Chem 40S Unit 5 Notes

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  • Now, when we think about acids donating protons, it isn't usually willingly.  You see, the hydrogen atoms involved are usually bonded to the rest of the molecule.  It then requires some "convincing" for the molecule to give up or "donate" its proton.  We can think of it in terms of the playground bully who "convinces" the children to "donate" their milk money.  The base is the playground bully.  Bases have a strong affinity or desire for protons.  So much so that, if they are strong enough, they will "tear" the proton from an acid and combine it with itself.   The stronger the bond between the hydrogen and the rest of the molecule, the more difficult it is for the base to take the proton.  So this accepting and donating business is just a nice way to talk about an acid being "mugged" by a base.
  • Adding Acid When an acid is dissolved in water, the acid produces a large amount of H3O+ ions. If the H3O+ ion concentration increases, the equilibrium will shift to the left to use up some of the added hydronium and maintain Kw at 1.0 x 10-14. Since equilibrium shifts left, the hydroxide ion concentration is reduced. Therefore, adding a strong acid to water increases the hydronium ion concentration and reduces the hydroxide ion concentration. Adding Base When a base is dissolved in water, the hydroxide ion concentration increases. According to Le Chatelier's Principle, the equilibrium shifts left to use up some of the added hydroxide and maintain KW at 1.0 x 10-14. Since equilibrium shifts left, the hydronium ion concentration is reduced. Therefore, adding a strong base to water increases the hydroxide ion concentration and reduces the hydronium ion concentration.
  • Adding Acid When an acid is dissolved in water, the acid produces a large amount of H3O+ ions. If the H3O+ ion concentration increases, the equilibrium will shift to the left to use up some of the added hydronium and maintain Kw at 1.0 x 10-14. Since equilibrium shifts left, the hydroxide ion concentration is reduced. Therefore, adding a strong acid to water increases the hydronium ion concentration and reduces the hydroxide ion concentration. Adding Base When a base is dissolved in water, the hydroxide ion concentration increases. According to Le Chatelier's Principle, the equilibrium shifts left to use up some of the added hydroxide and maintain KW at 1.0 x 10-14. Since equilibrium shifts left, the hydronium ion concentration is reduced. Therefore, adding a strong base to water increases the hydroxide ion concentration and reduces the hydronium ion concentration.
  • Adding Acid When an acid is dissolved in water, the acid produces a large amount of H3O+ ions. If the H3O+ ion concentration increases, the equilibrium will shift to the left to use up some of the added hydronium and maintain Kw at 1.0 x 10-14. Since equilibrium shifts left, the hydroxide ion concentration is reduced. Therefore, adding a strong acid to water increases the hydronium ion concentration and reduces the hydroxide ion concentration. Adding Base When a base is dissolved in water, the hydroxide ion concentration increases. According to Le Chatelier's Principle, the equilibrium shifts left to use up some of the added hydroxide and maintain KW at 1.0 x 10-14. Since equilibrium shifts left, the hydronium ion concentration is reduced. Therefore, adding a strong base to water increases the hydroxide ion concentration and reduces the hydronium ion concentration.
  • Adding Acid When an acid is dissolved in water, the acid produces a large amount of H3O+ ions. If the H3O+ ion concentration increases, the equilibrium will shift to the left to use up some of the added hydronium and maintain Kw at 1.0 x 10-14. Since equilibrium shifts left, the hydroxide ion concentration is reduced. Therefore, adding a strong acid to water increases the hydronium ion concentration and reduces the hydroxide ion concentration. Adding Base When a base is dissolved in water, the hydroxide ion concentration increases. According to Le Chatelier's Principle, the equilibrium shifts left to use up some of the added hydroxide and maintain KW at 1.0 x 10-14. Since equilibrium shifts left, the hydronium ion concentration is reduced. Therefore, adding a strong base to water increases the hydroxide ion concentration and reduces the hydronium ion concentration.
  • Chem 40S Unit 5 Notes

    1. 1. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases
    2. 2. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Properties of Acidic Solutions • Taste sour. For example, lemons (citric acid) and vinegar (acetic acid). • Burn when touching skin. • Turn blue litmus red. • Neutralize basic solutions. • React with carbonates to produce carbon dioxide gas. For example, when you add vinegar to baking soda (sodium bicarbonate), fizzing occurs. This fizzing is the production of carbon dioxide. • Corrosive to metals. Many acids react with active metals to produce hydrogen gas. • Another property of acids is that they are electrolytes
    3. 3. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Properties of Basic Solutions • Taste bitter. • Feel slippery. • Turn red litmus blue. • Neutralize acidic solutions. • Bases are electrolytes
    4. 4. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Arrhenius Definition vs Bronstead-Lowry Definition
    5. 5. Chemistry 40S Unit 4 – Acids & Bases Arrhenius Definition
    6. 6. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Arrhenius Definition of Acids & Bases • Arrhenius noticed that acidic and basic solutions were electrolytes. He determined that acids and bases must ionize or dissociate in water. • According to Arrhenius, an acid is defined as a substance which releases hydrogen ions in water. Example: hydrochloric acid: HCl(aq) → H+(aq) + Cl¯(aq)
    7. 7. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Arrhenius Definition of Acids & Bases • According to Arrhenius, a base is defined as a substance which releases hydroxide ions in water. Example 3, sodium hydroxide: NaOH(s) → Na+(aq) + OH¯(aq)
    8. 8. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Problems with Arrhenius Definition of Acids & Bases • The Arrhenius definition says that acids and bases can only occur in water solutions – not true. • There are many substances which are acidic or basic but do not have a hydrogen ion or a hydroxide ion. • The Arrhenius Theory is not a complete loss, however, since it was important in establishing the concept of dissociation.
    9. 9. Chemistry 40S Unit 4 – Acids & Bases Bronstead-Lowry Definition
    10. 10. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Brønsted-Lowry Definition of Acids & Bases • In 1923 a Danish chemist named Johannes Brønsted (1879-1947) and an English chemist Thomas Lowry (1874-1936) independently developed a more general definition of acids and bases within months of each other. • According to the Brønsted-Lowry definition, an acid is a proton or H+ ion donor. A base is defined as a proton acceptor.
    11. 11. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Brønsted-Lowry Definition of Acids & Bases • In this example, hydrogen chloride (hydrochloric acid) reacts with water by donating a proton. Water acts as the base, accepting the proton. The result is the H3O+ ion called the hydronium ion.
    12. 12. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Brønsted-Lowry Definition of Acids & Bases • In this example, ammonias properties as a base are better explained by the Brønsted-Lowry definition. Ammonia accepts a proton from water, making ammonia a base and water the acid. The result is the ammonium ion and the hydroxide ion.
    13. 13. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Brønsted-Lowry Definition of Acids & Bases • In the Brønsted-Lowry definition of acids and bases, substances like water can act as BOTH an acid and a base. These types of substances are called amphoteric (the root "amph" is similar to the root of amphibian which means "having two lives", on land and in water).
    14. 14. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Conjugate Acids and Bases • The general form of a Brønsted-Lowry acid-base reaction is Acid + Base → Conjugate Acid + Conjugate Base • The conjugate acid is what remains after a base has accepted a proton and the conjugate base is what remains after the acid has donated its proton.
    15. 15. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Conjugate Acids and Bases • Lets take another look at the reaction of ammonia with water. NH3(g) + H2O(l) → NH4+(aq) + OH¯(aq) Base Acid • The reverse reaction would be NH4+(aq) + OH¯(aq) → NH3(g) + H2O(l) Acid Base →
    16. 16. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Conjugate Acids and Bases • Notice an acid results form the ammonia accepting a proton from water. The ammonium ion can donate a proton to the hydroxide ion. The hydroxide ion accepts the proton making it a base. • In the first reaction, ammonia is the base and the ammonium ion is its conjugate acid. Water is the acid in the first reaction and the hydroxide ion is its conjugate base. • NH3(g) and NH4+(aq) are called a conjugate acid- base pair, as are H2O(l) and OH¯(aq).
    17. 17. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Conjugate Acids and Bases • So, for this reversible reaction: • In general for the acid HA, and for the base
    18. 18. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Hydrolysis • The Brønsted-Lowry definition can explain why a solution of sodium hydrogen carbonate (baking soda) is basic and can react with either acidic or basic solutions. • In water, NaHCO3 produces two ions: NaHCO3(s) → Na+(aq) + HCO3¯(aq)
    19. 19. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Advantages of The Brønsted-Lowry Theory • The Arrhenius Theory of acids and bases was limited to aqueous solutions. The Brønsted-Lowry Theory expands the definition of an acid and base to a proton donor or acceptor. This means the acid or base can be in any state. • The Brønsted-Lowry Theory is able to explain why substances such as the hydrogen carbonate ion can act as an acid and a base.
    20. 20. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Advantages of The Brønsted-Lowry Theory • All Arrhenius acids are Brønsted-Lowry acids, but NOT all Brønsted-Lowry acids are Arrhenius acids. • All Arrhenius bases are Brønsted-Lowry bases, but NOT all Brønsted-Lowry bases are Arrhenius bases.
    21. 21. Chemistry 40S Unit 4 – Acids & Bases Acids & Bases Your task is to write an analogy describing what a Bronstead- Lowry acid or Base is.
    22. 22. Chemistry 40S Unit 4 – Acids & Bases Strong & Weak Acids & Bases
    23. 23. Chemistry 40S Unit 4 – Acids & Bases Strong Acids • Strong acids are substances that easily donate protons. This means they must easily dissociate or ionize in water. The bond holding the proton to the rest of the molecule is not very strong. • Strong acids are usually strong electrolytes. • Strong acids are considered to COMPLETELY dissociate in water.
    24. 24. Chemistry 40S Unit 4 – Acids & Bases Strong Acids • For example, HCl(g) + H2O(l) → H3O+(aq) + Cl¯(aq) • The following are examples of strong acids. You must memorize the names and chemical formulas of these acids! Chemical Formula Name of Acid HClO4 perchloric acid HCl hydrochloric acid H2SO4 sulfuric acid HNO3 nitric acid HBr hydrobromic acid
    25. 25. Chemistry 40S Unit 4 – Acids & Bases Weak Acids • Weak acids are poor or weak electrolytes. They do not conduct an electric current in water very well. • Weak acids are poor electrolytes because they ionize incompletely. • Weak acids have a high affinity for their proton(s). The hydrogens are bound quite strongly to the rest of the molecule, so it takes a very strong base to “rip” the proton from a weak acid.
    26. 26. Chemistry 40S Unit 4 – Acids & Bases Weak Acids • An example of a weak acid is acetic acid. Only a small number of aqueous acetic acid molecules react with water to form an acetate ion and a hydronium ion. The rest exist as acetic acid molecules. • This establishes an equilibrium between the ions and the molecules of acetic acid. We show this using the double arrows: HC2H3O2(aq) + H2O(l) ↔ H3O+(aq) + C2H3O2¯(aq)
    27. 27. Chemistry 40S Unit 4 – Acids & Bases Weak Acids • Here are more examples of weak acids. Their names and formulas should be memorized! Chemical Formula Name of Acid H2CO3 carbonic acid HNO2 nitrous acid HF hydrofluoric acid
    28. 28. Chemistry 40S Unit 4 – Acids & Bases Strong Bases • Strong bases are substances have a very high affinity for protons, H+ ions. Strong bases are strong electrolytes, that is, they completely dissociate when dissolved in water. • The most common strong bases are sodium hydroxide, NaOH, and potassium hydroxide, KOH. NaOH(s) → Na+(aq) + OH¯(aq) KOH(s) → K+(aq) + OH¯(aq)
    29. 29. Chemistry 40S Unit 4 – Acids & Bases Strong Bases • Other strong bases are Chemical Formula Name of Base Mg(OH)2 Magnesium hydroxide Ca(OH)2 Calcium hydroxide CaO Calcium oxide (lime)
    30. 30. Chemistry 40S Unit 4 – Acids & Bases Weak Bases • Just as with weak acids, weak bases are poor proton acceptors. These do not have a large affinity for protons. • For example, ammonia is a weak base. Its dissociation should be memorized. NH3(g) + H2O(l) ↔ NH4+(aq) + OH¯(aq)
    31. 31. Chemistry 40S Unit 4 – Acids & Bases Weak Bases • Other examples of molecular compounds which exist as bases are Chemical Formula Name of Base CH3NH2 methylamine C6H5NH2 aniline C5H5N pyridine
    32. 32. Chemistry 40S Unit 4 – Acids & Bases Weak Bases • The most common weak bases are the conjugate bases of strong acids. • For example, the carbonate ion: CO32¯(aq) + H2O(l) ↔ HCO3¯(aq) + OH¯(aq)
    33. 33. Chemistry 40S Unit 4 – Acids & Bases Acid Strength Table
    34. 34. Chemistry 40S Unit 4 – Acids & Bases Using the Acid Strength Chart Example 1 • Arrange the following in decreasing strength as an acid and decreasing strength as bases. Note some may be used in both groups. HNO2, OH¯, HCO3¯, HPO42¯, HTe¯, C2H3O2¯
    35. 35. Chemistry 40S Unit 4 – Acids & Bases Acid-Base Reactions • Recall, an acid donates a proton and a base accepts a proton. • The stronger acid of the two tends to donate its proton, so the reaction favours the direction AWAY from the strongest acid. For example, if Acid1 was the stronger acid, the forward reaction would be favoured, away from the stronger acid.
    36. 36. Chemistry 40S Unit 4 – Acids & Bases Acid-Base Reactions Example 2 • Complete the reaction below, indicating the acids and bases. which direction is favoured and why? • HCO3¯ + PO43¯ ↔
    37. 37. Chemistry 40S Unit 4 – Acids & Bases The Equilibrium Law
    38. 38. Chemistry 40S Unit 4 – Acids & Bases Acid Dissociation Constant • Strong acids dissociate completely, and therefore do not establish an equilibrium. However, weak bases do establish equilibrium. This means, for weak acids we can make an equilibrium law. • In general, for the weak acid HA HA + H2O ↔ H3O+ + A¯
    39. 39. Chemistry 40S Unit 4 – Acids & Bases Acid Dissociation Constant • The equilibrium law would be • but water is a liquid and its concentration does not change, so we remove it and replace Kc with Ka, the acid dissociation constant or ionization constant. • The acid dissociation constant reflects the equilibrium that exists for an acid in solution.
    40. 40. Chemistry 40S Unit 4 – Acids & Bases Acid Dissociation Constant • The larger the Ka, the more product, so the greater the dissociation. The larger the Ka, the stronger the acid. • For example, hydrochloric acid completely dissociates according to the equation HCl(g) + H2O(l) ↔ H3O+(aq) + Cl¯(aq) • For a 1.0 mol/L solution, H3O+ and Cl¯ will both be 1.0 mol/L and there will be no HCl. • The smaller the Ka, the less product, so the weaker the acid. For example, the weak acid, acetic acid has a Ka equal to 1.8 x 10-5.
    41. 41. Chemistry 40S Unit 4 – Acids & Bases Base Dissociation Constant • Just as with acids, the base dissociation constant, Kb, reflects the strength of a base. The higher the value of Kb, the stronger the base. • For example, the weak base ammonia ionizes according to the equation NH3(g) + H2O(l) ↔ NH4+(aq) + OH¯(aq) • The equilibrium law is
    42. 42. Chemistry 40S Unit 4 – Acids & Bases Base Dissociation Constant • The Kb for ammonia at 25°C is 1.8 x 10-5. In general, the Kb for the weak base, B, whose dissociation is like that of ammonia is B + H2O(l) ↔ BH+(aq) + OH¯(aq)
    43. 43. Chemistry 40S Unit 4 – Acids & Bases Base Dissociation Constant • In general, for the weak base BOH, the equilibrium law would be BOH(aq) ↔ B+(aq) + OH¯(aq) • The strong base, sodium hydroxide, dissociates completely according to the equation NaOH(s) ↔ Na+(aq) + OH¯(aq) • Since no NaOH is present in solution, the Kb is very large.
    44. 44. Chemistry 40S Unit 4 – Acids & Bases Calculating the Dissociation Constant Example 1 • A 0.10 mol/L solution of acetic acid is only partly ionized. If at equilibrium, the hydronium ion concentration is 1.3 x 10-3 mol/L, what is the acid dissociation constant, Ka?
    45. 45. Chemistry 40S Unit 4 – Acids & Bases Calculating the Concentration of Dissociated Species • If we know the acid or base dissociation constant, we can calculate all the species in the solution. This requires the use of the “ICE” table. Example 2 • HA is a weak acid with a Ka of 7.3 x 10-8. What is the concentration of all species (HA, H3O+ and A¯) if the initial concentration of HA is 0.50 mol/L?
    46. 46. Chemistry 40S Unit 4 – Acids & Bases Percent Dissociation • Acids and bases can be described in terms of strong or weak, concentration and degree of dissociation. The acid and base dissociation constants represent the acids or bases degree of dissociation. Another way to describe the amount of dissociation is by percent dissociation. • Percent dissociation is calculated in a similar manner as calculating your percentage on a test or assignment.
    47. 47. Chemistry 40S Unit 4 – Acids & Bases Percent Dissociation • For percent dissociation, we use the hydronium ion, or hydroxide ion concentrations for reasons which will become more clear later in this module. In general, for the acid HA • For the acid HA, • Or, for the base BOH,
    48. 48. Chemistry 40S Unit 4 – Acids & Bases Percent Dissociation • Where [HA] and [BOH] are the initial concentrations of the acid and base, respectively. Example 3 • Calculate the percent dissociation of a 0.100 mol/L solution of formic acid (HCH2O2) if the hydronium ion concentration is 4.21 x 10-3 mol/L.
    49. 49. Chemistry 40S Unit 4 – Acids & Bases Calculating Dissociation Constant Given Percent Dissociation Example 4 • Calculate the Kb of the hydrogen phosphate ion (HPO42¯) if a 0.25 mol/L solution of sodium hydrogen phosphate is dissociated is 0.080%.
    50. 50. Chemistry 40S Unit 4 – Acids & Bases Ionization of Water • Recall, that water is amphoteric. That is, it can act as both an acid and a base. HA + H2O(l) ↔ H3O+(aq) + A¯(aq) or B + H2O(l) ↔ BH+(aq) + OH¯(aq) • Very sensitive instruments have shown that pure water actually dissociates into ions, or ionizes, slightly. We call this process self- ionization or autoionization.
    51. 51. Chemistry 40S Unit 4 – Acids & Bases Ionization of Water • The equation for self-ionization is written as: H2O(l) + H2O(l) ↔ H3O+(aq) + OH¯(aq) or H2O(l) ↔ H+(aq) + OH¯(aq) • This indicates there is an equilibrium established between hydronium and hydroxide ions
    52. 52. Chemistry 40S Unit 4 – Acids & Bases Ion Product of Water • If an equilibrium is established between hydronium ions, hydroxide ions and water molecules, an equilibrium law can be written: • Since water is a liquid, the product of Ka and water results in the ion product for water, Kw.
    53. 53. Chemistry 40S Unit 4 – Acids & Bases Ion Product of Water • The equilibrium law for water becomes: • KW = [H3O+][OH¯]At 25°C, the concentration of the hydronium and hydroxide ions are equal at 1.0 x 10-7 mol/L. • Therefore, at 25°C, the value of KW is constant at 1.0 x 10-14.
    54. 54. Chemistry 40S Unit 4 – Acids & Bases Ion Product of Water • The equilibrium law for water becomes: KW = [H3O+][OH¯] • At 25°C, the concentration of the hydronium and hydroxide ions are equal at 1.0 x 10-7 mol/ L. • Therefore, at 25°C, the value of KW is constant at 1.0 x 10-14.
    55. 55. Chemistry 40S Unit 4 – Acids & Bases Ionization of Water • If the ionization of water occurs by the equation H2O(l) + H2O(l) ↔ H3O+(aq) + OH¯(aq) • we can predict the effect of dissolving an acid or base on hydronium and hydroxide ion concentrations by using Le Chateliers Principle. • Adding Acid ? • Adding Base ?
    56. 56. Chemistry 40S Unit 4 – Acids & Bases Ionization of Water
    57. 57. Chemistry 40S Unit 4 – Acids & Bases Ionization of Water Adding Acid • When an acid is dissolved in water, the acid produces a large amount of H3O+ ions. If the H3O+ ion concentration increases, the equilibrium will shift to the left to use up some of the added hydronium and maintain Kw at 1.0 x 10-14. • Since equilibrium shifts left, the hydroxide ion concentration is reduced. Therefore, adding a strong acid to water increases the hydronium ion concentration and reduces the hydroxide ion concentration.
    58. 58. Chemistry 40S Unit 4 – Acids & Bases Ionization of Water Adding Base • When a base is dissolved in water, the hydroxide ion concentration increases. According to Le Chateliers Principle, the equilibrium shifts left to use up some of the added hydroxide and maintain KW at 1.0 x 10-14. • Since equilibrium shifts left, the hydronium ion concentration is reduced. Therefore, adding a strong base to water increases the hydroxide ion concentration and reduces the hydronium ion concentration.
    59. 59. Chemistry 40S Unit 4 – Acids & Bases Calculating Hydroxide Ion Concentration • Example 1 If 2.5 moles of hydrochloric acid is dissolved in 5.0 L of water, what is the concentration of the hydroxide ions? Assume the volume remains unchanged. Calculating Hydronium Ion Concentration • Example 2 O.40 g of NaOH is dissolved in water to make a solution with a volume of 1.0 L. What is the hydronium ion concentration in this solution?
    60. 60. Chemistry 40S Unit 4 – Acids & Bases The pH Scale
    61. 61. Chemistry 40S Unit 4 – Acids & Bases Defining pH • In 1909, a Danish chemist, named Soren Sorensen (1868 - 1939), developed a simplified system for referring to the degree of acidity of a solution. He used the pH or the potenz (power) of hydrogen. Therefore, pH describes the concentration of the hydronium or hydrogen ions in solution. • pH is defined as the negative logarithm of the hydronium ion concentration, or pH = -log[H3O+] or pH = -log[H+] (We will use the H3O+ion and H+ ion interchangeably.)
    62. 62. Chemistry 40S Unit 4 – Acids & Bases The pH Scale • Values for pH in most solutions range from 0.0 to 14.0. Pure water is considered to be neutral, or a pH of 7.0. The lower the pH, the more acidic the solution. The higher the pH the more alkaline or basic the solution. – pH < 7 acidic – pH = 7 neutral – pH > 7 basic (alkaline)
    63. 63. Chemistry 40S Unit 4 – Acids & Bases The pH Scale • Below is a pH scale with the pH values of some common solutions.
    64. 64. Chemistry 40S Unit 4 – Acids & Bases pH Calculations Example 1 • Calculate the pH of an HCl solution whose concentration is 5.0 x 10-6 mol/L.
    65. 65. Chemistry 40S Unit 4 – Acids & Bases pH Calculations Example 2 • The pH of a solution is 3.25. Calculate the hydronium ion concentration in the solution.
    66. 66. Chemistry 40S Unit 4 – Acids & Bases Defining pOH • Recall, the ion product for water: KW = [H3O+][OH¯] • If we take the negative log of each term, we get -log(KW) = pH + pOH • According to the rules of logs, when multiplying terms is equivalent to adding their logs. • If we calculate the negative log kW, -log(KW) = -log(1.0 x 10-14) = 14.00 • So, pH + pOH = 14.00
    67. 67. Chemistry 40S Unit 4 – Acids & Bases pOH Calculations Example 3 • The pH of a solution is 10.30, what is the hydroxide ion concentration?
    68. 68. Chemistry 40S Unit 4 – Acids & Bases pOH Calculations Example 4 • What is the pH of a 5.0 x 10-5 mol/L Mg(OH)2 solution?
    69. 69. Chemistry 40S Unit 4 – Acids & Bases Calculating pH, Given Ka • We can now calculate the pH of a weak acid(or base) solution, given the percent dissociation or Ka (Kb)and the acid concentration.
    70. 70. Chemistry 40S Unit 4 – Acids & Bases Calculating pH, Given Ka Example 5 • Calculate the pH of a 0.10 mol/L hydrogen sulfide solution. (Ka=1.0 x 10-7)
    71. 71. Chemistry 40S Unit 4 – Acids & Bases Calculating pH, Given Ka Example 5 - continued
    72. 72. Chemistry 40S Unit 4 – Acids & Bases Calculating pH, Given Ka Example 5 - continued

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