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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

Topic 11 – SolutionsTopics: 3.2.3 – Infer the quantitative nature of a solution (molarity, dilution, and titration with a 1:1 molar ratio).3.2.4 – Summarize the properties of solutions.3.2.5 – Interpret solubility diagrams.3.2.6 – Explain the solution process.

I. Solution Properties Solubility is an example of a physical property that is important in chemistry.

To say that something is soluble means that it can be dissolved. Let’s use sweet tea as an example. In sweet tea: sugar is a solute (the substance that is dissolved). water is the solvent (the substance that does the dissolving). the sweet tea that is produced is the solution (a homogeneous [uniform]

mixture of two or more substances in a single phase). Sweet tea is homogeneous because the sugar is evenly distributed throughout the tea.

There are four different types of solutions with which you need to be familiar:

1. Solid solutions – mixtures of solids, such as metal alloys Examples include: brass, bronze, steel

2. Liquid solutions – one liquid dissolved in another liquid Examples include:

vinegar (acetic acid dissolved in water), gasoline (hydrocarbons such as octane and pentane)

3. Gaseous solutions – mixtures of gases

Examples include: air (mostly oxygen dissolved in nitrogen), scuba tank (typically a mixture of nitrogen, oxygen and helium)

4. Aqueous solutions – solutes (solid, liquid or gas) dissolved in water Examples include: sweet tea (sugar and tannins dissolved in water),

vinegar (acetic acid dissolved in water), soft drinks (carbon dioxide dissolved in water)

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

♦ An electrolyte is a solute that conducts electricity when dissolved in water (in solution). Solutions are electrolytic because they contain ions, which carry electric charges in the solution.

Most covalent compounds (such as sugar) are not electrolytes, because they are not made up of ions and they do not react with water to form ions. Acids and some bases (such as ammonia and amines) are among the few covalent compounds

that are actually electrolytes, because they will react with water to form ions such as hydronium and hydroxide.

Ionic compounds (such as NaCl) are electrolytes, because they are made up of ions, and their interactions with water pull the ions apart in solution. All ionic compounds are at least very slightly soluble in water, and most are very soluble in water. Once dissolved, the ions are free to roam around in solution, thereby carrying an electric current.

Here is an animation of NaCl dissolving and acting as an electrolyte: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/015_ELECTANDNON.MOV

II. The Solution Process A solute will dissolve in a solvent if the solute particles have greater attraction for the solvent than

they do for their fellow solute particles. Think of it as a “tug-of-war” where the solvent has to pull hard enough on a solute particle to overcome the attractions it has to its neighboring solute particles. A material is insoluble in a particular solvent due to a lack of attraction between the material and

the solvent (or at least the attraction between the material and the solvent is weaker than the attractions within the bonds of the material).

The polarity of a compound can significantly affect a molecule’s solubility. Recall that “like dissolves like” – solvents tend to dissolve solutes with similar polarities. For example, polar solvents tend to dissolve polar solutes. Water dissolves NaCl, because the

sodium and chloride ions are attracted to the very polar water molecules, which play “tug-of-war” with the ionic bonds holding the Na+ and Cl– ions together. NaCl dissolves in water because the pull of many H2O molecules is stronger than the ionic forces holding Na+ and Cl– ions together. These attractions between the ions and the water are called ion-dipole attractions.

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

ENERGETICS OF THE SOLUTION PROCESS

heat of solution – the amount of heat energy absorbed or released when a solute dissolves in a specific amount of solvent NH4NO3 dissolved in water is cold – an endothermic reaction (absorbs heat) CaCl2 dissolved in water is hot – an exothermic reaction (releases heat)

endothermic = positive heat of solution solute + solvent + heat solution Heat is absorbed by the reactants; heat is added to the reactants to get the reaction going.

exothermic = negative heat of solution solute + solvent solution + heat Heat is given off (lost) as a product.

NaCl dissolves in water because of the fact that water molecules have an attraction sufficient for and ions to overcome attraction of the ions for one another in the crystal. To form an aqueous solution of NaCl, water molecules also have to separate in order to create spaces that will be occupied by the ions. This process is shown in the picture below. Because of these three steps, we are able to think of solution formation as having three components. Overall enthalpy change in a solution is, therefore, a sum of three terms:

The process of separating solute particles is endothermic because there is an input of energy required in order to overcome attractive interactions. The separation of solvent molecules also requires energy. The third part comes from the attractive interactions between solvent and solute and is exothermic. Thus, the three enthalpy terms can add to positive or negative sums.

+ or – + + –

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

III. Solutions & Solubility Curves It is also important that you be able to

read and interpret solubility curves that you might come across:

♦ The solubility of ~95% of all salts (ionic compounds) increases as temperature increases. Most of the rest have solubilities that stay constant with increased temperature. A very few, such as Li2SO4 here, decrease in solubility with increasing temperature.

♦ Note carefully the axes! Solute solubility is usually measured in grams of solute that will dissolve in 100 g of water.

♦ Consider the sodium chlorate (NaClO3) solubility curve as an example: At 40C, you can dissolve

approximately 115 g of NaClO3 in 100 g of water.

At this point, the solution is said to be saturated (it holds the maximum amount of solute that can be dissolved in 100 g of water at that temperature). If you tried to add any more solute, it would simply settle out at the bottom of the beaker and not dissolve.

♦ Use these solubility curves to answer the following questions:

1. How many grams of KNO3 will dissolve in 100g of water at 60C?2. Would 60g of NH4Cl dissolved in 100g of water produce a saturated solution at 60C?3. At what temperature would 20g of KClO3 dissolved in 100g of water produce a saturated

solution?4. Which is more soluble at 70C: Ba(NO3)2 or NaCl?

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

IV. Gas Solubility♦ Gases are different from solids when it comes to characterizing their solubility in water. As the

diagrams below indicate, gases are more soluble at low temperatures and high pressures.

Gases are usually more soluble in water at lower temperatures. The hotter it gets, the less likely gases are to stay dissolved in solution. To understand why, we need to revisit the energetics of solubility, this time in terms of gases: For a gas to dissolve in solvent such as water, a “pocket” has to open in the solvent into

which the gas molecule can fit. Energy is absorbed to open a pocket in the solvent. Solvent molecules attract each other.

Pulling them apart to make a cavity will require energy, and heat is absorbed in this step for most solvents. Water is a special case – it already contains open pockets in its network of loose hydrogen bonds at or near room temperature. For water, very little heat is required to create pockets that can hold gas molecules.

Energy is released when a gas molecule is popped into the pocket. Intermolecular attractions between the gas molecule and the surrounding solvent molecules lower its energy, and heat is released. The stronger the attractions are, the more heat is released. Water is capable of forming hydrogen bonds with some gases, while organic solvents often can't. A larger amount of heat is released when a gas molecule is placed in the pocket in water than in organic solvents.

+ or – + –

There is usually net absorption of heat when gases are dissolved in organic solvents, because the pocket-making contribution is bigger. Le Chatelier's principle predicts that when heat is absorbed by the dissolution process, it will be favored at higher temperature. Solubility is expected to increase when temperature rises.

There is usually net release of heat when gases are dissolved in water, because the pocket-filling contribution is biggest. Solubility is expected to decrease when temperature rises.

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

Gases also get more soluble at higher pressures. Pressure is not a factor for solubility of solids in solution.

To understand this, consider the following diagram, where the gas starts off in equilibrium with its liquid, and gases enter and leave the liquid phase at a constant rate.

Now, if we increase the pressure on the gas by decreasing the container volume, suddenly the gas molecules above the liquid are much more concentrated.

Gas molecules begin to enter the solution at a higher rate than they leave the solution (think Le Chatelier’s Principle), until a new equilibrium is reached, where the solution contains more dissolved gas than before.

This relationship can quantified by Henry’s Law: C = kP where: C = concentration of the dissolved gask = a constant unique to that solutionP = partial pressure of the gas above the solution

In other words, Henry’s Law states that the amount of a gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.

Henry’s Law is obeyed most accurately for dilute solutions of gases that do not dissociate in or react with the solvent (usually water).

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

V. Concentration & Molarity concentration – how much solute is present in a given amount of solution

One way to measure concentration is molarity.

molarity =

molesof soluteliter of solution (symbol = M or M)

What is the molarity of 3.50 L of solution that contains 90.0 g of sodium chloride?

What mass of KNO3 must be added to 425 mL of water to create a 0.250 M solution?(Assume no change in volume due to addition of the solute to the water.)

What mass of sodium hydroxide is present in 166 mL of a 0.308 M solution?

How many liters of water must be used to dissolve 14.6g of HCl to produce a 0.500 M HCl solution (assuming that adding the solute to the water does not change the solution volume)?

Keep in mind how to do dilution problems as well:

For dilutions, use this formula:

M1V1 = M2V2

You need to dilute 110.0 mL of a 0.333 M NaOH solution down to a 0.100 M solution. What will be your final solution volume?

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

How much water must you add in order to dilute 25.0 mL of 14.6 M HCl down to a 0.500 M solution?

Assume you add 250.0 mL of water to 100.0 mL of a 0.200 M KNO3 solution. What is the final molarity of your diluted solution?

VI. Colligative Properties Colligative properties are properties of solutions that depend only on the number of solute particles,

and NOT the identity of the solute. In other words, the more solute molecules you have (the greater the solute concentration), the

greater their effect on that property.

The formulas for equations that describe the effects of these colligative properties measure concentration by molality, instead of molarity.

molality =

moles of solutekilogram of solvent (symbol = m or m)

Calculate the molality when 75.0 grams of MgCl2 is dissolved in 500.0 g of water.

How many grams CaBr2 must you add to 1.50 kg of water to prepare a 0.810 molal solution?

How many grams of water must be used to prepare a 0.500 m solution using 18.5 g of NaCl?

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

There are 4 colligative properties we specifically need to know: (1) vapor pressure reduction, (2) boiling point elevation, (3) freezing point depression, and (4) osmotic pressure

(1) Vapor Pressure Reduction A non-volatile solute lowers the vapor pressure of a solvent. Non-volatile simply means that the solute itself has no tendency to evaporate at room

temperature (or the temperature of the reaction conditions).

At the surface of a liquid there is a competition between two forces:1. the KE of the molecules, which (if the KE is high enough) will allow them to launch off

the surface of the liquid into the gas phase; and 2. the intermolecular forces in the liquid, which are trying to keep the liquid molecules on

the surface. Some of the molecules on the surface with sufficient KE will overcome the intermolecular

forces and escape into the gas phase. These gas molecules collide with the container walls and exert pressure. As more molecules evaporate, some of these gas molecules collide with the surface and stick, reentering the liquid phase.

At equilibrium the number of molecules leaving the surface just balances those returning to the surface. The measured pressure at this point is the vapor pressure.

The solute molecules decrease the vapor pressure because some of the solvent molecules on the surface have now been replaced by solute molecules.

Remember that these solute molecules are non-volatile – they don’t evaporate – so they block other solvent molecules from being able to evaporate from the liquid surface.

Since there are fewer solvent molecules on the surface to escape, vapor pressure goes down. The more solute molecules there are in the solution (more concentrated), the more they inhibit evaporation of the water, further reducing vapor pressure.

Also, keep in mind that some of the solvent molecules may be attracted to the solute, which would further inhibit evaporation & lower the vapor pressure. (Think IMFs, dipole-ion attractions, etc.)

Note that we did not need to identify the nature of the solvent or the solute (except for its lack of volatility) to derive that the vapor pressure should be lower for a solution relative to the pure solvent. That is what makes vapor pressure lowering a colligative property – it only depends on the number of dissolved solute particles.

(2) Boiling Point Elevation

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

To boil, a liquid’s vapor pressure must equal atmospheric pressure.

A liquid with weak intermolecular forces will vaporize more easily, because less energy (heat) is needed to completely break the intermolecular forces of the liquid and turn it into a vapor. More vapor = higher vapor pressure, so you won't have to increase the temperature very much until the vapor pressure = external pressure. Therefore, a liquid with a high vapour pressure at a particular temperature (due to weak intermolecular forces) will have a lower boiling point.

For 2 liquids at the same temperature, the liquid with the higher vapor pressure is the liquid with the lower boiling point.

For a solution, the vapor pressure of the pure solvent is lower than that of the solution at any given temperature. (Refer to the previous section on vapor pressure reduction to remember why.)

Therefore, a higher temperature is required to boil the solution than the pure solvent. Because both pure solvent and solution need to reach the same pressure to boil (equal to atmospheric pressure), the solution requires a higher temperature to boil.

Boiling point elevation is a colligative property – the more solute you add, more it lowers the solution’s vapor pressure, which means more heat is needed to get the solution’s vapor pressure equal to atmospheric pressure (a higher boiling point).

Real world example of boiling point elevation: adding salt to water when cooking. You then have to heat up the water more to get it to boil, which cooks food faster.

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

If we represent the difference in boiling point between the pure solvent and a solution as ΔTb, we can calculate the change in boiling point as:

The term “i” is the van't Hoff factor. It represents the number of dissociated moles of particles per mole of solute. The van't Hoff factor is 1 for all non-electrolyte solutes and equals the total number of ions released for electrolytes. (Ex: The value of i for Na2CO3 is 3: 2 mol Na+ and 1 mol CO3

2– per mole of Na2CO3.)

Calculate the boiling point of 188 g of K2SO4 in 1.0 L (= 1.0 kg) of water.

(3) Freezing Point Depression Just as solutes dissolved in a solvent raise the boiling point of the solution, they also

lower the freezing point of a solution lower than that of the pure solvent. One way to understand freezing point depression is to consider the freezing process.

For a liquid to freeze, it must achieve a very ordered state that results in the formation of a crystal. If there are impurities in the liquid (i.e. solutes) the liquid is inherently less ordered. Therefore, a solution requires a lower temperature to freeze than the pure solvent.

The more detailed explanation involves vapor pressure and solids-liquid equilibrium, but this is complicated so just make sure you understand the previous point.

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

Real world examples: antifreeze keeps water in your car from freezing in the winter (and overheating in the summer); salting icy roads keeps from freezing.

If we represent the difference in freezing point between the pure solvent and a solution as ΔTf, we can calculate the change in freezing point as:

31.65 g of sodium chloride is added to 220.0 mL of water. What is the freezing point of this solution?

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

(4) Osmotic Pressure

Osmosis refers to the flow of solvent molecules (usually water) past a semipermeable membrane that stops the flow of solute molecules only. That flow of solvent from the pure solvent side makes the volume of the solution rise.

When the height difference between the two sides becomes large enough, the net flow through the membrane ceases due to the extra pressure exerted by the excess height of the solution chamber. This pressure is called osmotic pressure.

More solvent molecules are at the membrane interface on the solvent side of the membrane than on the solution side. Therefore, it is more likely that a solvent molecule will pass from the solvent side to the solution side than vice versa.

That difference in flow rate causes the solution volume to rise. As the solution rises, it exerts a larger pressure on the membrane's surface. As that pressure rises, it forces more solvent molecules to flow from the solution side to the solvent side.

When the flow from both sides of the membrane is equal, the solution height stops rising and remains at a height reflecting the osmotic pressure of the solution.

The relationship between a solution’s osmotic pressure and its concentration is very similar to the ideal gas law:

PV = inRT

where “i” again is the van’t Hoff factor.

Realizing that n / V gives the solute concentration in units of molarity, M:

P = iMRT or = iMRT

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