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Acids and Bases – Unit 13

Acids and Bases – Unit 13. 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

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Acids and Bases – Unit 13

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  1. Acids and Bases – Unit 13

  2. Chemistry of Acids and Bases1. Watch video and complete worksheet2. Gallery walk to complete notes on pages 3-5 in packet3. Homework is on page 6 in packet • Standard Deviants Teaching Systems: Chemistry: Module 05: Acids and Bases • http://app.discoveryeducation.com/player/view/assetGuid/DBD191DB-A10E-43C2-8DDE-A73858F12FE2

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

  4. 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 hydrogengas • Single replacement reaction • Acid + metal → hydrogen gas + a“salt” • Double replacement reaction • Acid + Base → water + a“salt” Unit 13 – Acids and BasesNotes #1: Intro

  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!!!!!

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

  8. Examples of Naming Ternary Acids Sulfate is the poly, so sulfuric acid • H2SO4 • H2CO3 • H2NO2 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. Properties of Bases: • bitter • slippery (soap) • electrolytic • Both strong and weak • Will cause an indicator to change colors Unit 13 – Acids and Bases

  11. Naming Bases • The easiest are the bases, since most of these are _metalhydroxides, 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

  12. 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.

  13. 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)

  14. Brønsted – Lowry Acids and Bases • Examples: • NH3 + H2O↔ NH4+ + OH- • H2O donated the H+ - Acid • NH3 accepted the H+ - Base • HCl + H2O ↔ H3O+ + Cl- • HCl donated the H+ - Acid • H2O accepted the H+ - Base

  15. 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. • Sodium oxide + water → Sodium hydroxide solution Na2O(s) + H2O(l) → NaOH(aq) • The soluble oxides of non-metals are acidic in nature. Examples include, carbon dioxide, sulfur dioxide and nitrogen dioxide. • Sulfur dioxide + water → Sulfurous acid SO2(g) + H2O(l) → H2SO3(aq) • 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.

  16. 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: NH3, Acetic Acid (vinegar) • Tooth decay is caused by the weak acid – lactic acid: C3H6O3

  17. Homework: pg 6

  18. Notes: pH and pOHpg 7

  19. pH Scale

  20. 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.

  21. 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

  22. Self-ionization of water • Self-ionization of water: • Reaction in which 2 water molecules produce ions • H2O + H2O → OH- + H3O+ • Also written as: H2O ↔ H+ + OH- • The H3O+ and H+ represent hydrogen ions in solution.

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

  24. 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 10-14M • Acidic Solution: • The [H+] will be greater than the [OH-]. • Therefore, the [H+] is greater than 1 x 10-7M. • 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 10-7M. • A.k.a. alkaline solutions

  25. NUMBER LINE and pH • Remember the number line • Which is greater? 0 or 3 • 3 • Which is greater? -7 or -4 • -4 • Which is less? -2 or -4 • -4 Increasing 2 3 4 -7 -6 -5 -4 -3 -2 -1 0 1 5 6 7 8

  26. Acids Bases

  27. Homework pg. 9

  28. 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)

  29. 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 10-2M b. [H+] = 1 x 10-9M c. [H+] = 1 x 10-5M pH = 2 acidic pH = 9 basic pH = 5 acidic

  30. 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

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

  32. Other Formulas and Problems EX: What is the pH of a solution with a [OH-] of 4.0 x 10-11M? • Use Kw 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

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

  34. 2. What is the pH of a solution that has a hydrogen ion concentration of 1.0 x 10-5M? 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

  35. 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

  36. 4. What is the pH of a solution that has a hydrogen ion concentration of 1.2 x 10-8M? 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

  37. 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

  38. HOMEWORK: pg 12

  39. Neutralization Notes pg. 15 • Acid-Base reactions will produce salt water when completely neutralized. • Salts are compounds consisting of a(n) anionfrom an acid and a(n) cationfrom 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.

  40. 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

  41. A strong acid + a strong base = neutral solutionExamples:HCl+ NaOH ↔ H2O + NaClHCl+ KOH ↔ H2O + KCl

  42. Practice: Don’t forget to balance them after you write them. • HCl + LiOH → • HNO3 + CsOH → • HBr + KOH → HOH + LiCl CsNO3 + H2O H2O + KBr

  43. 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 ofknown concentration.

  44. Performing Titrations • Steps in a neutralization reaction: • 1. A measured volume of an acidsolution 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.

  45. 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=RHTxIYDJ730

  46. Performing Titrations, cont. • Example: A 25 mL solution of H2SO4 is completely neutralized by 18 mL of 1.0 M NaOH. What is the concentration of H2SO4 solution? • Step 1: Balanced equation • ____H2SO4 + ____NaOH ↔ ____Na2SO4 + ____H2O • Step 2: Use formula to solve for unknown. • MaVa = MbVb nanb • na= Number of moles of your Acid (coefficient) • nb = Number of moles of your Base (coefficient) • M = Molarity of acid or base • V = Volume of acid or base (in Liters)

  47. MaVa = MbVb nanb Ma ( 25 mL) = (1.0 M)( 18 mL) 1 mol 2 mol Molarity = 0.36 M

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