Reactions in Aqueous (aq) Solutions by HC120519112450

VIEWS: 32 PAGES: 84

									Alex Van Name
Index
 CA Standards for this topic
 13-1Types of Mixtures
 13-2 The Solution Process
 13-3 Concentration of Solutions
CA Standard #5
Acids, bases, and salts are three classes of
compounds that form ions in water solutions.
 a. Students know the observable properties of acids,
  bases, and salt solutions.
 b. Students know acids are hydrogen-ion-donating
  and bases are hydrogen-ion-accepting substances.
 c. Students know strong acids and bases fully
  dissociate and weak acids and bases partially
  dissociate.
 d. Students know how to use the pH scale to
  characterize acid and base solutions.
CA Standard #5
Acids, bases, and salts are three classes of
compounds that form ions in water solutions.
 e.* Students know the Arrhenius, Brønsted-Lowry,
  and Lewis acid–base definitions.
 f.* Students know how to calculate pH from the
  hydrogen-ion concentration.
 g.* Students know buffers stabilize pH in acid–base
  reactions.
CA Standard #6
Solutions are homogeneous mixtures of two
or more substances.
a. Students know the definitions of solute and
  solvent.
b. Students know how to describe the dissolving
  process at the molecular level by using the
  concept of random molecular motion.
c. Students know temperature, pressure, and
  surface area affect the dissolving process.
CA Standard #6
Solutions are homogeneous mixtures of two
or more substances.
d. Students know how to calculate the concentration
   of a solute in terms of grams per liter, molarity,
   parts per million, and percent composition.
e.* Students know the relationship between the
   molality of a solute in a solution and the solution’s
   depressed freezing point or elevated boiling point.
f.*Students know how molecules in a solution are
   separated or purified by the methods of
   chromatography and distillation.
CA Standard #5
Acids, bases, and salts are three classes of
compounds that form ions in water solutions.
 a. Students know the observable properties of acids,
  bases, and salt solutions.
 b. Students know acids are hydrogen-ion-donating
  and bases are hydrogen-ion-accepting substances.
 c. Students know strong acids and bases fully
  dissociate and weak acids and bases partially
  dissociate.
 d. Students know how to use the pH scale to
  characterize acid and base solutions.
CA Standard #6
Solutions are homogeneous mixtures of two
or more substances.
a. Students know the definitions of solute and
  solvent.
b. Students know how to describe the dissolving
  process at the molecular level by using the
  concept of random molecular motion.
c. Students know temperature, pressure, and
  surface area affect the dissolving process.
Labs and Activities
 Light bulb to test electrolytic properties
 Precipitation Reactions
    Predict products
    Carry out experiment
    Write net ionic equations
 Acid Base
    Sea shell plus acid
    Base plus Egg white
 Redox Demo
    Zn + copper sulfate
    Demo a day
Types of Solutions
 The most common solutions are
  solutions that are a combination
  of a solid dissolved in a liquid.
 While solutions can be made
  using any combination of
  different states of matter
  (solid, liquid, gas)
Special Solutions
   Steel          Bronze         Brass
Iron + Carbon   Copper +Tin   Copper + Zinc
Special Solutions: Air
Solution Vocabulary
 Soluble = the substance can be dissolved
    Ex: Salt is soluble in water
 Solute = the substance dissolved in the solvent, less
  of this one…example: Salt
 Solvent = the substance that dissolves the solute,
  more of this substance….example: Water
 Solution = homogenous mixture of two or more
  substances that coexist in a single phase
   Ex: Saltwater
 Water is the universal solvent
Water is a molecular substance that is polar,
with a (+) side and a (-) side



             +              -
Solute + Solvent = Solution
Solute + Solvent = Solution
Most solutions are made with water as the solvent.
Aqueous (aq) means the substance is dissolved in
water.
 C6H12O6(s) + H2O(l) = C6H12O6 (aq)
   Sugar + Water = a sugar solution
 NaCl(s) + H2O(l) = NaCl(aq)
    Salt + Water = salt water solution
 CuSO4(s)+ H2O(l) = CuSO4(aq)
    Copper Sulfate + Water = Copper sulfate solution
Salt Dissolving Link
 Salt Dissolving
Ionic Solution(aq)
Ionic Compounds
 Ionic compounds dissociate in water and breakup
 into ions that are positive or negative.
   Positive ions (+) = Cations (pronounced Cat-ions)
   Negative ions (-) = Anions (pronounced An-ions)
  Example:
     NaCl(s)    
                
                 Water
                          Na+ (aq) + Cl-(aq)
 (aq) or aqueous means that the substance has been
 dissolved in water.
 Dissociation and Ionization results in ions being
  formed in solution
Ionic Crystals
 The unit cell is the shape that is repeated throughout
  the structure of the crystal.
 The unit cell defines the “formula unit” [Ex: CsCl] and
  calculate the mass of ionic substances. ⅛ x 8 Cs
Alloys
 Alloys are evenly
  distributed mixtures of 2
  or more different metals.
 Iron-Platinum Alloy is
  shown to the right
Amalgam
 Mercury (Hg) + Metal = Amalgam
Hydration
 Hydration is the process of water surrounding
 molecules arranged in a specific manner.
    Copper(II) sulfate
  Anhydrous vs. Hydrated
CuSO4 (Anhydrous)   CuSO4• 5H2O
 Tyndall Effect
 Light Scatters
  when light
  bounces off
  molecules are
  large.
 Light will NOT
  scatter in a true
  homogeneous
  solution
Solutions, Colloids, Suspensions
Solutions                  Colloids                    Suspensions
Homogeneous                Heterogeneous               Heterogeneous
Particles are small,       Large molecules, like       Polymers with particle
atoms or ions or small     proteins, 1-1000nm          size over 1000nm
molecules less than 1nm
Particles don’t separate   Particles don’t separate    Particles WILL separate
with time                  with time                   with time
Does not scatter light     Scatter’s light (Tyndall)   Scatter’s light (Tyndall)
Cannot filter out parts    Cannot filter out parts     Can filter and separate
                                                       parts of the suspension
Kool Aid                   Milk                        Muddy water
Opposites Attract, even in the
matrix!!
  Electrolyte vs. Non-electrolyte
 Electrolytes conduct electricity,
  Non-electrolytes don’t
 Electricity is defined as the
  movement of charge
   Can be ions moving in
    solution
   Can be electrons moving in a
    solid metal
 Ions are attracted to electrons
Dissolving
Dissolving can be influenced
3 factors that influence dissolving:

  Temperature
  Agitation
  Surface area of the solute
     Powder        vs. Solid Crystal
  Pressure
Dissolving (Temperature)
 Increasing the temperature increases the
  molecular motion of the solvent (and
  solute) particles.
 This increase in motion (kinetic energy)
  results in a greater frequency of collisions
  between particles.
 Lower temperatures have the exact
  opposite effect.
Dissolving (Agitation)
 This physically moves the
  molecules around
 Increases kinetic energy (slightly)
 Prevents “pockets” of high
  concentration in the solution and
  makes the solution heterogeneous.
Dissolving (Surface Area)
 Increasing the surface area
  of the solute (by grinding in a
  mortar and pestle) will
  increase the accessible
  surfaces for the solvent molecules to collide.
 A solid block only allows access to the outside
  surfaces of the block.
 If you grind or break apart the solid before
  dissolving, the solvent molecules will be able to
  contact places that are blocked by molecules in
  the solid piece.
Dissolving (Pressure)
 Increasing the pressure
  of the system forces
  more gas to dissolve in
  the liquid.
 Gases can dissolve in
  liquids. Carbonated
  soda is an example of
  a gas(CO2) dissolved
  in water.
  Dissolving saves space
 This generally happens because a
  system under pressure will try to
  conserve space under increased
  pressure.
 Since the fewer number of particles
  is favored, increasing pressure
  decreases the number of moles of
  particles and saves space.
 When you open the cap on a soda
  bottle, the pressure is decreased and
  the bubbles separate from the
  solvent
Solubility Defined
 Solubility: the maximum amount of
  solute that will dissolve in a given
  quantity of solvent at a specific
  temperature*
 Substances (solutes) can be…
    soluble (dissolve completely)
    insoluble (does not dissolve) or
    slightly soluble (a little bit dissolves)
Saturated vs. Unsaturated
 Saturated                   Unsaturated solutions
    No more room for            The solvent can
     solute molecules to          dissolve more solute
     dissolve.                   There is “more room”
    Tastes Strong, super         for solute to dissolve
     sweet, or super salty
Saturated vs. Unsaturated
Solubility
Curves
Practice: Solubility Curve
  Solute   Temperature                   Concentration

  NaCl       10C                ? (A)
  NaCl       90C                ? (B)
 K2Cr2O7     ? (C)       30 g K2Cr2O7 / 100g H2O
 K2Cr2O7     ? (D)       70 g K2Cr2O7 / 100g H2O
   KCl       90C                ? (E)
  ? (F)      90C          47g KClO3 / 100g H2O
Practice: Solubility Curve
  Solute   Temperature                   Concentration

  NaCl       10C          35 g NaCl / 100g H2O
  NaCl       90C                ? (B)
 K2Cr2O7     ? (C)       30 g K2Cr2O7 / 100g H2O
 K2Cr2O7     ? (D)       70 g K2Cr2O7 / 100g H2O
   KCl       90C                ? (E)
  ? (F)      90C          47g KClO3 / 100g H2O
Practice: Solubility Curve
  Solute   Temperature                   Concentration

  NaCl       10C          35 g NaCl / 100g H2O
  NaCl       90C          40 g NaCl / 100g H2O
 K2Cr2O7     ? (C)       30 g K2Cr2O7 / 100g H2O
 K2Cr2O7     ? (D)       70 g K2Cr2O7 / 100g H2O
   KCl       90C                ? (E)
  ? (F)      90C          47g KClO3 / 100g H2O
Practice: Solubility Curve
  Solute   Temperature                   Concentration

  NaCl       10C          35 g NaCl / 100g H2O
  NaCl       90C          40 g NaCl / 100g H2O
 K2Cr2O7     50C        30 g K2Cr2O7 / 100g H2O
 K2Cr2O7     ? (D)       70 g K2Cr2O7 / 100g H2O
   KCl       90C                ? (E)
  ? (F)      90C          47g KClO3 / 100g H2O
Practice: Solubility Curve
  Solute   Temperature                   Concentration

  NaCl       10C          35 g NaCl / 100g H2O
  NaCl       90C          40 g NaCl / 100g H2O
 K2Cr2O7     50C        30 g K2Cr2O7 / 100g H2O
 K2Cr2O7     90C        70 g K2Cr2O7 / 100g H2O
   KCl       90C                ? (E)
  ? (F)      90C          47g KClO3 / 100g H2O
Practice: Solubility Curve
  Solute   Temperature                 Concentration

  NaCl       10C          35 g NaCl / 100g H2O
  NaCl       90C          40 g NaCl / 100g H2O
 K2Cr2O7     50C        30 g K2Cr2O7 / 100g H2O
 K2Cr2O7     90C        70 g K2Cr2O7 / 100g H2O
   KCl       90C            55g KCl / 100g H2O
  ? (F)      90C          47g KClO3 / 100g H2O
Practice: Solubility Curve
  Solute   Temperature                 Concentration

  NaCl       10C          35 g NaCl / 100g H2O
  NaCl       90C          40 g NaCl / 100g H2O
 K2Cr2O7     50C        30 g K2Cr2O7 / 100g H2O
 K2Cr2O7     90C        70 g K2Cr2O7 / 100g H2O
   KCl       90C            55g KCl / 100g H2O
  KClO3      90C          47g KClO3 / 100g H2O
Solubility Curve
Saturated = on the line
Unsaturated = below line
Super Saturated = above line
Saturated: Organic Chem
Saturated: Organic Chem.

                     Fewer “H’s”
                     means the
                     molecule is
                     unsaturated
Monday, November 23, 2009
 Quiz: Solubility
 Warm up:
    Define electrolyte, give an example of an electrolyte
    What are cations and what is the charge?
    What are anions and what is the charge?
 Review Lab
 Topic: Precipitates and Net Ionic Equation
 HW: p 124 # 4.2, 4.10, 4.14, 4.20, 4.22, 4.26, 4.28
Super Saturated Solution
Solutions that have more solute than the solvent can
normally take are unstable. These overstuffed
solutions are super saturated.
Ex: Sodium Acetate (aq)

 http://www.youtube.com/watch?v=XSGvy2FPfCw
“Like dissolves Like”
 Substances that are polar will dissolve other
 substances that are polar.
CA Standard #5
Acids, bases, and salts are three classes of
compounds that form ions in water solutions.
 a. Students know the observable properties of acids,
  bases, and salt solutions.
 b. Students know acids are hydrogen-ion-donating
  and bases are hydrogen-ion-accepting substances.
 c. Students know strong acids and bases fully
  dissociate and weak acids and bases partially
  dissociate.
 d. Students know how to use the pH scale to
  characterize acid and base solutions.
CA Standard #6
Solutions are homogeneous mixtures of two
or more substances.
a. Students know the definitions of solute and
  solvent.
b. Students know how to describe the dissolving
  process at the molecular level by using the
  concept of random molecular motion.
c. Students know temperature, pressure, and
  surface area affect the dissolving process.
CA Standard #6
Solutions are homogeneous mixtures of two
or more substances.
d. Students know how to calculate the concentration
   of a solute in terms of grams per liter, molarity,
   parts per million, and percent composition.
e.* Students know the relationship between the
   molality of a solute in a solution and the solution’s
   depressed freezing point or elevated boiling point.
f.*Students know how molecules in a solution are
   separated or purified by the methods of
   chromatography and distillation.
• Measuring small quantities in large volumes of water
parts per million (ppm)
 What is ppm? One part per million ( 1 ppm) denotes
  one part per 1,000,000 parts, one part in 106, and a
  value of 1 × 10–6.
 Why ppm? ppm is used when discussing small
  concentrations, like the concentration of mercury in
  ocean fish, or pollution in drinking water
Drinking Water, 1 ppm
 1 mg/L = 1 ppm for dilute aqueous solutions.
    For example, a chlorine concentration of 1.8 mg/L
     chlorine is equivalent to 1.8 ppm chlorine.
 Chlorine is used to kill harmful bacteria in drinking
 water
   The EPA has set a limit for drinking water of
    4 milligrams of chlorine per liter of water (4 mg/L)
  Emergency Water Purification
  using 5.25% bleach
 Source http://www.epa.gov/safewater/faq/emerg.html
     Available            Drops per Quart/Gallon of Clear                         Drops per Liter of
     Chlorine                         Water                                          Clear Water
        4-6%                2 drops per Quart – 8 drops per                             2 per Liter
                                         Gallon
                               (1/8 teaspoon, or 8 drops)
(If the strength of the bleach is unknown, add ten drops per quart or liter of filtered and settled water.
Double the amount of chlorine for cloudy, murky or colored water or water that is extremely cold.)
Mix the treated water thoroughly and allow it to stand,
preferably covered, for 30 minutes. The water should have a slight chlorine odor.
If not, repeat the dosage and allow the water to stand for an additional 15 minutes. If the treated water has too
strong a chlorine taste, allow the water to stand exposed to the air for a few hours or pour it from one clean
container to another several times.
Why Bleach Kills Bacteria




  Essentially, bleach degrades to simple salt water
  which is not toxic to humans, animals, or our water
  supply.
Why Bleach Kills Bacteria
When this regular bleach is diluted in water at
the time of use, the excess water drives the
dissociation of the sodium hypochlorite into its
parts, while combining with some of the free H+
ions to form HOCl and NaOH.

The HOCl further breaks down spontaneously in
solution into HCl and O (oxygen).
Why Bleach Kills Bacteria
This free oxygen is very reactive, and when it
comes into contact with bacteria, it forces open their
cell walls, killing them. As this disruption is a
chemical/physical process (not mediated by protein
or gene expression), bacteria do not generally
acquire resistance to this mode of cell death,
making it one of the safest products to use for
disinfecting because it eliminates the concern of
antibiotic resistance acquisition by bacteria.
ppm for Mercury in Fish
 The majority of the mercury is
  present in an organic form, usually
  as methyl mercury, which can be
  toxic to humans in sufficient quantity.
 Fish in the U.S. diet average about
  0.3 parts per million (ppm) mercury.
 Food and Drug Administration (FDA)
  consider poisonous fish to have
  1.0 ppm mercury or higher.
Which fish are poisonous?
SOURCE:
http://www.oehha.ca.gov/fish/pdf/DiscAdvyUpdates032309.pdf

Common               Number          Mean          Mean Total
Name                 of Fish        Mercury       Length (mm)
                                     (ppm)
Sunfish                 10            0.54              171
Largemouth
                        77             1.37             419
bass
Rainbow trout           14             0.20             333
Sacramento
                        28             1.36             390
pikeminnow
Safe Fish
 Atlantic Salmon
 Shellfish
 Flat fish & Flounder
 Hake, Haddock,
  Pollock, Cod
 Canned Tuna
 Tuna steak
 Halibut
Evil = Chilean Sea Bass
Chilean sea bass                ECO-WORST CHOICE
(toothfish)
                                Avoid or eat infrequently until
HEALTH ALERT
                                improvements are made
Elevated levels of mercury.
•Adults and kids should



            http://www.edf.org/page.cfm?tagID=16319

ECO-BEST                         ECO-WORST
•Farmed striped bass             •All Chilean sea
•Sablefish from Alaska or Canada bass
M = n/V
Neutralization MaVa=MbVb
Dilution M1V1=M2V2
Calculations using the Molarity (M) formula
Molarity (M)
    moles solute
M
   Liters Solvent

   mol                 n
M             or   M
    L                  V
Molarity (M)
           moles solute
       M
          Liters Solvent
                  moles
          n
       M
          V         volume
Molarity, easy example
 What is the concentration of a solution that is made
 from 0.5 mol NaOH dissolved in 0.250 L of water?

     moles solute
 M
    Liters Solvent
     0.50 mol NaOH
 M                    2.0 M NaOH
    0.250 Liter Water
Molarity, harder example
 What is the concentration of a solution that is made
 from 72.9 g HCl dissolved in 400 mL of water?



    grams  moles             mL  L



            M=          
 Molarity (M) example 2
   What is the concentration of a solution that is made
   from 72.9 g HCl dissolved in 400 mL of water?


           1 mol HCl                  1L
72.9g HCl  mol HCl  2.0 mol HCl 
72.9g   2.0                 400mL           0.400 L
           36.5g HCl                1000mL

        2.0 mol HCl
  M                    5.0 M HCl
     0.400 Liter Water
 Dilutions, MiVi = MfVf

   initial
                MiVi = MfVf                       final
concentration                                volume of the
                   initial   final              solution
                  volume     concentration       (total)
 Dilutions, MiVi = MfVf
HOW to Dilute safely:
1. Calculate the final volume (Vf)
   using the above equation.
2. Add water to the larger container,
   to prevent accidents like acid
   splatter!
3. Add the concentrated solution (Vi)
4. Adjust the level to reach the final
   volume (Vf)
 Dilutions, MiVi = MfVf

   initial
                MiVi = MfVf                          final
concentration                                   volume of the
                   initial     final               solution
                  volume       concentration        (total)

Final volume, Vf, is the ending volume, it doesn’t tell you
how much water to add.
You need to subtract Vf - Vi = Volume of Water to add
Dilution Sample Problem
 A Chem Lab uses a 1.00 M HCl Solution. How much
  water is needed to dilute 50mL of 6M HCl?
1. Use MiVi=MfVf to calculate Vf
2. Subtract: Vf – Vi to calculate the water needed
Dilution Sample Problem Step 1
 A Chem Lab uses a 1.0 M HCl Solution. How much
 water is needed to dilute 50mL of 6.0 M HCl?
             M i Vi  M f Vf
(6.0M)(.05 0L)  (1.0M)(V f )
                                               final
                                          volume of the
     (6.0)(.050 )
                   Vf                       solution
                                              (total)
         (1.0)
             Vf  0.30L or 300mL
Dilution Sample Problem Step 1
 A Chem Lab uses a 1.0 M HCl Solution. How much
 water is needed to dilute 50mL of 6.0 M HCl?
M i Vi  M f Vf           Vf  0.30L or 300mL

    300 mL          Final Volume

                    Concentrated 6M Acid
    - 50 mL               Volume


    250 mL of water needed
Dilutions, MiVi = MfVf
 Laboratory chemicals are diluted from
  concentrated stock solutions
 Special Note for diluting ACIDS and
  BASES!
   Always add the acid/base to water
   The dissolving process releases heat,
    causing the acid/base to splatter and
    heat the container also!!

      MiVi = MfVf

								
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