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Chapter 19: Acids and Bases Sections 19.1 to 19.4

Chapter 19: Acids and Bases Sections 19.1 to 19.4. Acid Properties. Sour taste (citrus fruits) Conduct electric current Change the color of indicators React with bases to produce salt and water: HCl + NaOH  H 2 O + NaCl Some react with metals to release H 2 gas:

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Chapter 19: Acids and Bases Sections 19.1 to 19.4

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  1. Chapter 19:Acids and BasesSections 19.1 to 19.4

  2. Acid Properties • Sour taste (citrus fruits) • Conduct electric current • Change the color of indicators • React with bases to produce salt and water: HCl + NaOH  H2O + NaCl • Some react with metals to release H2 gas: Mg + HCl  MgCl2 + H2

  3. Naming Acids Review • Ternary Acids (oxyacids) • HClO3 • Chloric Acid • HClO2 • Chlorous Acid • Binary Acids: • Hydroiodic Acid • HI • HF • Hydrofluoric Acid

  4. CATEGORIES OF ACIDS • Monoprotic • ONE ionizable H+ Example: HNO3 • Diprotic • TWO ionizable H+ Example: H2SO4 • Triprotic • THREE ionizable H+ Example: H3PO4

  5. Base Properties • Bitter taste (coffee) • Feel slippery (soap) • Change the color of indicators • Caustic- attack the skin, cause severe burns • Conduct electric current

  6. Arrhenius Acids and Bases • Arrhenius Acid: A compound that produces H+ in solution. Ex:HCl (g) -------- H+ (aq) + Cl- (aq) • Arrhenius Base: A compound that produces OH- in solution. Ex: NaOH (s) --------- Na+ (aq) + OH- (aq) H2O H2O

  7. Acid/Base Strength • Strong Acid: Ionizes completely in aq. soln. HCl H2SO4 HBr HNO3 HI HClO4 HClO3 Strong Bases: Group 1 and 2 hydroxides

  8. Acid-Base Theories • Bronsted-Lowry: expands Arrhenius definition of acids and bases. • Bronsted-Lowry Acid: proton donor • Bronsted-Lowry Base: proton acceptor ex:HCl + NH3 NH4+ + Cl- Monoprotic B-L Acid B-L Base ex2: H3PO4 + H2O  H3O+ + H2PO4- Which is the B-L Acid? B-L Base?

  9. Lewis Acids and Bases • Based on bonding and structure and include substances that may not include Hydrogen. • Lewis Acid: electron pair acceptor • Lewis Base: electron pair donor ex: BF3 (aq) + F- (aq) BF4- (aq) Draw the dot structure for these substances and classify as a Lewis Acid or Base. Lewis Base Lewis Acid

  10. Conjugate Acids and Bases(Based on Bronsted-Lowry Classification) • Conjugate Base: The substance that remains after an B-L acid has given up a proton (H+). • Conjugate Acid: The substance formed when a B-L base has gained a proton. ex: HCl (aq) + H2O (l) Cl- (aq) + H3O+ (aq) Conjugate Base Conjugate Acid Acid Base • Table 19.6: The stronger the acid, the weaker its conjugate base. Equilibrium favors weak acid/base formation.

  11. Acid Reactions • Neutralization: HCl (aq) + NaOH(aq)  NaCl (aq) + H2O (l) • Acid Formation from Acid Anhydrides: SO3 (g) + H2O (l)  H2SO4 (aq) • Base Formation from Basic Anhydrides: Na2O + H2O(l)  2NaOH Acid Rain Acid Anhydride Basic Anhydride

  12. Aqueous Solutions and pH • Self Ionization of Water • Water also supplies H3O+ and OH- ions. H2O (l) + H2O (l) H3O+ (aq) + OH- (aq) Conductivity Experiments show the concentrations of ions at 25 °C: 1.0 x 10-7 M 1.0 x 10-7 M • Ionization Constant of Water, Kw Kw = [H3O+][OH-] = [1.0 x 10-7 M][1.0 x 10-7 M] Kw = 1.0 x 10-14 M2 Constant at a given temperature

  13. Neutral, Acidic, and Basic Solutions • Neutral: [H3O+] = [OH-] • Acids: [H3O+] > [OH-] • Bases: [H3O+] < [OH-] Determine the hydronium and hydroxide ion concentrations in a 1 x 10-5 M HCl solution.

  14. pH Scale A more convenient way to express acidity • pH = -log[H3O+] • pOH = -log[OH-] • pH + pOH = 14.0 • Find the pH and pOH of a 1x10-10 M solution of HBr.

  15. Find the pH and pOH of a 1x10-10 M solution of HBr. 1. HBr is an acid so we are looking at the [H+] concentration. [H+] = 1x10-10 M 2. pH = -log[1x10-10] pH = 10 3. pH + pOH = 14 10 + pOH = 14 pOH = 4

  16. Calculating pH pH = -log [H+] Ex: A solution has a H+ concentration of 1x10-5 pH = -log [1x10-5] pH = 5 and we have an acidic solution Ex: A solution has a pH=8 8 = -log [H+] [H+] = 1 x 10-8

  17. Indicators and Titration • Acid-Base Indicators: Compounds whose colors are sensitive to pH. • Titration: Method used to determine an unknown concentration of solution.

  18. Equivalence Point • The point at which the 2 solutions used in a titration are present in equal amounts. • End Point: The point in a titration during which an indicator changes color.

  19. NEUTRALIZATION OF ACIDS & BASES If we need to neutralize an acid or a base, we use the following formula: MaVaCb = MbVbCa Mx=Molarity (mol/L) of the acid or base Vx=Volume (L) of the acid or base Cx =Coefficient (balanced eq.) of the acid or base

  20. MaVaCb = MbVbCa Example: A 25mL solution of H2SO4 is neutralized by 18mL of a 1.0M NaOH using phenolphthalein as an indicator. What is the concentration (M) of the H2SO4? Step 1: Write the neutralization rxn and balance H2SO4 + 2NaOH Na2SO4 + 2H2O Step 2: Solve for the unknown. Ma = MbVbCa (1.0M)(0.018L)(1) VaCb (0.025L)(2) Ma = 0.36M

  21. MOLARITY vs NORMALITY • Molarity—moles of solute contained in 1 liter of solution (moles/Liter) • Normality—moles of reactive units for each liter of solution • IN ACID-BASE REACTIONS: • CaVa = CbVb • Ex #1—H3PO4 (N=H+ ions so N=3) • Ex #2---Mg(OH)2 (N=OH- ions so N=2)

  22. AMPHOTERIC SUBSTANCES • SUBSTANCES THAT CAN BE BOTH ACIDIC AND BASIC • Common amphoteric substances are the oxides of the elements and beryllium, aluminum, zinc, tin and lead. • TRENDS ON PERIODIC TABLE: • Across row: Metallic oxides → non-metallic oxides • Across row: Basic oxides → Amphoteric oxides → Acidic Oxides

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