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Chapter 15 Acid-Base Theories

Chapter 15 Acid-Base Theories

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Chapter 15 Acid-Base Theories

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  1. Chapter 15Acid-Base Theories

  2. Properties of Acids and Bases • Acids • Give foods a tart or sour taste • What acidic foods might you eat? • Aqueous solutions of acids are electrolytes • Conduct Electricity • Some are strong electrolytes (strong acids) • Some are weak electrolytes (weak acids) • Cause indicator dyes to change colors • Many metals react with acids producing hydrogen gas • React with compounds containing hydroxide ions to form water and a salt

  3. Properties of Acids and Bases • Bases • Have bitter taste, and slippery feel • Aqueous solutions of bases are also electrolytes • Conduct Electricity • Some are strong electrolytes (strong bases) • Some are weak electrolytes (weak bases) • Cause indicator dyes to change colors • Water and salt are formed when a base that contains hydroxide ions react with an acid

  4. Arrhenious Acids and Bases • Acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution • Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution

  5. Nonionizable Hydrogen H O H C C O- H+ Ionizable Hydrogen H Arrhenious Acids • Can be monoprotic, diprotic, or triprotic • Monoportic: HNO3 → H+ + NO3- • Ionization yields one hydrogen ion • Diprotic: H2SO4 → 2H+ + SO42- • Complete ionization yields 2 hydrogen ions • Triprotic: H3PO4 → 3H+ + PO43- • Complete ionization yields 3 hydrogen ions • Not all the hydrogens in an acid may be released as hydrogen ions • Not all hydrogen containing compounds are acids • Only hydrogens joined to very electronegative elements, and thus have very polar bonds, are ionizable in water Ethanoic Acid

  6. Arrhenious Bases • NaOH → Na+(aq) + OH-(aq) • KOH → K+(aq) + OH-(aq) • Bases formed with group one metals are very soluble and caustic • Can be made by reacting group one metals with water • Na + H2O → Na+(aq) + OH-(aq) H2 (g) • Bases of group 2 metals are very weak resulting low solubility • Examples are Ca(OH)2 and Mg(OH)2

  7. Bronsted-Lowry Acids and Bases • Arrhenious definition of acids and bases is not very comprehensive and does not explain why certain substances have basic or acidic properties • Ammonia (NH3) is a base, but there is no hydroxide (OH-) in the compound to ionize • The Bronsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogen-ion acceptor • Why ammonia is a base NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq) Hydrogen ion aceptor, Bronsted-Lowry Base Hydrogen ion donar, Bronsted-Lowry Acid Makes the solution basic

  8. conjugate acid-base pair NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Conjugate Acid Conjugate Base Base Acid conjugate acid-base pair conjugate acid-base pair HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) Conjugate Acid Conjugate Base Acid Base conjugate acid-base pair Conjugate Acids and Bases • A conjugant acid is the particle formed when a base gains a hydrogen ion • A conjugant base is the particle that remains when an acid has donated a hydrogen ion • A conjugate acid-base pair consists of two substances related by the loss or gain of a single hydrogen ion

  9. conjugate acid-base pair HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) Conjugate Acid Conjugate Base Acid Base conjugate acid-base pair Conjugate Acids and Bases • A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion H3O+ • In above equation, what is the hydrogen ion donor (acid) and which is the hydrogen ion acceptor (base) • Notice, water can both accept and donate a hydrogen ion and thus act as an acid and a base • A substance that can act as both an acid and a base is said to be amphoteric • Amino Acids as an example

  10. Lewis Acids and Bases • Acids accept a pair of electrons during a reaction while a base donates a pair of electrons • Lewis acid – a substance that can accept a pair of electrons to form a covalent bond • Lewis base – a substance that can donate a pair of electrons to form a covalent bond NH3 + BF3 → NH3BF3 Identify the Lewis Acid and the Lewis Base in the above equation

  11. Acid-Base Definitions Review

  12. Occasionally collusions between water molecules cause them to react forming hydroxide ions and hydronium ions The reaction in which water molecules produce ions is called the self ionization of water In aqueous solution, hydrogen ions H+ are always joined to a water molecule as hydronium ions In pure (neutral) water, the self-ionization of water results in 1 x 10-7 M of H+ ions and 1 x 10-7 M of OH- ions Any aqueous solution in which H+ and OH- ions are equal is described as a neutral solution Hydrogen Ions and Acidity H2O(l) H+(aq) + OH-(aq) H2O (l) + H2O (l) H3O+(aq) + OH-(aq)

  13. Ion Product Constant for Water • For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0x10-14 Kw = [H+] x [OH-] = 1.0x10-14M • The product of the concentrations of the hydrogen ions and hydroxide ions in water is called the ion-product constant • Aqueous acids and bases sift the ratio of hydrogen ions to hydroxide ions in solution causing it to become either acidic or basic • In a basic solution aka alkaline solution, the hydroxide ion (OH-) is greater than 1x10-7M and the hydrogen ion (H+) is less 1x10-7M • In a acidic solution, the hydrogen ion (H+) is greater than 1x10-7M and the hydroxide ion (OH-) is less 1x10-7M • Regardless of the acidity or alkalinity of the solution, the product of the Molarity (M) concentration of H+ and OH- always equals 1x10-14 at 25ºC • If the hydrogen ion H+ concentration in a soft drink is 1 x 10-5M, what is the concentration of the hydroxide ion OH-? • Is the solution basic, neutral, or acidic?

  14. The pH Concept • Expressing hydrogen-ion concentrations in molarity is cumbersome • Soren Sorensen suggested that the hydrogen ion concentration be expressed as the negative log of the hydrogen-ion concentration giving us much smaller numbers to work with pH = -log[H+] • The pH of a solution is the negative logarithm of the hydrogen ion concentration • A neutral solution H+ = 1x10-7 has a pH = -log[1x10-7]= 7 • A solution in which the [H+] is greater than 1x10-7 M and has a pH less than 7.0 is acidic • The pH of pure water or a neutral solution has a pH of 7 • A solution with a pH greater than 7 is basic and has a [H+] concentration of less than 1x10-7M • You can also calculate pOH which is the negative logarithm of hydroxide ion concentration pOH = -log[OH-] Work some pH problems (pH of 1 x10-5M H+?)

  15. pH + pOH = 14pH = 14 – pOHpOH = 14 - pH Relationship between pH and pOH

  16. pH and significant figures • Hydrogen ion concentrations should always be reported to two significant figures • pH and pOH calculations should always be reported to two decimal places • Rules are due to the sensitivity of pH meters

  17. An indicator (HIn) is an acid or base that undergoes dissociation in a known pH range An indicator is a valuable tool for measuring pH because its acid form and base form have different colors in solution The acid form of the indicator dominates the disassociation equilibrium at low pH The basic form of the indicator dominates the disassociation equilibrium at high pH Color change of an indicator occurs in a narrow pH range ≈ 2 pH units Thus it takes many indicators to span the entire pH spectrum Indicator dyes have limitations H+ H+ (aq) + In- (aq) HIn (aq) OH- Acid Form Base Form Acid-Base IndicatorsDyes

  18. Acid-Base IndicatorspH Meter • Makes rapid, accurate pH measurements • Can record pH continuously over time when performing a reactions • Measures pH to two decimal places • Color and cloudiness of solution does not interfere with reading • Are many different types specialized for different jobs were pH measurements are required

  19. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq) <1% ionized Strengths of Acids and BasesStrong and Weak Acids and Bases • Acids are classified as strong or weak depending on the degree to which they ionize in water • A strong acid completely ionizes in water • Weak acids ionize only slightly in aqueous solution • What are some example of strong acids and weak acids HCl(q) + H2O(l) H3O+(aq) + Cl-(aq) 100% ionized

  20. pH of 0.10 M Solutions of Common Acids and Bases • CompoundpH • HCl (hydrochloric acid)  1.1 • H2SO4 (sulfuric acid)  1.2 • NaHSO4 (sodium hydrogen sulfate)  1.4 • H2SO3 (sulfurous acid)  1.5 • H3PO4 (phosphoric acid)  1.5 • HF (hydrofluoric acid)  2.1 • CH3CO2H (acetic acid) 2.9 • H2CO3 (carbonic acid) 3.8 (saturated solution) • H2S (hydrogen sulfide)  4.1 • NaH2PO4 (sodium dihydrogen phosphate)  4.4 • NH4Cl (ammonium chloride)  4.6 • HCN (hydrocyanic acid)  5.1 • Na2SO4 (sodium sulfate)  6.1 • NaCl (sodium chloride)  6.4 • NaCH3CO2 (sodium acetate) 8.4 • NaHCO3 (sodium bicarbonate)  8.4 • Na2HPO4 (sodium hydrogen phosphate)  9.3 • Na2SO3 (sodium sulfite)  9.8 • NaCN (sodium cyanide)  11.0 • NH3 (aqueous ammonia)  11.1 • Na2CO3 (sodium carbonate)  11.6 • Na3PO4 (sodium phosphate)  12.0 • NaOH (sodium hydroxide, lye)  13.0

  21. The equilibrium constant for weak acids (HA) can be written as: For dilute solutions, the concentration of water is a constant, and can be combined with Keq to give the acid dissociation constant (Ka) HA(aq) + H2O(l) H3O+(aq) + A-(aq) [H3O+] X [A-] Keq = [HA] X [H2O] [H3O+] X [A-] Keq X H2O = Ka = [HA] Acid Disassociation Constant Acid Conjugate base • Ka reflects the fraction of an acid in the ionized form and thus is sometimes referred to as the ionization constant • Weak acids have small Ka values, while stronger acids have larger Ka values; why?

  22. The equilibrium constant for weak Bases (B) can be written as: For dilute solutions, the concentration of water is a constant, and can be combined with Keq to give the base dissociation constant (Kb) B(aq) + H2O(l) BH+(aq) + HO-(aq) [BH+] X [HO-] Keq = [B] X [H2O] [BH+] X [OH-] Keq X H2O = Kb = [B] Base Disassociation Constant base Conjugate acid • Kb is the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion to the concentration of the base • The magnitue of Kb indicates the ability of a weak base to compete with the very strong base OH- for hydrogen ions • The smaller the Kb the weaker the base

  23. Concentration and Strength • Remember, the word strong and weak acids and bases refers to the number particles of the acid or base that completely dissociate into their respective ions in solution • Concentration and dilute refer to how many moles of an acid or base is diluted in a constant volume of solution • Even though an acid may be “weak”, if it is highly concentrated, it will result in much lower pH of the solution it is dissolved in than a dilute solution of the same weak acid

  24. Calculating Dissociation Constants • Disassociation constants are calculated from experimental data • To find the Ka of weak acid or the Kb of a weak base, substitute the measured concentrations of all the substances present at equilibrium into the expression for Ka or Kb • A 0.1000M solution of ethanoic acid is only partially ionized and has a pH of 2.87. What is the acid dissociation constant (Ka) or ethanoic acid

  25. Neutralization ReactionsAcid Base Reactions • Reactions in which an acid and a base react in an aqueous solution to produce a salt and water are called neutralization reactions HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)

  26. Titrations • Titration - the process used to determine the concentration of solution (often an acid or base) in which a solution of known concentration (the standard) is added to a measured amount of the solution of unknown concentration until an indicator signals the end point • In titrations, it is important to know the mole ratios that the acid and base in question react • When an acid and base mix, the equivalence point is when the number of moles of hydrogen ions equals the number of moles hydroxide ions giving a pH of 7 • HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) Mole ratio 1:1 • H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l) Mole ratio 1:2 • 2HCl(aq) + CaOH(aq) → CaCl2(aq) + 2H2O(l) Mole ratio 2:1

  27. Steps in Titrating a Neutralization Reaction • A measured volume of an acid solution of unknown concentration is added to a flask • Several drops of the indicator are added to the solution while the flask is gently swirled • Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color • The solution of known concentration is called the standard solution • The point at which the indicator changes color is the end point • The point of neutralization is the end point aka equivalence point of the titration

  28. Titration of Strong Acid with a Strong Base

  29. Problem • A 25 ml solution of H2SO4 is completely neutralized by 18 ml of 1.0M NaOH. What is the concentration of the H2SO4 solution?

  30. Salts in SolutionSalt hydrolysis • Salt consist of anion from an acid and a cation from a base • Solutions of many salts are neutral while other salt solutions are not