SOLUTIONS Homogeneous Mixtures

Solution - Definition A solution is a homogeneous mixture of two or more substances (elements or compounds) in no definite ratio by mass. All homogeneous mixtures are solutions. Solutions can exist in any phase Liquid solutions must be clear (light can pass through) and can be colored.

Solvents and Solutes An aqueous solution contains dissolved substances. 2 components of a solution: solvent and solute. Solvent dissolves the solute; solute is what is being dissolved.

Question Which statement describes KCl(aq)? KCl is the solute in a homogeneous mixture KCl is the solute in a heterogeneous mixture KCl is the solvent in a homogeneous mixture KCl is the solvent in a heterogeneous mixture

Solution process http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

Electrolytes and Nonelectrolytes An electrolyte is a substance that conducts an electrical current when in an aqueous solution or a molten state. All ionic compounds are electrolytes. Strong electrolytes are good conductors of electricity. Weak electrolytes are weak conductors of electricity. Nonelectrolytes do not conduct electricity when in an aqueous solution or in a molten state.

Question Explain why CH4 is a nonelectrolyte.

Terminology Solute - substance that is being dissolved (lesser amount) Solvent - substance that is doing the dissolving (greater amount) Aqueous solutions - water solutions (water is the solvent) KCl(s) -> K+(aq) + Cl-(aq) Note the (aq) to indicate a water solution

Terminology (cont.) Solubility - the amount of solute that will dissolve in a given amount of solvent at a certain temperature. Dissociation - breaking apart into its ions KCl(s)  K+(aq) + Cl-(aq) Miscible - mixable (water & alcohol) Immiscible - will not mix (oil & water)

Solubility Factors The following have an effect on how much of something can be dissolved- Nature of the solute and solvent Temperature Pressure (gases)

“Like Dissolves Like” In chemistry the phrase is used to describe how solutes and solvents will interact with each other. In Both cases the substances are soluble water CCl4 liquid ethanol I2 solid Both polar Both nonpolar

Nature of the Substances Solute Type Non Polar Solvent Polar Solvent Nonpolar Soluble Insoluble Ionic

Temperature For most solids solubility in water increases with and increase in temperature. Ex: dissolving sugar in tea For gases the opposite is true; solubility in a liquid decreases with an increase with temperature. Why? Increasing temp allows the gases to escape the liquid.

Pressure Pressure has little no affect on the solubility of solids or liquids Pressure affects solubility of a gas in a liquid. As pressure increases, the solubility of a gas increases. Ex: Opening up a soda can.

Solubility Factors The following will effect how fast something dissolves- Particle size Stirring

Particle Size The rate at which a solute dissolves depends on the particle size of the solute. The bigger the particle the longer it takes to dissolve it.

Stirring Stirring increases the rate at which a solute dissolves because during stirring the fresh solvent is continuously in contact with the solute.

Question Describe two methods by which you could remove gases that are dissolved in water.

TEMPERATURE Refer to Table G of your Reference Tables for Chemistry The solubility of most solids increases with an increase in temperature The solubility of gases decreases with an increase in temperature

Temperature & Solubility These substances are gases Notice that the solubility of the gases decreases with an increase in temperature The solubility of most solids (those on the table) increases with an increase in temperature The solubility is given in grams of solute per 100 grams of water

Question Which compound’s solubility decreases most rapidly as the temperature changes from 10oC to 70oC? NH4Cl NH3 HCl KCl

Solution Conditions Supersaturated Grams solute/100 g H2O Unsaturated Temperature(°C)

Solution Types Unsaturated - the solution can dissolve more solute in the solvent at the specified temperature.

Solution Types Saturated - the solution is holding as much solute as it can hold at the given temperature ( the rate of solution equals the rate of dissolution).

Solution Types Supersaturated - the solution is holding more solute than it normally can hold at the specified temperature. These solutions are unstable and will seek to reach saturation when disturbed and the excess solute will precipitate out. Example: rock candy http://chemed.chem.purdue.edu/demos/main_pages/15.2.html

Question How many grams of NaNO3 would be needed to saturate 200 grams of H2O at 40oC?

Units of Concentration (Table T) ways to express the amount of solute in solution (concentration) Mass percent of solute Parts per million (ppm) Molarity

Mass Percent of Solute On Table T you will see the following relationship – Mass % solute = (mass solute/total solution mass) x (100) Example – In a solution prepared by dissolving 24 g of NaCl in 152 g of water, what is the % by mass of NaCl in solution? Solution – Mass % NaCl = (24 g/176 g) x 100 = 14

ppm On Table T you will see the following relationships ppm solute = (mass solute/total solution mass) x 106 There is a relationship between mass percent and ppm ppm = mass percent x 104

ppm Example Example – In the United States and Canada, drinking water cannot contain more than 5 x 10-4 mg of mercury per gram of sample of solution. In parts per million what would that be? Solution Ppm Hg = (5 x 10-4 mg Hg/1 x 103 mg) x 106 = 0.5 Remember there are 1000 mg = 1 gram)

Questions An aqueous solution has 0.0070 gram of oxygen dissolved in 1000. grams of water. In the space in your answer booklet, calculate the dissolved oxygen concentration of this solution in parts per million. Your response must include both a correct numerical setup and the calculated result.

Molarity (M) Remember moles (n) = given mass/molar mass Formula from Table G M = mols of solute /Liters of solution Example – if 40 g of NaOH is dissolved in water to prepare 500 mL of solution, what is the molarity? M = 1 mol/0.5 L = 2M (or a 2 molar solution) Molar mass of NaOH = 40 g/ mol

# grams used (convert to moles first) *Moles = mass/ molar mass *Mass = mole X molar mass Molarity Problems Fill-in the following blanks Substance Molar Mass (g/mol) # grams used (convert to moles first) L of solution Molarity (M) KOH 56.1 100 2.0 CaCl2 111 0.5 NH4NO3 80 160 1.0 Ca(OH)2 74.1 1.5

Question How many moles of solute are contained in 200 ml of 1M solution?

Percent by Volume % by volume (v/v) = 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑥 100 Example: If a 10 mL of propanone is diluted with water to a total solution volume 200 mL, what is the percent by volume of propanone in the solution?

Question A bottle of the antiseptic hydrogen peroxide H2O2 is labeled 3.0% (v/v). How many mL H2O2 are in 400.0-,L bottle of this solution?

Molarity Dilution Problems The number of mols of solute in a sample remains constant when we add water to dilute the sample, the concentration changes. M1V1 = M2V2 If a 200 mL of a 3.0 M HCl solution is diluted to 400 mL what is the new molarity?

Question How many milliliters of a solution of 4.00M KI are needed to prepare 250.0 ml of 0.760M KI?

Molecules vs. Ionic Compounds in Aqueous solutions Some molecular compounds dissolve but do not dissociate into ions. C6H12O6(s) (glucose)  C6H12O6 (aq) 1 mole of sugar gives 1 mole of sugar when dissolved in water

Molecules vs. Ionic Compounds in Aqueous solutions Many ionic compounds dissociate into independent ions when dissolved in water NaCl (s)  Na+(aq) + Cl-(aq) 1 mole of NaCl solid gives 1 mole of sodium ions and 1 mole of chloride ions when dissolves (total of 2 moles of dissolved particles). MgCl2 (s)  Mg+2(aq) + 2Cl-(aq) Total of 3 moles of dissolved particles

Colligative Properties Colligative properties refers to properties of a solution that depend on the concentration of particles. Vapor pressure Boiling point Freezing point

Vapor Pressure Lowering The presence of any solute (salt or sugar) lowers the vapor pressure of the solvent. The more moles of dissolved particles, the lower the vapor pressure. The greater the concentration of a solute the more it lowers the vapor pressure.

Boiling Point Elevation The presence of a nonvolatile solute (salt or sugar) raises the boiling point of the solvent. Nonvolatile: does not easily vaporize The greater the concentration of the solute, the more it raises the boiling point. The more moles of dissolved particles, the higher the boiling point. Boiling point elevates by 0.52 oC for each mole of particles in a kg of water.

Boiling Point (cont.) The elevation of the boiling point depends on the number of mols of particles present. For each mol of particles in a kg of water the boiling point of the aqueous solution is elevated by 0.52°C. That is called the boiling point elevation constant for water (0.52°C/m). Molecular compounds such as C6H12O6 do not break up into ions, so one mol of C6H12O6 would raise the boiling point 0.52°C. Ionic Compounds such as NaCl and CaCl2 break up into ions in aqueous solution. One mol of NaCl actually yields two mols of ions, one mol of CaCl2 yields three mols of ions. They would elevate the boiling point more than one mol of C6H12O6.

Freezing Point Depression The presence of any solute (salt or sugar) lowers the freezing point of the solvent. The more moles of dissolved particles, the lower the freezing point. The freezing point of an aqueous solution is lowered 1.86°C for each mol of particles per kg of water. The greater the concentration of a solute the more it lowers the freezing point.

Question How are the boiling and freezing points of a sample of water affected when a slat is dissolve in the water? Boiling point decreases and freezing point depression Boiling point decreases and freezing point increases Boiling point increases and freezing point decreases Boiling point increases and freezing point increases

Question Which solution containing 1 mole of solute dissolved in100 grams of water has the lowest freezing point? KOH (aq) C6H12O6 (aq) C2H5OH (aq) C12H12O11 (aq)

Summary The addition of solutes of water: Lowers vapor pressure Raises boiling point Lowers freezing point The more particles dissolved the greater change it has on the colligative properties.

Other things to remember Low vapor pressure=high boiling point High vapor pressure = low boiling point The stronger the intermolecular forces the lower the vapor pressure/higher the boiling point.