SOLUTIONS Solution – homogeneous mixture made up of very small particles; the size of individual molecules, atoms, or ions Parts of a solution solute- the substance that is considered to be the dissolved substance 2. solvent – the substance in which the solute dissolves

Aqueous solution (aq) water is the solvent Tincture alcohol is the solvent “like dissolves like” meaning that substances that contain similar bond types will dissolve in each other. Polar solvents will dissolve polar and ionic substances (molecule to ion attraction) but will not dissolve non polar substances (oil and water)

Typical properties: 1)solutions are homogeneous mixtures (aq) 2)dissolved particles will not come out of solution no matter how long it stands as long as it is covered to prevent evaporation 3) Clear and transparent- dissolved particles are too small to see. It will not disperse light

4) Because of extremely small particle size the solution will pass through the finest filters and leave no residue 5)A solution is considered to be a single phase even though the components may be of different phases before the mixture was formed

Types of solutions 1)gas solutions- gases or vapors dissolve in one another Example: Air 2)liquid solutions- liquid solvent in which a gas, liquid or solid is dissolved Ex: sugar water , carbonated beverages, antifreeze (ethylene glycol in water)

a. Miscible- two liquids that dissolve in one another- must have similar bonding b. Immiscible- two liquids that do not dissolve in one another. Like oil and water must have dissimilar bonding 3)Solid solutions – mixture of solids uniformly spread throughout each other at the atomic or molecular level.

EX: alloy –2 or more metals - brass is copper and zinc Amalgam-alloy when one of the metals is Hg- fillings are a Ag-Hg amalgam

Solubility- maximum amount of solute that can dissolve in a certain quantity of a solvent at a certain temperature Factors affecting solubility 1. Nature of the solute and solvent Example: 1 g of PbCl2 dissolves in 100g of water at room temp 200g of ZnCl2 dissolves in 100 g of water at room temp

Temperature. Generally solubility increases as temp. increases EXCEPTION: When the solute is a gas Increasing temp decreases solubility and vice versa.

3. Pressure For solid and liquid solutes no changes in solubility with changes in pressure, but for gases solubility increases with increasing pressure and vice versa Carbonated beverages are kept under pressure and when you decrease the pressure by opening them they lose the carbonation

Saturated – a solution that has dissolved in it all the solute it can normally hold at a given temperature. If any more solute is added it will fall to the bottom Unsaturated- a solution that contains less solute than it can hold at a certain temperature. If more solute is added it will dissolve

Supersaturated- Under special conditions, there are some solutions that can actually hold more solute than is present in their saturated solutions at a given temperature. If more solute is added all “extra” solute immediately falls out of solution Examples sodium acetate NaC2H3O2 and sodium thiosulfate Na2S2O3

Molecule-ion attraction a Molecule-ion attraction a. the term molecule refers to a polar solvent while ions refers to an ionic compound b. An ionic compound is placed in liquid ammonia, water or alcohol

c. Attraction between the positive and negative ions of the ionic compound and the oppositely charged poles of the solvent molecule. d. This results in a hydrated ion if the solvent is water

Preparation of a supersaturated solution 1 Preparation of a supersaturated solution 1. Add more solute to the solvent( usually water) than will dissolve at room temperature 2. Heat the solution until all the solute dissolves, do not boil 3. Gently place the hot solution in an ice bath and allow it to cool undisturbed. 4. Once the solution is cooled to room temperature a supersaturated solution exists.

Dilute- amount of solute dissolved is small in relation to the amount of solvent. It is considered a weak solution Concentrated – amount of solute dissolved is large in relation to the amount of solvent. It is considered a strong solution.

Table F- Solubility Guidelines Used to predict if various compounds dissolve in water.

Solubility curves 1. Graph of amount of solute in 100g of water versus temperature 2. TABLE G 3. Tells the number of grams that needs to be dissolved in 100g of water at a given temperature to make a SATURATED SOLUTION.

Examples: How many grams of KNO3 needs be dissolved to produce a saturated solution in 100g of water at 50C? How many grams of KNO3 needs be dissolved to produce a saturated solution in 50g of water at 50C? How many grams of KNO3 needs be dissolved to produce a saturated solution in 500g of water at 50C?

5. Can be used to predict if a solution is saturated, unsaturated or supersaturated. a. saturated solutions fall on the curve b. unsaturated solutions fall below the curve c. supersaturated solutions fall above the curve

6. Examples: A solution is prepared by dissolving 80 g of NaNO3 in 100 g of water at 30 C. Is this solution unsaturated, saturated or super saturated? A solution is prepared by dissolving 70g of NH4Cl in 100g of water at 80 C. Is this solution unsaturated, saturated or supersaturated?

Concentration 1. Definition: amount of solute dissolved in a given amount of solvent 2. The higher the value the more concentrated the solution 4 Methods for expressing concentrations Molarity a. number of moles of solute per liter of solution

b. Formula found on Table T Molarity = moles of solute Liter of solution Units: M or moles/liters 2. Percent by mass- a. Used when a solid is dissolved in a liquid % by mass = mass of solute x 100 mass of solution

3. Percent by volume a. Used when two liquids are mixed to form a solution % by volume = volume of solute x 100 volume of solution 4. Parts per million a. Specifies the number of parts of solute in 1 million parts of solution

b. Used when the amount of solute is small compared to the amount of solvent. c. Formula on Table T Ppm = mass of solute x 1,000,000 Mass of solution d. units ppm

Preparing a solution 1. Calculate the number of grams of solute that must be added using the correct formula from the information given 2. Add this much solute to a volumetric flask of the correct volume. 3. Fill the flask about ¾ of the way with solvent (usually water) 4. Cap the flask and invert to mix. (repeat until all solute is dissolve) 5. Fill the flask to the line with solvent. 6. Invert to finish mixing.

Dilutions: M1V1 = M2V2 M1 = molarity of original sample V1 = volume of original sample M2 = molarity after dilution V2 = volume of solution after dilution

Electrolytes and nonelectrolytes a. -compounds that conduct electricity in solution b. contain mobile charged particles(ions) in the solution

2. Strong electrolyte-electrolytes that form large numbers of ions in solution they cause the bulb of a conductivity apparatus to glow brightly. Examples: Ionic Solids, strong acids and bases

3. Weak electrolyte- electrolytes that are poor conductors- they form few ions in solution and cause the bulb of the conductivity apparatus to glow faintly. Examples weak acids and bases such as acetic acid and ammonia

4. Nonelectrolytes- non conductors of electricity they do not produce ions when placed in water and the bulb does not light at all Examples: sugars, glycerine Colligative properties 1. Properties of solutions such as vapor pressure, boiling point and freezing points that depend upon the number of particles of solute dissolved in a solution.

2. Therefore properties such as vapor pressure, boiling point and freezing point will be most affected if the solute is strong electrolyte and least affected if the solute is a nonelectrolyte. BOTH WILL AFFECT THE PROPERTIES JUST TO DIFFERENT EXTENTS

3. Solutes that are strong electrolyte break up into many ions therefore there are more dissolved particles in the solution than for a nonelectrolyte.

Boiling Point Elevations and Freezing point depressions 1. When any substance is added to a solvent the boiling point increases and the freezing point decreases. The increase or decrease is greatest if the substance added is ionic (strong electrolyte) and least if the substance is molecular (weak or nonelectrolyte)

Which of the following compounds when dissolved in water would lower the freezing point the most and why? C6H12O6 NaCl CH3OH MgCl2 AlCl3