3.19.12

  What is the difference between heterogeneous and homogenous mixtures?

  TYGAGT

  Distinguish between types of mixtures

  Compare properties of suspensions, colloids, and solutions

  Distinguish between electrolytes and non-electrolytes

Lots of terms

  Soluble – adj. capable of being dissolved

  Solution – a homogenous mixture of two or more substances in a single phase – does it have to be a liquid?

Components of solutions

  Solute – the substance dissolved in a solution

  Solvent – the dissolving medium

  Precipitate – undissolved/insoluble solid

  If you are not part of the solution you are part of the precipitate

Examples of solutions

Solute state Solvent state Example

Gas Gas Oxygen in nitrogen (our atmosphere)

Gas Liquid Carbon dioxide in water

Liquid Liquid Alcohol in water

Liquid Solid Mercury in tin (or silver) used in dental amalgam

Solid Liquid Sugar in water

Solid Solid Copper in nickel, brass, steel

Other mixtures

  Suspensions – if particles are so large that they settle out (precipitate) due to density unless constantly stirred or agitated. Example, sand in water

  Colloids – particles that are intermediate in size between those in solutions and suspensions

  Tyndall Effect – in colloids, the particles are too small to be seen, but large enough to scatter light. Can be used to distinguish between a solution and a colloid

Solutes

  Water has a high dielectric constant – is a poor conductor of electrical current

  Conductivity – ability to allow an electrical current (or ions) to pass unimpeded

  Water’s conductivity is similar to that of silicon’s – but only pure (de-ionized) water

  Adding certain kinds of solutes changes the conductivity

Particle mobility

  Dielectric constant – permitivity of free space pertaining to charged particles

  High constant = solvents charge to permitivity ratio very high

  Water is a very polar molecule that packs in very close (+ other reasons we’ll see later)

  Low constant = uncharged/non-polar solvent meaning low intermolecular interaction, easy for charged particles to pass between solvent molecules without interacting with them

electrolytes

  In general, a substance that when dissolved in water increases its conductivity (and consequently decreases the dielectric constant)

  In general, any salt (metal – nonmetal, ionic substance) will behave as an electrolyte

  Strongly polar substances will also behave as an electrolyte

  So if you are drinking “electrolytes”, what are you drinking?

Non-electrolytes

  A substance that dissolves in water and does not increase the conductivity (or even increases the dielectric constant)

  Non-polar covalents and even some

Rest of class

  Think about your lab reports, plan/outline your introduction

  Tutorial PP for citations is on my website

  I would like a draft of your introduction by Friday

  You will need to share data within the class so that you are basing your calculations on an average rather than a single trial

Criteria of introduction

  1st paragraph(s) - explain what a double replacement reaction is; how do you know if product is soluble or not? Is this impotant?

  Next paragraphs - summarize the reaction you performed. Here you give a little background on each of the reactants used (properties, toxicity, other info such as molecular weight, melting point, etc.).

  Last paragraphs - predict the yield of your reaction as we did in the last unit. This essentially is your hypothesis. We will be comparing your ACTUAL yield to your PREDICTED yield and determining if, within a given statistical range, we can accept our data.

3.20.12

  First day of spring!

  I like to add about 2 tsp of sugar to my coffee in the morning. What if I wanted to get 4 tsp of sugar dissolved in my coffee? What would I have to do?

  HW – problems on p 416

  Today – explain the factors that affect solubility of solutes in various kinds of solvents

Surface area

  Which is going to dissolve faster: finely ground sugar or whole sugar cubes?

  The dissolution process occurs at the surface of the solute, so increasing the surface area of the solute works

  Which has larger surface area relative to volume: small particles or large particles?

Agitation

  If you add sugar to coffee and let it just sit in there, will it dissolve?

  Stirring/shaking disperses the solute allowing interactions between greater numbers of solute particles and the solvent

Heat of solvent

  If you add sugar to iced tea and the same amount of sugar to hot tea, in which tea is the sugar going to dissolve faster?

  Heat is kinetic energy – dissolution requires collisions between solvent and solute, and at cold temperatures the forces of impact may be insufficient for the solute to dissolve

Solubility

  Despite heating, agitating, increasing surface area, eventually a solution will not be able to contain any more dissolved solute

  Solution equilibrium is the physical state in which the opposing process of dissolution and crystallization occur at equal rates

Saturation

  A solution that contains the maximum amount of dissolved solute is referred to as being saturated

  Example: at 20°C, 46.4 g of NaCH3COO is the maximum amount that will dissolve in 100.0 g of water.

  Alternatively, a solution that may be able to dissolve additional solute is unsaturated

Incredible!

  In general, if saturation is attained by heating, allowing a solution to cool will result in some crystallization of excess solute

  Crystallization is NOT precipitation!

  In certain solutions, the excess solute will not crystallize out, so we say the solution is supersaturated

  Adding a small amount of additional solute to a supersaturated solution will result in rapid crystallization of excess

Solubility values

  Solubility is the amount of substance required to form a saturated solution with a specific amount of solvent at a specified temperature

  Because solubility varies with temperature (kinetic energy)

Solute – solvent interaction

  “like dissolves like”

  Ionic/polar substances are soluble only in polar solvents

  Non-polar substances are only soluble in non-polar solvents

  Why?

Ionic solute in aqueous solution

Solubility in water

  If water can ionize a substance or form hydrogen bonds to it, the substance will dissolve in water

  From structure, which of the following would be soluble in water?

  CH4, CH3CH2OH, NH3, NH4+, C6H6

Miscibility

  Deals with liquid solutes and solvents

  Two liquids that are soluble in one another are referred to as being miscible, such as water and ethanol (CH3CH2OH)

  Immiscible means two liquids will not dissolve in one another; example oil and water

Effects of pressure on solubility

  In general, gases are not soluble in water (liquid solvents)

  Why? Why are they gases in the first place? (think kinetic energy)

  Also – most molecules that are gases at room temperature do not have molecular structures that lend themselves to Hydrogen bonding or ionization

Equilibrium of gases and solutions

  In general, this equation is true for gases

gas + solvent solution

  Increasing the pressure on the gas drives the equilibrium towards the solution side

  What will releasing pressure do?

  This brings us to…

Henry’s Law

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

  This law assumes constant temperature

  This is the law of carbonated beverages and effervescence

  That is not an official title!

Temperature and solubility

  What does increasing temperature do to gas solubility? Make it more or less soluble?

  Increasing temperature (in general) allows gas molecules to escape the solvent molecules and return to gas phase (after all, what does temperature do to volume of gases?)

Enthalpy

  Formation of a solution is accompanied by an energy change

  The net amount of energy absorbed as heat by dissolution of a solute in a solvent is the enthalpy of solution

  This may be positive or negative + values indicate an absorption of energy - values indicate a release of energy

  Which kind will feel hot/cold to the touch?

questions

  If adding heat decreases the solubility of gases, so will dissolving a gas involve the a positive or negative enthalpy change?

  What kind of intermolecular force exists between molecules of a gas (assume ideal gas)?

  Increasing temperature of water has little effect on solubility of NaCl. Why might this be the case?

3.21.12

  What does concentration mean?

  HW – p 424 1-3

  Objectives –

  Calculate molarity, molality of a solution

  Calculate changes to molarity due to dilution

concentration

  The measure of the amount of solute in a given amount of solvent or solution

  Two ways to describe concentration: molarity and molality

molarity

  Molarity (M): the moles of solute divided per liter of solution

  Example: 1.0 moles of NaOH dissolved into 1.0 L of water would be a “1.0 M solution of NaOH”

  Remember: moles = mass of solute/molar mass of solute

Example problems

  You have 3.50 L of a solution that contains 90.0 g of NaCl. What is the molarity of the solution?

  Strategy:

  you have 90.0 g – convert to moles

  M is moles per liter: divide your #moles by #Liters

  Answer: 0.440 M NaCl

Example 2

  You have 0.8 L of a 0.5 M HCl solution. How many moles of HCl does this solution contain?

  Solution:

0.5 mol HCl X 0.8 L = 0.4 mol HCl

1.0 L solution

Example

  To produce 40.0 g of silver chromate, you need 23.4 g of potassium chromate in solution as a reactant. You have 5.0 L of a 6.0 M potassium chromate. What volume of the solution is needed to give you 23.4 g potassium chromate?

  Strategy – what information is needed to answer this?

  How many moles of K2CrO4 is 23.4 g?

  What fraction of your solution will give you this many moles?

  Cont…

solution

  1 mol K2CrO4 = 194.2 g

23.4 g K2CrO4 X 1 mol K2CrO4 = 0.120 mol

194.2 g K2CrO4

6.0 M K2CrO4 = 0.120 mol K2CrO4

x L K2CrO4 solution

x = 0.020 L of the solution

example

  What volume of 3.00 M NaCl is needed for a reaction that requires 146.3 g of NaCl?

Molality

  Molality (m) is the concentration of a solution expressed in moles of solute per kilogram of solvent

  Example: 0.500 mol NaOH in 1 kg water would be a 0.500 m NaOH

molalities

  Less commonly used than molarity

  Concentrations expressed as molalities when studying properties of solutions related to vapor pressure and temperature changes

  Why does this make sense? What happens to volume when we consider pressure and temperature changes? Does mass change?

example

  A solution was prepared by dissolved 17.1 g of sucrose (table sugar, C12H22O11) in 125 g of water. What is the molal concentration of this solution?

  Strategy – moles of solute/kg solvent

example

  A solution of iodine (I2) in carbon tetrachloride is used when Iodine is needed in certain chemical tests. How much iodine must be added to prepare a 0.480 m solution of iodine in carbon tetrachloride if 100.0 g of CCl4 is used?

  What do you know? What are you solving for?

example

  What is the molality of acetone in a solution composed of 255 g of acetone CH3COCH3 dissolved in 200. g of water?

  Remember: moles solute/kg solvent

dilutions

  Dilution is diminishing the molarity of a solution by adding more solvent but keeping solute concentration constant

  Example – if you have 1.0 L a 1.0 M solution of NaCl and add an additional liter of water, what is the molarity now?

Dilution strategy

  Recall: mol = M X volume (V)

  What do the abbreviations mean?

  What if we change volume?

  M1V1 = M2V2

  This formula works whenever “solution 2” is made by diluting “solution 1”

example

  What is the molarity of a solution that is made by diluting a 50.00 ml of a 4.74 M solution of HCl to 250.00 ml?

  What is M1? V1? V2? Solve for M2!

Example 2

  What volume of water would add to a 15.00 ml of a 6.77 M solution of nitric acid in order to get a 1.50 M solution?

  What is M1? What is M2? What is V1?

End of chapter

  We will not test on this chapter by itself, this material will be included in a unit with the material in ch 13