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Acids and Bases: Arrhenius and Brønsted-Lowry Definitions

Explore the Arrhenius and Brønsted-Lowry definitions of acids and bases, their limitations, and the equilibrium in acid-base reactions. Learn about acid dissociation constants, writing and naming acids, and the pH scale.

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Acids and Bases: Arrhenius and Brønsted-Lowry Definitions

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  1. Chapter 14 Acids and Bases

  2. Arrhenius: 1st to Define Acids • Acids H+(aq) • Bases  OH-(aq) • Limitations: • Only aqueous solutions • Only one kind of base: OH- • NH3 ammonia could not be an Arrhenius base. 1883

  3. Brønsted-Lowry Definitions • Acid is proton (H+) donor, must have a removable (acidic) proton • Base is proton acceptor, must have a pair of nonbonding electrons • Acids and bases always come in pairs. • Example HCl(g) + H2O(l)⇌H3O+ + Cl- • Water as base becomes an Hydronium ion • Remember – when acids are added to water the process is Dissociation

  4. Bronsted Lowry • This is equilibrium • HA & A- are generic term for Bronsted Lowry acid & base • Competition for H+ between H2O and A- • In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. • If H2O is a stronger base it takes the H+ • p. 624 2nd paragraph • Equilibrium moves to right. Acid Base Conjugate acid Conjugate base

  5. Acid Dissociation Constant Ka • HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) • The equilibrium expression is: • Ka is the acid dissociation constant • H3O+ is often written H+ • Ignore the water in equation because it’s [l] remains constant (p. 625)

  6. Acid Dissociation Constant Ka HA(aq) ⇆ H+(aq) + A-(aq) • We can write the expression for any acid. • Strong acids dissociate completely. • Equilibrium far to right in strong acids • Conjugate base must be weak.

  7. Write the dissociation Reactionomitting water for the following: • Example Hydrochloric acid: HCl(aq)⇔ H+(aq) + Cl-(aq) • Acetic Acid HC2H3O2 • Ammonium ion • Hydrated aluminum (III) ion [Al(H2O)6]+3

  8. Various Ways to Describe Acid Strength Strength of acid is defined by the equilibrium position of its dissociated reaction. See figure 14.4 p. 626

  9. Acid and Base Strength • Strong acids are completely dissociated in water. • Their conjugate bases are quite weak. • Weak acids only dissociate partially in water. • Their conjugate bases are weak bases. • Substances with negligible acidity do not dissociate in water. • Their conjugate bases are exceedingly strong.

  10. Strong Acids Strong SolubleBases • Perchloric Acid >3.2 x 109 • Hydroiodic Acid 3.2 x 109 • Hydrobromic Acid 1.0 x 109 • Hydrochloric Acid 1.3 x 106 • Sulfuric Acid 1.0 x 103 • Chloric Acid 1.0 x 103 • Nitric Acid 2.4 x 101 • Lithium Hydroxide • Sodium Hydroxide • Potassium Hydroxide • Rubidium Hydroxide • Cesium Hydroxide • Calcium Hydroxide • Strontium Hydroxide • Barium Hydroxide KNOW THESE

  11. Rules for Writing and Naming Acidsreview p.67 • All inorganic acids starting with H • Binary Acids • Two elements or H plus non-oxy polyatomic anion • Name comes from root of anion • Begins w/ “hydro-” & ends “-ic” • Ex. HCl: hydrochloric acid • Ex. HCN: hydrocyanic acid

  12. Rules for Writing and Naming Acids • Oxyacids • Anion contains Oxygen • Named from root name of anion • Use suffix “ic” or “-ous” • –ate Anion becomes -ic acid • –ite Anion becomes -ous acid • Ex. Nitrate⇒ HNO3 is Nitric Acid • Ex. Nitrite⇒ HNO2 is Nitrous Acid

  13. Types of Acids • Polyprotic Acids- more than 1 acidic hydrogen (diprotic, triprotic) • Oxyacids - Proton attached to oxygen of ion • Organic acids contain the Carboxyl group -COOH with the H attached to O See Table 22.4 p 1011 • Generally very weak • Table 14.2 Common Monoprotic Acids – note strong acids not listed

  14. Amphoteric: Behave As Acid & Base • Example: Autoionization of Water (most common) 2H2O(l)⇆ H3O+(aq) + OH-(aq) Kw= [H3O+][OH-]=[H+][OH-] KW is dissociation constant for water In pure water, @ 25ºC KW = 1.0 x10-14 • In EVERY aqueous solution @ 25oC, no matter what it contains, the product of [H+][OH-] is always 1.0 x10-14

  15. Meaning of Kw • In any aqueous solution, we know KW = 1.0X10-14 • There are 3 possible situations: • Neutral solution [H+] = [OH-]= 1.0 x10-7 • Acidic solution [H+] > [OH-] • Basic solution [H+] < [OH-] • In each case: KW = 1.0X10-14

  16. Calculate [H+] or [OH-] Calculate [H+] for the following solution at 25oC and state whether the solution is neutral, acidic or basic: [OH-] = 1.0X10-5M Answer: KW=[H+] [OH-]=1.0X10-14 , since [OH-] = 1.0X10-5M, solving for [H+] gives: Since: 1.0X10-5 [OH-] > 1.0X10-9 [H+] the solution is basic Carnegie Time #38 p.674

  17. The pH Scale • pH= -log[H+] • Use pH scale because [H+] is usually very small Sig Figs Rule: Number of decimal places in log = number of sig figs in original number • [H+] = 1.0 x 10-8 pH= 8.00 (2 decimals in pH) • Similar for other quantities: pOH= -log[OH-] pK = -log K 2 sig figs

  18. These are the pH values of several common substances.

  19. Lots of Relationships • pH = -log[H+] • pOH = -log[OH-] • pK = -log K • KW = [H+][OH-] • log KW = log [H+]+ log[OH-] • -log KW = -log[H+] - log[OH-] • pKW = pH + pOH • Since KW = 1.0 x10-14 • thenpKW = -log (1.0 x10-14 )=14.00 • pH + pOH = 14.00 Given any one of these values: [H+], [OH-] ,pH, and pOH we can find the other three.

  20. [H+] 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 pH 0 1 3 5 7 9 11 13 14 14 13 11 9 7 5 3 1 0 pOH 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [OH-] Acidic Neutral Basic

  21. Calculating pH, pOH, [H+], & [OH-] The pH of a sample of blood was measured to be 7.41 at 25oC. Calculate the pOH, [H+], and [OH-] for the sample. Since pH + pOH = 14.00 pOH = 14.00 – pH = 14.00 – 7.41 = 6.59 Since pH = -log[H+] 7.41 = -log[H+] take the antilog of -7.41=3.9x10-8 Similarly for pOH = -log[OH-] 6.59 = -log[OH-] take the antilog of -6.59=2.6x10-7 Do #44 p. 674

  22. Calculating pH of Solutions • Alwayswrite down the major species in solution, this is first step in all problems • Remember these are at equilibrium. • Remember the chemistry. • Don’t just memorize

  23. Calculate pH of Strong Acid Solutions Common Strong Acids: HCl, HNO3, H2SO4, HClO4 • When working with a Strong Acid (SA) solution, focus on solution species • Determine which species are significant and which can be ignored. MAJOR SPECIES = COMPONENTS PRESENT IN RELATIVELY LARGE AMOUNTS. Ex. 1.0M HCl Calculate pH

  24. Calculate pH of Strong Acid Solutions No HCl in solution because all SA completely dissociate Focus on Major Species: H+ Cl- H2O Because want pH, focus on H+ Consider H+ from H2O, the dissociation of HCl drives H+ contribution from water to left, so H+ from H2O is negligible Therefore [H+] is 1.0 and pH = -log[1.0] = 0

  25. Calc. pH of 0.10 M HNO3 • Solution: • HNO3 strong acid, Major species: H+, NO3-, H2O • [HNO3] virtually 0 because SA • [OH-] very small because H2O(l) [H+](aq) + [OH-](aq) is driven to left by H+ from acid (Le Chatelier’s Principle) • Source of H+ are: • HNO3 • H2O, negligible from H2O (see #3 above) • Leaves HNO3 as source of H+ so: • [H+]=0.10 M and pH = -log(0.10) = 1.00 CHT #48 p.674

  26. Solving Weak Acid Equilibrium Problems • List major species • Choose the species that can produce H+ & write balanced equation for reactions producing H+ • Ka vs. Kw decide which equilibrium will dominate in producing H+ • Write Equilibrium Expression for dominate equilibrium. • BuildICETable: List the Initial conc. of species participating in dominate equilibrium • Define change needed to achieve equilibrium in terms of x • Initial conc. + Change = Equilibrium Concentration • Sub. Equil. Conc. into equilibrium expression #4 • Solve for x using Approximation Rule • Calculate [H+] & pH

  27. Example • Study Sample Exercise 14.8 p. 638 • Pay close attention to comments between equations and calculations • Step 5-7 is nICE • Step 9-10 is the Approximation Rule (Ch. 13 p.604) • You will use these steps in ALL problems in next couple of chapters, so learn them well Do # 53 p.675

  28. Finding The pH Of A Weak Acid Solution in Cases Where “x is small Approximation” Does Not Work Find the pH of a 0.100 M HClO2 solution See Solution Handout

  29. The initial concentration and Ka’s of several weak acid solutions are listed below. For which of these is the “x is small approximation” least likely to work in finding the pH of the solution? • Initial [HA] = 0.100M Ka = 1.0 x 10-5 • Initial [HA] = 1.00M Ka = 1.0 x 10-6 • Initial [HA] = 0.0100M Ka = 1.0 x 10-3 • Initial [HA] = 1.0M Ka = 1.5 x 10-3 The validity of the x is small approximation depends on both the value of the equilibrium constant and the initial concentration- the closer that these are to one another, the less likely the approximation will be valid

  30. pH of a Mixture of Weak Acids • Calc. pH of a solution that contains M HCN (Ka=6.2x10-10) and 5.00 M HNO2 (Ka=4.0x10-4). Also calc. the concentration of the (CN)- at equilibrium. • Major Species: • HCN, HNO2, andH2O • All 3 produce H+ • HCN Ka=6.2x10-10 • HNO2Ka=4.0x10-4 • H2O Kw=1.0x10-14 • Pick the strongest of the weak acids

  31. pH of a Mixture of Weak Acids • HNO2 is much stronger than the other 2 species. Thus HNO2 is assumed to be the dominant producer of H+ • Set up ICETable and write the Equilibrium Expression • Plug Equilibrium Concentrations into Expression and using the Approximation Rule, solve for x • x is the [H+] and you can use it to find [CN-]

  32. Percent Dissociation Weak acid: percent dissociation increases as acid becomes more dilute. Calculate the percent dissociation of acetic acid (Ka=1.8x10-5) for a 1.00M solution. Solution: • List major species: • HC2H3O2 H2O • Choose species that produces most H+ : HC2H3O2& write balanced equation: HC2H3O2(aq) H+(aq) + C2H3O2-(aq)

  33. Percent Dissociation • Write Equilibrium Expression • Create ICE chart & fill in • Plug equilibrium concentrations into equilibrium expression & solve for X (use “x is small approximation rule”) • X = 4.2 X 10-3 M

  34. Percent Dissociation • Calculate percent dissociation: General Trend: For weak acid solutions Fig 14.10 • [acid] and [H+] decrease as % dissociation increases • % dissociation increase as acid is diluted Calculating Ka from percent dissociation is illustrated in Sample Exercise 14.11 on p.643 Study this sample & CHT #65 p.675

  35. Bases • Remember the OH-list of strong bases. • Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. • The hydroxides of alkaline earths (Ex. Ca(OH)2) are strong dibasic (2 OH-) bases, but they don’t dissolve well in water. Not very soluble. • Many metal hydroxides are used as antacids because [OH-] can’t build up. Study Sample Exercise 14.12 p.645 Carnegie #77 p.675

  36. Bases without OH- • Bases are proton acceptors (Brønsted-Lowry). • Example: NH3 + H2O ⇆ NH4+ + OH- • NH3 is a base. It is the lone pair on nitrogen that accepts the proton, H2O is acid • Many weak bases contain N • General Reaction: B(aq) + H2O(l) ⇆ BH+(aq) + OH- (aq) BaseAcid Conj. Acid Conj. Base

  37. Strength of Bases • Equilibrium constant is Kb • Hydroxides bases are strong bases. • Others are weak. • Smaller Kb = weaker base. • Calculate the pH for 15.0 M solution of NH3 (Kb=1.8x10-5) • Refer to steps for solving weak acid handout • Create ICE table • Substitute conc. Into equilibrium expression

  38. Polyprotic Acids • Polyprotic acids furnish more than 1 proton • Always dissociate stepwise i.e. 1 proton at a time • The first H+ comes off much easier than the second. • Denoted Ka1, Ka2, Ka3 • Ka1 is much bigger than Ka2 • For typical weak acid Ka1 > Ka2 > Ka3 • Look at Table 14.4 on p. 651 for Ka values

  39. Polyprotic Acids Calculate the pH of a 5.0 M H3PO4 solution and the equilibrium concentration of the species H3PO4, H2PO4-, HPO4-2, and PO4-3 • List major species, write balanced equations • Use Ka & Kwto decide dominant equilibrium • Write equilibrium expression • Create ICE table • Solve for x • Repeat for each successive species Carnegie #95 p. 676

  40. Acid-Base Properties of Salts • Salts are ionic compounds. • Cations of strong bases(Na+ & K+) have no effect on [H+] • Anion of strong acids(Cl- & NO3-) have no effect on [H+] • Thus these salts NaCl, KNO3 have no effect on [H+] • Aqueous solutions of these salts are neutral (pH = 7) • Anion of weak acidyield strong conjugate base, solution will be basic • Alkali metal and Alkaline Earth Metal cations have no effect • Metal cations with +2 or +3 charges will produce weakly acidic solutions

  41. Acid-Base Properties of Salts

  42. Acid-Base Properties of Salts • Practice: Decide whether each of the following will give rise to an acidic, basic or neutral solution with water. • NaNO3 • K3PO4 • FeCl2 • NaHCO3 • NH4F

  43. Basic Salts • If salt solution: • Has neutral cation • And anion is conjugate base of a weak acid is basic • Reverse is true for acidic salts • If you know one (Ka or Kb ) you can calculate the other because you always know KW Ka x Kb = KW Add this to your formula list

  44. Salts as Weak Base • Study Sample Exercise 14.18 p. 656 • The anion of a weak acid is a weak base. • Calculate the pH of a 0.30 M NaF solution. • KaHF = 7.2 x 10-4 • List major species • Write equilibrium reaction • Calc. Kb expression from Ka x Kb = Kw • Create ICE table • Solve for x • Calculate pH Do #103 p. 676

  45. Acidic salts • Salts in which anion is not a base and cation is conjugate acid of weak base produce acidic solution • The same development as bases leads to • Ka x Kb = KW • Sample 14.19 Salt as Weak Acid 0.10 M NH4Cl , Kb of NH3 1.8 x 10-5 • Other acidic salts are those of highly charged metal ions, Example Al+3

  46. Hydrated metals • Highly charged metal ions pull the electrons of surrounding water molecules toward them. • Make it easier for H+ to come off. H O Al+3 H

  47. Qualitative PredictionsSalts w/ Acidic & Basic PropertiesCompare Ka for acidic ion with Kb for basic ion • IF Ka > Kb acidic • IF Ka < Kb basic • IF Ka = Kb Neutral

  48. Effect of Structure on Acid-base Properties 14.9 • Any molecule with an H atom is a potential acid • Two Factors will determine if molecule w/ H-X is acid: • Strength of Bond - The stronger the H-X bond the less acidic (Table 14.7) • Polarity: The more polar the H-X bond the stronger the acid (electronegativity) The more polar H-O-X bond -stronger acid.

  49. Strength of Oxyacids • The more oxygen hooked to the central atom, the more acidic the hydrogen. HClO4 > HClO3 > HClO2 > HClO • Remember that the H is attached to an oxygen atom. (Table 14.8) • Oxygen is electronegative & pull electrons away from hydrogen allowing it to dissociate more completely

  50. Cl O H Strength of oxyacids Electron Density

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