Solution ChemistryUnit 8

General ChemistrySpring ’13

Objectives (Ch. 15) Understand and describe the basic properties of water

and ice and how they effect the world around you. Explain the high surface tension and low vapor pressure

of water in terms of the structure of the water molecule and hydrogen bonding (15.1.1)

Distinguish between solvent and solute (15.2.1) Describe what happens in the solution process (15.2.2) Explain why all ionic compounds are electrolytes (15.2.3) Distinguish between suspension and solution (15.3.2) Identify the distinguishing characteristic of a colloid

(15.3.2)

Simply Water

Ice has a low density. (Does ice float?)

It’s a polar molecule Slightly positive (+) on one end Slightly negative (-) on another

Look what it does to salt! It also easily bonds to itself and

easily pulls compounds apart

Polar vs. Non-polar?

It’s like a tug of war… Even pulling of

electrons means it’s non-polar

Uneven pulling means it’s polar

Water’s Hydrogen Bond

Water molecules are not connected by full covalent bonds, but they’re pretty strong

The formation between the hydrogen atoms on one molecule and a highly electronegative atom on another is called a hydrogen bond.

Atoms that can do this are hydrogen, oxygen, fluorine, and nitrogen

Keep reading…

Hydrogen Bonds cont.

Any molecule with O-H bonds has the potential to form hydrogen bonds.

Alcohols (molecules with O-H bonds) also form hydrogen bonds. Have similar properties to water

Proteins, nucleic acids, and carbohydrates also can form hydrogen bonds. How they form and shape

determines how they’re used biologically

Water revisited

Boils at 100oC Freezes at 0oC

Expands due to hydrogen bonding Solid state is highly organized

One drop = ~2x1021 molecules

Water in the Solid State

The structure of ice is a regular open framework of water molecules arranged like a honeycomb. Add energy and the

framework collapses The molecules are closer

together making water more dense than ice

Implications on aquatic life?

Surface Tension

Surface tension: the inward force, or pull, that tends to minimize the surface area of a liquid Causes drop to pull together All liquids have a surface tension

Mercury has high surface tension

Surfactants Interfere with H-bonding and reduce

surface tension Soaps and detergents

Surface Tension at work

Capillarity

Results from the competition between the attractive forces between the molecules of the liquid and the attractive forces between the liquid and the tube that contains it

Vapor Pressure of Water

Results from the molecules escaping from the surface of the liquid and entering the vapor phase.

In water, H-bonds hold on to water molecules tightly Tendency for molecules to escape is low Evaporation of water is low What would happen if it was fast?!

Specific Heat of water

Specific heat: measures the amount of heat, in Joules, needed to raise the temperature of 1g of substance by 1oC.

For water it’s 4.18, pretty high. Why it takes so long to boil water It takes a long time to absorb or release more heat

for its temperature to change 1oC than a lot of other substances.

Think of a pool in the summer time.

Specific Heat

Substance Specificheat

Water 4.18

Chloroform 0.96

Aluminum 0.90

Mercury 0.14

Water: The Universal Solvent

Almost always found in solution A very good solvent due to its polar abilities Examples of aqueous solutions

Milk Soda pop Coffee and tea Tap water Look at the ingredient list of a liquidy beverage.

Water is probably there.

Solutions

When one substance dissolves into another, that is called a SOLUTION Example: sugar water, Kool-Aid

There are two main parts of a solution: SOLUTE= the dissolved material

Example: sugar, salt, oxygen (air) SOLVENT= the substance that is doing the

dissolving (usually a liquid) Usually present in the highest amount Example: Water, nitrogen (air)

Dissolving

Water, a polar molecule, is capable of dissolving a range of compounds

Many ionic compounds (like NaCl) are soluble in water

When dissolved, ionic solutions are very good conductors of electricity

Electrolytes vs. Nonelectrolytes

Electrolyte A compound that conducts an electric

current when it is in an aqueous solution All ionic compounds are electrolytes

because they dissociate into ions Nonelectrolyte

A compound that does not conduct an electric current in aqueous solution

Many covalents are this because they are not composed of ions

The electric pickle

How does dissolving happen?

Ionic solids are composed of positive and negative ions.

Water has a positive and negative end (it’s polar) Opposites attract and the ionic compounds

separate into ions. The process by which charged particles in an

ionic solid separate from one another is solvation

Dissolving Covalent substances

Sugar is the best example Almost 200 grams can dissolve in 100 mL of H2O! It has O-H bonds, which makes it polar, so it’s easily

dissolvable in water If a molecule contains O-H bonds, it will tend to be

polar and it can form hydrogen bonds with water.

Dissolving Covalent substances

Covalent molecules are simply separated from one another by water molecules.

They don’t solvate into separate ions Covalent solutions can’t conduct electricity

“Like Dissolves Like”

This means that dissolving occurs when similarities exist between the solvent and the solute.

Sugar is a polar molecule, so is water, and water tends to dissolve substances that are polar or that form hydrogen bonds.

Oil and water don’t mix. Oil is nonpolar But different oils are “like” enough to mix and stay

mixed. Oil and gasoline.

Suspensions

A (heterogeneous) mixture from which particles settle out upon standing

A suspension differs from a solution because the particles are much larger and do not stay suspended indefinitely Cornstarch mixed with water thickens

sauces

Suspensions

Two substances are clearly identified Dispersed phase

Ex) clay, dirt, sand, flour, CO2

Dispersion medium Water, ethanol, Milk or Cream

Think of a glass of water with sand or mud in it. Typically easy to separate

Colloids

A heterogeneous mixture containing particles that range in size from 1nm to 1000nm

Particles spread throughout the dispersion medium (s, l, g) Glues Gelatin Paint Aerosol sprays smoke

Colloid Examples (Table 15.3)

Tyndall Effect, Coagulation

Tyndall Effect The scattering of visible light by colloidal particles Suspensions also do this but solutions don’t.

Coagulation The clumping of particles in a colloid

Emulsions

A colloidal dispersion of a liquid in a liquid

Must have an “emulsifying agent” To form the emulsion To maintain stability Ex) soap, detergent, egg yolk

Objectives (Ch. 16)

Identify the factors that determine the rate at which a solute dissolves (16.1.1)

Identify the units usually used to express the solubility of a solute (16.1.2)

Identify the factors that determine the mass of solute that will dissolve in a given mass of solute (16.1.3)

Solution Formation

Determining factors Composition of solvent Composition of the solute

Speed of dissolving factors Stirring (agitation) Temperature Surface area

All involve contact between solvent and solute

Rate of Dissolving

Stirring the Solution Increases the interaction

between water molecules and the solute.

Solute and solvent interact more often, the rate of dissolution is faster.

Does not influence the amount of solute that will dissolve

Oil will never mix with water not matter how long you stir or shake that Italian dressing

Rate of Dissolving

Heating the Solution Increases kinetic energy

of the water molecules Increases frequency and

force of the collisions between solute and solvent

Rate of Dissolving

Grinding the solute Creates more surface area

(remember the big fireball demo?)

Solvent molecules attack the edged surfaces of solute crystals.

The more surface area exposed, the faster the rate of dissolving

Solubility

When solute enters the solvent… Particles move from the solid into the

solution Other dissolved particles move from

solution back to the solid Occurs at the same rate Called a saturated solution

Will stay this way as long as temperature stays the same

Solubility

The amount of solute that dissolves in a given amount of solvent at a specified temperature and pressure to produce a saturated solution

Units: grams per liter (g/L) Miscible

two liquids that dissolve in each other Immiscible

Two liquids not soluble in each other

Types of Solutions

Saturated solution Solution holding the max. amt. Of solute

per amt. Of solution under given conditions.

Add more solute it won’t dissolve Unsaturated solution

The amt. of solute is less than the max that could be dissolved.

Add more solute it will dissolve

Solutions

Supersaturated solution Contain more solute than the usual max.

amt. And are unstable. Add a crystal and it fills the container with

crystals

Effects of Pressure Huge effect on gases, very little on solids and

liquids Gas solubility increases as the partial pressure

of the gas above the solution increases. (direct relationship)

Ex) Soda bottle has lots of dissolved CO2 in it which is forced in at the plant.

When you open the bottle you hear a hiss and CO2 starts escaping from the bottle decreasing the concentration of CO2 in the bottle

Solubility Curves

The solubility of substances changes with temperature For example, is it easier to dissolve sugar

in hot or cold coffee? Solids become more soluble at higher

temperatures Gases become less soluble at higher

temperatures

Solubility Curves (cont.)

Scientist have studied many substances solubility at different temperatures They created

graphs which show this data

Solubility Curves (cont.) Let’s simplify the graph with all the

substances down to just one substanceSolubility of KCl in 100 g of water

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Temperature (degrees Celcius)

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Solubility Curves (cont.)

Is KCl a solid or

gas in this graph?

Solubility of KCl in 100 g of water

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Solubility Curves (cont.) How many grams of KCl will dissolve in

500g of water at 80°C? 260g of KCl (52g x 5 = 260g)

Solubility of KCl in 100 g of water

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Solubility Curves (cont.) How many grams of water will it take to dissolve

26 g of KCl at 80°C? 50g of H2O (1/2 of what dissolves in 100g H2O)

(% of 100g: 26g/52g=.50)

Solubility of KCl in 100 g of water

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Solubility Curves (cont.) If one dissolves 95 grams of KCl in 250

grams of water at 80°C, what kind of solution will they have? Unsaturated

Solubility of KCl in 100 g of water

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You need to determine the saturation point before you can decide the type of solution.

(Sat. pt. from graph)x(%H2O)

52 g x 2.5 = 130 g

Saturation point in 250g of water

Practice Problem #1

How many grams of NH4Cl will dissolve in 300 grams of water at 70°C?

If one dissolves 137.5 grams of NaNO3 in 125 grams of water at 45°C, what kind of solution will they have?

186g NH4Cl (62g x 3)

Saturated110g x 1.25= 137.5 The same as what it asks

Objectives

Solve problems involving the Molarity of a solution (16.2.1)

Describe the effect of dilution on the total moles of solute in solution (16.2.2)

Define percent by volume solutions (16.2.3)

Solution Concentration

What does concentration mean? It tells us how much solute is dissolved in a

given volume of solution “dilute solution” has a small amount of

solute “concentrated solution” has a large amount

of solute Both are relative and are very imprecise Qualitative… not quantitative

Molarity

You can describe the precise concentration quantitatively with Molarity.

Molarity (symbol M). Relates the amount of solute to a given volume of

solution. The number of moles of solute dissolved in one

liter of solution. Ex) a solution labeled 3M NaCl is read “three molar

sodium chloride solution”

Molarity Expressed in this manner:

How do you convert from grams to moles?Divide by molar mass foo!

Hint: what is the equation for density?

mass

Density volume

Molarity

Drain cleaner is made with caustic sodium hydroxide, NaOH. The Dow company prepares a bottle of drain cleaner from 24.0 g of NaOH dissolved in 0.100 L of solution. What is the molarity?

Molar mass of NaOH (40.00 g/mol)

24.0g NaOH 1 mol NaOH

mol NaOH

= L solution = 6.00M NaOH0.100 L solution 40.00g NaOH

We want Liters, leave this alone Units for molarity

Molarity- Writing Unit Factors

We can solve molarity calculations by using the solution concentration as a unit factor.

Example: 6.00M solution of NaOH contains 6.00 mol of solute in each liter of solution. Written as:

6.00 mol NaOH or 1 L solution

1 L solution 6.00 mol NaOH

There’s 1000mL in 1L

Preparing Solutions

How would you prepare 1.0L of a 0.15M sodium chloride solution? I.O.W. How many grams NaCl are needed?

Think: First, determine the mass of NaCl to add to a 1.0 L container. The 0.15M solution must contain 0.15 moles of NaCl per liter of solution (definition of molarity).

1. Use molarity to convert to moles.

2. Then use molar mass to go from moles to grams.

1.0 L solution 0.15 mol NaCl 58.5 g NaCl = 8.8 g NaCl 1 1 L solution 1 mol NaCl

Making that solution

1. Obtain a volumetric flask

2. Measure 8.8 g of NaCl

3. Add solute to a small amount of water, about 300 mL, to dissolve

4. Add enough additional water to bring the total volume to 1.0 L, to the etched line on flask

Preparing a Different Volume of Solution

How would you prepare 5.0 L of a 1.5M solution of glucose, C6H12O6

Think: You need to determine the number of grams of glucose to add to a 5-L container. The 1.5M solution has 1.5 mol of glucose (use this to convert to grams.

5.0 L solution 1.5 mol glucose 180 g glucose

= 1400 g glucose 1 1 L solution 1 mol glucose

MOLARITY MOLAR MASS

Making Dilutions

Stock solutions of acids are very concentrated HNO3 comes in 15.8M

H2SO4 comes in 18.0M HCl comes in 12M

This does not tell how “nasty” these are…that’s a different unit

Using an acid in this concentration is incredibly dangerous

I dilute them for lab purposes… how?

Acid burn(why you don’t wear flip flops in the lab)

Dilutions

Diluting a solution reduces the # of moles of solute per unit volume The total number of moles of solute in solution does

not change mol solute before dilution = mol solute after dilution

Equation for dilutions: M1 x V1 = M2 x V2

M1 and V1 are initial readings M2 x V2 are for the diluted solution

Units of volume must match (mL or L)

Dilution example How many milliliters of 2.00M MgSO4 must be diluted with

water to prepare 100.0mL of 0.400M MgSO4? M1= 2.00M M2=0.400M

V1= ? V2= 100mL M1 x V1 = M2 x V2

Dilution answer

Substitute and solve for V1

20 mL Take 20mL of initial solution and dilute with

enough water to increase volume to 100 mL

DO NOT ADD 100mL of water, this will give 120mL of solution, not 100mL

Dilution Example #2

If you take 10mL of a 3.42M solution and dilute to 100mL, what is your new concentration? M1 x V1 = M2 x V2

(10mL)(3.42) = (M2)(100mL)

M2 = 0.342M

Dilution Example #3

You need a 2.10M solution. You need 1500mL of it. How much of a 12M solution do you need to use? M1 x V1 = M2 x V2

(12M)(V1) = (2.10M)(1500mL)

V1 = 262.5mL

Percent Solutions

How many mL of isopropyl alcohol are in 100mL of a 70% solution? 70 mL are mixed with enough water to

make 100 mL

% by Volume = volume of solute x 100% volume of solution