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  1. Acids and Bases Main Concepts: Chapters 14 and 15 • Properties of Acids and Bases • Types of Acids and Bases • Arrhenius (traditional) Acids and Bases • Bronsted-Lowry Acids and Bases • Lewis Acids and Bases • Acid-Base Reactions • Molarity, normality, and equivalence calculations • pH and pH calculations, H+, OH-, pH, pOH • Neutralization and titration calculations 1 1

  2. Aqueous Solutions Electrolytes: substances whose water solutions form ions and conduct electricity including most acids, bases, and salts. Strong Electrolytes: very good conductors of electricity; nearly 100% of the solute is dissolved (dissociated into ions). Weak Electrolytes: very weak conductors of electricity; solute only partially dissolved. Non-electrolyte: a substance whose water solution does not form ions and does not conduct electricity (sugar).

  3. Aqueous Solutions Types of Electrolytes: • Ionic Substances (salts): compounds of a metal and nonmetal; bonded by ionic bonds, most ionic substances are soluble dissociate in water (table of solubility). NaCl(s) + H2O  Na+ + Cl- • Molecular (Covalent) Substances: compounds of nonmetals:bonded by covalent bonds, molecular substances ionize in water (acids). • HCl + H2O  H+(H3O+) + Cl- • HC2H3O2(l) + H2O(l)  H3O+(aq) + C2H3O2-(aq)

  4. Properties of Aqueous Acids • Acids are electrolytes: aqueous solutions of acids ionize when added to water and conduct electric current. • Traditional acids contain hydrogen and react with active metals to produce hydrogen gas (H2) by single replacement reactions. Mg + HCl  MgCl2 + H2

  5. Properties of Aqueous Acids • Acids change the color of dyes known as acid base indicators: acids turn blue litmus - red; turns bromothymol blue - yellow, and leaves phenolphthalein clear. • Neutralizes bases: acids react with bases to produce salts (ionic compounds) and water in neutralization reactions. NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

  6. Properties of Aqueous Acids Acids: • Sour Taste: Weak acids taste sour. • citrus fruits contain citric acid. • apples contain malic acid • vinegar contains acetic acid • sour milk contains lactic acid • wine contains tartaric acid • carbonated beverages contain phosphoric acid.

  7. Properties of Aqueous Acids • Acids have a pH less than 7.0 • Strong Acids: Acids that are completely or nearly completely ionized in water. HCl + H2O → H+(H3O+) + Cl- • Weak Acids:  Acids that are not completely ionized in water. • Strong acids in concentrated solutions are very corrosive and powerful poisons.

  8. Common Acids Strong acids: an acid that is a strong electrolyte and dissolves completely or nearly completely in water (aqueous solutions). • HCl HNO3 HClO3 HBr H2SO4 Hl HClO4 Weak acids: an acid that is a weak electrolyte and only partially dissolves in water yielding a relatively low concentration of ions in solution. • CH3COOHHF HNO2 H3PO4 H2S H3BO3 HCN HClO H2SO3

  9. Properties of Aqueous Bases • Bases are electrolytes: aqueous solutions of bases conduct electricity. • Traditional bases contain OH-(hydroxide ion). Group-1 metals react with water by single replacement reactions to form strong bases. K + H2O → KOH + H2 • Bases change the color of acid base indicators: bases turns red litmus blue, turns bromothymol blue - blue, and turns phenolphthalein pink.

  10. Properties of Aqueous Bases • Bases neutralize acids in neutralization reactions to produce salts and water. KOH(aq) + HCl(aq) KCl(aq) + H2O(l) • Bases have a bitter taste. • Dilute aqueous solutions of bases feel slippery to the skin (soap) • Bases have a pH greater than 7.0

  11. Properties of Aqueous Bases • Strong Bases: bases that are completely or nearly completely ionized including all group-1 metal hydroxides (KOH, NaOH,…) and the heavy group-2 metal hydroxides [Ca(OH)2, Sr(OH)2, and Ba(OH)2]. Ba(OH)2 + H2O → Ba2+ + 2OH- • Weak Bases: those bases that are only slightly ionized including NH3, NH4OH, C6H5NH2 (organic amines), and all others.

  12. Strong Acids and Bases Strong AcidsStrong Bases 12

  13. Aqueous Acids Monoprotic acids: an acid that can donate only one proton per molecule; one mole H+/mole acid (HCl) Polyprotic acid: an acid that can donate more than one proton per molecule. Diprotic acid: an acid that can donate two protons per molecule. (H2SO4) Triprotic acid: an acid that can donate three protons per molecule. (H3PO4)

  14. Aqueous Acids • The strength of an acid depends on the degree of ionization in aqueous (water) solution and not on the amount of hydrogen in the molecule. • Any acid that dissolves completely or nearly completely in water is strong and any base that dissolves completely or nearly completely in water is strong regardless of the concentration. • Organic acid: an organic compound that contains an acid group, the carboxyl group (-COOH). Acetic acid: CH3COOH

  15. Types of Acids and Bases • Arrhenius Acids and Bases (traditional) • Bronsted – Lowry Acids and Bases • Lewis Acids and Bases These different types or definitions of acids and bases have evolved over the years allowing more substances to be classified as an acid or base beyond the traditional definition.

  16. Arrhenius Acids • Arrhenius acid (traditional): any substance that produces or increases the hydrogen ion (H+,H3O+)ionconcentration of an aqueous solution (HCl, HC2H3O2). Dissolving acids in water: • HCl + H2O H+ + Cl- monoprotic, or HCl + H2OH3O+ + Cl- • H2SO42H+ + SO4-2 diprotic • H3PO43H+ + PO4-3 triprotic

  17. Arrhenius Bases • Arrhenius base (traditional): any substance that increases the hydroxide ion (OH-) concentration of an aqueous solution. (NaOH, Ba(OH)2, NH4OH) Dissolving bases in water: • NaOH + H2O  Na+ + OH- Ba(OH)2 + H2O  Ba+2 + 2OH-

  18. Arrhenius Acid-Base Reactions Arrhenius acid-base reactions (neutralization): acid + base  salt + water • H2SO4 + 2NaOH  Na2SO4 + 2H2O(l) • 2HCl + Ba(OH)2 BaCl2 + 2H2O(l) • Net ionic equation: H+(aq) + OH-(aq)  H2O(l)

  19. Arrhenius Acid-Base Reactions • To calculate concentrations or volumes in a [neutralization] reaction use: MaVa = MbVbor nMaVa = nMbVb 1) Calculate the concentration of KOH if 800 mL of KOH is required to neutralize 265 mL of 1.75M HCl. nMaVa= nMbVb (1.75M)(265 mL) = Mb(800 mL) Mb=(1.75M)(265 mL) = 0.58M 800 mL

  20. Arrhenius Acid-Base Reactions 2) Calculate the volume of 0.25M HCl required to neutralize 250 mL of 0.45M Ca(OH)2 nMaVa = nMbVb (1)(.25)(Va) = (2)(250)(.45) = 900 mL 3) Calculate the concentration of H3PO4 if 375 mL is needed to neutralize 0.14 L of 2.75M Sr(OH)2. • nMaVa = nMbVb (3)(Ma)(375 mL) = (2)(140 mL)(2.75M) = 0.68M

  21. Bronsted – Lowry Acids and Bases • A major problem with the Arrhenius definition of acids and bases is that it is limited to acids and bases in water solutions in which all acids contain the “H+” ion and all bases contained “OH-” ion in the formula. Ammonia was not included as an Arrhenius base: NH3+ H2O NH4+ + OH- Johannes Brønsted and Thomas Lowry felt the Arrhenius definition was too limiting because there are many non-aqueous systems (no water is present). They came up with the following definitions for acids and bases.

  22. Bronsted – Lowry Acids and Bases Bronsted – Lowry acid: aproton donor: any substance capable of donating a proton. The proton is exactly the same as the H+ (H3O+) ion in the Arrhenius acid. All Bronsted acids are also Arrhenius acids Bronsted – Lowry base: a proton acceptor: any substance capable of accepting a proton (not required to have OH- in the formula). A Bronsted – Lowry base must have an unshared pair of electrons to accept the proton forming a coordinate covalent bond.

  23. Bronsted – Lowry Acids and Bases A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor base acid conjugateacid conjugatebase

  24. Bronsted – Lowry Acids and Bases Bronsted – Lowry acid-base reaction: The products that result from a Bronsted–Lowry acid-base reaction are called the conjugate acid and the conjugate base or the conjugate acid-base pair. The conjugate baseis the particle that remains after a proton is donated by an acid. The conjugate acid is formed when a base accepts a proton from an acid. HF(aq) + HCO3-(aq)  H2CO3(aq) + F-(aq) acid base conj. acid conj. base

  25. Conjugate Pairs

  26. Conjugate Pairs

  27. Bronsted – Lowry Acids and Bases • The difference between an acid and its conjugate base or a base and its conjugate acid is one proton. • HF(aq) + HCO3-(aq)  H2CO3(aq) + F-(aq) acid base conj. acid conj. Base • The stronger the acid the weaker the conjugate base. The stronger the base, the weaker the conjugate acid.

  28. Bronsted Practice • Which of the following represent conjugate acid-base pairs? For those pairs that are not conjugates, write the correct conjugate acid or base for each species in the pair. • H2SO4, SO42- • H2PO4- , HPO42- • HClO4, Cl- • NH4+, NH2-

  29. Bronsted Practice • HSO4-, SO42- • HBr, BrO- • H2PO4- , PO43- • HNO3 , NO2-

  30. Bronsted Practice • Identify the conjugate acid-base pairs in each of the following equations: • HSO4- + H2O  SO42- + H3O+ • NH2- + H2O  NH3 + OH- • HCl + CH3OH  Cl- + CH3OH2+ • NH3 + H2O  NH4+ + OH-

  31. Bronsted Practice • Identify the conjugate acid-base pairs in each of the following equations: • PO43- + H2O  HPO42- + OH- • C2H3O2- + H2O  HC2H3O2 + OH-

  32. Lewis Acids and Bases • Lewis acid: any substance that can accept a pair of electrons to form a coordinate covalent bond (AlCl3). • Lewis base: any substance that can donate a pair of electrons ( OH-, HSO4-, O2-) to form a coordinate covalent bond: i.e. substances that have one or more unshared pairs of electrons (H2O, NH3). • Adduct: the product of a Lewis acid-base reaction • Review Lewis structures/molecular geometry.

  33. Lewis Acids and Bases Lewis acid: electron pair acceptor (Lewis structure). Lewis base: electron pair donor (Lewis structure).

  34. Lewis Acids and Bases Formation of the hydronium ion is also an example. Electron pair of the new O-H bond originates on the Lewis base.

  35. Lewis Acids and Bases

  36. Lewis Acid-Base Reactions Lewis expands acid/base reactions to include many substances without H in the formula that can react without being in aqueous solution. H+ + NH3 NH4+ LALB adduct NH3 + BF3 F3BNH3 LB LA adduct Na2O + SO3 Na2SO4 LB LA adduct

  37. Acid-Base Types: Practice • Identify each as an Arrhenius, Bronsted, or Lewis acid or base and the type of overall reaction where applicable. • HBr  H+ + Br- • HNO3 + H2O  H3O+ + NO3- • LiOH  Li+ + OH- • AlF3 + NH3 F3AlNH3

  38. Acid-Base Types: Practice • KOH + H2SO4 H2O + K2SO4 • HBr + H2O  H3O+ + Br- • HBr + NH3 NH4+ + Br-

  39. Bronsted/Lewis Bases • All Bronsted-Lowry bases are Lewis bases, but not all Lewis bases are Bronsted-Lowry bases. • All Bronsted-Lowry bases must have a lone pair of electrons to accept a hydrogen, and all Lewis bases have lone pairs. However, Lewis bases may act as nucleophiles (electron pair donors) for any number of atoms; B, Al, C, etc; not just H.

  40. Amphoteric (amphiprotic) Substances Amphoteric (amphiprotic): substances such as HCO3-, HSO4-, H2O, and NH3 which can act as either proton donors (acids) or proton acceptors (bases) depending upon what is present. • In the presence of a stronger acid, the HSO4- ion acts as a base (proton acceptor). H3O+ + HSO4- H2SO4 + H2O • In the presence of a stronger base, the HSO4- ion acts as a proton donor (acid). HSO4- + OH- H2O + SO42-

  41. Autoionization of Water Autoionization (self Ionization ) of Water: water spontaneously forms low concentrations of H+ (H3O+)(aq) and OH-(aq) by proton transfer from one water molecule to another. • H2O(l) + H2O(l) H3O+(aq) + OH-(aq)or… H2O(l) H+(aq) + OH-(aq) • Water is very slightly ionized yielding one mole of H+ ions and one mole of OH- ions per 107 liters. • H+ = OH- = 1 mole ion = 10-7 mol H+ = 10-7M 107 liter H2O liter H2O

  42. Equilibrium Product Constant • The equilibrium product constant for water is: H2O(l) H+(aq) + OH-(aq) KW = H+OH- = 10-7 x 10-7 = 10-14 Brackets   represent molar concentration (molarity) • KW = 10-14 remains constant in water and dilute aqueous solutions at constant temperature. An increase in the concentration of one ion will cause a decrease in concentration of the other. • This is the basis for the pH scale.

  43. pH Scale

  44. Normality and Equivalents Chemical Equivalents: the quantity of solutes that have equivalent combining capacity. HCl + KOH  KCl + H2O 1 mole 1 mole  1 mole 1 mole H2SO4 + 2KOH  K2SO4 + 2H2O 1 mole 2 mole  1 mole 2 mole ½ mole 1 mole  ½ mole 1 mole 1/2 mole H2SO4 and 1 mole HCl are chemical equivalents in neutralizing 1 mole of KOH.

  45. Normality and Equivalents Equivalent mass of an acid: the quantity, in grams, that supplies one mole of protons (H+). Equivalent mass of a base: the quantity, in grams, that supplies one mole of hydroxide ions (OH-). • The mass of one equivalent of HCl is 36.5 g and the mass of one equivalent of Be(OH)2 is 21.5 g. • One equivalent of HCl will neutralize one equivalent of Be(OH)2 in a neutralization reaction. • One equivalent of any acid will neutralize one equivalent of any base.

  46. Normality and Equivalents Normality: the number of equivalents of solute per liter of solution. Molarity = moles of solute/liter solution Normality = equivalents of solute/liter solution N = nM (n = # H’s or OH’s in the formula) MaVa = MbVb neutralization/titration nMaVa = nMbVb NaVa = NbVb One Na = one Nb

  47. Formulas • H+OH- = 1 x 10-14 • H+ = 1 x 10-14OH- • OH-  = 1 x 10-14H+ • pH = -logH+ and H+ = antilog(-pH) = 10-pH • pOH = -logOH- and OH- = antilog(-pOH) = 10-pOH • pH + pOH = 14

  48. Calculations: Strong Acids and Bases Strong acids and strong bases ionize completely and no equilibrium is established because all reactant is converted to products. • H2SO4(aq) + NaOH(aq)  Na2SO4(aq) + H2O(l) • 2H+ + SO42- + Na+ + OH- 2Na+ + SO42- + H2O(l) • Net ionic equation: H+ + OH- H2O(l) • The original solution concentration is the concentration for the acid or base (if monoprotic). • For 0.015 M HNO3 calculate [H+] and pH. • [H+] = 0.015 M, pH = 1.82

  49. Calculations 1. Calculate the pH of a solution whose hydrogen ion concentration H+ (H3O+) is 1.0 x 10-4 M. • pH = -logH3O+ • pH = -log(1.0 x 10-4) on calculator • pH = 1.0 EE 4  log  • pH = 4.0

  50. Calculations 2. Calculate the pH and pOH of a solution that contains a H+ concentration H3O+ of 3.5 x 10-3 M. • pH = -logH+ • pH = -log(3.5 x 10-3) on calculator • pH = 3.5 EE 3  log  • pH = 2.46 • and since pH + pOH = 14 • pOH = 14.00 - 2.46 = 11.54