C2 (Lessons 5­10) E.notebook
January 21, 2014
Lesson 5
Concentration
• concentration: a ratio that compares the quantity of solute to the quantity of solution
• A dilute solution has a relatively small quantity of solute per unit volume of solution compared to a concentrated solution.
Percentage Concentration
• A solution in which a pure liquid is dissolved in water may have a percentage volume by volume (% V/V) reported.
• Ex. Vinegar is labeled 5% V/V acetic acid
• EX. What is the concentration (in % V/V) of a solution in which 225 mL of pure sulfuric acid is found in 2.00 L of solution?
• EX. An ammonia solution in a beaker is labeled 13.0 % V/V. What volume of pure ammonia has been dissolved in making 350 mL of this solution?
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• Another common concentration ratio used for consumer products is “percentage weight by volume” or % W/V • Ex. Hydrogen peroxide is labeled 3% W/V
• EX. What is the concentration (in % W/V) of a solution in which 74.0 g of sugar is found in 1500 mL of solution?
• EX. A sodium chloride solution in a beaker is labeled 3.5 % W/V. Calculate the volume of this solution that contains 750 mg of dissolved solute.
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• A third concentration unit is the “percentage weight by weight”, or % W/W.
• Ex. A stainless steel alloy contains 11.5% W/W chromium.
• EX. A sterling silver ring has a mass of 12.0 g and contains 11.1 g of pure silver. What is the concentration (in % W/W) of silver in the ring?
• EX. A bronze medal is 88% W/W copper and 12% W/W tin. What mass of each metal is present in a 282 g bronze medal?
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C2 (Lessons 5­10) E.notebook
January 21, 2014
Parts Per Million Concentration
• In solutions found in the environment (ex. iron in well water), we often find very low concentrations.
• Very small concentrations are often expressed in parts per million (ppm) or even parts per billion (ppb).
At these very low concentrations, the solution is extremely dilute. Dilute aqueous solutions are assumed to have the same density as pure water (exactly 1 g/mL).
• EX. Well water is not recommended for drinking if dissolved lead has a concentration of 15 ppb or higher. What is the maximum mass of lead that can be dissolved in 20.0 L of water that is safe to drink?
• EX. A sample of hard water contains 300 ppm calcium carbonate. What mass of calcium carbonate will be found in 500 mL of hard water?
• EX. A mountain spring water source is analyzed to contain 485 ppb of dissolved iron. What volume of spring water must be collected in order to contain 1.00 g of dissolved iron? 4
C2 (Lessons 5­10) E.notebook
January 21, 2014
Lesson 6
Molar Concentration
• The most commonly used unit of concentration is molar concentration – the amount of solute dissolved in 1 L of solution.
• Chemists often use square brackets as shorthand for molar concentration because it is so commonly used.
• EX. [NaCl(aq)] = 0.175 mol/L means “the molar concentration of sodium chloride is 0.175mol/L”
• EX. Calculate the molar concentration of a solution that has 0.870 mol of sodium bromide dissolved in 1375 mL of solution.
• EX. A solution of magnesium chloride is prepared by dissolving 2.57 g of solute into 500 mL of water. Calculate the molar concentration.
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• EX. Calculate the mass of dissolved solute in 3.00 L of 0.200 mol/L sodium oxalate solution.
• EX. Determine the volume of 7.18 x 10­2 mol/L ammonium phosphate that would contain 500 mg of dissolved solute.
Molar Concentration of Individual Ions
• In solution, ionic compounds separate by dissociation into individual ions.
• The coefficient ratio found in the dissociation equation can be used to compare the concentration of ions to the concentration of the solute.
• For example, suppose the concentration of a magnesium chlorate solution is 0.225 mol/L. The [Mg2+(aq)] and [ClO3­(aq)] can be found by:
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• EX. Calculate the concentration of each ion in a 0.100 mol/L solution of niobium(V) sulfate.
• EX. A solution is prepared by dissolving 2.75 g of barium nitrate into 280.0 mL of water. Calculate the concentration of each ion in this solution.
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• EX. 500 mL of aluminium chloride solution has a [Cl­(aq)] = 0.336 mol/L. Calculate the mass of AlCl3
(s) that was dissolved to prepare this solution.
Lesson 7:
Dilution
• A standard solution is any solution whose concentration is accurately known.
• Dilution involves the addition of solvent to a solution in order to decrease its concentration.
• For aqueous solutions, water is added to dilute the solution. This increases the overall volume without affecting the amount of dissolved solute.
• Since the amount of solute remains constant throughout dilution, the following relationship exists:
• The stock solution is the more concentrated solution that is being diluted. It is generally a standard solution.
• By diluting a solution carefully and with accurate measurements, the dilute solution will also be a standard solution.
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• QUESTION: 300 mL of water is added to 17.2 mL of 5.14 mol/L HBr(aq). Calculate the resulting concentration.
• QUESTION: For a class experiment, a teacher must prepare 1.00 L of 0.170 mol/L sulfuric acid, H2SO4(aq). This acid is usually sold as a 17.8 mol/L concentrated solution. How much of the concentrated solution should be used to make a new solution with the correct concentration?
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• QUESTION: 30.0 mL of a stock solution of CaCl2(aq) is diluted to a total volume of 500.0 mL. The dilute solution is analyzed and found to have a concentration of 15.8 mmol/L. What was the molarity of the stock solution?
Preparing Standard Solutions
• The two common ways of preparing a standard solution include:
a. dissolving a measured mass of pure solute in a certain volume of solution
b. diluting a stock solution by adding a known volume of solvent
• When dissolving a measured mass of pure solute, the steps to follow are: 1. Calculate the mass of solute that will need to be measured out.
2. Measure out this mass of solute accurately with an electronic balance.
3. Dissolve the solute in a minimum volume of water.
4. Transfer solution into a volumetric flask that is half­filled with water.
5. Fill the volumetric flask with water up to the calibration line.
6. Stopper the flask and mix the solution.
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• QUESTION: How would a 0.180 mol/L solution of lithium nitrate be properly prepared by dissolving the solid chemical?
• When diluting a stock solution, the steps to follow are: 1. Calculate the volume of stock solution that will need to be measured.
2. Measure out this volume of solution accurately with a pipette.
3. Transfer solution into a volumetric flask that is half­filled with water.
4. Fill the volumetric flask with water up to the calibration line.
5. Stopper the flask and mix the solution.
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• QUESTION: How would a 0.220 mol/L solution of acetic acid be properly prepared by dilution of a 3.50 mol/L stock solution?
Lesson 9:
Naming Acids and Bases
• Most bases are ionic compounds that contain the hydroxide ion (OH­). They are named in the same manner as other ionic compounds – the cation is named before the anion without the use of prefixes.
• Acids can be named according to two different systems – the classical system and the International Union of Pure and Applied Chemistry (IUPAC) guidelines.
• In the IUPAC system, an acid is named the same as any other solution in water. The term “aqueous” is added to the beginning of the name of the compound in solution.
• The classical system names acids differently, according to whether they contain oxygen or not.
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C2 (Lessons 5­10) E.notebook
January 21, 2014
• Naming acids containing sulfur or phosphorus is analogous to naming acids containing chlorine. Add the –ic or –ous ending to sulfur­ and phosphor­.
• For example, H2SO2(aq) is known as either aqueous hydrogen hyposulfite (IUPAC system) or hyposulfurous acid (classical system).
• For example, H2PO5(aq) is known as either aqueous hydrogen perphosphate (IUPAC system) or perphosphoric acid (classical system).
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C2 (Lessons 5­10) E.notebook
January 21, 2014
Empirical Definitions of Acids and Bases
• Many solutions of acids and bases are clear and colourless. Many neutral solutions are also clear and colourless.
• These three types of solutions can be observed to behave chemically in very different ways when testing some of their characteristic properties.
• An empirical definition/explanation is based on directly observable evidence.
• Some observable properties of acids and bases are shown below:
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C2 (Lessons 5­10) E.notebook
January 21, 2014
Arrhenius Theory of Acids and Bases
• Arrhenius proposed that certain substances break apart in water to form ions, producing a solution that can then conduct electric current.
• Knowing that both acids and bases conduct electricity, Arrhenius suggested that acids contained hydrogen atoms that formed hydrogen ions (H+(aq)) in solution. He also proposed that bases contained hydroxide ions (OH­(aq)) that could dissociate in solution.
• Hydrogen chloride, HCl(g), is a molecular substance. According to Arrhenius, when dissolved in water this molecular substance will form H+(aq) and Cl­(aq):
• Sodium hydroxide, NaOH(s), is an ionic substance. When dissolved in water, this ions break apart to form Na+
­
(aq) and OH (aq):
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