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S
Solutions•I will describe and categorize solutions•I will calculate concentrations of solutions•I will analyze the colligative properties of solutions•I will compare and contrast heterogeneous mixtures
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What are Solutions?
I will describe the characteristics of solutions and identify the various types
I will relate intermolecular forces and the process of solvation
I will define solubility and identify factors affecting it
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Parts of a Solution
Solute Substance that dissolves
Solvent Dissolves the solute Water is the most common
Solutions are also known as homogeneous mixtures!
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Solution Vocabulary
Soluble A substance that dissolves in a solvent Sugar is soluble in water
Insoluble A substance that does NOT dissolve in a solvent Sand is insoluble in water
Immiscible Any two liquids that can be mixed together but separate
shortly after you cease mixing them Oil and Vinegar are immiscible
Miscible Any two liquids that are soluble in each other Water and ethylene glycol are miscible (antifreeze)
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Characteristics of Solutions
It is NOT possible to distinguish the solvent from solute
Exist as solid, liquid (most common), or gas Depends on state of solvent
Examples Gas = air (oxygen, nitrogen, argon, carbon dioxide, etc) Solid = braces (titanium and nickel) Liquids = sugar water
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Types of Solutions
Gas Gas in Gas—air (nitrogen + oxygen)
Liquid Gas in Liquid—carbonated soda ( soda + carbon dioxide) Liquid in Liquid—vinegar (water + acetic acid) Solid in Liquid—ocean water (water + sodium chloride)
Solid Liquid in Solid—dental amalgam (silver + mercury) Solid in Solid—steel (iron + carbon)
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Solvation
The process of surrounding solute particles with solvent particles to form a solution This process in water is called hydration Attractive forces
solvent + solute > solute + solute
“like dissolves like” Based on polarity and bonding of particles as well as
intermolecular forces
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Aqueous Solutions
Of ionic compounds
Of molecular compounds
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Factors that Affect Rates of Solvation
Increased solvation if increased collision of particles Because solvation ONLY occurs when and where the
solute and solvent touch!
3 ways to increase collisions Agitating the mixture—stirring and shaking Increasing the surface area of the solute—breaking
solute into smaller pieces Increasing the temperature of the solvent
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Heat of Solution
Overall energy change that occurs during the solution formation process
Energy required to overcome attraction within solute or solvent (endothermic)—feels cold
Energy released when solute and solvent particles mix (exothermic)---feels hot
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Solubility
The MAXIUMUM amount of solute that will dissolve in a given amount of solvent at a specified temperature and pressure
Expressed in grams of solute per 100g of solvent
Unsaturated Solution Solvation > than recrystallization Contains LESS dissolved solute for a given temperature and
pressure than a saturated solution (more solute could still be added)
Saturated Solution Solvation = recrystallization Contains MAXIMUM amount of dissolved solute for a given
amount of solvent at a specific temperature and pressure
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Factors that Affect Solubility
Pressure
The nature of the solute and solvent
Temperature Many (not all)
substances are MORE soluble at high temps, than low temps
Gas solutes in liquid solvents are LESS soluble at high temps fast gas particles
escape liquid faster when heated
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Supersaturated Solutions
Contains MORE dissolved solute than a saturated solution at the same temperature. Formed at high temperatures Cooled slowly Ex Boiling water and adding lots of sugar
Unstable Stirring or tapping container = crystallization Ex. Rock Candy
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Pressure
Affects solubility of gaseous solutes Pressure Solubility
Explains how we make carbonated soda pressure to solubility of carbon dioxide
(supersaturated) Cap the soda = trapped carbon dioxide in bottle Uncap the soda = carbon dioxide can escape
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Henry’s Law
Explains the decreased solubility of the carbon dioxide contained in the soda after the cap is removed
At a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid S = g/L P = varies
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Solution Concentration
I will state the concentrations of solutions in different ways
I will calculate the concentrations of solutions
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Solution Concentration
A measure of how much solute is dissolved ins specific amount of solvent or solution.
May be used as a qualitative description Concentrated- contains a large amount of solute Diluted- contains a small amount of solute
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Expressing Concentration
May be used as a quantitative description Percent by mass Percent by volume Molarity Molality
Express concentration as a ratio of measured amounts of solute and solvent or solution
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Concentration Ratios
Concentration Ratios
Concentration Description Ratio
Percent by Mass
Percent by Volume
Molarity
Dilutions
Molality
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Percent by Mass
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Percent by Volume
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Molarity (M)
Molar Concentration M = Molar Examples:
1M solution = A liter of solution containing one mole of solute 0.1M solution = A liter of solution containing 0.1 mole of solute
MUST know Volume of the solution Amount of dissolved solute
1𝐿=1000𝑚𝐿
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Molarity (M)
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Diluting a Solution
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Diluting a Solution
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Diluting a Solution
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Colligative Properties of Solutions
I will explain the nature of colligative properties
I will describe four colligative properties of solutions
I will calculate the boiling point elevation and the freezing point depression of a solution
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Molality (m)
Temperature causes volume of a solution to change, causing Molarity to change.
To prevent running into that problem, scientists sometimes uses Molality instead since that doesn’t change with temperature.
1𝑘𝑔=1000𝑔
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Molality (m)
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Colligative Properties
Physical properties of solutions that are affected by the number of particles but NOT the identity of dissolved solute particles “depending on the collection”
Example: Vapor pressure lowering Boiling point elevation Freezing point depression Osmotic pressure (NOT on test)
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Electrolytes
Ionic compounds (some molecular compounds) Dissociate in water Conduct an electric current
Strong Produce many ions in solution Ex NaCl(s) Na+(aq) + Cl-(aq)
1 mole NaCl (aq) = 1 mole Na ions + 1mole Cl ions 1 mole NaCl (aq) = 2 moles solute particles in solution
Weak Produce only a few ions in solution
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Nonelectrolytes
Many molecular compounds Dissolve in solvents Do NOT ionize (do NOT dissociate
when dissolved….stay as one particle) Do NOT conduct electric current
Example Sucrose 1m sucrose solution = 1 mole sucrose
particles
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Vapor Pressure Lowering
Vapor pressure The pressure exerted in a closed container by liquid
particles that have escaped the liquid’s surface and entered the gaseous state
LOWERS due to the number of nonvolatile solute particles in solution Solute particles occupy some surface area Solvent particles have less surface area to escape to a
gaseous state
Nonvolatile solute One that has little tendency to become a gas
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Predicting Effect of Solute on Vapor Pressure
Based on electrolyte or nonelectrolyte
Examples 1 mole nonelectrolytes = same relative effect on Vp
Glucose1mole Sucrose1mole Ethanol 1mole
1 mole electrolytes = increasingly greater affect on Vp Sodium chloride (NaCl) 1 mole Na+ 1mole Cl-
Sodium sulfate (Na2SO4) 2 mole Na+ 1 mole SO4-
Aluminum Chloride (AlCl3) 1 mole Al+ 3 mole Cl-
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Boiling Point Elevation
Boiling vapor pressure = atmospheric pressure
What happens when a nonvolatile solute is dissolved in a solvent? At normal boiling point, vapor pressure is still LESS than
atmospheric pressure Must be heated to raise vapor pressure
Boiling point elevation The temperature difference between a solution’s boiling point and
pure solvent’s boiling point The greater the number of solute particles in the solution, the
greater the boiling point elevation
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Freezing Point Depression
Freezing Point Temperature (solvent) Particles NO longer have sufficient kinetic energy to overcome
the interparticle attractive forces The particles form into a more organized structure (solid)
What happens when a solute is dissolved in a solvent? Solute interferes with attractive forces among solvent particles PREVENTS solvent from entering solid state at normal freezing
point Freezing points of solutions are always LOWER than that of the
pure solvent
Freezing Point Depression The difference in temperature between its freezing point and the
freezing point of its pure solvent
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Calculate Solution Boiling/Freezing Point
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Pure Solute vs Solution
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Osmosis and Osmotic Pressure
Osmosis the diffusion of solvent particles across a semipermeable membrane Higher solvent concentration Lower solvent concentration EX. Kidney dialysis or uptake of nutrients by plants
Semipermeable membranes Barriers with tiny pores Allow some (NOT all) kinds of particles to cross Ex. Surrounding ALL living cells
Osmotic Pressure Amount of additional pressure caused by the water molecules that
moved into the solution Depends on # of solute particles in a given volume of solution
NOT ON TEST
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Heterogeneous Mixtures
I will identify the properties of suspensions and colloids
I will describe different types of colloids
I will explain electrostatic forces in colloids
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Heterogeneous Mixtures
Contain substances that exist in distinct phases Regions with uniform composition and properties
Different substances remain physically separate
Types Suspensions Colloids
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Heterogeneous Mixtures
Example: Blood
The blood cells are physically separate from the blood plasma
The cells have different properties than the plasma.
The cells can be separated from the plasma by centrifuging physical change
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Suspensions
A mixture containing particles that settle out if left undisturbed
Can be separated with a filter or by settling out
Suspended Particles Diameter > 1000 nm
Examples Cornstarch and water Sand and water Muddy water
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Suspensions
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Colloids
Heterogeneous mixture of intermediate size particles (between suspensions and solutions)
Can NOT separate with a filter or settling
Particles 1nm< diameter < 1000 nm
Example Homogenized Milk
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Colloids
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Brownian Motion
Erratic movement of colloid particles Dispersed particles make jerky, random movements Results from collisions of particles with dispersion
medium Prevents particles from settling out
Destroying a colloid (gives colliding particles enough energy to overcome electrostatic forces) Stirring in an electrolyte heating
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Brownian Motion
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Tyndall Effect
Particles are large enough to scatter light
Both concentrated & diluted colloids exhibit this effect Concentrated colloids look cloudy Diluted colloids may look clear
Suspensions also exhibit this effect
Solutions NEVER exhibit this effect
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Tyndall Effect
The beam of light is visible in the colloid because of light scattering
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