1. Important – Read Before Using Slides in Class Instructor: This PowerPoint presentation contains photos and figures from the text, as well as selected animations and videos. For animations and videos to run properly, we recommend that you run this PowerPoint presentation from the PowerLecture disc inserted in your computer. To run the animations on these slides please click on the icon on each still. If you prefer to customize the presentation or run it without the PowerLecture disc inserted, the animations and videos will only run properly if you also copy the associated animation and video files for each chapter onto your computer. Follow these steps: 1.Go to the disc drive directory containing the PowerLecture disc, and then to the “Media” folder, and then to the “PowerPoint_Lectures” folder. 2.In the “PowerPoint_Lectures” folder, copy the entire chapter folder to your computer. Chapter folders are named “chapter1”, “chapter2”, etc. Each chapter folder contains the PowerPoint Lecture file as well as the animation and video files. For assistance with copying the slide and video files, please visit our Technical Support at cengage.com/support or call (800) 423-0563. Thank you.cengage.com/support

2. Chapter 9 Acids, Bases, and Salts Spencer L. Seager Michael R. Slabaugh www.cengage.com/chemistry/seager Jennifer P. Harris

3. ARRHENIUS ACIDS & BASES ARRHENIUS ACID An Arrhenius acid is any substance that provides hydrogen ions, H +, when dissolved in water. ARRHENIUS BASE An Arrhenius base is any substance that provides hydroxide ions, OH -, when dissolved in water. EXAMPLES OF AN ARRHENIUS ACID AND BASE HNO 3 is an acid: HNO 3 (aq) H + (aq) + NO 3 - (aq) KOH is a base: KOH(aq) K + (aq) + OH - (aq)

4. BRØNSTED ACIDS & BASES BRØNSTED ACID A Brønsted acid is any hydrogen-containing substance that is capable of donating a proton (H + ) to another substance. BRØNSTED BASE A Brønsted base is any substance capable of accepting a proton from another substance. EXAMPLE OF A BRØNSTED ACID AND BASE HNO 2 (aq) + H 2 O(l) H 3 O + (aq) + NO 2 - (aq) In this reaction, HNO 2 behaves as a Brønsted acid by donating a proton to the H 2 O. The H 2 O behaves as a Brønsted base by accepting the proton.

5. EXAMPLE OF A BRØNSTED ACID & BASE (continued) The reaction was written using a double arrow that points to both the right and the left, signifying that the reaction can go in either direction and establish an equilibrium. When the reaction is read from right to left, the H 3 O + behaves as a Brønsted acid by donating a proton to the NO 2 - ion, which behaves as a Brønsted base by accepting the proton. The behavior noted above is characteristic of Brønsted acids and bases. When a substance such as HNO 2 behaves as an acid and donates a proton, the species that remains (the NO 2 - ion in this case) is a Brønsted base. Similarly, the H 2 O that behaved as a base and accepted the proton from the HNO 2 was converted into the acid H 3 O +. HNO 2 (aq) + H 2 O (l) ⇆ H 3 O + (aq) + NO 2 − (aq)

6. CONJUGATE ACIDS & BASES CONJUGATE ACIDS AND BASES The base formed (NO 2 - ) when a substance (HNO 2 ) acts as a Brønsted acid is called the conjugate base of the acid. Similarly, the acid formed (H 3 O + ) when a substance (H 2 O) acts as a Brønsted base is called the conjugate acid of the base. CONJUGATE ACID-BASE PAIRS A Brønsted acid (such as HNO 2 ) and its conjugate base (NO 2 - ) form what is called a conjugate acid-base pair. The same name is given to a Brønsted base (such as H 2 O) and its conjugate acid (H 3 O + ).

7. BINARY ACIDS Solutions of binary acids such as HCl (aq) are formed by dissolving binary compounds such as HCl gas in water. The HCl gas before being dissolved in water is said to be anhydrous (meaning without water). The anhydrous gas is named hydrogen chloride. The water solution of the gas is called hydrochloric acid. Some hydrogen-containing compounds such as HCl, HI, HBr, and H 2 S form acidic solutions when they are dissolved in water. The names of the acid solutions, such as the hydrochloric acid given above, can be obtained by following four rules.

8. RULES FOR NAMING BINARY ACIDS Rule 1: Drop the word hydrogen from the anhydrous compound name. For example, HI, called hydrogen iodide becomes "iodide". Rule 2: Add the prefix hydro- to the result of step 1. "Iodide" becomes "hydroiodide". Rule 3: Drop the suffix -ide from the result of step 2 and replace it with the suffix -ic. "Hydroiodide" becomes "hydroiodic". Rule 4: Add the word acid to the end of the name as a separate word. The final name is "hydroiodic acid".

9. RULES FOR NAMING ACIDS CONTAINING POLYATOMIC IONS Rule 1: All hydrogen atoms that are written as the first part of the formula of the acid are removed. The hydrogens are removed in the form of H + ions. Rule 2: The polyatomic ion that remains after the H + ions are removed is named by referring to sources such as Table 4.7. Rule 3: When the remaining polyatomic ion has a name ending in the suffix -ate, the suffix is replaced by the suffix –ic, and the word acid is added. Rule 4: When the remaining polyatomic ion has a name ending in the suffix -ite, the suffix is replaced by the suffix -ous, and the word acid is added. Rule 5: If the polyatomic ion contains sulfur or phosphorus, the stems -sulf or -phosph that remain when the suffixes -ate or -ite are replaced, are expanded to -sulfur and -phosphor before the –ic or –ous suffixes are added.

10. THE SELF-IONIZATION OF WATER A sample of absolutely pure water does not contain only H 2 O molecules. In addition, small but equal amounts of H 3 O + and OH - ions are also present. The reason for this is that in one liter of pure water 1.0 x 10 -7 moles of water molecules behave as Brønsted acids and donate protons to another 1.0 x 10 -7 moles of water molecules, which act as Brønsted bases. The reaction is: As a result, absolutely pure water contains 1.0 x 10 -7 mol/L of both H 3 O + and OH -. The term neutral is used to describe any water solution in which the concentrations of H 3 O + and OH - are equal. Thus, pure water is neutral because each of the ions is present at a concentration of 1.0 x 10 -7 M. H 2 O (l) + H 2 O (l) ⇆ H 3 O + (aq) + OH − (aq)

11. THE ION PRODUCT OF WATER (continued) Even though this equilibrium equation was derived on the basis of pure water, it is true for any solution in which water is the solvent. ACIDIC SOLUTION An acidic solution is a solution in which the concentration of H 3 O + is greater than the concentration of OH -. It is also a solution in which the pH is less than 7. BASIC OR ALKALINE SOLUTION A basic or alkaline solution is a solution in which the concentration of OH - is greater than the concentration of H 3 O +. It is also a solution in which the pH is greater than 7.

12. EXAMPLE OF ACID-BASE CALCULATION Calculate the molar concentration of OH - in a solution that has an H 3 O + concentration of 1.0 x 10 -5 M. Classify the solution as acidic or basic. Solution: The molar concentration of H 3 O + will be substituted into the equilibrium expression for water, the resulting equation will be solved for [OH - ]: The molar concentration of OH - is seen to be smaller than the molar concentration of H 3 O +, so the solution is classified as being acidic.

13. THE pH CONCEPT It is often the practice to express the concentration of H 3 O + in an abbreviated form called the pH rather than to use scientific notation. It is also a common practice to represent the H 3 O + ion by the simpler H + ion. The pH notation is defined below, using H + in place of H 3 O + : pH = -log[H + ], or in alternate form [H + ]= 1x10 -pH Thus, the pH is seen to be the negative of the exponent used to express the molar concentration of H + using scientific notation.

14. EXAMPLES OF pH CALCULATIONS Example 1: Calculate the pH of a solution in which [H + ]= 1.0x10 -9 M. Solution: Because the pH is the negative of the exponent on 10 used to express [H + ] using scientific notation, pH = -log (1.0x10 -9 ) = -(-9) = 9.00.

15. EXAMPLES OF pH CALCULATIONS (continued) Example 2: Calculate the [OH - ] for a solution with a pH = 4.0. Solution: Because pH is the negative of the exponent on 10 used to express [H + ] in scientific notation, the exponent must be -4. Then, [H + ]= 1.0 x 10 -4. This value is substituted into the equilibrium expression for water, and the equation is solved for [OH - ]:

16. EXAMPLES OF pH CALCULATIONS (continued) Example 3: Calculate the pH of a solution in which [H + ]= 3.6x10 -4 M. Solution: Use the pH equation, pH = -log [H + ], to find pH= -log [3.6 x 10 -4 ], then evaluate with a calculator. The pH of the solution is 3.44.

17. EXAMPLES OF pH CALCULATIONS (continued) Example 4: Calculate the [H + ] of a solution in which pH = 5.92. Solution: Use the alternate pH equation, [H + ] = 1 x 10 -pH, to find [H + ] = 1.0 x 10 -5.92, then evaluate with a calculator. The [H + ] of the solution is 1.2 x 10 -6 M.

18. PROPERTIES OF ACIDS All acids have certain properties in common such as tasting sour and producing H 3 O + ions when dissolved in water. In addition, all acids undergo characteristic double- replacement reactions with solid oxides, hydroxides, carbonates, and bicarbonates. These reactions are given below, using hydrochloric acid, HCl(aq), as a representative acid. Reaction with metal oxide: 2HCl(aq) + MgO(s) → MgCl 2 (aq) + H 2 O(l) Reaction with metal hydroxide: 2HCl(aq) + Mg(OH) 2 (s) → MgCl 2 (aq) + 2H 2 O(l)

19. PROPERTIES OF ACIDS (continued) Marble, a naturally occurring form of CaCO 3, reacts with hydrochloric acid, HCl. Eggshells are also made of CaCO 3.

20. PROPERTIES OF ACIDS (continued) Acids can react with and dissolve certain metals to yield hydrogen gas in a redox reaction. The activity series is a tabular representation of the tendencies of metals to react with H +.

21. PROPERTIES OF ACIDS (continued) The reaction of zinc metal with hydrochloric acid can be written as follows: The chloride ion (Cl - ) is a spectator ion. The hydrogen ion gains an electron to be reduced, and therefore, the HCl is the oxidizing agent. The zinc metal loses electrons to be oxidized, and therefore, the zinc metal is the reducing agent. This reaction occurs because zinc is above the reactivity line that divides lead (reactive) from copper (unreactive) in the activity series. Molecular equation : Zn (s) + 2HCl (aq) → ZnCl 2 (aq) + H 2 (g) + 2H + (aq) Zn 2+ (aq) + 2Cl − (aq) 2Cl − (aq) Net ionic equation: Zn (s) + 2H + (aq) → Zn 2+ (aq) + H 2 (g) Total ionic equation : Zn (s) → + H 2 (g)

22. PROPERTIES OF ACIDS (continued) Metals vary in their ability to reduce hydrogen ions (H + ) to hydrogen gas (H 2 ). The difference is apparent when iron, zinc, and magnesium (left to right) are put into hydrochloric acid (HCl) of the same molarity.

23. PROPERTIES OF BASES Basic solutions feel soapy or slippery to the touch and contain the OH - ion. Basic solutions also change the color of litmus from red to blue. Their most characteristic chemical property is their ability to react readily with acids in what is called a neutralization reaction. Bases also react with fats and oils and convert them into smaller, soluble molecules. Most household cleaning products contain basic substances (e.g. lye (NaOH) in drain cleaner and ammonia (NH 3 ) in liquid household cleaners).

24. CLASSIFICATION OF HOUSEHOLD PRODUCTS Weak BasesWeak Acids

25. NEUTRALIZATION REACTIONS In neutralization reactions, an acid reacts with a base to produce a salt and water. The following are typical neutralization reactions involving the base sodium hydroxide, NaOH, which is also known commercially as lye. Reaction with hydrochloric acid: NaOH(aq) + HCl(aq) → NaCl(aq) + H 2 O(l) The salt produced in this reaction is sodium chloride, commonly called table salt. Reaction with nitric acid: NaOH(aq) + HNO 3 (aq) → NaNO 3 (aq) + H 2 O(l) The salt produced in this reaction is sodium nitrate.

26. NEUTRALIZATION REACTIONS (continued) Reaction with sulfuric acid: 2NaOH(aq) + H 2 SO 4 (aq) → Na 2 SO 4 (aq) + 2H 2 O(l) The salt produced in this reaction is sodium sulfate.

27. SALTS At room temperature, salts are solid crystalline ionic compounds that contain the cation (positive ion) of a base and the anion (negative ion) of an acid in their formulas. Sodium chloride, NaCl, contains one Na + cation from the base NaOH, and one Cl - anion from the acid HCl in its formula. Sodium nitrate, NaNO 3, contains one Na + cation from the base NaOH, and one NO 3 - anion from the acid HNO 3 in its formula. Sodium sulfate, Na 2 SO 4, contains two Na + cations from the base NaOH, and one SO 4 2- anion from the acid H 2 SO 4 in its formula. The cation of a salt can be any positive ion, except H +, and it will usually be a simple metal ion or NH 4 +. The anion of a salt can be any negative ion, except OH -.

28. SALTS (continued) Salts can be formed in a number of reactions: SALT HYDRATES When salts are obtained from water solutions by evaporating the water away, specific numbers of water molecules are retained, in some cases, as a part of the recovered solid salt. These solids, called hydrates, have formulas that indicate the number of water molecules retained by the solids called water of hydration. Acid + metal → salt + H 2 Acid + metal oxide → salt + H 2 O Acid + metal hydroxide → salt + H 2 O Acid + metal carbonate → salt + H 2 O + CO 2 Acid + metal bicarbonate → salt + H 2 O + CO 2

29. THE STRENGTH OF ACIDS & BASES The strength of an acid or base is determined by the extent to which dissolved acid or base dissociates to form ions. A strong acid or base dissociates 100%, while a weak or moderately weak one dissociates less than 100%.

30. MONOPROTIC, DIPROTIC & TRIPROTIC ACIDS Monoprotic acids give up only one proton per molecule when dissolved in water. Diprotic acids give up a maximum of two protons per molecule when dissolved in water. Triprotic acids give up a maximum of three protons per molecule when dissolved in water.

31. COMMON BASES Ammonia (NH 3 ) is the weak base most often encountered in addition to the anions of strong acids. The most common strong bases are the hydroxides of group IA(1) metals (NaOH, KOH, etc.) and the hydroxides of group IIA(2) metals (Mg(OH) 2, Ca(OH) 2, etc.). NH 3 (aq) + H 2 O (l) ⇆ NH 4 + (aq) + OH − (aq)

32. Indicators An indicator changes color with changes in pH (the numbers on the tubes). Methyl red goes from red at low pH to orange. Bromthymol blue from low pH to high pH. Phenolphthalein goes from colorless to pink.

33. ANALYZING ACIDS AND BASES The analysis of acid solutions to determine the amount of acid they contain is an important procedure done in many laboratories. An acid-base titration is one commonly-used method of analysis. When a titration is done, an accurately-measured volume of acid is put into a flask using a pipet. A few drops of indicator solution is added, then a base solution of known concentration is carefully added from a buret until all the acid has been reacted (equivalence point). The point at which all the acid has reacted is shown by a color change (endpoint) in the indicator. The concentration of the base and the volume required in the titration allow the concentration of acid to be determined.

34. TITRATION TECHNIQUE

35. pH METER A pH meter can also be used to detect the equivalence point of a titration. At the beginning, the pH meter gives the pH of the acid solution being titrated. Partway through the titration, the pH meter reading is of a solution of unreacted acid and the salt produced by the reaction. At the end of the titration, the pH meter gives the pH of the salt solution formed by the complete reaction of acid with base.

36. TITRATION CALCULATIONS Titration calculations are dependent upon knowledge of two things: the stoichiometry of the reaction that occurs between the acid and base, and the equation defining molarity. An example of a reaction equation is: H 2 SO 4 (aq) + 2NaOH(aq) → Na 2 SO 4 (aq) + 2H 2 O(l) Such an equation provides the relationship between the number of moles of acid and base that react. In this reaction it is seen that 1 mole of H 2 SO 4 acid reacts with 2 moles of NaOH base. The molarity equation may be rearranged to allow the calculation of the number of moles of solute contained in a specific volume of solution or the volume of solution that contains a specific number of moles of solute.

37. TITRATION CALCULATIONS M x liters of solution = moles of solute

38. HYDROLYSIS REACTIONS OF SALTS Salts consist of the cation of a base and the anion of an acid. The cation of a base is the conjugate acid of the base from which it came. Similarly, the anion of an acid is the conjugate base of the acid from which it came. The strength of a conjugate acid or base depends upon the strength of the base or acid from which they came. The stronger an acid is, the weaker its conjugate base is. Similarly, the stronger a base is, the weaker its conjugate acid is. The pH of a water solution of a salt depends on the strength of the salt cation as an acid and the strength of the salt anion as a base.

39. HYDROLYSIS REACTIONS OF SALTS (continued) Example 1: A solution containing the dissolved salt NaCl has a pH the same as the water used as a solvent for the solution. This is because the Na + ion is the conjugate acid of the strong base NaOH and is a very weak acid. Similarly, the Cl - ion is the conjugate base of the strong acid HCl and is a very weak base. Neither the Na + cation nor the Cl - anion will react appreciably with water to produce OH - or H +.

40. HYDROLYSIS REACTIONS OF SALTS (continued) Example 2: A solution containing the dissolved salt sodium carbonate, Na 2 CO 3, has a pH significantly higher than that of the water used as a solvent for the solution. The Na + ion is a weak acid as was discussed on the previous slide. The CO 3 2- ion is the conjugate base of the weak acid HCO 3 - and as a result is a significant base that will react with water as follows: This reaction, called a salt hydrolysis reaction, is seen to produce OH - ions which causes the pH to be higher than water and the solution is basic. CO 3 2− (aq) + H 2 O (l) ⇆ HCO 3 − (aq) + OH − (aq)

41. PURE WATER vs. SODIUM ACETATE Samples of pure water (left) and sodium acetate dissolved in water (right) behave differently when phenolphthalein indicator is added. The acetate ion hydrolyzes in water to form a basic solution that turns phenolphthalein to a pink color.

42. BUFFERS Buffers are solutions with the ability to resist changing pH when acids (H + ) or bases (OH - ) are added to them. Many useful buffers consist of a solution containing a mixture of a weak acid and a salt of the acid (e.g. acetic acid and sodium acetate). Any added acid (H + ions) react with the anion from the salt, which also happens to be the conjugate base of the weak acid. Any added base (OH - ions) react with the nonionized weak acid. The buffer capacity is the amount of acid (H + ) or base (OH - ) that can be absorbed by a buffer without causing a significant change in pH. C 2 H 3 O 2 − (aq) + H + (aq) ⇌ HC 2 H 3 O 2 (aq) HC 2 H 3 O 2 (aq) + OH − (aq) ⇌ C 2 H 3 O 2 − (aq) + H 2 O (l)

43. UNBUFFERED vs. BUFFERED SOLUTIONS The solution on the left is not buffered; the one on the right is; universal indicator has been added to each solution. Sodium hydroxide has been added to each solution Hydrochloric acid has been added to two fresh samples that originally looked like the first pair of samples.

44. pH and BUFFERS The pH of buffers made this way can be calculated using the Henderson-Hasselbalch equation: In this equation, pH = -log[H + ], pK a = -log K a, [B - ] is the molar concentration of the salt of the weak acid HB, and [HB] is the molar concentration of the weak acid. If [B - ] is equal to [HB], then the pH is equal to the pK a.