insoluble, write the
   formulas for the products after the arrow
   writing (s) after the product that is insoluble
   and will precipitate, and (aq) after products
   that are soluble and will not precipitate
2. Balance the equation

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Example 4.10 Write the equation for the
precipitation reaction between an aqueous
solution of potassium carbonate and an aqueous
solution of nickel(II) chloride

• Write the formulas of the reactants
• K2CO3(aq) NiCl2(aq) ?
• Determine the possible products
• Determine the ions present
• (K CO32-) (Ni2 Cl-) ?
• Exchange the Ions
• (K CO32-) (Ni2 Cl-) ? (K Cl-) (Ni2
  CO32-)
• Write the formulas of the products
• cross charges and reduce
• K2CO3(aq) NiCl2(aq) ? KCl NiCO3

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Example 4.10 Write the equation for the
precipitation reaction between an aqueous
solution of potassium carbonate and an aqueous
solution of nickel(II) chloride

• Determine the solubility of each product
• KCl is soluble
• NiCO3 is insoluble
• If both products soluble, write no reaction
• does not apply since NiCO3 is insoluble

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Example 4.10 Write the equation for the
precipitation reaction between an aqueous
solution of potassium carbonate and an aqueous
solution of nickel(II) chloride

• Write (aq) next to soluble products and (s) next
  to insoluble products
• K2CO3(aq) NiCl2(aq) ? KCl(aq) NiCO3(s)
• Balance the Equation
• K2CO3(aq) NiCl2(aq) ? 2 KCl(aq) NiCO3(s)

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Ionic Equations

• equations which describe the chemicals put into
  the water and the product molecules are called
  molecular equations
• 2 KOH(aq) Mg(NO3)2(aq) 2 KNO3(aq)
  Mg(OH)2(s)
• equations which describe the actual dissolved
  species are called complete ionic equations
• aqueous strong electrolytes are written as ions
• soluble salts, strong acids, strong bases
• insoluble substances, weak electrolytes, and
  nonelectrolytes written in molecule form
• solids, liquids, and gases are not dissolved,
  therefore molecule form
• 2K1(aq) 2OH-1(aq) Mg2(aq) 2NO3-1(aq)
  2K1(aq) 2NO3-1(aq) Mg(OH)2(s)

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Ionic Equations

• ions that are both reactants and products are
  called spectator ions
• 2K1(aq) 2OH-1(aq) Mg2(aq) 2NO3-1(aq)
  2K1(aq) 2NO3-1(aq) Mg(OH)2(s)

• an ionic equation in which the spectator ions
  are
• removed is called a net ionic equation
• 2OH-1(aq) Mg2(aq) Mg(OH)2(s)

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Acid-Base Reactions

• also called neutralization reactions because the
  acid and base neutralize each others properties
• 2 HNO3(aq) Ca(OH)2(aq) ? Ca(NO3)2(aq) 2
  H2O(l)
• the net ionic equation for an acid-base reaction
  is
• H(aq) OH?(aq) ? H2O(l)
• as long as the salt that forms is soluble in water

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Acids and Bases in Solution

• acids ionize in water to form H ions
• more precisely, the H from the acid molecule is
  donated to a water molecule to form hydronium
  ion, H3O
• most chemists use H and H3O interchangeably
• bases dissociate in water to form OH? ions
• bases, like NH3, that do not contain OH? ions,
  produce OH? by pulling H off water molecules
• in the reaction of an acid with a base, the H
  from the acid combines with the OH? from the base
  to make water
• the cation from the base combines with the anion
  from the acid to make the salt
• acid base ???salt water

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Common Acids
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Common Bases
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HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)
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Example - Write the molecular, ionic, and
net-ionic equation for the reaction of aqueous
nitric acid with aqueous calcium hydroxide

• Write the formulas of the reactants
• HNO3(aq) Ca(OH)2(aq) ?
• Determine the possible products
• Determine the ions present when each reactant
  dissociates
• (H NO3-) (Ca2 OH-) ?
• Exchange the ions, H1 combines with OH-1 to make
  H2O(l)
• (H NO3-) (Ca2 OH-) ? (Ca2 NO3-)
  H2O(l)
• Write the formula of the salt
• cross the charges
• (H NO3-) (Ca2 OH-) ? Ca(NO3)2 H2O(l)

53
Example - Write the molecular, ionic, and
net-ionic equation for the reaction of aqueous
nitric acid with aqueous calcium hydroxide

• Determine the solubility of the salt
• Ca(NO3)2 is soluble
• Write an (s) after the insoluble products and a
  (aq) after the soluble products
• HNO3(aq) Ca(OH)2(aq) ? Ca(NO3)2(aq) H2O(l)
• Balance the equation
• 2 HNO3(aq) Ca(OH)2(aq) ? Ca(NO3)2(aq) 2
  H2O(l)

54
Example - Write the molecular, ionic, and
net-ionic equation for the reaction of aqueous
nitric acid with aqueous calcium hydroxide

• Dissociate all aqueous strong electrolytes to get
  complete ionic equation
• not H2O
• 2 H(aq) 2 NO3-(aq) Ca2(aq) 2 OH-(aq) ?
  Ca2(aq) 2 NO3-(aq) H2O(l)
• Eliminate spectator ions to get net-ionic
  equation
• 2 H1(aq) 2 OH-1(aq) ? 2 H2O(l)
• H1(aq) OH-1(aq) ? H2O(l)

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Titration

• often in the lab, a solutions concentration is
  determined by reacting it with another material
  and using stoichiometry this process is called
  titration
• in the titration, the unknown solution is added
  to a known amount of another reactant until the
  reaction is just completed, at this point, called
  the endpoint, the reactants are in their
  stoichiometric ratio
• the unknown solution is added slowly from an
  instrument called a burette
• a long glass tube with precise volume markings
  that allows small additions of solution

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Acid-Base Titrations

• the difficulty is determining when there has been
  just enough titrant added to complete the
  reaction
• the titrant is the solution in the burette
• in acid-base titrations, because both the
  reactant and product solutions are colorless, a
  chemical is added that changes color when the
  solution undergoes large changes in
  acidity/alkalinity
• the chemical is called an indicator
• at the endpoint of an acid-base titration, the
  number of moles of H equals the number of moles
  of OH?
• aka the equivalence point

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Titration
58
Titration
The base solution is the titrant in the burette.
As the base is added to the acid, the H reacts
with the OH to form water. But there is still
excess acid present so the color does not change.
At the titrations endpoint, just enough base has
been added to neutralize all the acid. At this
point the indicator changes color.
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Example 4.14The titration of 10.00 mL of HCl
solution of unknown concentration requires 12.54
mL of 0.100 M NaOH solution to reach the end
point. What is the concentration of the unknown
HCl solution?
60
Gas Evolving Reactions

• Some reactions form a gas directly from the ion
  exchange
• K2S(aq) H2SO4(aq) ? K2SO4(aq) H2S(g)
• Other reactions form a gas by the decomposition
  of one of the ion exchange products into a gas
  and water
• K2SO3(aq) H2SO4(aq) ? K2SO4(aq) H2SO3(aq)
• H2SO3 ? H2O(l) SO2(g)

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NaHCO3(aq) HCl(aq) ? NaCl(aq) CO2(g) H2O(l)
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Compounds that UndergoGas Evolving Reactions
Reactant Type Reacting With Ion Exchange Product Decom-pose? Gas Formed Example
metalnS, metal HS acid H2S no H2S K2S(aq) 2HCl(aq) ? 2KCl(aq) H2S(g)
metalnCO3, metal HCO3 acid H2CO3 yes CO2 K2CO3(aq) 2HCl(aq) ? 2KCl(aq) CO2(g) H2O(l)
metalnSO3 metal HSO3 acid H2SO3 yes SO2 K2SO3(aq) 2HCl(aq) ? 2KCl(aq) SO2(g) H2O(l)
(NH4)nanion base NH4OH yes NH3 KOH(aq) NH4Cl(aq) ? KCl(aq) NH3(g) H2O(l)
63
Example 4.15 - When an aqueous solution of sodium
carbonate is added to an aqueous solution of
nitric acid, a gas evolves

• Write the formulas of the reactants
• Na2CO3(aq) HNO3(aq) ?
• Determine the possible products
• Determine the ions present when each reactant
  dissociates
• (Na1 CO3-2) (H1 NO3-1) ?
• Exchange the anions
• (Na1 CO3-2) (H1 NO3-1) ? (Na1 NO3-1)
  (H1 CO3-2)
• Write the formula of compounds
• cross the charges
• Na2CO3(aq) HNO3(aq) ? NaNO3 H2CO3

64
Example 4.15 - When an aqueous solution of sodium
carbonate is added to an aqueous solution of
nitric acid, a gas evolves

• Check to see either product H2S - No
• Check to see if either product decomposes Yes
• H2CO3 decomposes into CO2(g) H2O(l)
• Na2CO3(aq) HNO3(aq) ? NaNO3 CO2(g) H2O(l)

65
Example 4.15 - When an aqueous solution of sodium
carbonate is added to an aqueous solution of
nitric acid, a gas evolves

• Determine the solubility of other product
• NaNO3 is soluble
• Write an (s) after the insoluble products and a
  (aq) after the soluble products
• Na2CO3(aq) 2 HNO3(aq) ? 2 NaNO3(aq) CO2(g)
  H2O(l)
• Balance the equation
• Na2CO3(aq) 2 HNO3(aq) ? 2 NaNO3 CO2(g)
  H2O(l)

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Other Patterns in Reactions

• the precipitation, acid-base, and gas evolving
  reactions all involved exchanging the ions in the
  solution
• other kinds of reactions involve transferring
  electrons from one atom to another these are
  called oxidation-reduction reactions
• also known as redox reactions
• many involve the reaction of a substance with
  O2(g)
• 4 Fe(s) 3 O2(g) ? 2 Fe2O3(s)

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Combustion as Redox2 H2(g) O2(g) ? 2 H2O(g)
68
Redox without Combustion2 Na(s) Cl2(g) ? 2
NaCl(s)
2 Na ? 2 Na 2 e?
Cl2 2 e? ? 2 Cl?
69
Reactions of Metals with Nonmetals

• consider the following reactions
• 4 Na(s) O2(g) ? 2 Na2O(s)
• 2 Na(s) Cl2(g) ? 2 NaCl(s)
• the reaction involves a metal reacting with a
  nonmetal
• in addition, both reactions involve the
  conversion of free elements into ions
• 4 Na(s) O2(g) ? 2 Na2O (s)
• 2 Na(s) Cl2(g) ? 2 NaCl(s)

70
Oxidation and Reduction

• in order to convert a free element into an ion,
  the atoms must gain or lose electrons
• of course, if one atom loses electrons, another
  must accept them
• reactions where electrons are transferred from
  one atom to another are redox reactions
• atoms that lose electrons are being oxidized,
  atoms that gain electrons are being reduced

2 Na(s) Cl2(g) ? 2 NaCl(s) Na ? Na 1 e
oxidation Cl2 2 e ? 2 Cl reduction
71
Electron Bookkeeping

• for reactions that are not metal nonmetal, or
  do not involve O2, we need a method for
  determining how the electrons are transferred
• chemists assign a number to each element in a
  reaction called an oxidation state that allows
  them to determine the electron flow in the
  reaction
• even though they look like them, oxidation states
  are not ion charges!
• oxidation states are imaginary charges assigned
  based on a set of rules
• ion charges are real, measurable charges

72
Rules for Assigning Oxidation States

• rules are in order of priority
• free elements have an oxidation state 0
• Na 0 and Cl2 0 in 2 Na(s) Cl2(g)
• monatomic ions have an oxidation state equal to
  their charge
• Na 1 and Cl -1 in NaCl
• (a) the sum of the oxidation states of all the
  atoms in a compound is 0
• Na 1 and Cl -1 in NaCl, (1) (-1) 0

73
Rules for Assigning Oxidation States

• (b) the sum of the oxidation states of all the
  atoms in a polyatomic ion equals the charge on
  the ion
• N 5 and O -2 in NO3, (5) 3(-2) -1
• (a) Group I metals have an oxidation state of 1
  in all their compounds
• Na 1 in NaCl
• (b) Group II metals have an oxidation state of
  2 in all their compounds
• Mg 2 in MgCl2

74
Rules for Assigning Oxidation States

• in their compounds, nonmetals have oxidation
  states according to the table below
• nonmetals higher on the table take priority

Nonmetal Oxidation State Example
F -1 CF4
H 1 CH4
O -2 CO2
Group 7A -1 CCl4
Group 6A -2 CS2
Group 5A -3 NH3
75
Practice Assign an Oxidation State to Each
Element in the following

• Br2
• K
• LiF
• CO2
• SO42-
• Na2O2

76
Oxidation and ReductionAnother Definition

• oxidation occurs when an atoms oxidation state
  increases during a reaction
• reduction occurs when an atoms oxidation state
  decreases during a reaction

CH4 2 O2 ? CO2 2 H2O -4 1
0 4 2 1 -2
77
OxidationReduction

• oxidation and reduction must occur simultaneously
  
• if an atom loses electrons another atom must take
  them
• the reactant that reduces an element in another
  reactant is called the reducing agent
• the reducing agent contains the element that is
  oxidized
• the reactant that oxidizes an element in another
  reactant is called the oxidizing agent
• the oxidizing agent contains the element that is
  reduced

2 Na(s) Cl2(g) ? 2 NaCl(s) Na is oxidized, Cl
is reduced Na is the reducing agent, Cl2 is the
oxidizing agent
78
Identify the Oxidizing and Reducing Agents in
Each of the Following

• 3 H2S 2 NO3 2 H 3 S 2 NO 4 H2O
• MnO2 4 HBr MnBr2 Br2 2 H2O

79
Combustion Reactions

• Reactions in which O2(g) is a reactant are called
  combustion reactions
• Combustion reactions release lots of energy
• Combustion reactions are a subclass of
  oxidation-reduction reactions

2 C8H18(g) 25 O2(g) ? 16 CO2(g) 18 H2O(g)
80
Combustion Products

• to predict the products of a combustion reaction,
  combine each element in the other reactant with
  oxygen

Reactant Combustion Product
contains C CO2(g)
contains H H2O(g)
contains S SO2(g)
contains N NO(g) or NO2(g)
contains metal M2On(s)
81
Practice Complete the Reactions

• combustion of C3H7OH(l)
• combustion of CH3NH2(g)