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3.  A solution consists of a solute or solutes dissolved in a solvent: › The substance that is present in the largest quantity (whether by volume, mass or amount) is usually called the solvent. › The substances that are dissolved in the solvent are called the solutes.

4.  Solutes and solvents can come in different states of matter.

5.  Substances that dissolve in water to form solutions that conduct electric current are called electrolytes.  Electrolytes include soluble ionic compounds (including bases) and acids.  Substances that do not conduct electric current when they dissolve in water are called non-electrolytes.  Non-electrolytes include molecular compounds.

6.  As a soluble ionic compound is dissolved in water, the ionic bonds between the cations (+) and anions (-) are broken and the ions dissociate apart from one another.  Examples of dissociation equations:  NaCl (s) in water Na + (aq) + Cl - (aq)  CaS (s) in water Ca 2+ (aq) + S 2- (aq)  (NH 4 ) 2 CO 3(s) in water 2 NH 4 + (aq) + CO 3 2- (aq)  Al 2 (SO 4 ) 3(s) in water 2 Al 3+ (aq) + 3 SO 4 2- (aq)

7.  Some ionic compounds do not dissolve to any great extent – they are insoluble (or only slightly soluble ).  When placed in water, very little of the actual compound (if any) will dissociate.  CuCl (s) in water CuCl (s)  Sn 3 (PO 4 ) 3(s) in water Sn 3 (PO 4 ) 3(s)

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9.  Most molecular compounds are non- electrolytes.  In pure form, the molecules of these compounds are held together by intermolecular forces such as London (dispersion) forces and dipole-dipole forces.  When these compounds dissolve, the intermolecular bonds break and the solute molecules disperse throughout the solvent.  While this process (known as separation ) happens, the individual molecules remain intact.

10. Examples of separation equations:  C 12 H 22 O 11(s) in water C 12 H 22 O 11(aq)  C 3 H 7 OH (l) in water C 3 H 7 OH (aq)  CO 2(g) in water CO 2(aq)

11.  As with ionic compounds, not all molecular compounds are soluble in water.  Paraffin wax, C 25 H 52(s), is used in wax paper and helps to keep the paper from falling apart when it becomes damp because the wax is insoluble in water. C 25 H 52(s) in water C 25 H 52(s)

12.  Some molecular compounds are found to conduct electricity after they have been dissolved.  For example, hydrogen chloride, HCl (g), is a molecular compound, yet an aqueous solution of hydrogen chloride conducts electric current.  HCl (g) in water H + (aq) + Cl - (aq)  Acids are electrolytes because they ionize (form ions) in water.

13.  An electric conductivity apparatus can be used to determine whether a solution is an electrolyte or non-electrolyte.  Non-electrolytes can be classified as molecular compounds.

14.  Ionic compounds and acids are electrolytes. Further testing with litmus paper can help to differentiate between ionic compounds and acids.  Some ionic compounds contain the hydroxide ion. For example, sodium hydroxide, NaOH (s). These compounds are bases.

15.  You have 4 solutions in unidentified beakers: › hydrochloric acid (HCl (aq) ) › sodium chloride (NaCl (aq) ) › sodium hydroxide(NaOH (aq) ) › sucrose (C 12 H 22 O 11(aq) ).  Test each solution with an electrical conductivity apparatus, red litmus paper and blue litmus paper.  Record your observations and the identification of each solute.

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17.  The solubility of a solute is the maximum quantity of solute that dissolves in a given quantity of solvent, at a given temperature.  For example, the solubility of sodium chloride, NaCl (s), in water at 20 o C is 36 g/100 mL.

18.  A saturated solution is a solution that contains the maximum amount of dissolved solute at a given temperature in the presence of undissolved solute.  For example, 100 mL of saturated NaCl (aq) would contain 36 g of dissolved NaCl at 20 o C. If more sodium chloride is added to the solution, it will not dissolve.

19.  An unsaturated solution is not saturated and, therefore, can dissolve more solute at that particular temperature.  For example, a solution containing 20 g of dissolved NaCl in 100 mL of water at 20 o C is unsaturated. This solution has the potential to dissolve an additional 16 g of NaCl before it becomes saturated.

20.  A supersaturated solution is a solution that contains more dissolved solute than its solubility at a given temperature.  For example, a solution containing 37 g of dissolved NaCl in 100 mL of water at 20 o C is supersaturated. Supersaturated solutions are never very stable and eventually the additional dissolved solute may crystallize.

21.  What mass of NaCl would be dissolved in 275 mL of saturated solution at 20 o C?  What volume of saturated solution would contain 5.52 g of dissolved NaCl at 20 o C?  A solution contains 21.4 g of dissolved NaCl in 150 mL of water at 20 o C. How much more NaCl could be dissolved into the water before it becomes saturated?

22.  What mass of NaCl would be dissolved in 275 mL of saturated solution at 20 o C? › 99 grams  What volume of saturated solution would contain 5.52 g of dissolved NaCl at 20 o C? › 15.3 mL  A solution contains 21.4 g of dissolved NaCl in 150 mL of water at 20 o C. How much more NaCl could be dissolved into the water before it becomes saturated? › 32.6 grams

23.  As an ionic compound is dissolved, it dissociates into the individual cations and anions.  For example, the dissociation of CuSO 4(s) can be shown as:  CuSO 4(s) in water Cu 2+ (aq) + SO 4 2- (aq)

24.  Crystallization is the reverse of dissolving and can be written as an equation such as: Cu 2+ (aq) + SO 4 2- (aq) in water CuSO 4(s)  In a saturated solution, which contains both dissolved and undissolved solute, both of these processes take place at the same time. in water  CuSO 4(s) Cu 2+ (aq) + SO 4 2- (aq)

25.  A saturated solution is considered to be in a state of equilibrium. Equilibrium occurs when two opposing processes take place at the same rate in a closed system.

26.  The solubility of a substance can vary greatly at different temperatures.  When a solid dissolves in a liquid, energy is needed to break the bonds holding the solid together. At higher temperatures, the particles of the solute and solvent have more energy.

27. In general, solids have a greater solubility at higher temperatures.  The bonds between particles in a liquid are not as strong as the bonds between particles in a solid. When a liquid dissolves in a liquid, additional energy is not needed.

28. In general, the solubility of liquids is not greatly affected by changes in temperature.  Gas particles move quickly and have a great deal of kinetic energy.  When a gas is dissolved in a liquid, the gas particles lose some of their energy. At higher temperatures, the dissolved gas gains energy again. As a result, the gas comes out of solution and is less soluble.

29.  In general, the solubility of gases decreases at higher temperatures.  The effect of temperature on solubility is illustrated by the problem of thermal pollution (industries releasing hot water into rivers/streams).

30.  As the temperature of the rivers/streams increase: › the dissolved oxygen comes out of the water and returns back into the atmosphere – fish and other aquatic wildlife may not have enough oxygen to breathe › harmful solutes such as mercury compounds and pesticides may become more soluble in the river/stream

31.  Pressure changes have very little effect on the solubility of liquids and solids.  Pressure changes do have a significant effect on the solubility of a gas in a liquid solvent.  The solubility of a gas is directly proportional to the partial pressure of that gas above the liquid.

32.  The solubility of oxygen in lake water depends on the partial pressure of the oxygen in the air above the lake.  If the pressure is suddenly decreased (such as when the lid of a carbonated drink is opened), the solubility of the gas will decrease, causing the gas to “bubble” out of the solution.

33.  Say I wanted to increase the amount of carbon dioxide dissolved in my pop. What are two strategies I could use to dissolve more?  Solubility practice questions

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35.  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.

36.  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

37.  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?

39.  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?

41.  Another common concentration ratio used for consumer products is “percentage weight by volume” or % W/V  Ex. Hydrogen peroxide is labeled 3% W/V

42.  What is the concentration (in % W/V) of a solution in which 74.0 g of sugar is found in 1500 mL of solution?

44.  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.

46.  A third concentration unit is the “percentage weight by weight”, or % W/W.  Ex. A stainless steel alloy contains 11.5% W/W chromium.

47.  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?

49.  A bronze medal is 88% W/W copper and 12% W/W tin. What mass of each metal is present (to the tenth) in a 282 g bronze medal?

51.  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).

52.  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).

53.  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?

55.  A sample of hard water contains 300 ppm calcium carbonate. What mass of calcium carbonate will be found in 500 mL of hard water?

56.  m = 0.150 g

57.  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?

59.  Concentration Practice Problems  Solubility Graphing Assignment due tomorrow!

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61.  The most commonly used unit of concentration is molar concentration – the amount of solute dissolved in 1 L of solution.

62.  Chemists often use square brackets as shorthand for molar concentration because it is so commonly used.  Example: [NaCl (aq) ] = 0.175 mol/L means “the molar concentration of sodium chloride is 0.175mol/L”

63.  Calculate the molar concentration of a solution that has 0.870 mol of sodium bromide dissolved in 1375 mL of solution.

64. sodium bromide

65.  A solution of magnesium chloride is prepared by dissolving 2.57 g of solute into 500 mL of water. Calculate the molar concentration.

67.  Calculate the mass of dissolved solute in 3.00 L of 0.200 mol/L sodium oxalate solution.

69.  Determine the volume of 7.18 x 10 -2 mol/L ammonium phosphate that would contain 500 mg of dissolved solute.

71.  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.

72.  For example, suppose the concentration of a magnesium chlorate solution is 0.225 mol/L. The [Mg 2+ (aq) ] and [ClO 3 - (aq) ] can be found by: Mg(ClO 3 ) 2(s) in water Mg 2+ (aq) + 2 ClO 3 - (aq) 0.225 mol/L 0.225 mol/L 0.450 mol/L

73.  Calculate the concentration of each ion in a 0.100 mol/L solution of niobium(V) sulfate.

74.  0.200 mol/L Nb 5+ (aq)  0.500 mol/L SO 4 2- (aq)

75.  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.

76.  3.76 x 10 -2 mol/L Ba 2+ (aq)  7.52 x 10 -2 mol/L NO 3 1- (aq)

77.  500 mL of aluminium chloride solution has a [Cl - (aq) ] = 0.336 mol/L. Calculate the mass of AlCl 3(s) that was dissolved to prepare this solution.

78.  m = 22.4 g

79.  Molar Concentration Practice Problems  Have an awesome break!

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81.  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.

83.  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.

84. 300 mL of water is added to 17.2 mL of 5.14 mol/L HBr (aq). Calculate the resulting concentration.

85. 0.279 mol/L

86. For a class experiment, a teacher must prepare 1.00 L of 0.170 mol/L sulfuric acid, H 2 SO 4(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?

87. 9.55 mL

88. 30.0 mL of a stock solution of CaCl 2(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?

89. 0.263 mol/L

90. The two common ways of preparing a standard solution include:  dissolving a measured mass of pure solute in a certain volume of solution  diluting a stock solution by adding a known volume of solvent

91. When dissolving a measured mass of pure solute, the steps to follow are:  Calculate the mass of solute that will need to be measured out.  Measure out this mass of solute accurately with an electronic balance.  Dissolve the solute in a minimum volume of water.  Transfer solution into a volumetric flask that is half-filled with water.  Fill the volumetric flask with water up to the calibration line.  Stopper the flask and mix the solution.  http://www.youtube.com/watch?v=XMtm4h VCGWg http://www.youtube.com/watch?v=XMtm4h VCGWg

92. How would a 0.180 mol/L solution of lithium nitrate be properly prepared by dissolving the solid chemical?

93. mass of LiNO 3 = 12.411 g in 1.00 L of water

94. When diluting a stock solution, the steps to follow are:  Calculate the volume of stock solution that will need to be measured.  Measure out this volume of solution accurately with a pipette.  Transfer solution into a volumetric flask that is half-filled with water.  Fill the volumetric flask with water up to the calibration line.  Stopper the flask and mix the solution.  http://www.youtube.com/watch?v=kMDC4 vNEoVo http://www.youtube.com/watch?v=kMDC4 vNEoVo

95. How would a 0.220 mol/L solution of acetic acid be properly prepared by dilution of a 3.50 mol/L stock solution?

96. 62.9 mL

97.  Dilution and Standard Solutions Practice Questions  Solutions Quiz Wednesday

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99.  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.

100.  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.

102. Naming acids containing sulfur or phosphorus is analogous to naming acids containing chlorine. Add the –ic or –ous ending to sulfur- and phosphor-.  Ex. H 2 SO 2(aq) is known as either aqueous hydrogen hyposulfite (IUPAC system) or hyposulfurous acid (classical system).  Ex. H 2 PO 5(aq) is known as either aqueous hydrogen perphosphate (IUPAC system) or perphosphoric acid (classical system).

103.  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.

104.  An empirical definition/explanation is based on directly observable evidence.  Some observable properties of acids and bases are shown below:

105.  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.

106.  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) :  HCl (g) in water H + (aq) + Cl - (aq)

107.  Sodium hydroxide, NaOH (s), is an ionic substance. When dissolved in water, this ions break apart to form Na + (aq) and OH (aq) :  NaOH (s) in water Na + (aq) + OH - (aq)

108.  An Arrhenius acid is a substance that ionizes to form hydrogen ions, H + (aq), in aqueous solution.  An Arrhenius base is a substance that dissociates to form hydroxide ions, OH - (aq), in aqueous solution.

109.  Pg 29 in your notes

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111.  An Arrhenius acid is a substance that ionizes to form hydrogen ions, H + (aq), in aqueous solution.  An Arrhenius base is a substance that dissociates to form hydroxide ions, OH - (aq), in aqueous solution.

112. The Arrhenius theory has two significant limitations:  It does not account for the strong attraction between an H + in solution with a neighboring H 2 O molecule.  It does not explain how some substances that do not contain OH - are able to display characteristic properties of a base.

113.  When an acid ionizes, the H + are strongly attracted to a lone pair found on the oxygen atom of a neighboring and highly polar water molecule. They bind with this oxygen to form a new substance – the hydronium ion (H 3 O + (aq) ).

114.  Arrhenius ionization of HCl: HCl (g)  H + (aq) + Cl - (aq)  Modified Arrhenius (Brønsted-Lowry) ionization of HCl: HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq)

115.  Substances such as NH 3(aq) and NaHCO 3(aq) demonstrate the characteristic properties of a base: › they conduct electricity › they turn red litmus blue › they have a pH above 7  They display these characteristic properties despite lacking the hydroxide ion in their chemical formula.

116.  These chemicals react with water to produce the hydroxide ion, OH - (aq).  Ammonia, NH 3(g) is a molecular compound. However, when it is dissolved in water, it can begin to ionize by reacting with water:

118.  An Arrhenius acid reacts with water to produce H 3 O + (aq) in aqueous solution.  An Arrhenius base dissociates or reacts with water to produce OH - (aq) in aqueous solution.  According to the modified theory, when neutral substances dissolve in water they produce neither H 3 O + (aq) nor OH - (aq) ions.

122.  Notes page 33  Need extra practice? Text pg 217 #1-5

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124.  The acidic properties of a solution are affected by both the concentration of the solution and the identity of the acid.

125.  According to the modified Arrhenius theory, all acids produce hydronium ions in water.  However, not all acids ionize to the same extent.

126.  Common strong acids: › perchloric acid (HClO 4(aq) › hydrobromic acid (HBr (aq) ) › sulfuric acid (H 2 SO 4(aq) )  hydroioidic acid (HI (aq) )  hydrochloric acid (HCl (aq) )  nitric acid (HNO 3(aq) ) A strong acid will ionize completely in water.

127.  Hydrochloric acid is a strong acid. Nearly 100% of the molecules of hydrochloric acid in an aqueous solution react with water to form H 3 O + (aq) and Cl - (aq). HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) (  100% reaction)

128.  A weak acid will only partly ionize in water. Most of the acid molecules remain intact in their original molecular form.  Ethanoic acid is a weak acid. Only about 1% of the molecules of ethanoic acid will react with water to form H 3 O + (aq) and CH 3 COO - (aq). CH 3 COOH (aq) + H 2 O (l)  H 3 O + (aq) + CH 3 COO - (aq) (  1% reaction)

129.  Note that the ionization of a weak acid is an equilibrium system.  This means that ethanoic acid molecules are constantly reacting with water to form CH 3 COO - (aq) and H 3 O + (aq), and that those ions are constantly reacting to form ethanoic acid molecules and water.

130.  On average, at any one time, only about 1% of ethanoic acid molecules are ionized.

131. Based on this difference in ionization, a strong acid will more clearly display acidic properties than a weak acid of equal concentration.  A 0.10 mol/L strong acid will have a higher concentration of H 3 O + (aq) ions than a 0.10 mol/L weak acid.  Therefore, a 0.10 mol/L strong acid will react more vigorously with magnesium and will be a better conductor of electricity.

133.  As energy is added to the magnesium (in the form of thermal energy from the lighter), the electrons in the magnesium strip become “excited” and jump up to a higher energy level.  As the metal cools down, the electrons return to their lower energy levels; as they return, they give off that extra energy in the form of light.

134.  A strong base (a metal hydroxide) dissociates completely into ions in water.  A dissociation equation demonstrates what happens when a strong base is dissolved in water. in water  Ba(OH) 2(s)  Ba 2+ (aq) + 2 OH - (aq)

136.  One common weak base is aqueous ammonia, NH 3(aq). Only about 1 percent of the ammonia molecules react in water to form hydroxide ions.

137.  Aqueous Solutions Practice Problems  Notes page 37

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140.  Some acids react in a 1:1 ratio of water. These monoprotic acids have only one hydrogen atom per molecule that ionizes in water.  The general formula of a monoprotic acid would be HA (aq).

141.  Examples of monoprotic strong acids: hydroioidic acid (HI (aq) ) hydrochloric acid (HCl (aq) ) nitric acid (HNO 3(aq) )

142. The ionization of a strong monoprotic acid: Form: HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) Ex. HNO 3(aq) + H 2 O (l)  H 3 O + (aq) + NO 3 - (aq)

143.  Examples of monoprotic weak acids: hydrofluoric acid (HF (aq) ) ethanoic acid (CH 3 COOH (aq) ) hypochlorous acid (HOCl (aq) )

145.  Many acids contain two or more hydrogen atoms that can ionize. These are called polyprotic acids.  The general formula of a polyprotic acid would be H 2 A (aq) or H 3 A (aq).

146.  For example, sulfuric acid, H 2 SO 4(aq), has two hydrogen atoms that can ionize. The first ionization takes place completely – sulfuric acid is a strong acid: H 2 SO 4(aq) + H 2 O (l)  H 3 O + (aq) + HSO 4 - (aq)  The hydrogen sulfate ion, HSO 4 - (aq), that has been produced is also an acid but it is a weak acid.

150.  Most bases are anions that react with water to produce hydroxide ions.  A monoprotic base dissociates or reacts with water in one step. The general formula is B - (aq).

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156.  On average, only about two water molecules in every billion are ionized at any given moment.  Since such a small number of ions are produced, very little electric current can be conducted.

157.  In pure water at 25 o C, the concentration of the hydronium ions is only 1.0 x 10 -7 mol/L.  The concentration of the hydroxide ions in pure water is also found to be 1.0 x 10 -7 mol/L.  Therefore, in pure water or any neutral solution:  [H 3 O + (aq) ] = [OH - (aq) ] = 1.0 x 10 -7 mol/L

158.  Acids produce hydronium ions in water. Therefore, acidic solutions contain a higher concentration of hydronium ions than hydroxide ions.  [H 3 O + (aq) ] > [OH - (aq) ]  [H 3 O + (aq) ] > 1.0 x 10 -7 mol/L  [OH - (aq) ] < 1.0 x 10 -7 mol/L

159.  Bases produce hydroxide ions in water. Therefore, basic solutions contain a higher concentration of hydroxide ions than hydronium ions.  [OH - (aq) ] > [H 3 O + (aq) ]  [OH - (aq) ] > 1.0 x 10 -7 mol/L  [H 3 O + (aq) ] < 1.0 x 10 -7 mol/L

160.  acidic solution: [H 3 O + ] > [OH – ]  neutral solution: [H 3 O + ] = [OH – ]  basic solution: [H 3 O + ] < [OH – ]

161.  pH is a convenient method of communicating the acidity or basicity of a solution.  The pH scale typically runs from 0 to 14 – however, pH values less than 0 and greater than 14 are possible. These extreme pH values are rare. › A pH of exactly 7 represents a neutral solution. › Any pH less than 7 represents an acidic solution. › Any pH greater than 7 represents a basic solution.

162.  The pH of a solution can be calculated by using the following equation:  pH = - log [H 3 O + (aq) ] (read: pH is the negative log of the hydronium ion concentration)

163.  pH is written as a unitless value.  When reporting a pH value, only the digits to the right of the decimal are considered significant.  For example, a pH of 7.18 has two significant figures. A pH of 12.4 has only one significant figure.

164.  A solution has a hydronium ion concentration of 3.50 x 10 -8 mol/L. Calculate its pH.

165. pH = - log [H 3 O + (aq) ] pH = - log (3.50 x 10 -8 mol/L) pH = 7.456

166.  A solution is prepared by dissolving 2.18 g of HBr (g) in enough water to make 500 mL of solution. Calculate the resulting pH.

168. Chemistry 20

169.  Acid-base indicators can be used to identify the approximate pH of a solution.  Indicators are chemicals that have the ability to change colour depending upon the pH of the solution into which they are placed.

171.  By testing an unidentified solution with one or more indicators, its approximate pH can be determined.  The more indicators used, the more accurately the pH can be determined.

172.  A solution of unknown pH is tested with phenol red. The indicator remains with a red colour. Identify a possible pH range for this solution.

173.  A solution of unknown pH is tested with methyl orange, which appears yellow, and with bromothymol blue, which appears yellow. Identify a possible pH range for this solution.

174.  A solution of unknown pH is tested with methyl orange, orange IV, phenol red and methyl red, all of which appear yellow. Identify a possible pH range for this solution.

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176.  To find pH, we use pH = - log [H 3 O + (aq) ]  So, it should surprise you that to calculate pOH, we use pOH = - log [OH - (aq) ] (read: pOH is the negative log of the hydroxide ion concentration)

177.  Calculate the pOH of a solution with [OH - (aq) ] = 6.8 x 10 -5 mol/L. Is this solution acidic or basic?

178.  Calculate the pOH of a solution prepared by dissolving 0.21 g of sodium hydroxide into enough water to make 12.5 L of solution. Is this solution acidic or basic?

179.  For pure water at 25 o C, [H 3 O + (aq) ] = 1.0 x10 -7 mol/L and [OH - (aq) ] = 1.0 x10 -7 mol/L.  Therefore, pure water has a pH of 7.00 and a pOH of 7.00.

180.  For all dilute aqueous solutions at 25 o C,  pH + pOH = 14.00  Therefore, the lower the pH value of a solution, the greater the pOH value (and vice versa).

181.  This concept can be added to the pH scale:

182.  Calculate the pH and pOH of a 3.25 x 10 -5 mol/L solution of hydroiodic acid (a strong acid).

183.  Calculate the pH and pOH of a solution that contains 6.18 x 10 19 dissolved hydroxide ions in a volume of 210 mL.