1. Acids and Bases – Unit 13

2. Chemistry of Acids and Bases 1. Watch video and complete worksheet 2. Gallery walk to complete notes on pages 3-5 in packet 3. Homework is on page 6 in packet  Standard Deviants Teaching Systems: Chemistry: Module 05: Acids and Bases  http://app.discoveryeducation.com/player/view/asset Guid/DBD191DB-A10E-43C2-8DDE-A73858F12FE2 http://app.discoveryeducation.com/player/view/asset Guid/DBD191DB-A10E-43C2-8DDE-A73858F12FE2

3. Unit 13 – Acids and Bases Notes #1: Intro Acids: Something that produces a hydrogen ion (H + ) in solution

4. Unit 13 – Acids and Bases Notes #1: Intro Properties of Acids: Tart or sour taste (lemon juice) Electrolytic Both strong and weak Will cause indicators to change colors A metal + an acid will produce hydrogen gas Single replacement reaction Acid + metal → hydrogen gas + a “salt” Double replacement reaction Acid + Base → water + a “salt”

5. Single replacement reaction Acid + Metal → __Hydrogen gas_ + a “_salt_” Double replacement reaction Acid + Base → _water__ + a “_salt_”

6. Acid Naming Rules “Handle acids carefully so you don’t get a case of “ate-ic- ite-ous.””  Polys ending in “-ate” are changed to “-ic”  Polys ending in “-ite” are charged to “-ous” Hydro- prefix is not used with poly containing acids!!!!! ION TYPEION ENDINGACID NAME BEGINNINGACID ENDING Polyatomic -iteNO hydro- beginning-ous -ateNO hydro- beginning-ic Monatomic-idehydro- beginning-ic

7. Examples of Naming Binary Acids  HCl  HF  HBr Hydrochloric acid Hydrofluoric acid Hydrobromic acid

8. Examples of Naming Ternary Acids  H 2 SO 4  H 2 CO 3  H 2 NO 2 Sulfate is the poly, so sulfuric acid carbonate is the poly, so carbonic acid Nitrite is the poly, so nitrous acid

9. Base: Something that produces a hydroxide ion (OH - ) in solution

10. Unit 13 – Acids and Bases Properties of Bases: bitter slippery (soap) electrolytic Both strong and weak Will cause an indicator to change colors

12. Naming Bases The easiest are the bases, since most of these are _metal hydroxides, compounds you already know how to name. Metal hydroxides are named in the same way any other ionic compound is named. First give the name of the _metal_ ion. Follow this with the name of the anion, which, in the case of bases, is “__hydroxide__”. KOH – Mg(OH) 2 – Potassium Hydroxide Magnesium Hydroxide

13. Other definitions of Acids and Bases  Arrhenius Acids and Bases:  Acid:  Hydrogen containing compound that ionize to yield a hydrogen ion in solution.  Base:  Compounds that ionize to yield a hydroxide ion in solution.

14. Brønsted – Lowry Acids and Bases  They felt the Arrhenius definition was too limiting.  Acids:  Hydrogen ion donor (Proton donor)  Bases:  Hydrogen ion acceptor (Proton acceptor)

15. Brønsted – Lowry Acids and Bases  Examples:  NH 3 + H 2 O ↔ NH 4 + + OH -  H 2 O donated the H + - Acid  NH3 accepted the H + - Base  HCl + H 2 O ↔ H 3 O + + Cl -  HCl donated the H + - Acid  H 2 O accepted the H + - Base

16.  Amphoteric:  Substance that can act as both an acid or a base.  Background Theory:  The oxides of metals are basic in nature. For example, the oxides of the alkali metals (Group I) form alkali or basic solutions. o Sodium oxide + water → Sodium hydroxide solution Na 2 O(s) + H 2 O(l) → NaOH(aq)  The soluble oxides of non-metals are acidic in nature. Examples include, carbon dioxide, sulfur dioxide and nitrogen dioxide. o Sulfur dioxide + water → Sulfurous acid SO 2 (g) + H 2 O(l) → H 2 SO 3 (aq) o Insoluble non-metallic oxides like carbon monoxide do not form acidic solutions. This is often the cause of acid rain.  Compounds such as the amino acids, which contain both acidic and basic groups in their molecules, can also be described as amphoteric.

17. Strong Acids and Bases  Strong Acids/Bases:  Those that ionize completely in solution.  Ex: HCl, NaOH  Weak Acids/Bases:  Those that only slightly ionize in solution.  Ex: NH 3, Acetic Acid (vinegar)  Tooth decay is caused by the weak acid – lactic acid: C 3 H 6 O 3

20. Homework: pg 6

21. Notes: pH and pOH pg 7

22. pH Scale

23. MEASURING pH Scientists use a pH scale to measure the strength of an acid or base. The term pH stands for “potential for hydrogen”. The amount of hydrogen in a substance determines its acidity or alkalinity. Alkaline is another term for base. A number on the pH scale is used to describe the strength of acidity or alkalinity. The most commonly used pH scale goes from 1 (very acidic) to 14 ( very basic). The number 7 on a pH scale means neutral – neither acid nor base. Acids play important roles in the chemistry of living things. Many of the foods you eat are acids in vitamins like ascorbic acid or vitamin C, and folic acid. Other acids help the body such as stomach acids and others are waste products of cell processes like lactic acid in working muscles. Acids also are used to make valuable products for homes, farms and industries. People often use dilute solutions of acids to clean brick and other surfaces. Hardware stores sell muriatic (hydrochloric ) acid, which is used to clean bricks and metals. Industry uses sulfuric acid in car batteries, to refine petroleum and to treat iron and steel. Farmers depend on the nitric acid and phosphoric acid to make fertilizers for crops, lawns, and gardens.

24. The concentration of hydrogen ions in a solution is described by its number on the pH scale. A low pH tells you that the concentration of hydrogen ion is high. EX: pH 2 By comparison, a high pH tells you that the concentration of hydrogen ion is low. EX: pH 12

25. Self-ionization of water  Self-ionization of water:  Reaction in which 2 water molecules produce ions  H 2 O + H 2 O → OH - + H 3 O +  Also written as: H 2 O ↔ H + + OH -  The H 3 O + and H + represent hydrogen ions in solution.

26. Neutral Solutions  In pure water, the concentration of hydrogen ions is equal to the concentration of hydroxide ions  1 x 10 -7 M or pH of 7  Remember M represents Molarity  [H + ] = [OH - ]  (brackets represent concentration)  This represents a neutral solution.

27. Solutions  In a solution, if the [H + ] increases, the [OH - ] decreases and vice versa.  Think back to a see-saw. As one person went up the other went down.  Ion-product constant of water, Kw:  Kw = [H + ] x [OH - ] = 1 x 10 -14 M  Acidic Solution:  The [H + ] will be greater than the [OH - ].  Therefore, the [H + ] is greater than 1 x 10 -7 M.  Think about the # line. -5 is GREATER than -7  Basic Solution:  The [H + ] will be less than [OH - ].  Therefore, the [H + ] is less than 1 x 10 -7 M.  A.k.a. alkaline solutions

28. NUMBER LINE and pH  Remember the number line  Which is greater? 0 or 3 33  Which is greater? -7 or -4  -4  Which is less? -2 or -4  -4 34 5678 2 10-2-3-4-5-6-7 Increasing

29. AcidsBases

30. Homework pg. 9

31. Notes: pH Calculations pg. 10  The pH scale ranges from 0-14.  0 = strongly acidic  7 = neutral  14 = strongly basic  pH = -log [H + ]  What is the pH of a neutral solution?  Calculate using the Logarithmic function on the calculator (see at right)

32. Sample Problems  As long as you have a 1 x 10 to some power, the pH is the exponent. 1. What is the pH of the following concentrations? a. [H + ] = 1 x 10 -2 M b. [H + ] = 1 x 10 -9 M c. [H + ] = 1 x 10 -5 M pH = 2 acidic pH = 9 basic pH = 5 acidic

33. Sample Problems  If you do not have 1 to the power then you MUST use our formulas. 2. What is the pH of the following? a. [H + ] = 2x10 -2  pH = -log(2x10 -2 ) = 1.7 pH b. [H + ] = 6x10 -9  pH = -log(6x10 -9 ) = 8.2 pH c. [H + ] = 3x10 -5  pH = -log(3x10 -5 ) = 4.5 pH

34. Other Formulas and Problems  pH 14 = pH + pOH (See example 1 in Example Problems))  Equilibrium constant labeled as Kw  Kw is 1x10 -14  K w = [OH - ] x [H + ] = 1x10 -14

35. Other Formulas and Problems EX: What is the pH of a solution with a [OH - ] of 4.0 x 10 -11 M? o Use K w to find [H + ] then find pH using –log function. Step1: Step 2: Kw = [OH - ] x [H + ] = 1x10 -14 [H+] = 1x10 -14 /4x10 -11 = 2.5x10 -4 pH = -log [H+] pH= -log(2.5x10 -4 ) = 3.6

36. 1. If pH = 5, pOH = Kw pH 14 = pH + pOH 14 = 5 + pOH 14 – 5 = 9 pOH Acid because pH = 5

37. 2. What is the pH of a solution that has a hydrogen ion concentration of 1.0 x 10 -5 M? Is this solution acidic, basic or neutral? Given: [H + ] Solving for: pH pH = - log [H + ] pH = - log(1.0 x 10 -5 M) pH = 5 pH < 7 ACIDIC

38. 3. What is the hydrogen ion concentration of a solution with a pH of 11? Which has a greater concentration: H + or OH - ? [H + ] = 1 x 10 -11 M more OH -, So basic

39. 4. What is the pH of a solution that has a hydrogen ion concentration of 1.2 x 10 -8 M? Is this solution acidic, basic or neutral? Given: [H + ]Solving for: pH pH = - log [H + ] pH = - log(1.2 x 10 -8 M) pH = 7.92 pH > 7 BASIC

40. 5. Assuming Kw = 1x10 -14, calculate the molarity of OH- in solutions at 25ºC when the H+ concentration is 0.2M At 25ºC, Kw = [OH-] [H+] = 1x10 -14 1x10 -14 = [OH-] 0.2M = 1x10 -14 /.2 [OH-] = 5x10 -14 M

41. HOMEWORK: pg 12

42. Neutralization Notes pg. 15  Acid-Base reactions will produce salt water when completely neutralized.  Salts are compounds consisting of a(n) anion from an acid and a(n) cation from a base.  In general, reactions in which an acid and a base react in an aqueous solution to produce a salt and water is called Neutralization Reactions.

43. Neutralization Reactions  Neutralization occurs when an Acid + Base ↔ water + salt  Salt: Anion from acid and the cation from the base join together to form a salt. Where do we see this process? Antacids Farmers controlling the pH of soil Formation of caves

44.  A strong acid + a strong base = neutral solution Examples: HCl + NaOH ↔ H 2 O + NaCl HCl + KOH ↔ H 2 O + KCl

45. Practice: Don’t forget to balance them after you write them.  HCl + LiOH →  HNO 3 + CsOH →  HBr + KOH → HOH + LiCl CsNO 3 + H 2 O H 2 O + KBr

46. Titrations  Titration: The process of adding a known amount of solution of known concentration to determine the concentration of the other solution.  If you don’t know the concentration of one solution, you can figure it out by performing a neutralization reaction, or titration, with a standard solution.  A standard solution is one of known concentration.

47. Performing Titrations  Steps in a neutralization reaction:  1. A measured volume of an acid solution of unknown concentration is added to a flask.  2. Several drops of indicator are added to the solution.  3. Measured volumes of a base with a known concentration are mixed into the acid until it barely changes color.

48. Performing Titrations, cont.  End Point: The point at which the indicator changes color.  Once you have reached the end point, you can perform calculations to find the unknown solution. Let’s show a video! http://app.discoveryeducation.com/search? Ntt=titration## http://www.youtube.com/watch?v=RHTxIY DJ730

49. Performing Titrations, cont.  Example: A 25 mL solution of H 2 SO 4 is completely neutralized by 18 mL of 1.0 M NaOH. What is the concentration of H 2 SO 4 solution?  Step 1: Balanced equation  ____H 2 SO 4 + ____NaOH ↔ ____Na 2 SO 4 + ____H 2 O  Step 2: Use formula to solve for unknown.  M a V a = M b V b n a n b  n a = Number of moles of your Acid (coefficient)  n b = Number of moles of your Base (coefficient)  M = Molarity of acid or base  V = Volume of acid or base (in Liters)

50. M a V a = M b V b n a n b Ma ( 25 mL) = (1.0 M)( 18 mL) 1 mol 2 mol Molarity = 0.36 M

51. 1. How many moles of HCl are needed to neutralize 6 mols of KOH?  1 st ask, what is the mol ratio and then set it up as a proportion. HCl + KOH  KCl + H 2 O This equation is balanced so 1 mole HCl = 1 mole KOH So 6 mols KOH will neutralize 6 moles HCl

52. 2. H 2 SO 4 + 2NaOH ↔ Na 2 SO 4 + 2H 2 O a. One mole of sulfuric acid is needed to neutralize moles of NaOH. b. How many moles of NaOH are needed to neutralize 4 moles of H 2 SO 4 ? 2 Given that 1 H 2 SO 4 = 2NaOH So if you have 4 mols H 2 SO 4 you will need 8 moles NaOH

53. Homework: Page 17