New chm 152 unit 3 power points sp13
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New chm 152 unit 3 power points sp13 New chm 152 unit 3 power points sp13 Presentation Transcript

  • ACIDS AND BASES Chapter 4.4 & 18.1-5 Silberberg
  •  See the Learning Objectives on page 821.  Understand these Concepts: 18.1-20, 23-25.  Master these Skills: 18.1-5, 6-10, 12-14, 17. 
  • Table 18.1 Some Common Acids and Bases and their Household Uses.
  •  The Savante Arrhenius Theory (1884)   Acid-ionizes in aqueous solution to produce H+ ions. Base-ionizes in aqueous solution to produce OHions.
  • Svante August Arrhenius  Svante Arrhenius was born on February 19, 1859 and he died October 2, 1927. He was a Swedish physical chemist best known for the development of his acidbase theory. In 1903 he was awarded the Nobel Prize for Chemistry.
  • Strong and Weak Acids A strong acid dissociates completely into ions in water: HA(g or l) + H2O(l) → H3O+(aq) + A-(aq) A dilute solution of a strong acid contains no HA molecules. A weak acid dissociates slightly to form ions in water: HA(aq) + H2O(l) H3O+(aq) + A-(aq) In a dilute solution of a weak acid, most HA molecules are undissociated. Kc = [H3O+][A-] has a very small value. [HA][H2O]
  • Classifying the Relative Strengths of Acids  Strong acids include    the hydrohalic acids (HCl, HBr, and HI) and oxoacids in which the number of O atoms exceeds the number of ionizable protons by two or more (eg., HNO3, H2SO4, HClO4.) Weak acids include  the hydrohalic acid HF,  acids in which H is not bonded to O or to a halogen (eg., HCN),  oxoacids in which the number of O atoms equals or exceeds the number of ionizable protons by one (eg., HClO, HNO2), and  carboxylic acids, which have the general formula RCOOH (eg., CH3COOH and C6H5COOH.)
  •  HBr + H2O   H3O+ + Br -1 HBr is a strong acid--it completely dissociates. HF + H2O  ------> = H3O+ + F HF is a weak acid--it only partially ionizes.
  •     H+ ions are bare protons. In aqueous solutions these protons are hydrated H(H2O)n+ where n is a small number. The hydrated hydrogen ion is normally represented as H3O+. In many reactions when it is obvious that aqueous solutions are involved, H+ will be used to represent H3O+.
  •        HCl HBr HI HNO3 HClO3 HClO4 H2SO4
  • Classifying the Relative Strengths of Bases  Strong bases include  water-soluble compounds containing O2- or OH- ions.  The cations are usually those of the most active metals:  M2O or MOH, where M = Group 1A(1) metal (Li, Na, K, Rb, Cs)  MO or M(OH)2 where M = group 2A(2) metal (Ca, Sr, Ba).  Weak bases include  ammonia (NH3),  amines, which have the general formula  The common structural feature is an N atom with a lone electron pair.
  •  NaOH ------> Na+ + OH-1   NaOH is a strong base--it completely dissociates. NH3 + H2O  = NH4+ + OH-1 NH3 is a weak base-- it only partially ionizes to produce OH- ions.
  •         LiOH NaOH KOH RbOH CsOH Ca(OH)2 Ba(OH)2 Sr(OH)2
  •    Neutralization is the combination of hydrogen ions(acid) with hydroxide ions(base) to produce water(neutral). The net ionic equation for the reaction of a strong acid with a strong base would be H+ + OH= H2O
  •       Acid- a proton (H+) donor Base- a proton(H+) acceptor An acid-base reaction is the transfer of a proton from an acid to a base. NH3 + H2O = NH4+ + OH-1 base acid conjugate conjugate acid base
  • Johannes Nicolaus Bronsted  Johannes Brønsted was born on February 22, 1879 and died on December 17, 1947. He was a Danish physical chemist known for a widely applicable acidbase concept introduced in 1923. His work was independent of Lowry.
  • Thomas Martin Lowry   Thomas Lowry was born on October 26, 1874 and died on September 2, 1936. He was an English chemist widely - known for an acid-base concept identical to that of Johannes N. Bronsted.
  • Brønsted-Lowry Acid-Base Definition An acid is a proton donor, any species that donates an H+ ion. • An acid must contain H in its formula. A base is a proton acceptor, any species that accepts an H+ ion. • A base must contain a lone pair of electrons to bond to H+. An acid-base reaction is a proton-transfer process.
  • Conjugate Acid-Base Pairs In the forward reaction: NH3 accepts a H+ to form NH4+. H2S + NH3 HS- + NH4+ H2S donates a H+ to form HS-. In the reverse reaction: NH4+ donates a H+ to form NH3. H2S + NH3 HS- + NH4+ HS- accepts a H+ to form H2S.
  • Conjugate Acid-Base Pairs H2S + NH3 HS- + NH4+ H2S and HS- are a conjugate acid-base pair: HS- is the conjugate base of the acid H2S. NH3 and NH4+ are a conjugate acid-base pair: NH4+ is the conjugate acid of the base NH3. A Brønsted-Lowry acid-base reaction occurs when an acid and a base react to form their conjugate base and conjugate acid, respectively. acid1 + base2 base1 + acid2
  • Table 18.4 The Conjugate Pairs in some Acid-Base Reactions Conjugate Pair Acid + Base Base + Acid Conjugate Pair Reaction 1 HF + H2O F- + H3O+ Reaction 2 HCOOH + CN- HCOO- + HCN Reaction 3 NH4+ + CO32- NH3 + HCO3- Reaction 4 H2PO4- + OH- HPO42- + H2O Reaction 5 H2SO4 + N2H5+ HSO4- + N2H62+ Reaction 6 HPO42- + SO32- PO43- + HSO3-
  • Sample Problem 18.4 Identifying Conjugate Acid-Base Pairs PROBLEM: The following reactions are important environmental processes. Identify the conjugate acid-base pairs. (a) H2PO4-(aq) + CO32-(aq) (b) H2O(l) + SO32-(aq) HPO42-(aq) + HCO3-(aq) OH-(aq) + HSO3-(aq) PLAN: To find the conjugate pairs, we find the species that donated an H+ (acid) and the species that accepted it (base). The acid donates an H+ to becomes its conjugate base, and the base accepts an H+ to becomes it conjugate acid. SOLUTION : (a) H2PO4-(aq) + CO32-(aq) acid1 base2 HPO42-(aq) + HCO3-(aq) base1 acid2 The conjugate acid-base pairs are H2PO4-/HPO42- and CO32-/HCO3-.
  • Sample Problem 18.4 (b) H2O(l) + SO32-(aq) acid1 base2 OH-(aq) + HSO3-(aq) base1 acid2 The conjugate acid-base pairs are H2O/OH- and SO32-/HSO3-.
  • Figure 18.8 Strengths of conjugate acid-base pairs. The stronger the acid is, the weaker its conjugate base. When an acid reacts with a base that is farther down the list, the reaction proceeds to the right (Kc > 1).
  •     Water can act as both an acid and a base. Water is said to be amphoteric (amphiprotic). H2O + H2O = H3O+ + OH -1 base acid conjugate conjugate acid base
  • Autoionization of Water Water dissociates very slightly into ions in an equilibrium process known as autoionization or self-ionization. 2H2O (l) H3O+ (aq) + OH- (aq)
  •     Acid--species that accepts a share in a pair of electrons--an electron pair acceptor Base--species that donates a share in a pair of electrons--an electron pair donor :NH3 + H2O base acid = H:NH3+ + OH-1 conj.acid conj.base
  • Gilbert Newton Lewis  G. N. Lewis was born on October 23, 1875 and died on March 23, 1946. He was a famous American physical chemist, who in 1923, developed the electron-pair theory of acid-base chemical reactions. He is also known for the creation of Lewis structures for drawing chemical molecules.
  • The Lewis Acid-Base Definition A Lewis base is any species that donates an electron pair to form a bond. A Lewis acid is any species that accepts an electron pair to form a bond. The Lewis definition views an acid-base reaction as the donation and acceptance of an electron pair to form a covalent bond.
  • Lewis Acids and Bases A Lewis base must have a lone pair of electrons to donate. Any substance that is a Brønsted-Lowry base is also a Lewis base. A Lewis acid must have a vacant orbital (or be able to rearrange its bonds to form one) to accept a lone pair and form a new bond. Many substances that are not Brønsted-Lowry acids are Lewis acids. The Lewis definition expands the classes of acids.
  • Electron-Deficient Molecules as Lewis Acids B and Al often form electron-deficient molecules, and these atoms have an unoccupied p orbital that can accept a pair of electrons: BF3 accepts an electron pair from ammonia to form a covalent bond.
  • Lewis Acids with Polar Multiple Bonds Molecules that contain a polar multiple bond often function as Lewis acids: The O atom of an H2O molecule donates a lone pair to the S of SO2, forming a new S‒O σ bond and breaking one of the S‒O p bonds.
  • Metal Cations as Lewis Acids A metal cation acts as a Lewis acid when it dissolves in water to form a hydrated ion: The O atom of an H2O molecule donates a lone pair to an available orbital on the metal cation.
  • Sample Problem 18.15 Identifying Lewis Acids and Bases PROBLEM: Identify the Lewis acids and Lewis bases in the following reactions: (a) H+ + OHH2O (b) Cl- + BCl3 BCl4- (c) K+ + 6H2O K(H2O6)+ PLAN: We examine the formulas to see which species accepts the electron pair (Lewis acid) and which donates it (Lewis base) in forming the adduct. SOLUTION : The H+ ion accepts the electron pair from OH-. H+ is the Lewis acid and (a) OH- is the Lewis base. (b) BCl3 accepts an electron pair from Cl-. Cl- is the Lewis base and BCl3 is the Lewis acid. (c) An O atom from each H2O molecule donates an electron pair to K+. H2O is therefore the Lewis base, and K+ is the Lewis acid.
  •   ionize or dissociate completely Strong acids   100% H+ + Cl-1 Strong bases   HCl NaOH 100% Na+ + OH-1 Soluble salts  NaCl 100% Na+ + Cl-1
  •   Determine the ion concentrations in 0.050M nitric acid solution. Determine the concentration of all ions in a 0.020M solution of calcium hydroxide, Ca(OH)2.
  •     H2O = H+ + OH-1 KW = [H+][OH-1] At 25OC KW = 1.0x10-14 Calculate the [H+] and [OH-1 ] in 0.050M HCl solution.
  • The pH Scale pH = -log[H3O+] The pH of a solution indicates its relative acidity: In an acidic solution, In a neutral solution, In basic solution, pH < 7.00 pH = 7.00 pH > 7.00 The higher the pH, the lower the [H3O+] and the less acidic the solution.
  •       a convenient way of expressing the acidity and basicity of dilute aqueous solutions. pH = -log[H+] This applies to other ion concentrations as well pOH = -log[OH -1 ] Another useful relationship pH + pOH = 14
  • Figure 18.7 Methods for measuring the pH of an aqueous solution. pH (indicator) paper pH meter
  • Figure 18.5 The pH values of some familiar aqueous solutions. pH = -log [H3O+]
  • Table 18.3 The Relationship between Ka and pKa Acid Name (Formula) Ka at 25°C pKa Hydrogen sulfate ion (HSO4-) 1.0x10-2 1.99 Nitrous acid (HNO2) 7.1x10-4 3.15 Acetic acid (CH3COOH) 1.8x10-5 4.75 Hypobromous acid (HBrO) 2.3x10-9 8.64 Phenol (C6H5OH) 1.0x10-10 10.00 pKa = -logKa A low pKa corresponds to a high Ka.
  • pH, pOH, and pKw Kw = [H3O+][OH-] = 1.0x10-14 at 25°C pH = -log[H3O+] pOH = -log[OH-] pKw = pH + pOH = 14.00 at 25°C pH + pOH = pKw for any aqueous solution at any temperature. Since Kw is a constant, the values of pH, pOH, [H3O+], and [OH-] are interrelated: • If [H3O+] increases, [OH-] decreases (and vice versa). • If pH increases, pOH decreases (and vice versa).
  • Figure 18.5 The relations among [H3O+], pH, [OH-], and pOH.
  • Sample Problem 18.3 Calculating [H3O+], pH, [OH-], and pOH PROBLEM: In an art restoration project, a conservator prepares copper-plate etching solutions by diluting concentrated HNO3 to 2.0 M, 0.30 M, and 0.0063 M HNO3. Calculate [H3O+], pH, [OH-], and pOH of the three solutions at 25°C. PLAN: HNO3 is a strong acid so it dissociates completely, and [H3O+] = [HNO3]init. We use the given concentrations and the value of Kw at 25°C to find [H3O+] and [OH-]. We can then calculate pH and pOH. SOLUTION : Calculating the values for 2.0 M HNO3: [H3O+] = 2.0 M [OH-] pH = -log[H3O+] = -log(2.0) = -0.30 Kw 1.0x10-14 = = = 5.0x10-15 M [H3O+] 2.0 pOH = -log[OH-] = -log(5.0x10-15) = 14.30
  • Sample Problem 18.3 Calculating the values for 0.30 M HNO3: [H3O+] = 0.30 M [OH-] pH = -log[H3O+] = -log(0.30) = 0.52 Kw 1.0x10-14 = = = 3.3x10-14 M [H3O+] 0.30 pOH = -log[OH-] = -log(3.3x10-14) = 13.48 Calculating the values for 0.0063 M HNO3: [H3O+] = 0.0063 M [OH-] pH = -log[H3O+] = -log(0.30) = 2.20 Kw 1.0x10-14 = = = 1.6x10-12 M [H3O+] 0.0063 pOH = -log[OH-] = -log(1.6x10-12) = 11.80
  •    Calculate the pH of a solution in which [H+] = 0.030M The pH of a solution is 4.597. Determine the [H+] of this solution. Determine the [H+], [OH-1 ], pH and pOH for a 0.020M HNO3 solution.
  • Solving Problems Involving Weak-Acid Equilibria The notation system • Molar concentrations are indicated by [ ]. • A bracketed formula with no subscript indicates an equilibrium concentration. The assumptions • [H3O+] from the autoionization of H2O is negligible. • A weak acid has a small Ka and its dissociation is negligible. [HA] ≈ [HA]init.
  •      Consider the reaction when the weak acid acetic acid is added to water. CH3COOH + H2O = H3O+ + CH3COO-1 Ka = [H+][CH3COO-1 ] [CH3COOH] Write the equation for the ionization of HCN in aqueous solution.
  •   In 0.12M solution, a weak acid HY is 5.0% ionized. Determine the value for the ionization constant for this weak acid. The pH of a 0.10M solution of a weak monoprotic acid HA is 2.97. Calculate the value for the ionization constant of this weak acid.
  •     Determine the concentrations of all species in 0.15M acetic acid , CH3COOH, solution. Ka = 1.8x10-5 Determine the concentrations of all species in 0.15M HCN solution. Ka = 4.0x10-10 Determine the concentrations of all species in 0.15M NH3. Kb=1.8x10-5
  •  The pH of an aqueous NH3 solution is 11.37. Determine the molarity of this aqueous ammonia solution. Kb=1.8x10-5