Solutions Notes:



Solutions Notes:

Because water is all over the place and most of the reactions we see in our daily lives take place in water, it’s handy to learn as much as we can about things that are dissolved in it.

There are three main different ways that stuff can be mixed with liquids:

1) Suspensions: Heterogeneous mixtures where small particles of a solid are floating around in a liquid.

• How you can tell if something is a suspension: over time, these particles settle to the bottom of the liquid. For example, if mud is left alone for a long period of time, the particles settle to the bottom.

• Examples of solutions: Mud, peanut butter, ketchup, etc.

2) Colloids: Heterogeneous mixtures in which mixtures consist of two phases of matter that are stable and don’t settle out no matter how much time you give them.

• Why they don’t settle out: Brownian motion

o Brownian motion is when the particles of the thing that’s dispersed is hit from all directions by the particles of the thing it’s dispersed in. Because the particles of the thing that is dispersed are so small, they keep it from settling.

• How you can tell if something is a colloid: It experiences the Tyndall effect: When light shines through a colloid, you can see the rays of light as they travel though the solid because the suspended particles reflect/scatter it. For example, smoke is colloidal because you can see light beams pass through it, and the smoke particles don’t settle to the bottom. This causes it to appear “milky” or “foggy.”

• Examples of colloids: Butter, milk, smoke, aerosols.

3) Solutions: Homogeneous mixtures, in which one material completely dissolves in another.

• All solutions have two parts:

o Solvent: The thing that does the dissolving.

o Solute: The thing that gets dissolved.

• How you can tell if something is a solution: If you let it sit for a long time, nothing settles to the bottom. Additionally, it doesn’t experience the Tyndall effect (i.e. it doesn’t appear “cloudy” and light beams are invisible as they pass through it.

• Give several examples of solutions and have students determine what the solvent and solute are.

Mini lab where they determine whether several unknown materials are colloids, suspensions, or solutions.

• Possibilities for the mixtures: Suspended silica gel, water/methyl orange, milk/water mixture, milk/water/dye mixture (just 1 drop of milk, 1 drop of dye), copper (II) sulfate solution.

Concentration:

It’s frequently handy to know how much stuff is dissolved in a solution.

concentration: Any measurement of how much solute is dissolved in a solution.

Qualitative ways of measuring concentration:

• Unsaturated: (define and explain)

• Saturated: (define and explain)

• Supersaturated: (define and explain). Some properties:

o Two ways of making them:

▪ Cool a saturated solution.

▪ Evaporate solvent from a saturated solution.

o Supersaturated solutions are never stable; when disturbed they will recrystallize.

Quantitative methods of measuring concentration (from most to least commonly used):

• molarity: moles of solute / liter of solution.

o Sample problems:

▪ What’s the molarity of a solution that contains 0.5 moles of NaCl dissolved to make 1.5 L of solution? (0.33 M)

▪ If I have 30 g LiOH dissolved in 300 mL of solution, what’s the molarity? (4.2 M)

o How to make a solution:

▪ Using the equation M = mol/L, determine how much solute will be required.

▪ Add solvent until the solution is at the desired volume.

▪ Stir until the solute is completely dissolved.

▪ Add more solvent until the solution is at the desired volume (this will require more solvent because the volume of the solution will decrease when the solute dissolves – explain this).

o How NOT to make a solution:

▪ Using the equation M = mol/L, determine how much solute will be required.

▪ Add a volume of solvent equal to the volume of the solution you’re trying to make.

▪ Stir and dissolve.

▪ The reason this doesn’t work well is that the solute has a volume of its own, so the final volume of the solution will be greater than desired, making it slightly less concentrated that you’d like.

Molarity practice sheet

Lab: Making a solution

▪ General: Have them make 100 mL of a 1.50 M NaCl solution.

▪ Honors: Have them make 100 mL of a 1.50 M MgSO4.7H2O solution.

Tell them ahead of time that this lab will be graded completely by how close their concentration is to the desired concentration – if the concentration is 90% of what it should be, they will receive a 90 on the lab.

Molarity Homework

Dilutions:

It’s frequently handy to take a very concentrated solution of a compound and dilute it so it’s less concentrated – for one thing, it’s much cheaper to do that than it is to always purchase exactly what you need. This is frequently done with acids and bases.

To do dilutions, use this equation:

M1V1 = M2V2 (Explain what the terms mean)

o Example: If I add 25 mL of water to 5 mL of 0.15 M NaCl, what’s the final concentration of the NaCl solution? (0.025 M)

o Example: If I dilute 5 mL of 0.15 M NaCl solution to a final volume of 25 mL, what is the final concentration of the NaCl solution? (0.030 M)

o [More examples that are aimed at making sure that students can tell the difference between these]

Dilutions practice worksheet

Dilutions lab:

o General chemistry: Have them make 100 mL of 0.150 M NaCl solution from 2 M NaCl.

o Honors chemistry: Have them explain how to make a 1.00 x 10-8 M NaCl solution using the equipment in our lab.

Dilutions homework

Other quantitative units of concentration:

• molality = moles of solute / kilograms of solvent

o The unit of molality is m, called “molal” – for example, a 0.25 m solution is said to be “0.25 molal.”

o For water, 1 mL = 1 gram and 1 L = 1 kg.

o Sample problems:

▪ What’s the molality of a solution made by adding 1.50 L of water to 0.75 moles of NH4OH? 0.50 m

▪ How many grams of NaOH should be added to 750 mL of water to make a 0.25 m solution? 7.5 grams

• [pic]

o This is sometimes used when a solid solute is dissolved in a liquid solvent.

o Example: What’s the mass percent of NaOH if I have added 45 grams of NaOH to 250 mL of water? 15% (note: make sure the denominator = 250 + 45 = 295 grams!)

• [pic]

o This is sometimes used when a liquid solute is dissolved in a liquid solvent (the “solute” is the less abundant liquid in the solution).

o Example: If I add 50 mL of rubbing alcohol to water and the final volume of the solution is 200 mL, what’s the percent by volume of rubbing alcohol in the solution? 25% (50/200)

• mole fraction (X):

o [pic], [pic], etc.

o The sum of the mole fractions in a solution is equal to 1.

o Mole fraction is one of the few numbers that has no units associated with it.

o Example: What’s the mole fraction of CH3OH in a solution that contains 0.75 moles of CH3OH and 6.0 moles of water? (Xa = 0.75 / 6.75 = 0.11)

Concentration practice sheet

In-class group activity: Describe how to make a solution that’s 0.5 in each of the units above.

Concentration homework

Solubility – Why stuff dissolves

The main rule: Things dissolve when it’s more stable for the solute particles to mix with the solvent particles than it is for the solute and solvent particles to stay separated.

• Polar solutes in polar solvents: When you put a polar solute in a polar solvent, the partial charges on the solvent molecules will pull on the partial charges of the solute molecules. Because the solvent and solute molecules interact well with each other, polar solvents dissolve polar solutes.

• Polar solutes in nonpolar solvents: When you put a polar solute in a nonpolar solvent, the solute molecules would rather stick to each other (because of the opposite charges) than the solvent molecules (which don’t have them). As a result, polar solutes don’t dissolve in nonpolar solvents.

• Nonpolar solutes don’t dissolve in polar solvents: Though the nonpolar solutes don’t want to stick to each other, the polar solvents do. Because the polar solvent molecules would rather stick to each other than to the nonpolar solutes, nonpolar solutes don’t dissolve in polar solvents.

• Nonpolar solutes dissolve in nonpolar solvents: Even though nonpolar solvent molecules don’t pull on nonpolar solutes, there’s also not much holding the solute together. As a result, nonpolar solute molecules will eventually mix with nonpolar solutes.

Because things that are polar dissolve in polar solvents and things that are nonpolar dissolve in nonpolar solvents, we use the saying “LIKE DISSOLVES LIKE” to describe the solubility of solids in liquids.

If we find that something does dissolve in a solvent, how can we make it dissolve more quickly?

• Grind the solute particles and make them smaller: Because solutes only dissolve at the point where their molecules hit the surface of the liquid, smaller particles will equal faster dissolving.

• Stir it: When you dissolve a solute the area immediately around the undissolved solute is in contact with very concentrated solution. By stirring it, you put the solute in contact with unconcentrated solution which makes it dissolve more quickly.

• Change the temperature:

o Solids are usually more soluble in hot solvents than in cold solvents because the solvent molecules have more energy to pull them apart.

▪ Show them the solubility chart on p. 458 and have them use it for practice.

o Gases are usually less soluble in hot solvents than in cold ones because when energy is added to the gas molecules they are more prone to vaporize than to dissolve.

• If it’s a gas, increase its partial pressure: If you want to dissolve a gas into a liquid, increasing the partial pressure of the gas will allow more gas molecules to come into contact with the liquid and dissolve. This relationship is expressed by Henry’s Law:

[pic]

o Sample: If a gas has a solubility of 0.38 g/L at 1.0 atm and we increase the partial pressure of the gas to 2.5 atm, what will the new solubility of the gas be? 0.95 g/L

Lab: The group that most quickly dissolves their solute wins.

• Give them 20.0 grams of various solutes (MgSO4.7H2O, sugar, salt, mothballs) and quantities of different solvents (isopropanol, water). The first group to dissolve their solute gets an additional 10 percent on their lab grade.

Like dissolves like homework sheet

Solubility homework sheet

Colligative Properties:

Not surprisingly, some of the properties of solutions change when you change the concentration. Concentrated solutions generally tend to be thicker, darker in color, and denser.

Colligative property: Any property of a solution that changes when the concentration of the solution changes.

Examples of colligative properties and how they work:

• Boiling point elevation:

o As you increase the concentration of a solution, the boiling point of the solution increases.

o The reason for this is that adding solute decreases the vapor pressure of a solution because there are fewer molecules at the surface of the solution that can vaporize. Because liquids boil when their vapor pressure is equal to the ambient atmospheric pressure, you need to heat them more to make them boil.

[pic]

More about BP elevation (

o The relationship between concentration and boiling point is described by the equation:

∆Tb = Kbm

Where:

▪ ∆Tb is the change in boiling point

▪ Kb is the ebullioscopic constant (0.520 C/m for water)

▪ m is the effective molality of the solute

o “Effective molality” refers to the molality of particles in the solution.

▪ For covalent compounds, the effective molality is the same as the regular molality because the molecules don’t break into smaller pieces when you put them in water.

• For example, if you have a 0.75 m solution of CS2, the effective molality is 0.75 m.

▪ For ionic compounds, the effective molality is equal to the molality times the number of ions in the compound.

• For example, if you have a 0.75 m solution of NaCl, the effective molality is 1.50 m (0.75 x 2)

• More examples, if needed.

o Sample problems:

• What is the boiling point of a 1.75 m solution of H2CO? 100.910 C

• What is the boiling point of a 1.75 m solution of NaCl? 101.820 C

• What is the boiling point of a 1.75 m solution of Ca3(PO4)2? 104.550 C

• Melting point depression:

o Solutions melt at lower temperatures than pure liquids because the solute molecules disturb the intermolecular forces between the solvent molecules. As a result, less energy is needed to break up the solid.

o The relationship between concentration and melting point is determined by the equation:

∆Tf = Kfm

Where:

▪ ∆Tf is the change in freezing point

▪ Kf is the cryoscopic constant (1.860 C/m for water)

▪ m is the effective molality of the solute

o The rules for determining effective molality are the same as for boiling point. The only thing different is the constant used.

▪ Example: What is the melting point of a 1.75 m solution of NaCl? -6.510 C

• Vapor pressure decrease

o The more concentrated a solution, the lower the vapor pressure (for the reasons we described above with the BP elevation chart, shown again here):

[pic]

• Osmotic pressure increase

o Osmotic pressure is the force with which a pure solvent moves across a semi-permeable barrier into a container that holds a solution (explain what a semi-permeable barrier is).

o This occurs because solvents tend to diffuse across barriers in ways that will decrease the difference in concentration. The bigger the difference in concentration (because of high solute concentration), the stronger the force of osmosis.

• Conductivity of electricity

o Electrolytes: Solutions that are able to conduct electricity.

▪ These are solutions in which ionic compounds are the solute – recall that ionic compounds are good at conducting electricity when dissolved in water.

o The more concentrated the solution of an ionic compound, the better it is able to conduct electricity.

▪ The larger the number of mobile ions present, the larger the number of electrons that can be moved from one place to another.

In class colligative property worksheet

Colligative property lab: Determine which solution is more concentrated and calculate its concentration from the boiling point.

Homework like this

Review sheet for quiz

Quiz on solutions

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