1. Acids & Bases

2. Models Arrhenius’ definition says that acids contain a hydrogen ion and bases contain a hydroxide ion. But Arrhenius’ model cannot explain ammonia, NH 3, which is a base and produces hydroxide ions in solution.

3. Bronsted-Lowry defines acids as proton donors and bases as proton acceptors. HX + H 2 O   H 3 O + + X - A conjugate acid is formed when a base accepts a proton. A conjugate base is formed when an acid donates a proton.

4. Identify the acid/base pairs: HF + H 2 O   H 3 O + F - NH 3 + H 2 O   NH 4 + + OH -

5. All substances classified as acids and bases by Arrhenius are also classified as such by Bronsted-Lowry. However, bases not classified by Arrhenius are classified by Bronsted-Lowry.

6. Substances that can act as acids and bases are amphoteric. Monoprotic acids are acids that can donate only one hydrogen ion. Polyprotic acids are acids that can donate more than one hydrogen ion.

7. Monoprotic or Polyprotic? HF HNO 3 H 2 SO 4 HBr H 2 CO 3 HI H 2 S H 3 PO 4

8. Acid Base Strength Acids & bases that ionize completely are strong acids and bases. This is because they produce the maximum amount of ions in solution. As such, they are excellent conductors. HCl  H + + Cl - NaOH  Na + + OH -

9. Acids and bases that partially ionize or disassociate are called weak acids and bases. Some of the acid and some of the base remains in molecular form. As such, they do not conduct electricity as well as a strong acid or a strong base.

10. Clarification Strong or Weak refers to the degree at which an acid/base ionizes. Concentrated or Dilute refers to the molarity, or the number of acid/base molecules in the solution. Do not confuse these descriptions.

11. pH

12. pH is derived from the French for pouvoir hydrogene, or hydrogen power. It is calculated as the negative log of the hydrogen ion concentration. pH = -log [H + ] There is also a formula to calculate the hydroxide power. pOH = -log [OH - ]

13. Ion Product Constant K w = [H + ] [OH - ] = 1.0 * 10 -14 This is used to find either [H + ] or [OH - ], so K w = [H + ] or K w = [OH - ] [OH - ] [H + ] Also, pH + pOH = 14.0.

14. pH Scale Strong Acids Weak Acids Weak Bases Strong Bases H + > OH - H + = OH - H + < OH -

15. How to calculate pH from [H] and [OH]. [H] = 1.0* 10 -7 [OH] = 6.1 * 10 -5 -(log 1 + log -7) -(log 6.1 + log -5) -(0 + -7)-(.79 + -5) 74.21

16. How to find [H] and [OH] from pH Given a pH of 2.37. Using your calculator, punch in this sequence: 2 nd LOG (-2.37) = 4.27*10 -3 M This value is [H]. To find [OH], divide K w by [H]. K w = 1.0 * 10 -14. You should get 2.34*10 -12 M for [OH].

17. Your turn! Calculate the pH of the following solutions: a. [H] = 1.0*10 -2 b. [OH] =8.2*10 -6 Calculate [H] and [OH]: a. pH =11.05b. pH= 6.50

18. Neutralization

19. When acids and bases react, they neutralize each other to form water and a salt. A salt is any ionic compound, not just sodium chloride. Neutralization reactions are double replacement reactions.

20. An example of a neutralization reaction: HCl + NaOH  NaCl + H 2 O Acid Base Salt Water Mg(OH) 2 + H 2 SO 4  MgSO 4 + 2H 2 O Base Acid Salt Water

21. Titration Stoichiometry of an acid-base neutralization is the same as any other reaction. It provides the basis for titration. Titration is a method for determining the concentration of a solution by reacting a known volume of the solution with a solution of known concentration.

22. Steps A measured volume of an acidic or alkaline solution is measured and placed into a beaker. A few drops of phenolphthalien are added. A buret is filled with the titrating solution, which is usually the opposite of what you are trying to find. Measured volumes of the titrating solution are added slowly and mixed into the beaker. The point at which the [H] = [OH] is called the equivalence point, which is the point you’re trying to reach.

23. Titration Demonstration

24. Salt Hydrolysis Why are some salt solutions acidic, some alkaline, and some neutral? The answer is salt hydrolysis. In salt hydrolysis, the anions of the disassociated salt accept H+ ions from water, or the cations of the disassociated salt donate H+ to water.

25. KOH + HF  KF + H 2 O; KF  K + + F - F - + H 2 O   HF + OH - The fluoride ion acts as a Bronsted-Lowry base and accepts a H+ ion from water. HF and OH- are produced, and because there is a greater concentration of hydroxide ions (because KOH is a strong base), the solution is basic.

26. NH 3 + HCl  NH 4 Cl + H 2 O; NH 4 Cl  NH 4 + + Cl - NH 4 + + H 2 O   NH 3 + H 3 O + The ammonium ion acts as a Bronsted-Lowry acid in the forward reaction, and because there are more hydronium ions in the solution (due to HCl being a strong acid), the solution is acidic. Sodium nitrate (NaNO 3 ) is the product of a strong acid (HNO 3 ) and a strong base (NaOH). No salt hydrolysis occurs and the solution is neutral.

27. Buffers Your blood pH must be maintained at an average of 7.4 in order to survive. Acidosis or alkalosis will occur if the pH decreases or increases. Your gastric juices must have a pH between 1.6 and 1.8 to promote digestion of certain foods. How does your body maintain these strict pH values? By producing buffers.

28. Buffers are solutions that resist changes in pH when limited amounts of acid or base are added. A buffer is a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid. The mixture of ions resist changes in pH by reacting with hydrogen or hydroxide ions added to the buffered solution.

29. If your blood pH falls more than.3 units, acidosis can occur. Cramps are the result of acidosis, or lactic acid buildup in muscle tissue. If your blood pH rises more than.3 units, alkalosis can occur. 3 ways the body maintains proper pH: excretion of acid through urine, elimination of CO2 through breathing, and a system of buffers. Read page 625 to find out what they are and what they do.

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