Solutions

Unit 13

• Mixtures

Composed of two or more substances that can be separated by physical means.

Separating a Mixture• Filtration: the process of separating

materials of a mixture by passing the materials through a filter and paper.

• Distillation:process of separating different parts of a mixture (usually all liquids) based on their boiling points.

• Chromatography: process of separating parts of a mixture by using another material that will diffuse – bringing parts with it.

Mixtures can be Heterogeneous or Homogeneous

• Heterogeneous: not uniform in appearance Ex:Soil, Salad, Chex Mix

• Homogeneous: uniform in appearance Ex:Milk, (aq)

Solutions

• homogeneous mixture of 2 or more substances, usually in unequal amounts.

Definitions• 1.      Solvent: (usually liquid) –

substance found in greater amount – does the dissolving.

• 2.      Solute: (often solid) – substance found in smaller amount – gets dissolved.

• 3.      Aqueous Solution: solution in which water is the solvent (aq)4. Concentration: ratio of solute to solvent in a solution (increase the solute, increase the concentration).

Summary

1. Solutions are homogeneous mixtures

• 2.  Solutions are clear and do not disperse light

• 3.  Solutions can have color

• 4.  Solutions will not settle on standing

5. Solutions will pass through a filter

Solubility

• The amount of solute that will dissolve into a solvent under set conditions of temperature and pressure.

• Factors that Affect Solubility

• 1. Nature: Polar substances vs. Nonpolar substances “like dissolves like”

Solute Type Nonpolar Solvent

Polar Solvent

Nonpolar Soluble Insoluble

Polar Insoluble Soluble

Ionic Insoluble Soluble

• 2. Temperature: as temperature increases, most solids become more soluble in water. Gases react opposite – as the temperature rises, the solubility decreases.

• 3. Pressure: *only effects the solubility of gases* As pressure increases, the solubility of gases increases.

• Precipitation: Process in which dissolved particles “fell out” of solution, usually settling on the bottom of the container.

• **Table F and Table G**

Solubility Graphs (Table G) and Tables (Table F)

• 1.      Saturated a solution that contains a maximum amount solute that will dissolve at a specific temperature.

2. Unsaturated a solution that contains less than the maximum amount of solute that will dissolve at a specific temperature.

• 3. Supersaturated a solution that holds more solute than is present in a saturated solution at that temperature. *Very Unstable*

• a. How to recognize saturation: If a solution contains some undissolved material it must be saturated. If you add some solute to the solution and it dissolves, it was unsaturated. If you add some solute to the solution and other crystals form, it was supersaturated.

4. Soluble material that dissolves (has a high solubility)

5. Insoluble material that does not dissolve (has a low solubility)

Methods of Indicating the Concentration of a Solution

• Dilute/Weak Concentrated/Strong

• 1. Molarity: equal to the of moles of solute dissolved in 1L of solution Abbreviation: M

• Formula: Molarity = #moles of solute/liter of solution Unit: moles/liter

• Examples: • How many grams of NaOH are in 1L of a 2M

solution?

How many moles of solute are in 2L of a 3M solution?

What is the molarity of a solution containing 2 moles of solute in 0.5L of a solvent?

How many moles are in 0.100L of a 0.75M solution?

Which of the following contains the least number of moles of solute?

a. 1L of 1M solution c. 1L of 0.5M solutionb. 2L of 1M solution d. 2L of 0.5M solution

Percent by Mass

• is calculated by dividing the mass of the solute by the mass of the solution (total mass of the solvent + solute) and multiplying it by 100%

• Formula: %mass = (mass of solute/mass of solution) x 100%

• Example: A mass of 25g of NaCl is dissolved in 200g of water. What is the concentration of this solution expressed as a percentage by mass?

Percent by Volume

• Usually used when the solvent is alcohol. Similar to percent by mass, but volumes are used.

• Formula: % Volume = (Volume of solute/Volume of solution) x 100%

• Example: What is the percent by volume concentration of a 25mL solution of ethanol dissolved in 75mL of water?

Parts Per Million

• Usually used when only a small amount of solute is dissolved in a significantly larger amount of solvent.

• Formula: ppm = (grams of solute/grams of solution) x 1,000,000

• Example: Find the concentration on ppm if 0.15g of sugar is dissolved in 500g of water.

Preparation of a Solution of Known Concentration

• In order to prepare a solution of known concentration, there is a simple equality for you to use:

• Formula: M1V1 = M2V2

Colligative Properties of Solutions Changes in boiling pt/freezing pt

• Boiling Point: Increase with the addition of dissolved particles to the solution.

• Increases 0.52 degrees celcius per one mole added.

• Example: 1000g of water at standard pressure boils at 100C 1000g of water + 1 mole of solute boils at 100.52C

• Freezing Point: decreases with the addition of dissolved particles to solution.

• Decreases 1.86 degrees celcius per one mole added.

• Example: 1000g of water at standard pressure freezes at 0C 1000g of water + 1 mole of solute freezes at –1.86C

• The most amount of solute in the least amount of water will have the highest boiling point and the lowest freezing point. Note: Electrolytic solutions have a much greater colligative effect than non-electrolytic solutions.

• Electrolyte: solute has to be able to break up into ions in solution. Capable of conducting electricity in aqueous forms

• Nonelectrolyte: do not “ionize” well in water Make very poor conductors of electricity

• Examples:

• Electrolytes Nonelectrolytes• Ionic Solids non ionic solids• NaCl C6H12O6

• KBr NH3

• CaCl2 CH4

• The more ions that a solution has, the more of a change it will cause!!!

• Example: Which 0.1M solution has the lowest freezing point?

a. C6H12O6 b. (NH4)2SO4

c. KBr d. CuSO4