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Acids and Bases

Acids and Bases. Chapter 19. Describing Acids and Bases. Mini-Project. Work with a partner. Organize the following formulas into two groups with four formulas in each group: HNO 3 , NaOH, H 2 SO 4 , H 2 CO 3 , Ca(OH) 2 , KOH, H 8 PO 4 , Mg(OH) 2. One way to organize them into groups is:

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Acids and Bases

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  1. Acids and Bases Chapter 19

  2. Describing Acids and Bases

  3. Mini-Project Work with a partner. Organize the following formulas into two groups with four formulas in each group: HNO3, NaOH, H2SO4, H2CO3, Ca(OH) 2, KOH, H8PO4, Mg(OH) 2 One way to organize them into groups is: Group One Group Two HNO3 NaOH H2SO4 Ca(OH) 2 H2CO3 KOH H8PO4 Mg(OH) 2 Group One formulas represent acids. Group Two formulas represent bases.

  4. ACIDS • Electrolytes in solution • Taste sour (lemon, vinegar) • React with metal (corrosion) • React with carbonates (makes bubbles of CO2 • Turns blue litmus RED • In Water forms Hydrogen ION Water HCl H+ + Cl-

  5. BASES • Electrolyte in solution • Taste Bitter (soap, tonic water) • Feel Slippery (soap) • Turns Red Litmus Blue • React with Acids to make water Water NaOH Na+ + OH-

  6. Why are Some Solutions Acid & Others Base? • Acid solutions contain more H+ ions than OH- ions. • Base solutions contain more OH- ions than H+ ions. • Water is the standard for Acid/Base and is defined as NEUTRAL • Water has equal amounts of H+ and OH- ions

  7. Arrhenius Model of Acids/Bases • Substance is an acid if it contains hydrogen and dissociation causes hydrogen ions to form in solution • Substance is a base if it contains a hydroxide and dissociates to produce hydroxide ions in solution

  8. Bronsted-Lowry Model • Acid is a proton (hydrogen ion) donor • Base is a proton (hydrogen ion) receptor • This is a broader definition than Arrhenius model because there are substances that cause donation or reception without having hydrogen in them.

  9. Example • When an acid dissolves in water, it donates an H+ ion to a water molecule forming H3O+. • The water molecule acts as a base and accepts the H+ ion HX + H2O ⇄ H3O+ + X- • Conjugate acid = species produced when a base accepts a hydrogen ion from an acid • Conjugate base = species produces when an acid donates a hydrogen ion to a base • Conjugate base pair = 2 substances related to each other by donating and accepting a single hydrogen ion

  10. Ionization + H H O H O Cl Cl H H H Historical views on acids • O (e.g.H2SO4) was originally thought to cause acidic properties. Later, H was implicated, but it was still not clear why CH4 was neutral. • Arrhenius made the revolutionary suggestion that some solutions contain ions & that acids produce H3O+ (hydronium) ions in solution. + + • The more recent Bronsted-Lowry concept is that acids are H+ (proton) donors and bases are proton acceptors

  11. + + + H H O H O Cl Cl H H H The Bronsted-Lowry concept • In this idea, the ionization of an acid by water is just one example of an acid-base reaction. conjugate acid conjugate base acid base conjugate acid-base pairs • Acids and bases are identified based on whether they donate or accept H+. • “Conjugate” acids and bases are found on the products side of the equation. A conjugate base is the same as the starting acid minus H+.

  12. Practice problems Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs: HC2H3O2(aq) + H2O(l)  C2H3O2–(aq) + H3O+(aq) conjugate base conjugate acid acid base conjugate acid-base pairs OH–(aq) + HCO3–(aq)  CO32–(aq) + H2O(l) base acid conjugate base conjugate acid conjugate acid-base pairs

  13. Answers: question 18 HF(aq) + SO32–(aq)  F–(aq) + HSO3–(aq) (a) conjugate base conjugate acid acid base conjugate acid-base pairs (b) CO32–(aq)+HC2H3O2(aq)C2H3O2–(aq)+HCO3–(aq) base acid conjugate base conjugate acid conjugate acid-base pairs (c) H3PO4(aq) + OCl–(aq)  H2PO4–(aq) + HOCl(aq) conjugate base conjugate acid acid base conjugate acid-base pairs

  14. More • Amphoteric = substances like water that can act like either an acid or a base

  15. Monoprotic and Polyprotic Acids • Monoprotic – acids based on formula that can donate only one hydrogen ions • CH3COOH + H2O ⇄ H3O+ + CH3COO- • Polyprotic – acids that can donate multiple hydrogen ions • H3PO4 + H2O ⇄ H3O+ + H2PO4+ • H2PO4+ + H2O ⇄ H3O+ + HPO4+2 • HPO4+2 + H2O ⇄ H3O+ + PO4+3 • Anhydride = oxides that can become acids or bases by adding elements contained in water

  16. Acid Rain • Acid rain comes from rain collecting gasses from the air to create acids: • Carbon Dioxide = carbonic acid • Sulfur oxides = sulfuric acid • Nitrogen oxides = nitric acid • Damages statues, buildings, kills forests, kills fish

  17. Acids and Bases in Solution Chapter 19.2

  18. Acid/Base Strength • In strong acids, almost all molecules ionize. • In weak acids, fewer molecules ionize.

  19. Conjugate Pairs Strength • If an acid is a strong acid, its conjugate pair base is a weak base • Why? • If HX is strong acid, it ionizes completely. • The conjugate base must be a weak base because it has a greater attraction to the H+ than HX • The reaction equilibrium lies far to the right of the equation.

  20. Conjugate Pair Strength • For a weak acid, the equation equilibrium lies to the right (reactant side) • Conjugate base (Y-) has a stronger attraction for the H+ ion than the base H2O • HY + H2O H3O+ + Y-

  21. Acid Ionization Constants • An “ionization constant” is the tendency of an item to make ions in solution. • Higher the constant, the higher the amount of ions. • Acid ionization constant is value of the equilibrium constant expression for a weak acid • Value Ka indicates whether reactants or products are favored at equilibrium • Weak acids have low Kavalues

  22. Acid Ionization Constants

  23. Base Ionization Constant • Same Basic Principle as Acid • Measures OH- concentrations

  24. pH Scale Chapter 19.3

  25. Ionization Constant for Water • The ionization constant for water is: • 1.0 x 10-14 • [Ka][Kb] • = [1.0x10-7] [1.0 x 10-7] • Experiments show that the product of [H+] and [OH-] always equals 1.0 x 10-14 at 298°C • pH scale is a way of showing this relationship of ionization constants

  26. The pH Scale • pH stands for ‘per Hydrion’ • Low pH is Acid • High pH is base • Water is neutral (7.0)

  27. pH • There are many ways to consider acids and bases. One of these is pH. • [H+] is critical in many chemical reactions. • A quick method of denoting [H+] is via pH. • By definition pH = –log [H+], [H+] = 10-pH • The pH scale, similar to the Richter scale, describes a wide range of values • An earthquake of “6” is 10 as violent as a “5” • Thus, the pH scale condenses possible values of [H+] to a 14 point scale (fig. 2, p370) • Also, it is easier to say pH=7 vs. [H+]=1x10–7

  28. pH • pH = -log [H+] • [H+] = 10-pH • pOH = -log [OH-] • pH + pOH = 14

  29. Calculations with pH Q: What is the pH if [H+]= 6.3 x 10–5? pH = –log [H+] ‘(-)’, ‘log’, ‘6.3’, ’10x’, ‘(-)’, ‘5’, ‘)”, ‘)”, ‘ENTER’) Ans: 4.2 Q: What is the [H+] if pH = 7.4? [H+] = 10–pH mol/L (’10x’, ‘(-)’, ‘7.4’, “)” ‘ENTER‘) 3.98x10–8 M

  30. Calculating pH from Strong Acid Solutions • Strong acids are 100% ionized • For monoprotic acids, concentration of the acid IS the concentration of the H+ ion • Use Acid concentration as substitute for H+ ion concentration. • Use Base concentration as substitute for OH- concentration

  31. Calculating pH from Strong Acid Solutions • Example: What is the pH of a 0.1M solution of HCl? • 0.1 M HCl = 1 x 10-1 M • Calculate pH = 1 • Example: What is pH of solution that is 7.5 x 10-4 M Ca(OH)2? • (7..5 x 10-4) x 2 = 1.5 x 10-3M • There are 2 OH- ions per molecule • Calculate pOH = -log[OH-] • = 2.8 • pH = 14-2.8 = 11.2

  32. Calculating Molarity from pH • Example: what is the molarity of an acid solution with a pH of 2.37? • [H+] = 10-pH • [H+] = 10-2.37 = 4.27 x 10-3 M

  33. Neutralization Chapter 15.4

  34. Acid-Base Reactions • Neutralization reaction is a reaction between an acid and a base • Makes Water + Salt • Solution becomes Neutral (not acid or base) HCl + NaOH H2O + Na+ + Cl-

  35. Acid-Base Reactions • Mg(OH)2 + 2 HCl → MgCl2 + 2H2O • Note: • Cation from base (Mg) is combined with anion from acid (Cl) • The salt is MgCl2 • The H+ and OH- always combine to form water

  36. Acid-Base Titration • Acid/Base Titration is the stoichiometry of acid/base reactions. • Titration is a method for determining the concentration of a solution by using another solution of known concentration • Uses an INDICATOR to show when the acid/base reaction is complete (neutral) • Indicator is a chemical that changes color as determined by acid or base conditions • There are many indicators with different pH points.

  37. Acid/Base Titration Curve

  38. pH Indicators

  39. Calculating Molarity from Titration • Write the balanced equation • Calculate the number of moles used in the ‘known’ solution • Use the mole ratio from the balanced equation to calculate moles of reactant in the ‘unknown’ solution • Calculate the molarity of the ‘unknown’ solution based on moles used and liters used.

  40. Salt Hydrolysis • When you put salts in water, the resulting solution can be either acid, base, or neutral • Salts will dissolve to form ions • The anions will accept hydrogens from water • The cations will accept hydroxides from water • Which way it goes depends upon the strength of the conjugate acids/bases • If conjugate acid is strong, it will be acid • If conjugate base is strong, it will be base • If both are strong, it will be neutral

  41. What are Buffers? • Buffers are solutions that resist changes in pH when limited amounts of acid or base are added. • Buffer is a weak acid and its conjugate base or a weak base and it’s conjugate acid • Buffer can accept or donate Hydrogen ions and shift its equilibrium point left or right • Buffers have limits, but the work for a while • Heavily used in human body, especially blood

  42. Acid Names 1. Binary or hydrohalic acids – HF, HCl, HBr, HI, etc. “hydro____ic acid” are usually strong acids • If name ends in ‘-ide’ • Acid name will be “hydro ____ic acid” • HF and H2S are weak hydrohalic acid. Although the H-F bond is very polar, the bond is so strong (due to the small F atom) that the acid does not completely ionize.

  43. Acid Naming 2. Oxyacids – contain a polyatomic ion • Most common form (MCF) “ic” ending– strong acids (contain 2 oxygen per hydrogen) • If chemical name ends in “-ate” • Acid name will be “___IC Acid” • HNO3 – nitric from nitrate • H3PO4 - phosphoric from phosphate • H2SO4 - sulfuric from sulfate • HClO3 - chloric from chlorate

  44. Acids with 1 less oxygen than the MCF “ous” ending- weaker acids • Chemical name ends in “-ite” • Acid name is “___OUS Acid” HNO2 – nitrous from nitrite H3PO3 - phosphorous from phosphite H2SO3 - sulfurous from sulfite HClO2 - chlorous from chlorite c. Acids with 2 less oxygen than the MCF “hypo___ous” – very weak acids HNO - hyponitrous H3PO2 - hypophosphorus HClO - hypochorous

  45. d. Acids with 1 more oxygen than the MCF “per______ic” – very strong acids HClO4 – perchloric acid HNO4 - pernitric acid • Organic acids – have carboxyl group -COOH - usually weak acids • Acid names are based on the base organic name or common name HC2H3O2 - acetic acid C7H5COOH - benzoic acid

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