chapter 12 solutions. objectives describe the components of a solution explain why water is a good...
TRANSCRIPT
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Chapter 12
Solutions
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Objectives
• Describe the components of a solution
• Explain why water is a good solvent
• Explain the expression “Likes Dissolve Likes”
• Define solubility
• Compare solubility of solids and gases in solution
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Solutions
• Homogeneous mixture of two or more substances
– Homogenous – Uniform
• Two components of solutions
– Solute – Substance being dissolved
• Present in lower amount
– Solvent – Substance that dissolves solute
• Present in larger amount
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Types of Solutions
• Solid dissolved in a liquid– Sugar in water, sodium chloride in water
• Liquid dissolved in a liquid– Ethanol in water (wine)
• Liquids are Miscible • Gas dissolved in a liquid
– Oxygen in water (Lake water)• Gas dissolved in a gas
– Oxygen in nitrogen (Air)
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What is the solute?
What is the solvent?
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Types of Solutions• Solid dissolved in a solid - Alloy
– Copper in Silver (Sterling Silver)
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Aqueous Solutions
• Solutions where water is the solvent
• Very common type of solution
• Water is an excellent solvent
– Dissolves a number of solutes
– Often referred to as a “universal solvent”
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Water as a Solvent
• Water is an excellent solvent because of its shape
– Water is a bent molecule
– Very Polar
• Oxygen end of molecule is very negative
• Hydrogen end is very positive
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Water as a Solvent
• These ends of the molecule seek out opposite charges to rip apart solutes
• Process called hydration
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Water as a Solvent
• This is how water dissolves all ionic compounds that are soluble
– Creates ions in solution
• Process is slightly different, but similar, for polar covalent compounds
– Water just dissolves groups of polar molecules into individual polar molecules
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“Likes Dissolve Likes”
• This saying tells us what kind of solute a given solvent will dissolve
• Polar solvents dissolve
– Polar and ionic solutes
• Alcohols, Sugars, Salts
• Nonpolar solvents dissolve
– Nonpolar solutes
• Oils, Plastics
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Polarity Demo
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Solubility
• The maximum amount of solute that can dissolve in a given amount of solvent at a given temperature
– Usually in units of
• grams solute / 100 grams solvent
• Solubility information exists for solids, liquids, and gases
• Solubility changes with temperature
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Solubility of Sucrose vs. Temp
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Think About This
• What are some solids that are very soluble in water?
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Types of Solutions
• Unsaturated Solution
– A solution that has less solute than the maximum solubility
• Saturated Solution
– A solution that has reached solubility
• Supersaturated Solution
– A solution that exceeds solubility
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Solubility Demo
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Solubility of Gases
• Gases can dissolve in liquids
– Oxygen or Carbon Dioxide in water
• Depends on Pressure and Temperature
– High pressure and low temp increases solubility of gases
• High pressure forces more gas into the liq.
• Low temp. slows down particles so fewer gas particles escape the surface.
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Implications of Gas Solubility
• Habitat for specific fish species
– Trout live in cold streams
• Water is cold so more oxygen dissolves
– Bullheads live in prairie potholes
• Water is warm so less oxygen dissolves
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Implications of Gas Solubility
From: http://tellus.ssec.wisc.edu/outreach/teach/ideas/kotoski/Minifact_Sheets/Minifact4_Dissolved_Oxygen.pdf
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Implications of Gas Solubility
• Scuba Diving
– Scuba Divers cant rise to the surface quickly
– Pressure in deep water is high
– More gas dissolves in their blood
– A quick rise will cause gas to bubble out
– Causes the Bends
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Homework
• p.479 37,38,39,40,43,86,87,93
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Objectives
• Explain molarity
• Use molarity in calculations
• Describe how to prepare a solution of a given concentration
• Describe the process of dilution
• Calculate dilution problems
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Concentration
• Concentration refers to the amount of solute in a given amount of solvent
– General Terms
• Dilute
• Concentrated
–Not very specific
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A Bad Time 4 General Concentrations
“Let’s give Mr. Hermansen a dilute solution of potassium chloride and see how he does.”
What kind of solution?!?!?!?
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Molarity
• Moles of solute per liter of solution
– Molarity = Moles/Liters
• Units of mol/L
• Also (M)
• Can be said “Molar”
• Ex. 3.0 mol/L = 3.0M = Three Molar
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Molarity
• Most common way of expressing concentration in chemistry b/c involves moles
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Simple Examples
• How many moles of NaCl are required to prepare 1.0L of a 2.0 M solution
– 2.0 mol
• How many moles of NaCl are required to prepare 500mL of a 2.0 M solution
– 1.0 mol
• What volume of a 2.0 M solution of NaCl contains .20 mol NaCl
– 0.10L
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Example• Calculate the molarity of a solution made by
dissolving 11.5 g NaOH in 1.5L of solution.
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Example• Calculate the molarity when 1.56 g of HCl is
dissolved in 26.8 mL of solution.
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Example• What mass of potassium dichromate is
required to prepare 200.0 mL of a .200M solution?
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Preparation of Solutions
• Determine the mass of solute necessary• Mass solute into a weight boat• Transfer solute to volumetric flask of proper
volume – Rinse weigh boat
• Add water to flask and dissolve solute• Fill to line• Insert stopper • Invert to stir
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Lets Make a Solution
• 250.00mL of a 0.018M solution of copper (II) nitrate
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Dilution• Process of making a solution less
concentrated by adding solvent
• Dilution equation
C1 V1 = C2 V2
• C is concentration, V is volume
• Concentration times volume = moles
• When a sample of solution is diluted the number of moles does not change
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Example• Explain how would you prepare 300.0 mL of
1.5 M HCl from 6.0 M HCl
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Lets Dilute Our Solution
• Explain how to prepare 100.00mL of 0.009 M Copper (II) Nitrate from 0.018 M stock solution.
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Better Concentration Information
“Let’s give Mr. Hermansen a 0.10M solution of potassium chloride and see how he does.”
Thanks Doc..10M
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Homework
• p. 482 #’s 109,110,111,114,115,117
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Solution Preparation
Sodium Hydrogen Carbonate
1. 0.025M2. 0.018M3. 0.036M4. 0.0095M5. 0.013M6. 0.0082M7. 0.010M8. 0.020M
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Objectives
• Complete stoichiometry calculations involving solutions
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Solution Stoichiometry
• Stoichiometry with solutions is VERY similar to other types of stoichiometry.
• Need to find moles based on volume and concentration
– Concentration times Volume = Moles
• General Process
Volume * Concentration * Mole Ratio * Molar Mass
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Example• What mass of Lead (II) Sulfate is produced
when 2.00L of 0.025M of Sodium Sulfate is mixed with an excess of Lead Nitrate?
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Example• What volume (mL) of 0.100M HCl is
required to neutralize 25.0 mL of 0.350M Sodium Hydroxide?
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Example• What mass of Silver Chloride is produced
when 23.0 mL of 0.20M Sodium Chloride is mixed with 50.0 mL of 0.11M Silver Nitrate?
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Homework
• p. 482 #’s 128,130ab,132
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Demo
• Start water boiling!
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Objectives
• Explain what is meant by freezing and boiling point
• Describe colligative properties
• Calculate changes in boiling and freezing points
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Boiling Point
• Temperature at which the vapor pressure of the liquid is equal to the atmospheric pressure– Water boils at 100ºC
• At this temperature the vapor pressure of water is 760mmHg
– Ethanol boils at 78.4ºC• At this temperature the vapor pressure
of ethanol is 760mmHg–Ethanol is more volatile than water
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Demo
• Boiling water with a vacuum
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When Liquids Boil
Q) What happens to the temperature of a liquid when it is boiling?
R) The temperature stays the same
- The heat energy is used to overcome the intermolecular forces present in the liquid.
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Freezing Point
• Temperature at which a solid and liquid have the same vapor pressure.
– Water freezes at 0ºC
– Ethanol freezes at -115ºC!
• Freezing Pt. and Melting Pt. are the same
Q) Is freezing exothermic or endothermic
R) Exothermic
- Energy must be lost to freeze
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Colligative Properties
• Properties that depend on the number of solute particles present but NOT their identity.
– Only the number of dissolved particles present matters
• Properties of solvents change when a solute is added
– Solutions have different properties than the solvent that is present
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Which Bottle Contains More Dissolved Particles?
0.10M KI
0.10M CH3OH
0.10M Na2S
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Colligative Properties
• When a solute is added to a solvent the vapor pressure decreases
• What will this do to the boiling point?– The boiling point increases
• Boiling Point Elevation– The freezing point decreases
• Freezing Point Depression– Can you think of any examples?
• Salt on Roads, Antifreeze
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Boiling Point Elevation
• ΔTb = Change in the boiling point• Kb = Boiling point elevation constant
Units of (ºC*kg/mol)– Constant for a given substance
• Data found in table 12.2 on page 474– If ΔTb for salt water is 2.2ºC then the new
B.P. is 102.2
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The van’t Hoff Factor
• i = The van’t Hoff factor– Tells you how many moles of particles a
solute would form if 1 mole dissolves– All covalent compounds have an i of 1
• They do not ionize– i for ionic compounds equals the # of ions
• NaCl = 2• MgCl2 = 3• AlCl3 = 4
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Freezing Point Depression
• ΔTf = Change in the freezing point
• Kf = Freezing point depression constant Units of (ºC*kg/mol)
– If ΔTf for salt water is 2.2ºC then the new F.P. is -2.2ºC
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Example• What is the freezing point of water if 0.25
mol. of glucose is dissolved in 500mL of water? Kf=1.86 ºC*kg/mol?
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Example• What is the new boiling point of water when
100. grams of NaCl is dissolved in 2.0L of water? Kb=0.51 ºC*kg/mol
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Example• Ethylene Glycol (C2H6O2 62.08 g/mol) is
placed in radiators to prevent the water from boiling. How many moles of EG must be added to 10.0L of water to make the water boil at 110.00 ºC? Kb=0.51 ºC*kg/mol
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Homework
• p. 483 #’s 135,137,143,144,145,148