Solutions

1

• Solutions
• Solution Composition
• A solution is a homogeneous mixture
• Solutions are prepared by dissolving components
  in a solvent
• The solvent is usually the component of the
  solution in highest quantity
• When solids are dissolved in a liquid, the liquid
  is the solvent
• If a liquid is dissolved in another liquid, the
  one in highest amount is the solvent
• When water is one of the components, it is the
  solvent regardless of how much is present
• The solutes are the other components in a
  solution they are what is dissolved
• Water is a most important solvent
• Dissolves a wide variety of substances
• Important in biological process

2
Solutions Concentration units for
stoichiometry Molarity(M) Example What is
the molarity of a solution made by dissolving
14.2 g NaCl in enough water to make 225 mL of
solution? Step 1 Find the number of moles of
solute, NaCl Step 2 Find the volume of solution
in liters Step 3 Find the molarity by
calculating the moleliter ratio
3

• Solutions
• Molarity, volume and moles
• From the definition of molarity the equation can
  be rearranged go calculate moles from molarity
  and volume of solution
• moles molarity x volume
• Example calculate the number of moles NaCl in
  31.5 mL 1.08 M NaCl
• moles NaCl
• From the definition of molarity the equation can
  be rearranged to calculate the volume from
  molarity and moles
• Example calculate the volume of 1.08 M NaCl
  solution that contains 2.00 mol NaCl

4

• Solutions
• Properties of Aqueous Solutions
• Ionic compounds dissolve by dissociating into
  ions
• NaCl(s) ?H2O Na(aq) Cl-(aq)
• Water molecules have a strong affinity for ions
• The end is attracted to anions the - end is
  attracted to cations
• This attraction overcomes the attraction of the
  cations and anions for each other
• Concentration of ions in solution
• A 1.00 M solution of KBr is 1.00 M in K and 1.00
  M in Br-
• A 1.00 M solution of K3PO4 is 3.00 M in K and
  1.00 M in PO43-

Water molecules have a - end water molecules
have a end
d-
electric dipole
d
5

• Solutions
• Ionic Compounds in Solution
• Ions in solution conduct electricity
• Electrolytes are ionic substances whose aqueous
  solutions conduct electricity.
• Non-electrolytes are substances whose aqueous
  solutions do not conduct electricity.
• Most Molecular compounds when dissolved to not
  produce ions
• Some important exceptions occur when molecules
  interact with water to produce ions
• HCl(aq) H(aq) Cl-(aq) there are
  no HCl molecules in water
• NH3(aq) H2O NH4(aq) OH-(aq)
• Strong Electrolytes are substances that dissolve
  by forming electrolytes.
• Almost all ionic compounds and some molecular
  compounds are strong electrolytes
• Some ionic compounds are not very soluble, yet,
  because the dissolve by forming electrolytes,
  they are strong electrolytes even though the
  concentration of ions in solution is small.
• HCl is a strong electrolyte in water because it
  completely ionizes when it dissolves

6

• Solutions
• Weak Electrolytes are substances that dissolve
  but do not form many ions
• Weak electrolyte solutions do not conduct
  electricity very well
• Example acetic acid, HC2H3O2
• Acetic acid is very soluble in water
• A 1 M solution of acetic acid is only about 1
  ionized
• HC2H3O2(aq) H(aq) C2H3O2-(aq)
• The double arrow represents and equilibrium the
  reaction is occurring in both directions
• A single arrow implies the reaction is complete
  in the direction of the arrow
• Acids, Bases and Salts
• Acids are substances that produce H in aqueous
  solution
• Examples HCl, HC2H3O2, H2SO4, H3PO4
• HCl, HC2H3O2 are monoprotic acids they produce
  only 1 H
• H2SO4 is a diprotic acid it produces 2 H, but
• H2SO4(aq) H(aq) HSO4-(aq)
• HSO4(aq) H(aq) SO42-(aq)

7

• Solutions
• Bases are substances that react with H
• OH- is a base H(aq) OH-(aq) H2O(l )
• Metal hydroxides are common bases NaOH, KOH,
  Ca(OH)2. Ba(OH)2
• Some non-hydroxides produce OH-(aq) NH3
• NH3(aq) H2O(l ) NH4(aq) OH-(aq)
• Strong and Weak Acids and Bases
• Strong acids and bases are acids or bases that
  are strong electrolytes
• Examples Acids HCl, HBr, HI, HNO3, HClO4,
  H2SO4
• Bases Group 2A hydroxides, Ca(OH)2, Sr(OH)2,
  Ba(OH)2
• Weak acids and bases are acids or bases that are
  weak electrolytes
• Figure 4.8, p. 118, can help assign strength of
  acid or base
• Strong acids are more reactive when the
  reactivity depends on the concentration of H
  sometimes its the anion
• HCl(aq) Cu(s) N.
  R.
• 4H(aq) 2NO3-(aq) Cu(s) Cu2
  2NO2(g) 2H2O(l )

8

• Solutions
• Neutralization reactions involve the reaction
  between an acid and a base to produce water and
  a salt
• HCl(aq) NaOH(aq) H2O(l ) NaCl(aq)
• A salt is an ionic compound whose cation comes
  from a base and whose anion comes from an acid
  NaCl
• H2SO4(aq) Ba(OH)2(aq) 2H2O(l )
  BaSO4(s)
• Note BaSO4 is not soluble in water
• NH3(aq) HCl(aq) NH4(aq) Cl-(aq)
• NH3(aq) H2O(l )
  NH4(aq) OH-(aq)
• NH4(aq) OH-(aq) HCl(aq) NH4(aq)
  Cl-(aq) H2O(l )
• add eqns NH3(aq) HCl(aq) NH4(aq)
  Cl-(aq)

9

• Ionic Reactions in Solution
• Ionic Equations
• Example AgNO3(aq) NaCl(aq) AgCl(s)
  NaNO3(aq)
• This is called a molecular equation even though
  all the reactants and products are ionic
  compounds
• The ionic character of the reactants and product
  are not shown in the equation
• Another way to write the equation
• Ag(aq) NO3-(aq) Na(aq) Cl-(aq)
  AgCl(s) Na(aq) NO3- (aq)
• This is called a complete ionic equation because
  it shows the ionic nature of all the reactants
  and products
• Notice AgCl forms as a solid and is not
  dissociated. It looks like a molecular
  compound but in fact is ionic but is not very
  soluble in water.
• Both Na(aq) and Cl-(aq) show up on both sides of
  the equation in the same form they can be
  canceled
• Ag(aq) Cl-(aq) AgCl(s)
• This is called a net ionic equation
• It shows the net chemical change that takes place

10

• Ionic Reactions in Solution
• Net Ionic Equations
• Another Example
• HNO3(aq) NaOH(aq)
  NaNO3(aq) H2O molecular equation
• H(aq) NO3-(aq) Na(aq) OH-(aq)
  H2O(l ) Na(aq) NO3-(aq)
• H(aq)
  OH-(aq) H2O(l )
  net ionic equation
• HCl(aq) NaOH(aq)
  NaCl(aq) H2O (l ) molecular
  equation
• H(aq)
  OH-(aq) H2O(l )
  net ionic equation
• The net ionic equations for the two
  neutralization reactions are the same
• The net ionic equations for any neutralization
  reaction between a strong acid and a strong base
  will be the same.
• Spectator ions are the ions not involved in the
  net ionic equations but show up in complete
  ionic equations.

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Ionic Reactions in Solution What species are
written in ionic form? Species that are soluble
and
write in ionic
form Species that are strong electrolytes What
species are not written in ionic form? Soluble
weak electrolytes Soluble non-electrolytes
write in molecular
form Insoluble substances - solids, liquids,
gases Metathesis reactions also called double
replacement reactions Cations in two compounds
exchange anion partners AX BY
AY BX Pb(NO3)2(aq) 2HCl(aq)
PbCl2(s) 2HNO3(aq)
12

• Metathesis Reactions
• Three processes lead to a metathesis reaction
• Formation of a precipitate, an insoluble product
• Formation of a weak electrolyte or
  non-electrolyte
• Formation of a gas that escapes from solution
• Precipitation Reactions
• Ag(aq) Cl-(aq) AgCl(s)
• AgCl is not very soluble in water
• Solubility is the amount of substance that
  dissolves in 1 L water
• Solubility of AgCl is about 10-5 mol/L
• Insoluble substances have solubilities lt about
  0.01 M

13
Solubility Guidelines
14
Solubility Guidelines Predict metathesis
precipitation reactions Na2SO4 CaCl2
2NaCl CaSO4 Cl-s are soluble SO42- s
are soluble except Ca2 Na2SO4(aq)
CaCl2(aq) 2NaCl(aq) CaSO4(s)
Ca2 (aq) SO42- (aq)
CaSO4(s) Metathesis Reactions Formation of a weak
electrolyte or non-electrolyte Water is a very
weak electrolyte
HNO3(aq) NaOH(aq) NaNO3(aq) H2O
H(aq) OH-(aq)
H2O(l ) net
Mg(OH)2(s) 2HCl(aq)
MgCl2(aq) 2H2O(l )
Mg(OH)2(s) 2H (aq) Mg2
(aq) 2H2O(l ) net
?
15
Metathesis Reactions Formation of a weak
electrolyte or non-electrolyte Other weak
electrolytes HC2H3O2, HNO2, and other weak
acids Formation of a gas Anion
Reaction S2- S2-(aq) 2H(aq)
H2S(g) CO32- CO32-(aq) 2H(aq)
H2CO3(aq) CO2 (g) H2O(l ) SO32-
SO32-(aq) 2H(aq) H2SO3(aq)
SO2 (g) H2O(l ) NO2- NO2-(aq) H(aq)
HNO2(aq)
2HNO2(aq) H2O(l) NO2(g) NO
(g) Cation Reaction NH4 NH4 (aq) OH-
(aq) NH3(g) H2O(l )
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Metathesis Reactions Formation of a
gas Examples Reaction between PbS(s) and
HNO3(aq) PbS(s) HNO3 (aq)
Pb(NO3)2 (aq) H2S(g)
PbS(s) 2H (aq) Pb (aq)
H2S(g) Reaction between Fe2(CO3)3(s) and
HCl(aq) Fe2(CO3)3(s)
6HCl(aq) 2FeCl3 3CO2(g) 3H2O(l )
Fe2(CO3)3(s) 6H (aq)
2Fe3 (aq) 3CO2 (g) 3H2O(l
) Introduction to Oxidation Reduction
Reactions Example 2Na(s) Cl2(g)
2NaCl(s) The Na atoms have lost electrons to
become Na The Cl atoms have gained electrons to
become Cl-
17

• Introduction to Oxidation Reduction Reactions
• Oxidation is the loss of electrons by a substance
• Na Na 1e-
• Reduction is the gain of electrons by a substance
• Cl 1e- Cl-
• Oxidation must be accompanied by reduction
• If a species loses electrons another species must
  pick them up
• The number of electrons gained must equal the
  number lost in a balanced chemical equation
• Oxidation of metals by acids
• Many metals react with acids to produce a metal
  ion (or a metal salt) and H2(g)
• Fe(s) 2HCl(aq) FeCl2(aq) H2(g)
• Fe(s) 2H (aq)
  Fe2 H2(g)

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Introduction to Oxidation Reduction
Reactions Activity Series of Metals Compare the
ability of one metal to be oxidized by another
metals ion Example Zn(s) Cu2(aq)
Cu(s) Zn2(aq) Cu2(aq) can oxidize
Zn metal Zn metal can reduce Cu2 See Table 4.4,
page 131, in Brown for a list of metals ordered
by their oxidizing abilities Metals high in the
table are more easily oxidized Metals high in
the table can reduce metal ions lower in the
table Metals above hydrogen can be oxidized
by H Zn 2H Zn H2(g) Metals
below hydrogen cannot be oxidized by H Cu H
no reaction Cu can be oxidized by
HNO3 because NO3- oxidizes Cu Cu(s)
4H(aq) 2NO3-(aq) Cu2(aq) 2NO2(g)
2H2O
19
Solution Stoichiometry When solutions of
reactants are used, concentrations and volumes
give moles of reactants To find moles of solute
in a solution of known volume Example how
many moles of NaOH are required to react with
25.0 mL of 0.250 M HCl? NaOH HCl
NaCl H2O Example how many mL of 0.300
M HCl are required to neutralize 6.25 x 10-3
mol NaOH?
20

• Titrations
• A titration is a method for finding the volume of
  one solution that contains a substance that
  reacts with another substance in a second
  solution.
• If we know
• the number of moles of the second substance
• the balanced equation for the reaction
• the volume of the first solution that exactly
  reacts with the amount of reactant in the
  second solution
• We can calculate the concentration of the first
  solution
• One of the solutions used in a titration is a
  standard solution whose concentration is
  accurately and precisely known.

21

• Titrations
• The titration procedure involves very carefully
  adding from a buret one of the solutions to the
  second solution.
• The solution in the buret is the titrant.
• The addition is stopped when the volume of
  solution added contains a stoichiometric
  quantity of reactant
• the amount of reactant in the titrant will
  be exactly the amount
  required to consume the substance in the second
  solution.
• The equivalence point is the point in a titration
  when a stoichiometric amount of titrant has
  been added to the second solution
• The titration procedure often uses an indicator.
• An indicator is a substance added to the solution
  titrated that will change color hopefully when
  the equivalence point is reached
• The end point in a titration is the point where
  the indicator changes color.

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Titrations Example If 45.3 mL of 0.100 M NaOH is
required to exactly neutralize 75.0 mL of
H2SO4, what is the molarity of the H2SO4? H2SO4
2NaOH Na2SO4 2H2O 2 mol NaOH?1 mol
H2SO4 Example Photo developing laboratories
recover silver from the solutions used to
process film. A 50.0 mL sample from a 5.00 x 103
liter tank is analyzed by titration with Cl-,
requiring 18.3 mL of 0.100 M Cl- for the
titration. How many grams of Ag are in the
tank? Ag(aq) Cl- AgCl(s) 1 mol
Cl-?1 mol Ag