Understanding Solutions: Types, Separation, Concentration, and Solubility

1. Ch 13Solutions

2. Solutions • Mixture: combination of 2 or more substance that ca be separated by physical means • Solution homogeneous mixture (even throughout) of 2 or more substances uniformly dispersed throughout a single phase • Suspension: mixture which particles are dispersed throughout a liquid or gas.

3. Solution/homogeneous mixture: • Particles evenly distributed. • Solvent: substance which the solute is dissolved • Solute: substance that is dissolved in the solvent. • Colloids/heterogeneous mixture: • Colloid mixture consisting of tiny particles that are intermediate in size between those in solution and suspension. Suspended in a liquid, solid or gas (milk) • Use Tyndell effect to test for colloids

4. Separating mixtures • Filtration, evaporation and centrifuges. Depends on the properties of the solution • Chromatography: separation due to difference of attraction of the components of the mixture. • Mobile phase: solvent move along a solid phase • Paper chromatography: separate dyes(some dyes more attracted to paper than others. • Distillation: separation due to different boiling pts. • Lower boiling pt are removed first (alcohol/water mix in fermentation process) alcohol boils first then collected by a distilling column. • Remove drinking water from salt water, fractional distillation of crude oil.

5. Concentration/molarity • Concentration: amount of a particular substance in a given quantity of a solution. • Calculating concentration/unit dimension analysis • Parts per million (ppm): # of grams of solute in 1 million grams of solution • Molarity (M): moles of solute per liter of solution • Molality (m): moles of solute per kg of solvent. • Calculate the following: • How many parts per million of mercury are in a sample of tap water with a mass of 750g contain 2.2 mg Hg? • A sample of water from a stream has a mass of 625g. The sample is found to contain 0.35 mg of arsenic.

6. Preparing solutions pg 463 • Solute has volume • Determine mass of solute needed •  moles to mass using molar mass • Add solvent (H2O) to flask then add solute to mix. • Add H2O until volume of solution is made • Calculating molarity: • 3 variables mass/mol solute, volume & [ ] • Know two of three and solve for the third • Determine the molarity of a solution by dissolving 16.9 g NaOH in enough water to make 250.0 mL of solution • What mass of ammonium dichromate is need in 500.0 mL of solution and 0.2384M [ ]

7. Molarity/stoichiometry • Know [ ] and volume of solution to predict mass of product or vise versa. • [ ] x vol<-> moles<->moles<->mass • mole ratio molar mass • What volume in mL of a 1.50M HCl solution would needed to to react completely with 28.4g of Na2CO3 to produce CO2, H2O & NaCl? • A zinc bar is placed in 435 mL of a 0.770M CuCl2 solution. What mass of zinc would be replaced by copper if all the copper ions were used up?

8. Solubility/polarity • Solubility: ability of one substance to dissolve into another at a given temperature and pressure in terms of [ ] in a saturated solution. • Polarity determines if a solute/solvent combination will form a solution. • Like dissolves like (polar/polar or nonpolar/nonpolar) • Vitamin C water soluble: • Cannot be stored in the body, synthesis collagen, scurvy (lack of vit C) • Vitamin A fat soluble (can be stored) • Yellow vegetables, good vision, skin & bone growth

9. Like dissolves like • Nonpolar molecules: LDF of attraction between molecules. • Nonpolar molecules soluble in other nonpolar molecules. • Miscible: 2 or more liquids that are able to dissolve into each other. • Polar molecules; opposite charges attract 2 molecules to each other. • Polar molecules soluble in polar molecules. • Mix polar/nonpolar molecules • Low solubility in each other • Immiscible: 2 or more liquids that do not mix

10. Solubility's of solids • Dissolving occurs at the surface between two molecules • Rate of dissolving effected by: • Mixing (stirring, more contact) • Increase surface area (crushing) • Increase temperature: increase speed of solvent. • Solubility of solid generally increases with temperature (pg 472 figure 12) • Depends on net energy flow during process • Net endothermic: increase temp increase solubility • Net exothermic: increase temp decrease solubility

11. Enthalpy/entropy & solubility • Dissolving ionic cpds • Separate ions from lattice to individual dissolved ions (dissociation) • When water is solvent its called hydration • Enthalpy (H) • Separate ion require +H • As ions attract to water -H • H changes nearly cancel • Entropy (S) • Increase S as ions scattered throughout solution • Solubility of ionic cpds: • Based on size and change of ions • Decrease size and charge increase solubility (pg 473 table 2)

12. Solubility cont. • Saturation: maximum amount of solute dissolved in an amount of solvent. • Any addition of solute will settle to bottom of container (dynamic equilibrium) • Unsaturated: more solute can still be dissolved. • Super saturated: a solution holding more dissolved solute than what is required to reach equilibrium at a given temperature.( pg 475 figure 16 hand warmer) • Solubility equilibrium: • Physical sate of opposing processes of dissolution & crystallization of a solute occur at equal rates • Amount of solute dissolved remains constant.

13. Gases dissolved in liquid • Solubility depends on pressure and temperature • Henry’s law: at constant temperature solubility of gases in liquid is directly proportional to the partial pressure of gases on the surface of the liquid • Gases are less solubility as the temperature of solution increases. • Increase solvent molecular motion allows gas molecules to escape.

14. Physical properties of solutions • Electric conductivity in solution • Ability to conduct electric current • Electrolytes: substance that dissolves in water conducts an electric current (mobile ions) • Strong electrolyte: completely dissociates into ions • Weak electrolyte: provide a few ions in solutions • Nonelectrolyte: polar molecules that do not produce ions in water, nonconductor. • Acid in water (polar molecule) produces hydronium ion (H3O+) & is an electrolyte • Tap water has electrolytes dissolved & is a conductor

15. Colligative properties • Property of a system that is determined by the number of particles present in the system but independent of the properties of the particles themselves. • When solute is added to solvent • Vapor pressure and freezing pt decrease and boiling pt increases. => extends the liquid phase of solvent • Nonvolatile solute: • Depends on [ ] • When ionic cpds effects colligative property by a whole # factor equal to the # of ions in solution. • Sugar (aq) factor = 1, NaCl(aq) factor = 2, CaCl2(aq) factor = 3

16. Solute(aq)/decrease VP • Vapor pressure : ability of a solvent to change phase. • Adding solute to solvent => less [ ] of solvent molecules and decrease VP (pg 483 figure 25) • Boiling pt elevated & freezing pt depressed (g) 1.0 Solvent freezes At a lower T & Boils at a higher T Vapor Pressure (atm) solvent (l) solution (s) Temperature (oC)

17. Surfactants • Cpd (like soap) that concentrates at the boundary surface between two immiscible phases (s/s, l/l, or g/g) • Detergent (surfactant): water soluble cleaner that can emulsify dirt & oil. • Soap: substance that is used as a cleaner and dissolves in water • Sodium & potassium salts of fatty acids.(pg 484, figure 26) • Emulsifying agent (soap): connect polar/nonpolar substance to make a mixture oil-soap-water • Emulsion: mixture of two or more immiscible liquids in which one liquid is dispersed in the other.

18. Hard water/soap • Hard water limits soap’s detergent ability • Some soap anions insoluble in water (Ca 2+ , Mg 2+ & Fe 2+ ) • Hard water has a high [ ] of these ions. => form ppt with soaps anions and form bath tub rings and soap scum. • Synthetic detergent (486, figure 28): • Will not form ppt. in hard water • Early form where not biodegradable => pollution problems. • Newer detergents are now biodegradable.