Slide 1

Slide 2

Types of mixtures Heterogeneous mixture: Heterogeneous mixtures do not blend smoothly. Individual substances can still be seen Suspension: A suspension is a heterogeneous mixture where particles will settle to the bottom if left undisturbed. Example: sand in water. Colloids: Colloids are heterogeneous mixtures where particles will not settle to the bottom. Example: Milk

Slide 3

Homogeneous Mixtures Homogeneous mixtures are mixtures where a substance is dissolved in another substance. These are also called solutions. Parts of a solution: Solute: The substance that is dissolved Solvent: The substance that does the dissolving When a solute readily dissolves in a solvent it is said to be soluble. Two liquids that can be combined to make a solution are said to be miscible.

Slide 4

A solute that does not readily dissolve in a solvent is said to be insoluble. Two liquids that do no mix to form a new solution are said to be immiscible.

Slide 5

Solution Concentration Concentration is the measure of how much solute is dissolved in the solution. There are many ways to measure concentration 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 Molality: Moles of solute/Kilograms of solvent Mole Fraction: Moles of solute/Moles of solute + Moles of solvent

Slide 6

Percent by mass In order to maintain a sodium chloride concentration similar to ocean water, an aquarium must contain 3.6g NaCl per 100.0 g of water. What is the percent by mass of NaCl in the solution? 3.5% What is the percent by mass of NaHCO 3 in a solution containing 20.0 g of NaHCO3 dissolved in 600.0 mL of H 2 O? 3%

Slide 7

Percent by Volume What is the percent by volume of ethanol in a solution that contains 35mL of ethanol dissolved in 155mL of water? 18%

Slide 8

Molarity A 100.5 mL intravenous solution contains 5.10 g of glucose. What is the molarity of the solution? The molar mass of glucoes is 180.6 g/mol 0.282M

Slide 9

If you have 1500 g of a bleach solution with a percent by mass of NaOCl of 3.62 % how many grams of NaCOl are in the solution? 54.3 g What is the percent by volume of isopropyl alcohol in a solution that contains 24 mL of isopropyl alcohol in 1.1 L of water? 2.1 % Calculate the molarity of 1.6 L of a solution containing 1.55 g of KBr. 8.13 x 10 -3

Slide 10

Molality In the lab a student adds 4.5 g of NaCl to 100.0 g of water. Calculate the molaity of the solution. 0.77 mol/kg

Slide 11

Preparing Solutions How would you prepare 1 L of a 1.5 molar CuSO 4 Solution? Diluting Solutions: Dilution equation M 1 V 1 = M 2 V 2 Calculate the new volume if you wanted to dilute the above solution to 0.5 molar.

Slide 12

Diluting Solutions Remember Molarity = (Moles of solute/Liters of solvent) The Dilution Equation: M 1 V 1 = M 2 V 2 Example: What volume of a 2.00M CaCl 2 solution would you use to make 0.50 L of a 0.300M solution? M 1 = 2.00M V 1 = ? M 2 = 0.300M V 2 = 0.5L (2.00M)(V 1 ) = (0.300M)(0.5L) V 1 = 0.075 L

Slide 13

What volume of a 3.00M KI solution would you use to make 0.300 L of a 1.25M KI solution? 125mL How many milliliters of a 5.0M H 2 SO 4 stock solution would you need to prepare 100.0mL of 0.25M H 2 SO 4 ? 5.0mL

Slide 14

Mole Fraction To express concentration as a mole fraction we need to know the number of moles of solute and the number of moles of solvent. The mole fraction is calculated by dividing the number of moles of either the solute or solvent by the total number of moles in the solution.

Slide 15

Example If we have a solution that contains 36 g of HCl and 64 g of H 2 O. What are the mole fractions of HCl and H 2 O? HCl: H 2 O:

Slide 16

Practice What is the mole fraction of NaOH in an aqueous solution that contains 22.8% NaOH by mass? If the mole fraction of H 2 SO 4 is an aqueous solution is 0.325, what is the percent by mass of H 2 SO 4

Slide 17

Factors Affecting Solvation Many physical factors affect the solubility of a solute. Such as: Temperature Pressure Polarity Solvation is the process of surrounding solute particles with solvent particles.

Slide 18

Aqueous ionic solutions Many ionic compounds are soluble in water. Why? Water is a polar molecule. Why might an ionic compound not be water soluble?

Slide 19

Heat of solution The overall energy change that occurs during the formation of a solution is called the heat of solution. certain solvation processes are exothermic while others are endothermic.

Slide 20

Factors that affect solvation Agitation: Stirring or shaking moves solvated particles away from the surface of the solid and allows more solute particles to move into solution. Surface Area: Breaking the solute into small pieces increases the available surface area and increase the rate of solvation. Temperature: As temperature increase so does the rate of solvation.

Slide 21

Types of solutions Unsaturated: A solution that more solute could be dissolved in. Saturated: In a saturated solution the rate of solvation and the rate of crystallization are in equilibrium. Supersaturated solution: A super saturated solution is a solution that has more solute dissolved in it that it would normally allow.

Slide 22

Solubility of Gases Unlike solid solutes the solubility of gases (CO 2, or O 2 for example) increases as the temperature of the solvent goes down. The solubility of gases is also affected by pressure. How? Think about soda.

Slide 23

Henrys Law Henrys Law states that at a constant temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure of the gas above the liquid. Henrys Law S 1 /P 1 = S 2 /P 2 Example: If 0.85g of a gas at 4.0 atm of pressure dissolves in 1.0 L of water at 25 o C, how much will dissolve in 1.0 L at 1.0 atm at the same temperature? 0.21 g

Slide 24

Decompression Sickness Decompression sickness is a term used to describe the effects of a drop in external pressure on a persons body. One common symptom of decompression sickness is AGE, Arterial Gas Embolism. The bubbles that form in the blood stream from gas leaving solution can restrict blood flow.

Slide 25

Slide 26

Boiling Point Elevation Dissolving a solute in a solution raises the boiling point of the solution. This is called boiling point elevation. We can calculate the difference in boiling point ( T b ) by multiplying a solutions molality by a constant we look up for our particular solvent (K b ). T b = K b m

Slide 27

What is the boiling point of a 0.625m aqueous solution of a nonelectrolyte?

Slide 28

What is the boiling point of a 0.4m aqueous solution of Ca(OH) 2 (A strong electrolyte) K b = 0.512 o C/m

Slide 29

Molal Boiling Point Elevation Constants. SolventBoiling Point ( o C)K b ( o C/m) Water100.00.512 Benzene80.12.53 Chloroform61.73.63 Ethanol78.51.22

Slide 30

Freezing Point Depression Adding a solute to a solvent also lowers the solvents freezing point. The equation is very similar to boiling point elevation. T f = K f m SolventFreezing Point ( o C) K f ( o C/m) Water0.001.86 Benzene5.55.12 Ethanol-114.11.99

Slide 31

Example Sodium chloride is often used to prevent icy roads and to freeze ice cream. What are the boiling point and freezing point of a 0.029m AQUEOUS solution of sodium chloride? BP: 100.03 0 C FP: -0.11 o C

Slide 32

Slide 33

Calculate the freezing and boiling point of a solution containing 6.42 g of sucrose (a nonelectrolyte) in 100.0 g of water. T f = -0.350 o C T b = 100.096 o C Calculate the freezing and boiling point of a solution containing 23.7 g of Copper (II) sulfate (A strong electrolyte) in 250 g of water T f = -2.02 o C T b = 100.606 o C Calculate the freezing and boiling point of a solution containing 0.15 moles of the compound naphthalene, a nonelectrolyte in 175 g of benzene (Normal F.P. = 5.5, Normal B.P. = 80.1, K f = 5.12, K b = 2.53) T f = 1.1 o C T b = 82.3 o C