Solutions

1. Solutions Unit 11

2. Water • Think back to the structure of water: • The charges on water can attract other “things” which makes water the universal solvent.

3. Water • Also, because of the different charges water molecules can bond to one another. • This hydrogen bonding gives water many of its unique properties.

4. Properties of Water • Hydrogen bonding causes water to have some unique properties: • high surface tension • high specific heat • low vapor pressure • high heat of vaporization

5. Liquid vs. Solid Water • Ice is less dense than liquid water • Hydrogen bonds hold the water molecules in place in the solid phase. • The structure of the ice is a regular open framework of water molecules arranged like a honeycomb. • Ice has “holes” in its structure. • Because it is less dense ice will float.

7. Water of Hydration • Hygroscopic • a compound that removes water from the environment (hydrates) • Desiccant • a hygroscopic substance that removes water from the atmosphere to keep the environment dry • used as a drying agent (shoes, electronics) • Deliquescent • substances that removes so much water from the air that it turns into a solution

8. Solutions • Solutions are homogeneous mixtures • Solute is the dissolved substance • Seems to “disappear” or “Takes on the state” of the solvent • Solvent is the substance the solute dissolves in • Does not appear to change state • Solutions in which the solvent is water are called aqueous solutions • Water is often called the universal solvent • When in question, the solvent is the substance that you have more of.

10. The Solution Process - Solvation • When ionic compounds dissolve in water they dissociate into ions • ions become surrounded by water molecules - hydrated • When solute particles are surrounded by solvent molecules we say they are solvated • Solvation of molecular compounds  dissolving • Solvation of ionic compounds  dissociation • Rate of solvation affected by: • Agitation • Temperature • Particle Size

11. Solubility • When one substance (solute) dissolves in another (solvent) it is said to be soluble • When one substance does not dissolve in another they are said to be insoluble • Rule of Thumb: “Likes dissolve likes” (polarity) • There is a limit as to how much solute can dissolve in a solvent. • this is called solubility • Ex.  at 20oC 64.2 g NiCl2 in 100 g H2O  g/L with a gas solute and liquid solvent

13. Solutions & Solubility • The solubility of the solute in the solvent depends on the temperature • Higher Temp = Larger solubility of solid in liquid • Lower Temp =Larger solubility of gas in liquid • The solubility of gases depends on the pressure of the gas above the solution • Higher pressure = Larger solubility • Henry’s Law: S1 P1 S2 P2 =

14. Describing Solutions - Qualitatively • A concentrated solution has a high proportion of solute to solution – lots of solute • A dilute solution has a low proportion of solute to solution – little solute • A saturated solution has the maximum amount of solute that will dissolve in the solvent • Depends on temp • An unsaturated solution has less than the saturation limit • A supersaturated solution has more than the saturation limit • adding a seed crystal will initiate the crystallization of this supersaturated solution

15. Review • If a solution can hold a max of 100 g of solute and the solution currently is holding 99 g of solute then it is said to be ___________________. • Water’s low vapor pressure is a result of the ____________ ___________ that occurs in water. • As the pressure of a gas above a liquid increases, the solubility ______________.

16. Describing Solutions Quantitatively • Solutions have variable composition • To describe a solution accurately, you need to describe the components and their relative amounts • Concentration = amount of solute in a given amount of solution

17. moles of solute liters of solution molarity = Solution ConcentrationMolarity (M) • moles of solute per 1 liter of solution • used because it describes how many particles of solute in each liter of solution • If a sugar solution concentration is 2.0 M , 1 liter of solution contains 2.0 moles of sugar, 2 liters = 4.0 moles sugar, 0.5 liters = 1.0 mole sugar, etc.

18. Examples - Molarity • An aqueous solution has a volume of 2.0 L and contains 36.0 g of glucose (C6H12O6). What is its molarity? • If you want to make 250. mL of a .500 M solution of copper (II) chloride in water, how many grams of solute will you need?

19. Dilution • Dilution is adding solvent to decrease the concentration of a solution • The amount of solute stays the same, but the concentration decreases • Dilution Formula Ms x Vs = Md x Vd Concentrations and Volumes can be most units as long as consistent

20. Examples - Dilution • How much stock NaCl (aq), which is 1.00 M, is required to make 100.0 mL of a .200 M NaCl(aq)? • How would you prepare 500. mL of a .100 M solution of MgSO4 from a stock solution of 2.00 M MgSO4?

21. Molality (m) • molality (m) = mol solute kg solvent • How many grams of KI must be dissolved in 500. g of H2O to produce a .0600 m solution? • But WHY???? • That’s so annoying!!!!! • Who would make that up???? • Molality: Is used in situations where temp. is an issue. • Why would you not want to use a volume based concentration, like Molarity, with highly varying temperatures?

22. Review • 1) Explain how you would prepare 500. mL of a 0.33 M solution of calcium chloride. • 2) Determine the molal concentration of a solution in which 320. grams of glucose C6H12O6 are dissolved in 4000 grams of water. • 3) If the solubility of Carbon Dioxide in water is .625g/L at 790. Torr and 25.0 degrees Celsius; what would be the solubility at constant temperature and 3.00 atm? • 4) If water and ethanol mix in all proportions they are said to be? • If however water and oil don’t mix at any proportion they are said to be?

23. Mole Fraction • mole fraction (χ) = mol solute mol solution • What is the mole fraction of solute and solvent in a .150 m solution of KCl in H2O?

24. Colligative Properties • Depend only on the number of solute particles present, not on the identity of the solute particles. • Among colligative properties are • Vapor pressure depression • Boiling point elevation • Freezing point depression otic pressure

25. Vapor Pressure Depression • Vapor pressure is caused by solvent evaporation • Adding a non-volatile solute ALWAYS lowers the vapor pressure • more solute = less solvent at surface = ↓ evaporation = ↓ vapor = ↓ vapor pressure • Amount of solute is the only thing that alters this property – colligative property • 1 mol glucose  1 mol glucose • 1 mol NaCl  1 mol Na+ and 1 mol Cl- • van’t hoff factor – mol of particles per mol of solute

26. K Cl K+ Cl- K+ SO4 K K SO42- K+ Cu SO4 Cu+2 SO42- Dissociation KCl(aq) = K+ (aq) + Cl- (aq) CuSO4(aq) = Cu+2(aq) + SO42-(aq) K2SO4(aq) = 2 K+ (aq) + SO42-(aq)

28. Boiling Point Elevation • For something to boil the vapor pressure = atmospheric pressure. • Adding solute = ↓ vapor pressure = ↑ boiling point • Example – adding salt to water before cooking spaghetti noodles causes that water to boil at a hotter temperature, which leads to the noodles cooking faster

29. Freezing Point Depression • During freezing, the particles of a solid take on an orderly pattern. • Adding solute disrupts this pattern, so even more energy must be removed for the solution to solidify. • Examples – applying salt to icy roads helps prevent the water from freezing and the use of antifreeze in vehicles.

31. BPE and FPD Calculations • BPE/FPD = kmi • k = a different constant for each solvent = positive for BP (.512 kg0C/mol for H2O) = negative for FP (-1.86 kg0C/mol) • m = molality • i = van’t hoff factor

32. Examples – Colligative Calculations • What is the boiling point of a solution that contains 1.25 mol of CaCl2 in 1400. mL of H2O? • What is the freezing point of 72.3 g of magnesium sulfate in 1350 mL of H2O • PRE-LAB CALCULATION: The freezing point for H2O is lowered to -0.3900C when 3.90 g of a non-volatile molecular solid is dissolved in 475 g H2O. Calculate the molar mass of the substance.

36. Heterogeneous Aqueous Systems • Two types: 1. Suspension - a mixture from which particles settle out - suspension differs from a solution because the particles of a suspension are much larger & do not stay suspended (> 1000 nm) - Ex. – muddy water, gravy 2. Colloid - a mixture containing intermediate-sized particles larger than those in solutions but smaller than those in a suspension (1-1000 nm) - Ex. – glue, paint, smoke, milk, fog

37. Comparison of Particle Size • Solution • Less than 1 nm • Colloid • Between 1 nm and 1000 nm • Suspension • Larger than 1000 nm

38. Tyndall Effect • The scattering of visible light by colloidal particles is called the Tyndall effect. • Good test to distinguish between a solution and a colloid.