Chapters 17 and 18: Water and Aqueous Solutions

1. Chapters 17 and 18: Water and Aqueous Solutions A. The Water Molecule (Ch. 17): • H2O is polar (has “dipoles”) ex. • H2O is bent due to its lone e– pairs • H2O covers about 75% of the earth’s surface

2. Intramolecular bonding – bonding within a molecule ex. covalent bonds • Intermolecular bonding – bonding between molecules ex. hydrogen bonds • Hydrogen bonding – when a hydrogen atom in one molecule bonds with a very electronegative atom (F, O, N) of another molecule ex. H2O

3. Solvation – how H2O molecules arrange themselves around ions that are dissolved in H2O ex.

4. Properties of H2O: • All due to hydrogen bonding (strong)! • High surface tension • High specific heat • High heat of evaporation • High boiling point

5. High surface tension – the inward force (or pull) that minimizes the surface area of a liquid ex.

6. High surface tension (cont.) • Surfactants – decreases surface tension by getting (H2O) molecules to “relax” and spread out ex. soaps or detergents

7. High specific heat – the amount of heat needed to raise the temperature of 1 g of a substance 1 C - it takes more energy (heat) to heat up H2O compared to other substances ex. H2O has a higher specific heat than sand and metals This explains why sand and metals get hot quickly when in the same amount of sun (energy/heat), as compared to H2O

8. High heat of evaporation – it takes a high amount of heat to evaporate H2O molecules - evaporation is the same as vaporization (l  g) - condensation (g  l) • When H2O vaporizes, bonds are broken • H2O has strong hydrogen bonds, which take a lot of heat/energy to break • High FP’s/MP’s/BP’s vs. other covalent bonds

9. Structure of H2O(l) vs. H2O(s) : ex. H2O(l) vs. H2O(s) • Liquid molecules are closer ; Solid farther apart • Ice has a bigger volume for the same mass • Ice has a smaller density (D = m / v) • Ice floats on liquid water

10. B. Solubility of Solutions (Ch. 18): • Solute – the substance being dissolved (s) • Solvent – what the solute is being dissolved in (l) - usually H2O • Solution – when the solute dissolves in the solvent (aq) ex. NaCl dissolves in water NaCl(s) = solute H2O(l) = solvent NaCl(aq) = solution

11. Solubility – the amount of solute that can dissolve in given amount of solvent at a given temperature (see solubility curves) • Miscible – liquids that dissolve each other • Immiscible – liquids that do not dissolve each other • “Likes dissolve likes”: • Polar/Polar = dissolves • Non-Polar/Non-Polar = dissolves • Polar/Non-Polar = does not dissolve

12. Factors affecting (to increase) solubility: • Stirring or shaking • Increase the temperature • Except gases (dec. temp to inc. solubility) • Increase the surface area (smaller particles) • Increase the amount of solvent (H2O) • Solubility curve – a graph that shows the different solubilities of many substances at different temperatures (see worksheet)

13. Saturated solution – contains the max. amount of solute per solvent at a given temp. - on the curve/line • Unsaturated solution – contains less than the max. amount of solute per solvent at a given temp. - below the curve/line • Supersaturated solution – contains more than the max. amount of solute per solvent at a given temp. - above the curve/line

14. Solubility Curve: ex.

15. C. Molarity (Ch. 18): • Molarity (M) – the concentration of a solution - formula: M = moles / L - units: moles/L or M (molar) - shows how dilute (small amount of solute per solvent) or concentrated (large amount of solute per solvent) - chemists need to know how much solute is in a solution for a reaction - “average” solutions are 1 M

16. ex. What is the molarity of a solution that contains 2 moles of sugar in 5 liters of water? M = moles / L M = 2 moles / 5 L M = 0.4 moles/L or 0.4 M ex. What is the molarity of a solution that contains 2 moles of NaCl in 5 liters of water? M = moles / L M = 2 moles / 5 L M = 0.4 moles/L or 0.4 M

17. ex. What is the concentration of a solution that contains 0.90 grams NaCl in 100 mL of H2O? M = moles / L 0.90 g NaCl x 1 mole = 0.015 moles 58.5 g 100 mL x __1 L__ = 0.1 L 1000 mL M = 0.015 moles / 0.1 L M = 0.15 M

18. Solving for other variables (moles or liters): • M = moles / L • moles = M x L • L = moles / M ex. How many grams are in 1500 mL of a 0.24 M solution of sodium sulfate? moles = M x L = (0.24 M) x (1.5 L) = 0.36 moles 0.36 moles Na2SO4 x 142.1 g = 51.16 g Na2SO4 1 mole

19. D. Dilutions (Ch. 18): • Dilutions – make the solution less concentrated - make the molarity lower - formula: M1 x V1 = M2 x V2 - the units of both volumes (V1 and V2) must be the same - cannot take out dissolved solute from a solution ; so add more solvent (H2O) - chemists make a “stock solution” of a high concentrated solution ; from a high M, they can add a specific amount of H2O to dilute the solution to the exact M they want

20. ex. What is the new molarity of 100 L of solution that was diluted from 50 L of a 5 M solution? M1 x V1 = M2 x V2 (5 M) x (50 L) = M2 x (100 L) 100 L 100 L M2 = (5 x 50) / 100 M2 = 2.5 M

21. ex. How many mL of a 3.0 M solution are needed to prepare 1 L of a 1.5 M solution ofMgSO4? M1 x V1 = M2 x V2 (3.0 M) x V1 = (1.5 M) x (1000 mL) 3.0 M 3.0 M V1 = (1.5 x 1000) / 3 V1 = 500 mL

22. “To how much…be added ” problems: • Solve for V, but subtract the 2 volumes to get the volume that should be added ex. To how much water should 25 mL of 8 M HCl be added to produce a 2 M solution? V2 = (25 x 8) / 2 = 100 mL = V2 - 25 mL = V1 Answer = 75 mL must be added

23. E. Colligative Properties (Ch. 18): • Colligative properties – depend on the number of particle units that are dissolved in a solvent • Particle unit – the number of “units” that each substance breaks up into when dissolved • Covalent compounds = 1 particle unit ex. sugar (C12H22O11) = 1 unit (no ions) • Ionic compounds = # of total ions ex. NaCl = 2 units (1 Na, 1 Cl) MgBr2 = 3 units (1 Mg, 2 Br) Al2O3 = 5 units (2 Al, 3 O)

24. Effects of the solute and solvent: • Solute • The more moles of solute, the greater the effect on the colligative property • The more particle units (assuming same number of moles), the greater the effect on the colligative property • Solvent • The higher the value of the solvent’s “constant”, the greater the effect on the colligative property (water has a very low constant)

25. Types of colligative properties: • Boiling point elevation • The boiling point is higher with added solute when compared to the pure solvent (just by itself) ex. Adding salt to water when boiling water to cook with (a myth?) • Freezing point depression • The freezing point is lower with added solute when compared to the pure solvent (just by itself) ex. Rock salt on icy roads

26. F. Suspensions and Colloids (Ch. 17): • Heterogeneous solutions – not uniform (not the same throughout) - 2 different phases - show the Tyndall effect • Tyndall effect – the scattering of light ex.

27. Suspensions – particles that settle out upon standing (stay at the bottom) - large particles (> 100 nm) - particles are visible - particles can be filtered out - exhibits the Tyndall effect ex. Pieces of clay in water Snow globes

28. Colloids – particles that do not settle out upon standing - small particles (1-100 nm) - particles are not visible - particles cannot be filtered out - exhibits the Tyndall effect - substances contain 2 states of matter that are dispersed evenly

29. ex. Colloids: (p. 491)

30. Solutions – particles that do not settle out upon standing - very small particles (< 1 nm) - particles are not visible - particles cannot be filtered out - do not exhibit the Tyndall effect - substances are homogeneous - look at p. 492 for a comparison of solutions, colloids, and suspensions