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Arrhenius Acids and Bases

Arrhenius Acids and Bases. Acid : A substance that produces H 3 O + ions in aqueous solution. Base : A substance that produces OH - ions in aqueous solution. H + reacts immediately with a water molecule to give a hydronium ion. Arrhenius Acids and Bases.

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Arrhenius Acids and Bases

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  1. Arrhenius Acids and Bases Acid: A substance that produces H3O+ ions in aqueous solution. Base: A substance that produces OH- ions in aqueous solution. H+reacts immediately with a water molecule to give a hydronium ion.

  2. Arrhenius Acids and Bases • When ACIDS dissolve in water, they react with water to give hydronium ion and chloride ion: • Use curved arrows to show bonding changes:

  3. Arrhenius Acids and Bases When BASES dissolve in water, the ions separate (ions become hydrated): Metal hydroxides such as NaOH, KOH, Mg(OH)2, and Ca(OH)2behave this way. Some basesalso react with waterto produce OH-ions, as shown for ammonia:

  4. Arrhenius Acids and Bases We use curved arrows to show the transfer of a proton from water to ammonia. The first arrow shows the formation of an N-H bond, the second shows breaking of an H-O bond.

  5. Acid and Base Strength Strong acid: One that reacts completely or almost completely with water to form H3O+ ions. Strong base: One that reacts completely or almost completely with water to form OH- ions. MEMORIZE THIS TABLE Table of Strong Acids and Strong Bases

  6. Acid and Base Strength Weak acid: A substance that dissociates only partially in water to produce H3O+ ions. • Acetic acid, for example, is a weak acid. In water, only 4 out every 1000 molecules are converted to acetate ions: Weak base: A substance that only partially reacts with water to produce OH- ions. • Ammonia, for example, is a weak base:

  7. Brønsted-Lowry Acids and Bases Acid: A proton donor Base: A proton acceptor. Acid-base reaction: A proton-transfer reaction. Conjugate acid-base pair: Any pair of molecules or ions that can be interconverted by transfer of a proton.

  8. Brønsted-Lowry Acids and Bases We can use curved arrows to show the transfer of a proton from acetic acid to ammonia:

  9. Acids and Conjugate Bases

  10. Brønsted-Lowry Acids and Bases 1. An acid can be positively charged, neutral, or negatively charged; examples of each type are H3O+, H2CO3, and H2PO4-. 2. A base can be negatively charged or neutral; examples are OH-, Cl-, and NH3. 3. Acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H2CO3, and H3PO4.

  11. Brønsted-Lowry Acids and Bases • Carbonic acid, for example, can give up one proton to become bicarbonate ion, and then the second proton to become carbonate ion: 4. Several molecules and ions can function as either an acid or a base (next screen).

  12. Brønsted-Lowry Acids and Bases The HCO3- ion, for example, can give up a proton to become CO32-, or it can accept a proton to become H2CO3. A substance that can act as either an acid or a base is said to be amphiprotic. H2O is amphiprotic; it can accept a proton to become H3O+, or lose a proton to become OH- . 5. Not all hydrogen atoms are acidic. • Acetic acid, CH3COOH, for example, gives up only one proton.

  13. Brønsted-Lowry Acids and Bases 6. There is an inverse relationship between the strength of an acid and the strength of its conjugate base. • The stronger the acid, the weaker its conjugate base. • HI, for example, is the strongest acid from Group 7 and its conjugate base, I-, is the weakest base. • CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3- (bicarbonate ion).

  14. Acid-Base Equilibria • We know that HCl is a strong acid, which means that the position of this equilibrium lies very far to the right. • In contrast, acetic acid is a weak acid, and the position of its equilibrium lies very far to the left. • But what if the base is not water? How can we determine which are the major species present?

  15. Acid-Base Equilibria To predict the position of an acid-base equilibrium such as this, we do the following: • Identify the two acids in the equilibrium; one on the left and one on the right. • Identify which is the stronger acid and which is the weaker acid. • Also determine which is the stronger base and which is the weaker base, Remember that the stronger acid gives the weaker conjugate base, and the weaker acid gives the stronger conjugate base. • The equilibrium lies on the side of the weaker acid and weaker base.

  16. Acid-Base Equilibria • Identify the two acids and bases, and their relative strengths. • The position of this equilibrium lies to the right, on the side of the weaker acid and weaker base.

  17. Acid-Base Equilibria Example: Predict the position of equilibrium in this acid-base reaction: Solution: The position of this equilibrium lies toward the right.

  18. Acid Ionization Constants When a weak acid, HA, dissolves in water: The equilibrium constant expression, Ka, is: What does K mean? The stronger the acid, the bigger value for K Big K means stronger acid Little K means weaker acid

  19. Acid Ionization Constants • Acid ionization constants for weak acids are small numbers, i. e., numbers with negative exponents • Easier to work with pKa values: pKa = -logKa • Values of Ka and pKa for some weak acids are given in the following Table • Note the inverse relationship between Ka values and pKavalues

  20. Acid Ionization Constants • Note the inverse relationship between Ka values and pKa values

  21. Properties of Acids and Bases Neutralization: • Acids and bases react with each other in a process called neutralization Reaction with metals: • Strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H2 • Reaction of a strong acid with a metal is a redox reaction; the metal is oxidized to a metal ion and H+ is reduced to H2.

  22. Properties of Acids and Bases Reaction with metal hydroxides: • Reaction of an acid with a metal hydroxide gives a salt plus water. • The reaction is more accurately written as follows: • Omitting spectator ions gives this net ionic equation:

  23. Properties of Acids and Bases Important Biological Reactions • Strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O. • Strong acids react similarly with bicarbonates:

  24. Self-Ionization of Water Pure water contains a very small number of H3O+ ions and OH- ions formed by proton transfer from one water molecule (the proton donor) to another (the proton acceptor). • The equilibrium constant for the ionization of water, Kw, is called the ion product of water.

  25. Self-Ionization of Water • In pure water, H3O+ and OH- are formed in equal amounts (remember the balanced equation for the self-ionization of water). • This means that in pure water:

  26. Self-Ionization of Water • The product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14. • Example: • 0.010 mol of HCl dissolved in 1.00 liter of pure water • HCl reacts completely to give 0.010 mole of H3O+. • [H3O+] is 0.010 or 1.0 x 10-2. • Also, the concentration of hydroxide ion is 1.0 x 10-12.

  27. pH and pOH • Because hydronium ion concentrations for most solutions are numbers with negative exponents, we commonly express these concentrations as pH, where: pH = -log [H30+] • We can now state the definitions of acidic and basic solutions in terms of pH: • Acidic solution: One whose pH is less than 7.0. • Basic solution: One whose pH is greater than 7.0. • Neutral solution: One whose pH is equal to 7.0.

  28. pH and pOH • Just as pH is a convenient way to designate the concentration of H3O+, pOH is a convenient way to designate the concentration of OH-. pOH = -log[OH-] • The ion product of water, Kw, is 1.0 x 10-14 • Taking the logarithm of this equation gives: pH + pOH = 14 • Thus, if we know the pH of an aqueous solution, we can easily calculate its pOH.

  29. pH Buffers pH buffer: A solution that resists change in pH when limited amounts of acid or base are added to it. • A pH buffer is an acid or base “shock absorber.” • A pH buffer is commonly referred to simply as a buffer. • The most common buffers consist of approximately equal molar amounts of a weak acid and a salt of the weak acid; that is, approximately equal molar amounts of a weak acid and a salt of its conjugate base. • For example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer.

  30. pH Buffers • How does an acetate buffer resist changes in pH? • If we add a strong acid, such as HCl, added H3O+ ions react with acetate ions and are removed from solution: • If we add a strong base, such as NaOH, added OH- ions react with acetic acid and are removed from solution:

  31. pH Buffers Buffer pH • If we mix equal molar amounts of a weak acid and a salt of its conjugate base, the pH of the solution will be equal to the pKa of the weak acid. • If we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H3BO3), pKa 9.14, and sodium dihydrogen borate (NaH2BO3), the salt of its conjugate base.

  32. pH Buffers Buffer capacity:The amount of hydronium or hydroxide ions that a buffer can absorb without a significant change in pH. • Buffer capacity depends both its pH and its concentration.

  33. Blood Buffers The average pH of human blood is close to 7.4. • Any change greater than 0.10 pH unit in either direction can cause illness. To maintain this pH, the body uses three buffer systems: • Carbonate buffer: H2CO3 and its conjugate base, HCO3- • Phosphate buffer: H2PO4- and its conjugate base, HPO42- • Proteins

  34. Henderson-Hasselbalch Eq. Henderson-Hasselbalch equation: A mathematical relationship between: • pH • pKa of the weak acid, HA • The concentrations of HA and its conjugate base A- . It is derived in the following way: • taking the logarithm of this equation gives:

  35. Henderson-Hasselbalch Eq. Multiplying by -1 gives: • -log Ka is by definition pKa, and -log [H3O+] is by definition pH. Making these substitutions gives: • rearranging terms gives:

  36. Henderson-Hasselbalch Eq. Example: What is the pH of a phosphate buffer solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 dissolved in enough water to make 1.0 liter of solution? Solution • The equilibrium we are dealing with and its pKa are: • Substituting the concentrations in the H-H equation gives:

  37. Biochemical Buffers The original laboratory buffers for use in biochemical studies were made from simple acids and bases, such as acetic acid, phosphoric acid, and citric acid and their conjugate bases. However, many of these have severe limitations: • ---They often change their pH too much if the solution is diluted or the temperature is changed. • ---They often permeate cells in the solution thereby changing the chemistry in the interior of the cells. To overcome these short comings, N.E. Good developed a series of buffers that consist of zwitterions, molecules that to not readily permeate cell membranes.

  38. Biochemical Buffers Table 8.6 Acid and Base Forms of Some Useful Biochemical Buffers

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