1. 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

2. 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

3. Parts of a Solution Solute Solvent
Substance that dissolves
Solvent
Dissolves the solute
Water is the most common
Solutions are also known as homogeneous mixtures!

4. Solution Vocabulary Soluble Insoluble Immiscible Miscible
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)

5. 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

6. Types of Solutions Gas Liquid Solid Gas in Gas—air (nitrogen + oxygen)
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)

7. 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

8. Aqueous Solutions Of ionic compounds Of molecular compounds
Ionic Compounds…Solid sodium chloride dissolves as its ions are surrounded by solvent water molecules. The polar water molecules orient themselves differently around positive and negative ions.
Molecular Compounds…polar molecules contain O—H bonds. Forces between polar water molecules and polar solute molecules break these o—H bonds and the solute dissolves.
Oil (nonpolar) + water (polar) DO NOT make a solution “not like!”

9. 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

10. 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

11. 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

12. 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

13. 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

14. 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

16. 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

17. Solution Concentration
I will state the concentrations of solutions in different ways
I will calculate the concentrations of solutions

18. 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

19. 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

20. Concentration Description
Concentration Ratios
Concentration Ratios
Concentration Description
Ratio
Percent by Mass
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑋 100
Percent by Volume
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑋 100
Molarity
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑙𝑖𝑡𝑒𝑟 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
Dilutions
𝑀 1 𝑉 1 = 𝑀 2 𝑉 2
Molality
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑘𝑖𝑙𝑜𝑔𝑟𝑎𝑚 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡

21. Percent by Mass

22. Percent by Volume

23. Molarity (M) Molar Concentration MUST know 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

24. Molarity (M)

25. Diluting a Solution

26. Diluting a Solution

27. Diluting a Solution

28. 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

29. 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

30. Molality (m)

31. 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)

32. Electrolytes Ionic compounds (some molecular compounds) Strong Weak
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

33. Nonelectrolytes Many molecular compounds Example 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

34. Vapor Pressure Lowering
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

35. Predicting Effect of Solute on Vapor Pressure
Based on electrolyte or nonelectrolyte
Examples
1 mole nonelectrolytes = same relative effect on Vp
Glucose 1mole
Sucrose 1mole
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-

36. Boiling Point Elevation
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

37. 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

38. Calculate Solution Boiling/Freezing Point

39. Pure Solute vs Solution

40. Osmosis and Osmotic Pressure
NOT ON TEST
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

41. 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

42. Heterogeneous Mixtures
Contain substances that exist in distinct phases
Regions with uniform composition and properties
Different substances remain physically separate
Types
Suspensions
Colloids

43. 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

44. 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

45. Suspensions

46. 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

47. Colloids

48. 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

49. Brownian Motion

50. 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

51. Tyndall Effect
The beam of light is visible in the colloid because of light scattering

52. Review