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Chapter 18: Properties of Acids & Bases

Chapter 18: Properties of Acids & Bases. pH < 7 ----------- Acidic solution pH = 7 ----------- Neutral solution pH >7 ------------ Basic solution. What Do You Think?. The pH of milk is 6.4. Based on this information, which of the following statements best describes milk?

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Chapter 18: Properties of Acids & Bases

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  1. Chapter 18: Properties of Acids & Bases

  2. pH < 7 ----------- Acidic solution pH = 7 ----------- Neutral solution pH >7 ------------ Basic solution

  3. What Do You Think? • The pH of milk is 6.4. Based on this information, which of the following statements best describes milk? A. It is very basic B. It is very acidic C. It is slightly basic D. It is slightly acidic

  4. Properties of Acids • Acids are electrolytes • Acids produce hydrogen ions (H+) when dissolved in water HCl (aq)  H+ (aq) + Cl- (aq) • Aqueous solutions of acids have a sour taste • Acids change the color of acid-base indicators Example: Litmus paper turns red in presence of acid

  5. Properties of Acids (cont.) • Some acids react with metals to release hydrogen gas (H2) Zn(s) + 2HCl(aq)  ZnCl2(s) + H2(g) • Acids react with bases to produce salts and water (neutralization) 7. Acids have pH below 7

  6. Properties of Bases • Bases are electrolytes • Bases produce hydroxide ions (OH-)when dissolved in water NaOH (aq)  Na+ (aq) + OH- (aq) • Bitter taste and slippery 4. Bases change the color of acid-base indicators. Example: Litmus paper turns blue in presence of base 5. Bases react with acids to produce salts and water (neutralization) • Bases have pH above 7

  7. Arrhenius Theory of Acids • Arrhenius theorized that acids and bases must produce ions in solution • Arrhenius Acid: Chemical compound that increases the number of hydrogen ions, H+, in aqueous solutions • Ionize in water to produce H+

  8. Arrhenius Theory of Bases • Arrhenius Base: Substance that increases the concentration of hydroxide ions, OH-, in aqueous solution • Most bases are ionic compounds containing metal cations and the hydroxide anion, OH- • Ionize in water to produce OH-

  9. Naming Acids

  10. Naming Acids: Binary Acids • A binary acid is an acid that contains only two different elements: hydrogen and one other element • Examples: HF, HCl, H2S • Naming Binary Acids • The name of a binary acid begins with the prefix hydro- • The root of the name of the second element follows this prefix • The name then ends with the suffix -ic

  11. Naming Acids: Binary Acids (cont.)

  12. Naming Acids: Oxyacids • An oxyacidis an acid that contains hydrogen, oxygen, and a third element, usually a nonmetal. • Written as one or more hydrogen atoms followed by an anion • Examples: HNO3, H2SO4, H3PO4 • The name of the oxyacid is based on the name of the anion contained in the acid • If the anion of an oxyacid ends with the suffix “–ate,” replace with “–ic acid” • HNO3---- anion is nitrate ----- acid is called nitric acid • HC2H3O2---- anion is acetate ---- acid is called acetic acid

  13. Naming Acids: Oxyacids (cont.) • If the anion of an oxyacid ends with the suffix “–ite,” replace with “–ous acid” • H2SO3 ----- anion is sulfite ----- acid is called sulfurous acid • Oxyacids do not begin with the prefix hydro-

  14. Strong vs. Weak Acids & Bases

  15. Strong Acids Example: Hydrochloric Acid HCl (aq) + H2O  H3O+ (aq) + Cl- (aq) • Strong Acid: An acid that completely ionizes in aqueous solution and is a strong electrolyte • HCl - hydrochloric acid • HNO3 - nitric acid • H2SO4 - sulfuric acid • HBr - hydrobromic acid • HI - hydroiodic acid • HClO4 - perchloric acid • HClO3 - chloric acid

  16. Weak Acids Example: Acetic acid HC2H3O2 (aq)  H+ (aq) + C2H3O2- (aq) • Mixture of moleculesandions at equilibrium (note the double arrows, indicating that the reaction does not go to completion) • Weak Acid:An acid that releases few hydrogen ions in aqueous solution • Not complete ionization • Weak electrolyte

  17. Strong Bases Example: Sodium hydroxide NaOH (aq)  Na+ (aq) + OH- (aq) • Strong Base: A base that completely ionizes in an aqueous solution and is a strong electrolyte • NaOH - sodium hydroxide • KOH - potassium hydroxide • RbOH - rubidium hydroxide • CsOH - cesium hydroxide • Ca(OH)2 - calcium hydroxide • Sr(OH)2 - strontium hydroxide • Ba(OH)2 - barium hydroxide

  18. Weak Bases Examples: Slightly soluble hydroxides Cu(OH)2   Cu+2 (aq) + 2 OH- (aq) • Weak Base: A base that produces very few hydroxide ions when added to water • Either not very soluble or poor electrolyte

  19. Concentrations of Acids & Bases

  20. Self-Ionization of Water [H+] = Concentration of hydrogen ions [OH-] = Concentration of hydroxide ions • As water molecules collide, sometimes a water molecule will break up into H+ and OH- • An equilibrium exists between the water molecules and ions • This process is called the self-ionization of water H2O H+ +OH-

  21. Self-Ionization of Water • At room temperature, [H+] = 1.0 ×10−7 M and [OH−] = 1.0 × 10−7 M • The ionization constant of water (KW) is expressed by the following equation: Kw = [H+] [OH−] Kw = [1.0 ×10−7 ] [1.0 ×10−7 ] = 1.0 ×10−14 • [H+] and [OH−] are interdependent • Altering the concentration of one ion affects the other • Increasing the concentration of one ion decreases the concentration of the other ion

  22. Neutral, Acidic, and Basic Solutions • Solutions are classified as acidic, basic, or neutral based on the [H+] and [OH−] in the solution • Solutions in which [H+] = [OH−] are neutral • [H+] =[OH−] =1.0 × 10−7 M • Solutions in which the [H+] > [OH−] are acidic • [H+] > 1.0 × 10−7 M • Solutions in which the [OH−] > [H+] are basic • [OH−] > 1.0 × 10−7 M

  23. Strength of Acidic or Basic Solutions • The greater the [H+] in an aqueous solution, the more acidic the solution • The lower the [H+] in an aqueous solution, the more basicthe solution • Using the equation for self-ionization of water, solve the following problem: At 298 K, [H+] in a cup of coffee is 1.0 x 10-5 M. What is [OH-] in this coffee? Is the coffee acidic, basic, or neutral?

  24. Strength of Acidic or Basic Solutions Practice Problems For each solution, calculate either [H+] or [OH-]. State whether the solution is acidic, basic, or neutral. 1.) [H+] = 1.0 x 10-13 M 2.) [OH-] = 1.0 x 10-7 M 3.) [H+] = 4.0 x 10-5 M

  25. pH of a Solution • Since [H+] values tend to be very small, another more convenient way to describe acidity is in terms of pH • pH measures the concentration of hydrogen ions • pOH measures the concentration of hydroxide ions • pH + pOH = 14

  26. pH and the Strength of Acids and Bases   • Which solution is more acidic: [H+] = 1.0 x 10-6 or [H+] = 1.0 x 10-5? • What is the pH of each solution? • The strength of an acid increases as the pH decreases • Which solution is more basic: [H+] = 1.0 x 10-8 or [H+] = 1.0 x 10-9? • What is the pH of each solution? • The strength of a base increases as the pH increases

  27. Neutralization Reactions

  28. Neutralization Reactions • Neutralization reactions are double replacement reactions that involve the reaction of an acid with a base HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) • Products of a neutralization reaction are a saltand water • Salts are electrolytes that disassociate into ions other than H+ and OH-

  29. Neutralization Reactions • In aqueous solutions, neutralization is the reaction of hydrogen ions and hydroxide ions to form water molecules • The cation of a salt formed in a neutralization reaction comes from the base and the anion comes from the acid • Complete the following neutralization reaction and identify the salt: Ba(OH)2 + HNO3

  30. Neutralization Reactions Practice Problems Complete each neutralization reaction by determining the products. 1.) Mg(OH)2 (aq) + 2HCl (aq)  2.) KOH (aq) + HBr (aq)  3.) KOH (aq) + H2CO3 (aq) 

  31. Titrations

  32. Titrations Titration: Method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration Standard Solution: A solution whose concentration is known Equivalence Point: The point at which moles of acid equals moles of base

  33. Titrations • An acid-base indicator is used to show when the equivalence point has been reached - Common indicator is phenolphthalein, which turns pink at the equivalence point • The point in a titration at which an indicator changes color is called the end point of the indicator

  34. Titrations MAVA = MBVB Remember, a titration is done to determine the concentration of a solution by reacting a known volume of that solution with a solution of known concentration MA= concentration of the acid VA= volume of the acid MB= concentration of the base VB = volume of the base

  35. Titrations: Example Problem #1 In a titration, 27.4 mL of 0.154 M NaOH is added to a 20.0 mL sample of HCl solution of unknown concentration until the equivalence point is reached. What is the molarity of the HCl solution? HCl + NaOH  NaCl + H2O MAVA = MBVB MA(0.02 L)= (0.154M)(0.0274 L) MA = .21M

  36. Titrations: Example Problem #2 18.3 mL of a standard solution of 0.1 M NaOH was used to neutralize 25.0 mL of a solution of methanoic acid (HCOOH). What is the molarity of the methanoic acid solution? MAVA = MBVB

  37. Titration Curve

  38. Titrations: Particulate Level Before addition of base ----- all H+ Aftersmall amount of base added -----H+ and H2O At the equivalence point ----- H2O Past the equivalence point ----- OH-

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