Unit 6 Solutions and Solubility

Unit 6 – Solutions and Solubility Text – Ch. 6 and 7

Introduction

Many of the substances we use in our lives and our chemistry class are dissolved in water. Why? Well, they are easier to react and take in our bodies. Because many substances can dissolve in water, water is usually described as the universal solvent. Not everything can dissolve in water, but many things do. The substances that do not dissolve in water are defined as insoluable. What is the chemical definition for insoluable and soluable? Precipitates are insoluable and sometimes form when ions from different solutions are combined. Why? How does our City of Winnipeg water supply ensure that there are no precipitates forming?

Definitions

Using your textbook and other resources, define the following:

1. solution 2. aqueous solutions 3. solute 4. solvent 5. electrolyte 6. ionization 7. dissociation 8. cation 9. anion 10. saturated solution 11. non-saturated solution 12. supersaturated solution

Dynamic Equilibrium and Solutions

For salts, when a solution is saturated, there is a point when the production of ions (dissolving) is in equilibrium with the ions coming back together to form the salt. This is called ______The more ions that form means the more soluable the substance is. The less ions that form means the substance is less soluable.

Diagram Animation Example - http://phet.colorado.edu/en/simulation/soluble-salts Liquids – Properties

Liquids have their particles further apart than a solid but closer than a liquid. The spaces allow the movement of particles. As a result, liquids can take the shape of their container

Since the particles are further apart, the intermolecular forces are weaker than most solids. This affects the boiling point and vapour pressure of a liquid.

Boiling Point –

Vapour Pressure –

Substances with high vapour pressures have ______boiling points.

 If the vapour pressure is high, this means that a lot of the liquid is turning into a gas. The intermolecular forces are weak and the energy needed to boil is low.

Substances with low vapour pressures have ______boiling points

 If the vapour pressure is high, this means that a lot of the liquid is turning into a gas. The intermolecular forces are weak and the energy needed to boil is low.

Ex. Vapour pressure of pure water o 0 C = 0.6 kPa o 100 C = 101.3 kPa

Notes and Diagram Factors that Affect Solubility (p. 273- 280)

In Unit 3, we discussed the roles of intermolecular forces and the molecular shapes that influence them. For dissolving of substances our rule was “like dissolves like”. What this means is polar substances will dissolve in polar substances and non-polar will dissolve non-polar. As we will see later, even though the solubility rules state it is soluable, every ionic substance is soluable to some degree.

Diagrams What are the factors that affect solubility?

1. Temperature

2. Nature of solute or solvent 3. Pressure (Gases only)

Diagrams Properties of Aqueous Solutions (p. 267 – 268)

1. One Phase

2. Conduct Electricity a. Solvation

3. Acidity or Basicity

4. Changes in the Boiling Point and Freezing Point of Water

Freezing Point Depression

Boiling Point Elevation Solution Concentrations (p. 281 – 299)

Of all the concentration units that are in general use, only four are likely to be encountered in an introductory chemistry course: mass percent, volume percent, mole fraction, and moles per liter (mol / L).

Moles per liter, or molar concentration, is called molarity (M). We will examine all four of these units in some detail, although mol/L is by far the most useful.

Mass Percent:

When we are trying to express a concentration of a solution, one of the methods chemists use is to state the solution in a “Mass Percent”. This calculation can help us to figure out the percent of certain substances in a solution. This can be practical for example when trying to find the percent of lead in our drinking water to see if it is safe to have.

To express a concentration as a “Mass Percentage” we use this equation. *

% Solute = Mass of solute x 100% Mass of solute + Mass of solvent

So we can see that in comparing the mass of the solute compared to the total mass of the solution multiplied by 100 gives us the percent mass.

Example - A solution of sodium hydroxide is made up by dissolving 10.0g of sodium hydroxide in 100.0 g of water. Find the mass percentage of sodium hydroxide in the solution.

Going back to the point about using mass percent to find iron in our water; in situations like this we use the term parts per million (ppm) or parts per billion (ppb). These terms are usually used in environmental situations caused by presence of quantities of toxic chemicals.

The equation to find ppm is very similar to the mass percentage equation. * ppm Solute = Mass of solute x 106 Mass of solute + Mass of solvent Note*: since the mass of the solute is small we can take the bottom “mass of solute” out of the equation simplyfing the equation to… ppm Solute = Mass of solute x 106 Mass of solvent

The equation for ppb is… * ppb = Mass of solute x 109 Mass of solvent

Example - Water containing more than 50 ppb of lead is unfit to drink. A certain sample of water is found to contain 8.5 x 10-5 g of lead in 1.0 L of water. Is it safe to drink? (Density of water is 1.0 g / mL)

Volume Percent:

When we are trying to measure the concentration of two liquids in a solution it is more practical to deal in terms of volume rather than mass. We normally determine that the liquid with the smaller volume is the solute and the equation for volume percent is… *

% Solute = Volume of solute x 100% Volume of solution

Example - A total of 25.0 mL of alcohol is placed in a container and sufficient water is added to bring the volume of the solution up to 125 mL. Find the percent alcohol by volume.

Other examples of volume percentages are when we label alcohol percents in drinks

U.S. Beer 4-4.6 % Alcohol Canadian Beer 4-6 % Wine 8-15% Fortified wine 20-22% Liquors 40-45% Mole Fraction:

The advantage of expressing a concentration as mole fraction is that it provides more information concerning the actual ratio of particles in the solution than the units previously discussed. The mole fraction of solute is the ratio of the amount of solute to the amount of solution.

Mole fraction of solute = Amount of Solute . Amount of solute + Amount of solvent

Example - A solution is made up of 23.0 g of ethanol (C2H5OH) and 18.0 g of water. What is the mole fraction of ethanol in the solution?

Homework Questions

1. What is the percentage by mass of the solute in each of the following solutions? a. 8.60 g of sodium chloride in 95.0 g of water b. 3.85 g of calcium chloride in 78.50 g of solution

2. An alloy of gold and copper is used for making jewellery. The composition of the alloy is normally expressed in karats, where 24 carats represents 100% gold. Calculate the composition (mass percent) of 14 karat gold.

3. Water is added to 45.0 mL of alcohol until the total volume of the mixture is 250.0 mL. Find the volume percent of alcohol in the mixture.

4. A certain bronze statue contains 80% copper and 20% zinc by mass. Calculate the mole fraction of copper in this alloy?

5. A California wine contains 11% alcohol by volume. The density of ethanol (pure alcohol) is 0.790 g/mL and that of water in 1.00 g/mL. Assuming that the volume of the solution is the sum of the volumes of alcohol and water, calculate the mass percent and the mole fraction of alcohol in the wine. (The formula of ethanol is C2H5OH) Molar Concentration:

By far the most useful way in which to express the concentration of a solution is to use moles per liter (mol/L or molL-1). We use molarity to compare all solutions as if they were 1L. Molarity is represented by the value in [ ]. An example is a 1 M solution could be represented as [1.00].

Concentration = Amount of solute (mol) Formula Volume of solution (L)

For example, what is the difference between the molarity of a solution if 0.5 mols of sodium chloride is dissolved in 500 ml of water and 1 mol of sodium chloride is dissolved in 1L of water?

Example - A solution contains 5.85 g of sodium chloride dissolved in 5.00 x 102 mL of water. What is the concentration of the sodium chloride in mol/L?

Example - What mass of potassium hydroxide is required to prepare 600 mL of a solution with a concentration of 0.225 mol/L?

Example - A solution of magnesium chloride of 0.125 mol/L is required. What is the maximum volume of solution that we can prepare if we have only 87.8 g of solid magnesium chloride? Concentrations of Ions

When certain substances are dissolved in water, the resulting solutions will conduct an electric current. This conductivity is due to the presence of ions in the solution. These ions are formed when certain covalent compounds (such as hydrogen chloride) are dissolved in water or from the dissociation of ionic solids (such as sodium chloride or magnesium chloride) when they dissolve in water. The concentration of the cations and anions in such solutions may be easily determined if we know the formula for the compound and the concentration of the solution.

Example –

Suppose we have a solution containing 15.6 g of magnesium chloride in 1.25 L of solution. How can we find the concentrations of the magnesium ions and the chloride ions?

Example – 15.50 g of aluminum sulfate are dissolved in 250 ml of water. What is the -2 [SO4 ]?

Example - A solution of volume 525 mL containing 6.78 g of calcium bromide (CaBr2) is mixed with 325 mL of a solution containing 11.4 g of potassium bromide (KBr). What is the concentration of bromide ions in the resulting solution? Assume the volumes are additive (i.e. the final volume is the sum of the two initial volumes). Homework Questions

1. What is the molarity of a solution of potassium nitrate if 16.00 g is dissolved in 500 ml of water?

2. What mass of potassium iodide would you dissolve in 200 ml of water to make a 0.5 M solution?

3. 15.00 g of sodium hydroxide is dissolved in 250 ml of water

a. What is the molarity of the solution? b. If another 15.00 g of solute is added to the water, what is the increase in molarity?

4. Using the graph on the right: a. Calculate the molarity of a sodium nitrate solution at 45oC. b. Calculate the molarity of a potassium nitrate solution at 90oC Dilution (p. 302 – 307)

We can prepare solutions by adding solid to a set volume of water. Occasionally you may find that your laboratory work calls for the use of a dilute solution of a substance, but you discover that the only solutions available are more concentrated than what you need. This situation calls for a dilution to be carried out. The thing to remember is that the number of moles of solute does not change as we are only adding more solvent to the existing solution.

Moles of solute = Concentration of solution (mol/L) x Volume of solution (L)

As a result we can set up a relationship of

Concentrationinitial x Volumeinital = Concentrationfinal x Volumefinal

Formula –

Example - We want to make 500 mL of a hydrochloric acid solution with a concentration of 3.0 mol/L starting with a concentrated solution with a concentration of 12.0 mol/L. What volume of the concentration solution should we start with and dilute to 500 mL to end up with the correct concentration of 3.0 mol/L? What volume of water needs to be added?

Example – You make a stock solution by adding 12.50 g of sodium chloride to 500 ml of water. You take 50 ml of the stock solution and add it to 100 ml of water. What is the new concentration of the sodium chloride solution?

Example – You have a 0.3 M solution of cobalt (II) nitrate. How many grams were dissolved in 250 ml of water to make the solution. If you add 50 ml of the stock solution to 50 ml of water, what is the new concentration? Homework Questions

1. How many grams of hydrochloric acid must be dissolved in 1 L of water to make a 1.0 M solution?

2. If 40.00 grams of sodium hydroxide is dissolved in 200 ml of water, what is the molarity of the solution?

3. If you want to make 0.013 M solution of potassium dichromate:

a. how many grams would you dissolve in 500 ml of water? b. If 50 ml of the above solution is added to 150 ml of water, what is the molarity of the new solution?

4. We need to prepare 500 ml of a 1.0 M acetic acid solution. The stock solution is 17.5 M. What volume of concentrated acetic acid is needed to make the solution? What volume of water is added?

5. Vinegar is usually 10% by mass of acetic acid.

a. What is the percentage by mass of acetic acid in a 1 L bottle bought at Safeway? b. What is the molarity of the vinegar? c. If you wanted to make the solution yourself by diluting the concentrated acetic acid, how much acid would you add to make the 1 L bottle? Solubility Rules

Earlier we said that if a substance is soluable it puts lots of ions in solution. It ionizes 100%. If a substance is insoluable very little ionizes and stays dissociated together.

Examples

Sodium hydroxide

Copper (II) sulfate

Aluminum carbonate

Magnesium hydroxide

How do we know what is soluable?

We use a list of Solubility Rules. (Handout)

Are the following soluable? Use the rules.

Salt Formula Soluable? Rule Sodium iodide

Potassium sulfate

Copper (II) chloride

Lead (II) acetate

Lead (II) Iodide

Mercury (II) nitrate

Sodium Carbonate

Calcium carbonate Precipitates and Net Ionic Equations

When two solutions combine and some of the ions re-associate, the resulting product is a precipitate

Precipitate –

The precipitate is insoluable based upon our solubility rules

If no precipitate forms, or no new substance forms, there is no reaction. In solutions, the ions stay dissociated and do not re-associate.

We can predict if a precipitate forms by ionizing the reactants and reforming the products by combining ions. If one of the products is insoluable, determined by the solubility rules then a precipitate forms.

Sodium chloride reacts with silver (I) nitrate

REMEMBER – PRECIPITATE REACTIONS ARE DOUBLE DISPLACEMENT REACTIONS Problems

Determine whether a precipitate occurs in the following reactions.

1. Lithium chloride reacts with copper (II) sulfate 2. cobalt (II) chloride reacts with lead (II) nitrate 3. aluminum sulfate reacts with barium chloride 4. sodium hydroxide reacts with tin (II) iodide 5. magnesium chloride reacts with potassium hydroxide

Net Ionic Equations

A Net Ionic Equation is a chemical equation for a reaction which lists only those species participating in the reaction. We eliminate the spectator ions.

Spectator Ions –

A net ionic equation just leaves us with the species that react, and form the new product in the end. **Net ionic equations are not only for precipitate reactions, but all reactions where there are species which do not take part in the formation of a new substance.

Identify the spectator ions and write the net ionic equation for the reaction of sodium chloride with silver (I) nitrate

1. Start by simply writing the overall balanced chemical reaction. This is also called the Molecular Equation.

2. Then, you break apart the soluble ionic compounds into the two ions from which it is formed (one positive and one negative). You will have to use the solubility rules to do this, they can be found online. If something is insoluble, it should not be broken apart. Write the reaction out with all of the separated ions. This is called the Total Ionic Equation.

3. Then, you simplify by cancelling things out if they appear on both sides of the reaction, resulting in the Net Ionic Equation. Ex.

Ex. Write the net ionic equation for the reaction of Copper (II) nitrate and sodium hydroxide

Ex. Write the net ionic equation for the reaction of calcium carbonate with hydrochloric acid to form calcium chloride, water, and carbon dioxide

Notice that the Cl- was cancelled out from the Net Ionic Equation, because it really isn't playing an important part of this reaction. It is just there to balance out the charge because you can't have an ion just by itself -- you must always pair an ion with another one of opposite charge so that the overall charge is zero.

Homework (I have given you an Octet to do)

Write the net ionic equations for the following. Name the precipitate if one forms. If no reaction occurs, state “NR”.

1. lead (II) acetate reacts with lithium chloride

2. sodium chromate reacts with barium nitrate

3. magnesium chloride reacts with sodium nitrate

4. sulfuric acid reacts with barium chloride

5. zinc metal reacts with hydrochloric acid to produce zinc (II) chloride and hydrogen gas

6. sodium hydroxide reacts with tin (II) iodide

7. sodium hydroxide reacts lithium chloride

8. sodium bicarbonate reacts with acetic acid to produce sodium acetate, water and carbon dioxide

Solution Stoichiometry (same steps, same stuff)

Quite often the best way to make two chemicals react with one another is to dissolve them in a suitable solvent and to mix the resulting solutions. It is particularly important to be able to perform stoichiometric calculations when situations such as this arise. As with the previous examples on stoichiometry, we usually have to determine the amount of reactant (or product) involved.

If 10 ml of a 0.1 M solution of sodium chloride reacts with excess silver nitrate solution, what mass of precipitate is formed?

Calculate the volume of silver nitrate solution [0.800] that is needed to react completely with 12.0 g of copper metal.

Cu (s) + AgNO3 (aq)  Cu(NO3)2 (aq) + Ag (s) Ex.

If 20 ml of a 0.5 M solution of barium chloride is reacted with a solution of aluminum sulfate, what mass of precipitate forms?

If 25 ml of a 0.2 M solution of silver (I) nitrate reacts with 5.00 g of copper wire, what mass of silver would be produced?

Cu (s) + AgNO3 (aq)  Cu(NO3)2 (aq) + Ag (s)

Homework

1. What mass of precipitate is produced if 10 ml of a 0.5 M solution of lead (II) nitrate solution is added to an excess amount of potassium iodide solution? State the name of the precipitate. Show the net ionic equation.

2. Calcium chloride reacts with sodium carbonate. 15.50 g of calcium chloride are dissolved in 250 ml of water.

a. What is the molarity of the calcium chloride solution? b. If 10 ml of the calcium chloride solution are needed to react completely with 10 ml of the sodium carbonate solution, what is the molarity of the sodium carbonate? c. What mass of precipitate is formed?

3. If 5 ml of a 0.1 M solution of sodium chloride is added to 10 ml of a 0.1 M solution of silver (I) nitrate, what mass of precipitate will form?

4. If 14 ml of a 0.5 M solution of copper (II) sulfate solution is added to excess sodium hydroxide, what mass of precipitate forms?

5. If 20 ml of a 0.2 M solution of sodium hydroxide is added to 20 ml of a 0.5 M solution of aluminum sulfate, what mass of precipitate is formed? Write the net ionic equation for the reaction.