Solution chemistry unit 8
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Solution Chemistry Unit 8. General Chemistry Spring ’09 Mr. Hoffman. Objectives (Ch. 15). Understand and describe the basic properties of water and ice and how they effect the world around you.

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Solution Chemistry Unit 8

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Solution chemistry unit 8

Solution ChemistryUnit 8

General ChemistrySpring ’09

Mr. Hoffman

Objectives ch 15

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

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

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’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

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

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

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

  • 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

Surface Tension at work



  • 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

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 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

Specific Heat

Water the universal solvent

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.



  • 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)



  • 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

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

How does dissolving happen

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

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 substances1

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

“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.



  • 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



  • Two substances are clearly identified

    • Dispersed phase

      • Ex) clay, dirt, sand, flour

    • Dispersion medium

      • Water, ethanol

  • Think of a glass of water with sand or mud in it.

  • Typically easy to separate



  • 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

Colloid Examples (Table 15.3)

Tyndall effect coagulation

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



  • 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

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

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

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 dissolving1

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 dissolving2

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 expose, the faster the rate of dissolving



  • 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



  • The solubility of a substance is 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

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



  • 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

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 on CO2 in the bottle

Solubility curves

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

Solubility Curves (cont.)

  • Scientist have studied many substances solubility at different temperatures

    • They created graphs which show this data

Solubility curves cont1

Solubility Curves (cont.)

  • Let’s simplify the graph with all the substances down to just one substance

Solubility curves cont2

Solubility Curves (cont.)

  • What does this graph tell you about KCl at 80°C?

    • 52g of KCl dissolve in 100g of water

Is KCl a solid or gas in this graph?

Solubility curves cont3

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 curves cont4

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 curves cont5

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

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

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)


110g x 1.25= 137.5

The same as what it asks



  • 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

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



  • 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”



Hint: what is the equation for density?


Density volume

  • Expressed in this manner:

  • How do you convert from grams to moles?

Divide by molar mass foo!



  • 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)

We want Liters, leave this alone

Molarity writing unit factors

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!! Remember this foo!

Preparing solutions

Preparing Solutions

  • How would you prepare 1.0L of a 0.15M sodium chloride solution?

  • 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).

  • Use molarity to convert to moles.

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

Making that solution

Making that solution

  • Obtain a volumetric flask

  • Measure 8.8 g of NaCl

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

  • 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

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.



Making dilutions

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?



  • 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

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

  • 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

Percent solutions

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

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