1. Name __________________________ Block_____ Chapter 17 Solutions and Molarity Some Definitions A solution is a homogeneous mixture of 2 or more substances. One constituent is usually regarded as the SOLVENT and the others as SOLUTES.

2. Parts of a Solution SOLUTE – the part of a solution that is being dissolved (usually the lesser amount) SOLVENT – the part of a solution that dissolves the solute (usually the greater amount) Solute + Solvent = Solution Water is called the Universal Solvent

3. Factors Affecting Solubility 1. Nature of Solute / Solvent Like dissolves like – polar substances dissolve polar while non-polar dissolve non-polar substances. 2. Temperature Solids/Liquids - Solubility increases with Temperature Gases - Solubility decreases with Temperature 3. Pressure Solids/Liquids - Very little effect Gases - Solubility increases with Pressure. (Henry’s law)

4. Intermolecular Forces or IMF Forces of attraction between molecules Intermolecular forces act between molecules holding them together –Ex. 2 molecules of H 2 O close together IMF determine what state of matter a substance will be at room temperature. –solid, liquid or gas Intramolecular Forces Bonds (Ionic and Molecular/Covalent) are examples of Intramolecular forces or forces within a molecule holding atoms together.

6. 3 types of Intermolecular Forces 1.Dispersion forces - Very weak IMF between adjacent nonpolar (equal sharing of electrons) molecules based on a momentary movement of electrons creating momentary partial charges. Ex.// CH 4 molecules Neon atoms 2.Dipole interactions - Attraction between adjacent polar molecules (having a polar molecules are attracted to each other. The positive dipole of one molecule is attracted to the negative dipole of another. Ex.// HCl molecules

7. Intermolecular Forces 3.Hydrogen bond: Attraction between polar molecules containing hydrogen bonded to very electronegative element (F, O, N) Hydrogen bonding is the strongest intermolecular force. It explains water’s high boiling point (100 0 C).

8. Temperature & the Solubility of Solids Solubility of most solids INCREASES at higher temperatures Solubilities of several ionic compounds (salts) as a function of temperature. MOST ionic compounds have greater solubility in hot water.

9. Temperature & the Solubility of Gases The solubility of gases DECREASES at higher temperatures

10. Phase Change Diagram Solid Liquid Gas temperature pressure melting freezing evaporating condensing sublimation deposition triple point critical point

11. Phase Change Diagram Triple Point – –where all 3 phases (solid, liquid and gas) coexist. Critical Point – –Beyond this point, it is no longer possible to distinguish between the gas and liquid phases—substance called a supercritical fluid. Substances can boil/freeze at different temperatures by changing the pressure.

12. Definitions Solutions can be classified as saturated or unsaturated. A saturated solution contains the maximum quantity of solute that dissolves at that temperature. An unsaturated solution contains less than the maximum amount of solute that can dissolve at a particular temperature.

13. Definitions Supersaturated Solutions contain more solute than is possible to be dissolved. Supersaturated solutions are unstable. The supersaturation is only temporary, and usually accomplished in the following way: Warm the solvent so that it will dissolve more, then cool the solution.

14. Electrolytes and Nonelectrolytes Electrolyte – A compound that conducts an electric current when it is in an aqueous solution or molten state. All Ionic Compounds are electrolytes because they dissociate into ions. Ex., NaCl, CuSO 4

15. Nonelectrolytes Non-electrolyte is a compound that does not conduct an electric current in either aqueous solution or molten state. Many molecular compounds are non-electrolytes because they are not composed of ions. Ex., Sucrose (C 12 H 22 O 11 ) Ethanol (C 2 H 5 OH)

16. Concentration of Solute The amount of solute in a solution is given by its concentration The amount of solute in a solution is given by its concentration. Molarity (M) = moles of solute liters of solution

17. Problem: 5.00 g of NaCl is dissolved in enough water to make 250 mL of solution. Calculate the Molarity. Step 1: Calculate moles of NaCl mm (NaCl) = (23 x 1) + (35 x 1) = 58 g/mole. 5.00 g x 1 mole = 0.08 mol NaCl 1 58 g Step 2: Convert 250 mL into L Since 1000 mL = 1 L, we will get 0.250 L. Step 3: Calculate Molarity Molarity = 0.08 moles = 0.34 M 0.250 L

18. When a solution is diluted, solvent is added to lower its concentration. The amount of solute remains constant before and after the dilution: moles BEFORE = moles AFTER M 1 V 1 = M 2 V 2 M 1 and V 1 are the beginning molarities and volumes M 2 and V 2 are the ending molarities and volumes Suppose you have 0.500 M sucrose stock solution. How do you prepare 250 mL of 0.348 M sucrose solution ? Concentratio n 0.500 M Sucrose 250 mL of 0.348 M sucrose Dilution What is the molarity of a 10 mL sample of 2.0 M aqueous HCl diluted to 40 mL Answer: (2.0)(10) = (M 2 )(40) so M 2 = 0.5 Molar HCl What is the molarity of a 10 mL sample of 2.0 M aqueous HCl diluted to 40 mL Answer: (2.0)(10) = (M 2 )(40) so M 2 = 0.5 Molar HCl

19. Evaporation and Vapor Pressure In a closed container  some particles at the surface of the liquid evaporate  these particles collide with the walls of the sealed container and produce a vapor pressure  over time, the number of particles entering the vapor increases and some particles condense and return to liquid state. Eventually – Rate of Evaporation becomes = Rate of Condensation Hence, a dynamic equilibrium is established between the vapor and liquid.

20. Vapor Pressure The Vapor Pressure of a liquid is the equilibrium pressure of a vapor above its liquid; that is, the pressure of the vapor resulting from evaporation of a liquid above a sample of the liquid in a closed container. An increase in the temperature of a contained liquid increases the vapor pressure. Boiling Point of a liquid is the temp. at which Vapor Pressure of liquid = Atmospheric Pressure NOTE: Normal Boiling point is Boiling point at 1 atm pressure.

21. Colligative Properties Physical properties of a solution differ from those of pure solvent used to make the solution. These properties are called Colligative Properties because they depend only on the NUMBER of solute particles in the solution, not on the KIND of solute particles. 3 important colligative properties of solutions are – 1. Vapor Pressure lowering 2. Freezing point depression 3. Boiling point elevation

22. 1. Vapor Pressure Lowering The Vapor Pressure of the solution is LOWER than that of the pure solvent. Why? The solute particles are surrounded by water molecules and this reduces the number of solvent molecules that can escape the liquid as vapor. Hence, A solution containing a solute that is nonvolatile (not easily vaporized) has a lower vapor pressure than that of the solvent. Nonvolatile solutes are Glucose, NaCl.

23. 2. Freezing Point Depression Pure Water Ethylene glycol/water solution The freezing point of the solution is LOWER than that of the pure solvent. 1. Ethylene glycol is used as an antifreeze and added to water in automobiles to depress freezing point of water below 0 0 C. 2. Salt is added to icy surfaces to make the ice melt.

24. 3. Boiling Point Elevation The boiling point of a solution is HIGHER than that of the pure solvent. Applications – The same antifreeze – Ethylene glycol – added to automobile engines to prevent freeze-ups in winter, protects the engine from boiling in summer.