Types of Reactions

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Chapter 4

• Types of Reactions Solution Stoichiometry

2
Unit essential Question

• How do chemicals react with one another in
  aqueous solutions?

3
Lesson essential questions (4.1-4.4)

• 1) How do water molecules interact with
  chemicals?
• 2) How is the concentration of a solution
  measured?

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Water, the Common Solvent

• Section 4.1

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Aqueous solutions
d

• Dissolved in water.
• Good solvent- polar molecules.
• Hydration ions in salts break apart due to
  attraction to polar water molecules.
• Example
• NH4NO3 (s) ? NH4 (aq) NO3- (aq)

d-
d
6
Hydration
7
Solubility

• Amount of substance that will dissolve in a given
  amount of water.
• If they do dissolve, ions are separated, and can
  move around.
• Water can also dissolve non-ionic compounds if
  they have polar bonds.

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Like dissolves like

• Polar substances generally dissolve other polar
  and ionic substances
• Alcohol is slightly polar and dissolves (mixes)
  in water
• Nonpolar substances dissolve other nonpolar
  substances
• Fat will not dissolve in water

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The Nature of Aqueous Solutions Strong Weak
Electrolytes

• Section 4.2

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Parts of Solutions

• Solution- homogeneous mixture.
• Solute- what gets dissolved.
• Solvent- what does the dissolving.
• Soluble- Able to be dissolved.
• Miscible- liquids dissolve in each other.

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Electrolytes

• Electrolytes- ionic compounds in solution that
  conduct electricity.
• Strong electrolytes- completely dissociate (fall
  apart into ions).
• Many ions conduct electricity well.
• Weak electrolytes- partially dissociate into
  ions.
• Few ions conduct electricity slightly.
• Non-electrolytes- dont dissociate at all.
• No ions dont conduct electricity.

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Acid/Base Electrolytes

• Arrhenius acid- forms H ions when dissolved.
• Strong acids dissociate completely.
• Ex H2SO4 HNO3 HCl HBr HI
• Weak acids do not dissociate completely.
• Ex HC2H3O2
• Arrhenius base - forms OH- ions when dissolved.
• Strong bases also dissociate completely.
• Ex KOH NaOH (Groups 1 2 hydroxides)

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Sections 12 Homework
Pg. 170-171 1,9,18,19
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Warm-Up
HNO3 is a strong acid. Write the chemical
equation for a solution of HNO3. Will it conduct
electricity?
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Composition of Solutions

• Section 4.3

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Measuring Composition of Solutions

• To do stoichiometry
• Need to know chemicals
• Need to know amounts (concentrations)
• Concentration- how much is dissolved.
• Molarity Moles of solute
• Liters of solution
• abbreviated M (molar)
• 1 M 1mol solute / 1 liter solution

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Molarity Calculations

• Can solve for
• Amount or mass of solid to dissolve
• Moles of solute
• Volume of solution
• Standard solution
• Solution whose concentration is accurately known.

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Examples

• Calculate the molarity of a solution prepared by
  dissolving 11.5g of solid NaOH in water to make
  1.50L of solution. (pg. 134)
• Give the concentration of each ion in 0.50 M
  Co(NO3)2. (pg. 135)
• 27 pg. 172

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Dilutions

• Stock solution a concentrated solution
• Dilution number of moles of solute stays the
  same, just adding more water
• M1V1 M2V2
• Example 30 (a) pg. 172

mol1 x V1 mol2 x V2
L1
L2
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Section 3 Homework
Pg. 171-172 21-23,28,31
21
Lesson essential questions (4.5-4.7)

• 1) How do we identify and work with precipitation
  reactions?

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Precipitation Reactions

• Section 4.5

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Precipitation Reactions

• Solid forms when two solutions of ionic compounds
  are mixed.
• Precipitate (ppt)
• To help you remember If youre not a part of
  the solution, your part of the precipitate!

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Precipitation reactions

• NaOH(aq) FeCl3(aq) NaCl(aq)
  Fe(OH)3(s)
• is really
• Na(aq)OH-(aq) Fe3 Cl-(aq) Na
  (aq) Cl- (aq) Fe(OH)3(s)
• So all that really happens is
• OH- (aq) Fe3 (aq) Fe(OH)3 (s)
• Also a double displacement reaction

net ionic equation!
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Precipitation reaction

• Can predict products, but can only be certain by
  experimenting.
• The anion and cation switch partners.
• Only occurs if a product is insoluble!
• Otherwise all the ions stay in solution- nothing
  has happened (spectators)
• Memorize solubility rules! Pg. 144

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Solubility Rules

• All nitrates, Na, K, NH4 are soluble.
• You must know this for the AP exam!
• Additional solubility rules on pg. 144.

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Describing Reactions in Solutions

• Section 4.6

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Three Types of Equations

• 1. Formula Equation- write formulas, not ions.
• K2CrO4(aq) Ba(NO3)2(aq)
• 2. Complete Ionic equation- show dissolved
  electrolytes as the ions.
• 2K CrO4-2 Ba2 2 NO3-
  BaCrO4(s) 2K 2 NO3-
• Spectator ions are those that dont react- appear
  as ions on both sides.

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Three Type of Equations

• 3. Net Ionic equation- show only ions that react,
  not spectator ions
• Ba2 CrO4-2 BaCrO4(s)
• If all species in a reaction are aqueous
  (soluble), write NR!

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Sections 56 Homework
Pg. 172-173 36,42,44
32
AP Practice Question

• How many moles of Na2SO4 must be added to 500
  milliliters of water to create a solution that
  has a 2- molar concentration of the Na ion?
  (Assume the volume of the solution does not
  change.)
• ? Think about what you need to answer this!
• ? Need to find moles Na. Then find moles Na2SO4
• 0.5 moles
• 1 mole
• 2 moles
• 5 moles

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Stoichiometry of Precipitation Reactions

• Section 4.7

34
Stoichiometry of Precipitation

• Steps for reference pg.148
• Similar to other stoichiometry problems weve
  done!
• Sample problem What volume of 0.15M KCl is
  needed to precipitate out all of the lead ions
  from 100.mL of 0.20M Pb(NO3)2?

270mL KCl needed
35
Section 7 Homework
Pg. 173 47,48,50,54
36
Acid-Base Reactions

• Section 4.8

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Lesson essential question (4.8)

• How do we classify acids and bases?
• What happens when acids and bases are mixed
  together?

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Acid-Base Reactions

• For our purposes an acid is a proton donor, H
  (BrØnsted-Lowry theory).
• A base is a proton acceptor, usually OH-
• acid base salt water
• H OH- H2O
• Practice Write the net ionic equation for the
  acid/base rxn. below
• HNO3(aq) NaOH(aq) ? ?
• Note H2CO3 always breaks down into CO2 H2O

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Acid-Base Reactions

• Follow same steps as precipitation reactions for
  stoichiometry problems.
• See p.149-150
• Practice What volume (in mL) of 0.100M HCl will
  react completely with 25.00mL of 0.200 M NaOH?
• (1) Write net ionic equation
• (2) Find moles youre starting with
• (3) Find moles needed
• (4) Find volume needed

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Acid-Base Reactions

• Also called neutralization reactions.
• Use titrations to determine concentrations.
• Titrant solution of known concentration
• Analyte solution of unknown concentration
• Equivalence Point when enough titrant has been
  added to exactly react with the analyte
  (neutralization is complete).
• Stoichiometric amounts come from balanced
  equation!
• Tells us how many moles of the titrant fully
  reacted with the analyte- then can solve for
  moles of analyte!

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Titration

• Solution of known concentration (titrant), is
  added to the unknown (analyte), until the
  equivalence point is reached.
• How do we know when the equivalence point has
  been reached?
• Add indicator to analyte at the beginning

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Titration

• Where the indicator changes color is the
  endpoint.
• Ex phenolphthalein used often
• Pink in base, colorless in acid
• As close as we can get to the equivalence point
  still assume complete neutralization.
• The solution will not turn pink until one drop
  after the equivalence point (when the solution is
  more basic).
• Can also use titration for non acid/base
  substances to find amounts/concentrations.

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AP Practice Question

• Which of the following best represents the
  balanced net ionic equation for the reaction of
  lead(II) carbonate concentrated hydrochloric
  acid? (All lead compounds are insoluble.)
• Pb2CO3 2H Cl- ? Pb2Cl CO2 H2O
• PbCO3 2H 2Cl- ? PbCl2 CO2 H2O
• PbCO3 2H ? Pb2 CO2 H2O
• PbCO3 2Cl- ? PbCl2 CO3-2

44
AP Practice Question
The conductivity of a solution of Ba(OH)2 is
monitored as the solution is titrated with 0.10 M
H2SO4. The original volume of the Ba(OH)2
solution is 25.0 mL. A precipitate of BaSO4 is
formed during the titration. The data collected
from the experiment is plotted in the graph above.
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Question Continued

• As the first 30.0 mL of 0.10 M H2SO4 are added to
  the Ba(OH)2 solution, two types of chemical
  reactions occur simultaneously. Write the
  balanced net-ionic equations for (i) the
  neutralization reaction and (ii) the
  precipitation reaction.
• (i) Equation for neutralization reaction
• (ii) Equation for precipitation reaction

OH- (aq) 2H (aq) ? H2O (l)
Ba2 (aq) SO4-2 (aq) ? BaSO4 (s)
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Question Continued

• 2) The conductivity of the Ba(OH)2 solution
  decreases as the volume of added 0.10 M H2SO4
  changes from 0.0 mL to 30.0 mL.
• Identify the chemical species that enable the
  solution to conduct electricity as the first 30.0
  mL of 0.10 M H2SO4 are added.
• (ii) On the basis of the equations you wrote in
  question 1, explain why the conductivity
  decreases.

OH- (aq) Ba2 (aq) (Cant be anything from
H2SO4 because the ions immediately react.)
Ba2 in sltn. decrease as they precipitate out,
and OH- in sltn. decrease as they react to form
H2O. Note be specific in your answers!!
Reference all species and reactions!
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Question Continued
Think about what information can be determined
from this point!
At equivalence point complete neutralization
3) Using the information in the graph, calculate
the molarity of the original Ba(OH)2 solution.
0.12M Ba(OH)2
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Section 8 Homework
Homework pg.173-174 56, 58, 60, 64, 66
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Oxidation Reduction Reactions

• Section 4.9

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Lesson essential questions (4.9-4.10)

• How can we identify redox reactions?
• How do we assign oxidation states?
• Why is balancing different for redox reactions?

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Redox Reactions

• Ionic compounds are formed through the transfer
  of electrons.
• An oxidation-reduction reaction involves the
  transfer of electrons.
• One element gains, one loses
• Non-ionic compounds can also undergo redox
  reactions.

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Oxidation States charge

• A way of keeping track of the electrons.
• Not necessarily true of what is in nature, but it
  works.
• Need to memorize rules for assigning (pg.156)
• The oxidation state of elements in their standard
  states is zero.
• Oxidation state for monatomic ions are the same
  as their charge.

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Oxidation states

• Oxygen is assigned an oxidation state of -2 in
  its covalent compounds except in peroxide (-1).
• In compounds with nonmetals hydrogen is assigned
  the oxidation state 1.
• In its compounds fluorine is always 1.
• The sum of the oxidation states must be zero in
  compounds or equal the charge of the ion.

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Practicing Oxidation States

• Determine the oxidation states in the following
• Cl2
• SO4-2
• CaBr2
• C6H12O6

Cl 0
S 6 O -2
Ca 2 Br -1
C 0 H 1 O -2
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Section 9 Homework
Pg. 174 67(c-e),68(a-c),72
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Balancing Redox Reactions

• Section 4.10

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Oxidation-Reduction

• e- transferred, so the oxidation states change.
• Oxidation is the loss of electrons.
• More positive oxidation state.
• Reduction is the gain of electrons.
• More negative oxidation state.
• OIL RIG
• LEO (the lion goes) GER

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Agents

• Oxidizing agent- substance that gets reduced
  (causes oxidation in another species).
• Gains electrons.
• More negative oxidation state.
• Reducing agent- substance that gets oxidized
  (causes reduction in another species).
• Loses electrons.
• More positive oxidation state.

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Identify the

• Oxidizing agent
• Reducing agent
• Substance oxidized
• Substance reduced
• 1 2Na Cl2 ? 2NaCl
• 2 CH4 2O2 ? CO2 2H2O

reducing agent, substance oxidized
oxidizing agent, substance reduced
oxidizing agent, substance reduced
reducing agent, substance oxidized
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Half-Reactions

• All redox reactions can be thought of as
  happening in two halves.
• One produces electrons - oxidation half.
• The other requires electrons - reduction half.
• Ex Fe (s) CuSO4 (aq) ? Cu (s) FeSO4 (aq)
• Net Ionic Fe (s) Cu2 (aq) ? Cu (s) Fe2
  (aq)
• Oxidation Fe (s) ? Fe2 (aq) 2e-
• Reduction Cu2 (aq) 2e- ? Cu (s)

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Balancing Redox Equations

• Redox reactions may involve an acid or base as a
  reactant.
• The number of electrons produced must be the same
  as those required.
• 8 step procedure for acidic solution, 10 step
  procedure for basic solution.
• Called the half reaction method.
• Balance each half reaction, then combine for
  total balanced reaction

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Balancing in Acidic Solution

• Write separate half reactions.
• For each half reaction balance all species except
  H and O.
• Balance O by adding H2O to one side.
• Balance H by adding H to one side.
• Balance charge by adding e- to the more positive
  side.

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Balancing in Acidic Solution

• Multiply equations by a number to make electrons
  equal.
• Add equations together and cancel identical
  species. Reduce coefficients to smallest whole
  numbers.
• Check that charges and elements are balanced.

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Balancing in Acidic Solution

• Ex Balance the following equation
• H (aq) Cr2O7-2 (aq) C2H5OH (l) ?
• Cr3 (aq) CO2 (g) H2O (l)
• Reduction 6e- 14H Cr2O7-2 ? 2Cr3 7H2O
• Oxidation C2H5OH 3H2O ? 2CO2 12H 12e-
• Final 16H 2Cr2O7-2 C2H5OH ? 4Cr3 11H2O
  2CO2
• Note there should NOT be any e- in the final
  balanced equation! If so, not balanced!

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Basic Solution

• Do everything you would with acid, but add one
  more step.
• Add enough OH- to both sides to neutralize the
  H.
• Any H and OH- on the same side form water.
  Cancel out any H2Os on both sides.
• Simplify coefficients, if necessary.

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Balancing in Basic Solution

• Assume previous example in acidic solution was
  actually in a basic solution.
• Had 16H 2Cr2O7-2 C2H5OH ? 4Cr3 11H2O
  2CO2
• For any H ions, add same number of OH- ions to
  both sides. This forms water with H. Cancel out
  waters on both sides.
• Now 16H2O 2Cr2O7-2 C2H5OH ? 4Cr3 11H2O
  2CO2 16OH-

16H, so add 16OH-
5 H2O
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Practice Balancing Redox Rxns.

• Pg. 174 74(b)
• Pg. 174 75(b)

Answer 6Cl- Cr2O7 14H ? 3Cl2 2Cr3 7H2O
Answer 2OH- Cl2 ? OCl- Cl- H2O
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Side Note Redox Titrations

• Same as titrations discussed before, just looking
  at redox reactions instead of acid/base
  reactions.
• Permanganate ion is used often because it is its
  own indicator MnO4- is purple, Mn2 is
  colorless. When reaction solution remains clear,
  MnO4- is gone.
• Chromate ion is also useful, but color change,
  orangish yellow to green, is harder to detect.

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Section 10 Homework
Pg. 174-175 73-76 ONLY letter a for each