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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
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