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“Water, water everywhere

The very boards did shrink.

Water, water everywhere

Nor, any drop to drink.”Rhyme of the Ancient Mariner

Samuel Coleridge

Heterogeneous matter:Mixtures

Review of suspensions and colloids with concentration on

solutions.

Introduction to Mixtures

A mixture is a combination of different substances where each substance retains its chemical properties.

Generally they can be separated by non-chemical means such as filtration, heating, or centrifugation.

3 Types of Mixtures

Heterogeneous Suspensions

Colloid Homogenous

solutions

Suspensions

A heterogeneous mixture in which some of the particles will settle (separate) upon standing. may be filtered

Example: oil/water clay/water

ColloidsA mixture containing particles that are intermediate in size between those of suspensions and true solutions.Example: hair spray, cold cream, milk

Are usually cloudy or milky in appearance, but look clear when very dilute.Particles CANNOT be separated by filtering and do not settle.

Solutions The same throughout. Best mixed. Examples

Kool-aide®

iced tea Air metal alloys

gold jewelry brass

Solutions

Introduction to solutions andcalculations for solutions

Introduction to solutions

A solution is a homogeneous mixture where all particles exist as individual molecules or ions.

Particles are so small (less than 1nm) that the mixture never settles out.

Will not scatter light. Does not display the Tyndall effect.

Tyndall effect The Tyndall Effect is caused by reflection

of light by very small particles in suspension or colloid. It is often seen from the dust in the air when sunlight comes in through a window, or comes down through holes in clouds. It is seen when headlight beams are visible on foggy nights, and in most X-File episodes when Moulder and Sculley check out some dark place with flashlights.

Tyndall effect

Components of a solution

Solvent the dissolving

medium greater in

amount typically a liquid usually water

aqueous

sometimes alcohol

tincture

Solute Substance

dissolved lesser in amount frequently a

solid

Solution Formation

If it will dissolve depends on the solute and solvent.

Water is called the universal solvent because it dissolves most substances.

Liquid solvent and solute

Miscible – liquids that dissolve freely in one another in any proportion are completely miscible.

Immiscible – liquids solutes and solvents that are not soluble in each other.

Like dissolves like

A polar solvent (water) dissolves most ionic solutes.

A polar solvent dissolves polar solutes.

A polar solvent does NOT dissolve nonpolar solutes.

Nonpolar solvents (oil) dissolve nonpolar solutes.

Salt(ionic)

Water(polar)

NaCl

Solubility practice problem:

Determine whether the following substances will dissolve in water (polar), carbon tetrachloride (nonpolar), and/or alcohol. (see worksheet 1)

MgCl2, F2, methanol (note: -ol ending indicates and alcohol)

Solubility practice problem ans

SolutesWater CCl4 alcohol

MgCl2

F2

MethanolCH3OH

X

X

X

X

X X

Solubility practice problems

At this time complete the worksheet solubility (polar vs. nonpolar)

Electrolytic Solutions Electrolytic solutions

are composed of substance that dissolves in water to give a solution that conducts electric current.

Ionic compounds dissolved in water.

Nonelectrolytic solutions substances dissolved in

water to give a solution that does NOT conduct electricity

solutions containing neutral solute molecules do not conduct electric current because there are no mobile charged particles available

covalent compounds dissolved in water

Electrolyte

At this time stop and complete the electrolyte worksheet.

You are familiar with electrolytic solutions. In the Ionic or Molecular lab you tested conductivity of solutions .

(You dipped the wires in solutions and the light lit up if the solution was ionic.)

What ionic solids dissolve?

To determine which ionic salts dissolve in water you must consult a Table of Solubility Rules.

Luckily for you one is included in you EOC reference tables. Consult them now (page 6).

NCDPI Reference tables for Chemistry (adopted 2000)SOLUBILITY RULES:

SOLUBLE: All Nitrates, Acetates, Ammonium and Group I salts All Chlorides, Bromides, and Iodides, except Silver, Lead, and Mercury(I) All Fluorides except Group II, Lead(II), and Iron(III) All Sulfates except Calcium, Strontium, Barium, Mercury, Lead(II), and

SilverINSOLUBLE: All Carbonates and Phosphates except Group I and Ammonium All Hydroxides except Group I, Strontium, and Barium All Sulfides except Group I, II, and Ammonium All Oxides except Group I INSOLUBLE means a precipitate forms when equal volumes of 0.10 M

solutions or greater are mixed

Solubility example problem #1

dentify the precipitate(s) and aqueous component(s) that form when aqueous solutions of zinc nitrate and ammonium sulfide are combined.

Solubility example #1: Plan

Write the possible double-replacement reaction between reactants and use the solubility rules to determine which products are soluble and which will precipitate (not soluble).

Solubility example #1: solution

U: state of products K: Zn(NO3)2 and (NH4)2S (I checked

charges)

P: Write equation and consult rules.

S: Zn(NO3)2 + (NH4)2S 2NH4NO3 + ZnS Consult chart

Solubility example #1: solution

S: Zn(NO3)2 + (NH4)2S 2NH4NO3(aq) + ZnS(s) Chart reveals that ammonium nitrate is soluble. Zinc

sulfide is not soluble and is therefore a precipitate according to the table.

All Nitrates, Acetates, Ammonium and Group I salts are soluble.

Solubility example #1: solution

S: Zn(NO3)2 + (NH4)2S 2NH4NO3(aq) + ZnS(s) Chart reveals that ammonium nitrate is soluble. Zinc

sulfide is not soluble and is therefore a precipitate according to the table.

All Nitrates, Acetates, Ammonium and Group I salts are soluble.

All Sulfides except Group I, II, and Ammonium are insoluble.

Note: You do not have to do a ukps for each of these problems. You do need to justify your answers.

Solubility rules practice

At this time stop and work the solubility rules practice problems.

You must complete the bold portion of worksheet 5 (#1, 2, 4, 8)

Check your answers with your lab partners.

Solubility

If you determine that the salt will dissolve, how much of it will dissolve?

Solubility

The amount of a substance that dissolves in a given quantity of a solvent at a given temperature to produce a saturated solution is its solubility.

Unit = grams solute / 100 g solvent

Once you determine that a substance will dissolve consult the solubility curve to determine how much of it will dissolve at a given temperature.

For example after adding 37 grams of NaCl to 100 ml of tap water no more will dissolve. Any additional salt settles to the bottom.

NCDPI Reference tables for Chemistry (adopted 2000) Solubility Curve

This is no longer on the reference tables. It will be included with a question as needed.

Saturated

Saturated solution – contains the maximum amount of solute for a given amount of solvent at constant temperature.

This is an equilibrium; the same amount changes into a solid as is dissolving in the liquid.

Same amount of solid solidifying as dissolving.

Saturation Unsaturated – a solution that contains

less solute than a saturated solution. (Less than it could hold.)

Supersaturated – a solution which contains more solute than it can theoretically hold at a given temperature. Additional solute will cause all the “extra”

solute to crystallize and precipitate. this is how rock candy is made

Practice problem using Solubility Curve

When 100 grams of water saturated with KNO3 at 70oC are cooled to 25oC, what is the total number of grams of KNO3 that will precipitate?

Practice problem using Solubility Curve

At 70oC 135g of KNO3 would be dissolved in a saturated solution. At 25oC only 40g of KNO3 could be dissolved.

Therefore 95grams of KNO3 will precipitate.

Henry’s law

The solubility of a gas in a liquid is directly proportional to the partial pressure of that has on the surface of the liquid.

Pressure has little effect on the solubility of liquid or solid solutes.

Rate of Solution

How fast a solute will dissolve depends on: Size of solute particles

Since dissolving occurs on the surface, increasing surface area (power) allows more solvent to reach more solute and speeds dissolving.

Agitation/stirring Moves solute away from solid so solvent can reach

more solute. Temperature

For solids a warmer solvent moves faster and farther apart thus increases rate of dissolving.

However, for gas solutes, as temperature of solvent increases the solubility of a gas decreases.

Solubility Curve problems

At this time stop and complete the solubility curve problems.

Concentration Refers to how much solute is dissolved Dilute means only a little solute is

dissolved Concentrated means a lot is dissolved

These are NOT quantitative values just words you should be familiar with.

Concentration Calculations

There are two major concentration words that are quantitative in nature: Molarity and Molality

Molarity

As is clear from its name, molarity involves moles.

It is calculated by taking the moles of solute and dividing by the liters of solution.

Molarity = moles of solute liters of solution

Molarity Example #1

Suppose we had 1.00 mole of sucrose (about 342.3 grams) and proceeded to mix it into some water. It would dissolve and make sugar water. We keep adding water, dissolving and stirring until all the solid was gone. We then make sure that when everything is well-mixed there is exactly 1.00 liter of solution.

What would by the molarity of this solution?

Example #1 answer

U: molarity K: 1.00 mol sucrose; 1.00 L solution p:

s: M= 1.00 mol 1.00 liter solution

=1.00 mol/L = 1.00 M

Molarity = moles of solute liters of solution

Molarity Example #1 cont’d

Note, some textbooks make the M using italics and some put in a dash, like this: 1.00-M

When you say it out loud, say “one point oh oh molar”

Never forget to replace M with mol/L when you do calculations. “M” is just shorthand.

Molarity Example #2 What is the molarity when

0.750 mol is dissolved in 2.50 L of solution?

U: molarity K: 0.750 mol; 2.50 L solution p: Molarity = moles of solute

liters of solution s: M= 0.750 mol

2.50 L solution =0.300 mol/L = 0.300 M

Molarity Example #3 (using grams which is much more common)

Suppose you had 58.44 g of NaCl and you dissolve it in exactly 2.00L of solution. What would be the molarity of the solution?

p: convert grams to moles Molarity = moles of solute

liters of solution

s: 58.44 g NaCl 1 mol NaCl =1.00 mol NaCl 58.44 g NaCl

M= 1.00 mol NaCl2.00 L solution =0.500 mol/L = 0.500 M

Example #4 As a class calculate the molarity of 25.0g

of KBr dissolved to 750.0 mL solution.

Work the above problem.

Answer: 0.280M

Example #5

80.0 grams of glucose (C6H12O6) is dissolved in 1.00L of solution. What is the molarity?

Answer: 0.444M

Molarity

Notice how the phrase “of solution” keeps showing up. The molarity definition is based on the volume of solution, NOT the volume of pure water used. For example, to say: “A one molar solution is prepared by adding one mole of solute to one liter of water” is totally INCORRECT.

Problems

At this time stop and complete the Molarity Worksheet.

Dilution

Definition and Calculations

Dilution

To dilute a solution means to add more solvent without the addition of more solute.

Since the amount of solute stays the same: moles before = moles after

Dilution equation From the definition of molarity we get that

moles of solute equals molarity times the volume.

If we are just diluting (adding water) the moles of solute remains the same.

Therefore: M1 V1 = M2 V2

Dilution example #1

53.4 mL of a 1.50M soln of NaCl is on hand, but you need some 0.800M solution. How many mL of 0.800M can you make?

S: (1.50mol/L)(53.4mL)= (0.800mol/L)(V2) V2 = 100. mL

Dilution Practice problems

Complete the dilution practice problems

Boiling point elevation Freezing point depression

When a solute is added to a solvent the freezing point of the solvent decreases and the boiling point increases.

For example, “salt” is added to the ice in an ice cream freezer so that the temperature of the ice (freezing point) will decrease and the milk can freeze (become solid)

Boiling point elevationFreezing point depression

Likewise salt (any solute) added to a liquid being heated raises the boiling point.

If salt is added to water then the water is heated, instead of the water boiling at 100oC it would boil at a higher temperature, maybe 104oC.