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Acid Base Definitions

Acid Base Definitions. Originally recognized by properties like taste, feel, reactions with indicators Acids taste sour and turn blue litmus red Bases are bitter, feel slippery, and turn red litmus blue. Arrenhius Definition. Very limited

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Acid Base Definitions

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  1. Acid Base Definitions • Originally recognized by properties like taste, feel, reactions with indicators • Acids taste sour and turn blue litmus red • Bases are bitter, feel slippery, and turn red litmus blue

  2. Arrenhius Definition • Very limited • Very few substances could be classified by this definition. • Acids produce excess Hydrogen ions when added to water • Bases produce excess Hydroxide ions when added to water

  3. Bronsted-Lowry Definition • Most frequently used • Acids are proton donors (give off an H+) • Bases are proton acceptors

  4. Strong Acids • Acid that dissociates completely in water • 100% of the sample breaks apart into ions • Seven strong acids • a.HClO4 - Perchloric acid • b.HClO3 - Chloric acid • c.H2SO4 - Sulfuric acid • d.HNO3 - Nitric acid • e.HCl - Hydrochloric acid • f.HBr - Hydrobromic acid • g.HI - Hydroiodic acid

  5. Strong Base • Base that dissociates completely in water • 100% of the sample breaks apart into ions • Strong bases • Hydroxides of the metals in group 1A and 2A (not Be or Mg)

  6. Acid-Base Equilibrium • A system at equilibrium can be described by its equilibrium constant. • For acidic systems, we call the equilibrium constant the acid dissociation constant and is symbolized as Ka. • For basic systems, the equilibrium constant is known as the base dissociation constant and is symbolized as Kb.

  7. Acid-Base Equilibrium • Let’s consider water in equilibrium with hydroxide and hydronium ions • The balanced equation is 2H2O (l)  H3O1+ (aq) + OH1- (aq) • The equilibrium constant would be written as Keq = [H3O1+] [OH1-]

  8. Acid-Base Equilibrium • The equilibrium constant for water has a set value, 1.0 x 10-14, and is referred to Kw • In pure water, the concentration of H3O1+ and OH1- are equal, and the product of the two concentrations is always equal to 10-14. • Thus in pure water, [H3O1+] = [OH1-] = 1.0 x 10-7 • If the concentration of one of these two ions changes, the concentration of the other ion must also change. For example, if the [H3O1+] = 10-3, then the [OH1-] = 10-11.

  9. Acid-Base Equilibrium • It is important to remember that when you are working with a system involving water, the water always contributes to the [H3O1+] and [OH1-] • The contribution of the water can be ignored when working with strong acids or if the [H3O1+] from the water is < or = to 1% of the contribution of the weak acid. • For strong acids, the [H3O1+] will be equal to the concentration of the acid. • For example, a 0.1 M solution of HCl would have [H3O1+] = 0.1 M • The same is true for strong bases and the [OH1-]

  10. Summary • Water • 2H2O  H3O1+ + OH1- • Kw = [H3O1+] [OH1-] = 1.0 x 10-14 • [H3O1+] = [OH1-] = 1.0 x 10-7 (pure water) • If [H3O1+] increases, [OH1-] will decrease • You must remember that water contributes to the [H3O1+] for equilibrium • Can be ignored for strong acids

  11. Weak acid or weak base • Do NOT dissociate completely in water • In water, establishes equilibrium between the molecular form and ionic form • Any acid or base that is not a strong acid or base is weak • Example: Acetic acid HC2H3O2 + H2O  C2H3O21- + H3O1+

  12. Conjugate base • Ion that is formed when an acid donates a Hydrogen ion (proton) • Examples: Acid Conjugate base H2SO4 HSO41- HNO3 NO31- HC2H3O2 C2H3O21-

  13. Conjugate base • If original acid is a strong acid, the conjugate base is so weak that is does not behave as a base. • It behaves as a neutral species • If original acid is a weak acid, the conjugate base behaves as a weak base • (Important for equilibrium considerations later)

  14. Conjugate acid • Ion that is formed when a base accepts a Hydrogen ion • Examples: Base Conjugate acid NaOH H2O NH3 NH41+

  15. Conjugate acid • If original base is a strong base, the conjugate acid is so weak that is does not behave as an acid. • It behaves as a neutral species • If original base is a weak base, the conjugate acid behaves as a weak acid

  16. Direction of acid-base reactions • The stronger acid and base will react with each other to yield the weaker acid and base Example: HClO4 + H2O  ClO41- + H3O1+ stronger acid stronger base weaker base weaker acid

  17. Acidic Hydrogen • Hydrogen that will be donated or removed • Not all hydrogen’s in an acid can be donated • This is especially true for organic acids • In oxy acids, the hydrogen attached to an oxygen is the acidic hydrogen • Example: Acetic acid • Most acids are monoprotic (donate one H) • Some acids are polyprotic (i.e. H2SO4)

  18. Autoionization of water • The transfer of a hydrogen ion from one water molecule to another water molecule, • results in the formation of a hydroxide ion and a hydronium ion. • Equation: 2 H2O  H3O1+ + OH1- • Equal amounts of hydroxide and hydronium are formed, • so water remains neutral.

  19. Amphoteric Substances • A species that can behave as either an acid or a base • Water is the best example of an amphoteric substance

  20. Amphoteric Substances • Many aluminum compounds are also amphoteric: • Base: Al(OH)3 + 3HCl → AlCl3 + 3H2O • Acid: Al(OH)3 + NaOH → NaAl(OH)4

  21. pH system • The pH of a system is an indication of the [H3O1+]. While it is based on the autoionization of water, it works for all acid-base systems.

  22. pH system • Definitions: • pH = -log [H3O1+] • pOH = -log [OH1-] • pKw = -log Kw Since Kw = 1.0 x 10-14, pKw = -log (1.0 x 10-14) = 14 • Kw = [H3O1+] [OH-1]  pKw = pH + pOH = 14

  23. pH Strong Acids and Bases • [H3O1+] = initial concentration of acid • To find the pH of a strong acid, use the initial concentration of the acid as the concentration of H3O1+ • [OH1-] = initial concentration of base • pH (strong acid) = -log (initial conc.)

  24. 14.5 pH of weak acids and bases • Weak acids/bases do NOT dissociate completely • To find pH, you must first find the [H3O1+] • This is done by setting up an ICE chart!!! • Weak acid equilibrium • HA  H1+ + A1- • Ka = [H+] [A-] / [HA]

  25. pH weak acids and bases • Weak base equilibrium • B + H20  BH1+ + OH1- • Kb = [BH+] [OH-] / [B]

  26. Weak acid equilibrium • Percent ionization = Degree of dissociation • Amount of substance that breaks apart into ions • If an acid has an initial concentration of 0.1 M and a 1 % ionization, [H+] = 1% of 0.1 M • We can calculate the value of the equilibrium constant, as well as equilibrium concentrations for weak acid systems. Let’s try some examples!

  27. Polyprotic Acids • Polyprotic Acids – more than one acidic hydrogen • Dissociate in a stepwise fashion • Ka values assigned for each dissociation step • The first step in the dissociation happens completely before the next step begins. • When solving problems involving polyprotic acids, you can sometimes ignore the contribution of the subsequent dissociations because it is ≤ 1% of the dissociation of the first step. • Let’s try a couple

  28. Acid/Base Properties of Salts • Review 1st Sem. Lab Notes • Ions are often modified when dissolved in solution. Examine the photos of Fe(III) salts and solutions. Why do they have different colors?

  29. Acid/Base Properties of Salts • Fe(NO3)3.6H2O contains pink Fe(H2O)63+ • Solutions may hydrolyze to give yellow Fe(H2O)5OH2+ or even reddish brown Fe(H2O)3(OH)3 • FeCl3.6H2O contains ions such as yellow Fe(H2O)5Cl2+

  30. Hydrolysis • Hydrolysis is more important for more highly charged ions

  31. Hydrolysis • Highly charged metal ions (> +3) cause pH shifts due to hydrolysis: Fe(H2O)63+ + H2O ⇌ Fe(H2O)5OH2+ + H3O+ • Many other salts also cause pH shifts when dissolved in water. These salts contain the conjugate acid or base of a weak base or acid. NH4Cl: NH4+ + H2O ⇌ NH3 + H3O+ pH<7 NaCH3CO2: CH3CO2- + H2O ⇌ CH3CO2H + OH - pH>7 • These reactions are called hydrolysis

  32. Hydrolysis • Hydrolysis is not observed with ions derived from strong acids or bases: Cations of group I and II (except Be2+) Anions: Cl-, Br-, I-, NO3-, ClO4- • Hydrolysis is observed for: Cations with charge > +3 Transition metal +2 ions Some post-transition metal ions with high charge • Common for Fe3+, Cr3+, Al3+, Zn2+, Cu2+, Bi3+, Pb4+

  33. Hydrolysis • See table below for values of Ka for metal ions. Na+ 95 pm 3.3 x 10-15 Li+ 60 pm 1.5 x 10-14 Be2+ 31 pm 3.2 x 10-7 Mg2+ 65 pm 3.8 x 10-12 Ba2+ 135 pm 1.5 x 10-14 Cr3+ 69 pm 9.8 x 10-5 Zr4+ 78 pm 6.0 x 10-1 • Greater values of Ka for ions with larger charge and smaller size.

  34. Hydrolysis of Salts • Cation/Anion from: • Strong base, NaCl strong acid no hydrolysis pH = 7 • Strong base, LiCN weak acid anion hydrolysis pH > 7 • Weak base, NH4Cl strong acid cation hydrolysis pH < 7 • Weak base, NH4CN weak acid cation and anion hydrolysis pH depends on relative Ka and Kb

  35. Hydrolysis of Salts • Calculate the pH of a solution the same as for any weak acid or weak base, using the appropriate Ka or Kb for the equilibrium constant.

  36. H3O+ OH- Hydrolysis of Salts • Is a solution of NH4OCN acidic or basic? • NH4+ + H2O ⇌ NH3 + H3O+ Ka = 1.0 x 10-14/1.76 x 10-5= 5.7 x 10-10 • OCN- + H2O ⇌ HOCN + OH- Kb = 1.0 x 10-14/3.46 x 10-4 = 2.9 x 10-11 • Produces more H3O+ than OH-, so the solution is acidic.

  37. Effect of Structure on Acid-Base Properties • 1. Bond Polarity (with H) • As polarity increases, acid strength increases • The more polar a molecule is, the easier it will be to remove the acidic hydrogen. • More of the sample will dissociate

  38. Effect of Structure on Acid-Base Properties • 2. Bond Strength • As bond strength increases, acid strength decreases • It is more difficult to remove the acidic hydrogen when the bond strength increases. • Less of the sample will dissociate if it is more difficult to remove the hydrogen, leading to a weaker acid.

  39. Effect of Structure on Acid-Base Properties • Atomic radius • Larger atoms form weaker bonds.stronger acids • As mentioned, weaker bonds result in stronger acids. • The bigger the atom, the weaker the bond • The weaker the bond, the easier to remove the H • Acids from the same family • Elements at the bottom of the column are stronger • HI > HBr > HCl > HF

  40. Effect of Structure on Acid-Base Properties • Electronegativity • As electronegativity increases, bond polarity also increases • Acids formed with highly electronegative elements will be stronger than acids formed with elements that have low electronegativity values. • Acids in the same period • Elements to the right will form stronger acids • HF > H2O > NH3

  41. Effect of Structure on Acid-Base Properties • Number of acidic hydrogens • The Neutral acid is always the strongest • With each hydrogen that is removed, the remaining acid is much weaker • H3PO4 > H2PO41- > HPO42- > PO43-

  42. Effect of Structure on Acid-Base Properties • Oxy acids • a) Electronegativity of nonmetal: • The greater the electronegativity, the stronger the acid • when comparing oxy acids of elements in the same family, the element at the top of the column will form the strongest oxy acid • HClO > HBrO > HIO

  43. Effect of Structure on Acid-Base Properties • Oxy acids • b) Number of oxygen atoms: • This is true because the addition of oxygen atoms increases the difference in electronegativity between the oxygen atoms and the central atom, which in turn causes the molecule to be more polar • More oxygen atoms = stronger acid • More oxygen atoms causes the molecule to be more polar • HClO4 > HClO3 > HClO2 > HClO

  44. Effect of Structure on Acid-Base Properties • Oxides that react with water to produce bases are called basic oxides or base anhydrides. • These are generally metal oxides

  45. Metal (basic) oxides • Sodium oxide reacts with water to produce the strong soluble base sodium hydroxide: Na2O (s) + H2O (l)  NaOH (aq) • Calcium oxide ("lime") reacts with water to produce the insoluble base calcium hydroxide ("slaked lime"): CaO (s) + H2O (l)  Ca(OH)2 (s)

  46. Acid-base properties of oxides • Oxides that react with water to produce acids are called acidic oxides or acid anhydrides. • These are generally non-metal oxides.

  47. Non-Metal (Acidic) oxides • Carbon dioxide and sulfur trioxide are two such compounds. Carbon dioxide reacts with water to produce carbonic acid: CO2 (g) + H2O (l) ---> H2CO3 (aq) • Sulfur trioxide reacts with water to produce sulfuric acid: SO3 (g) + H2O (l) ---> H2SO4 (aq)

  48. Non-Metal (Acidic) oxides • Most acidic oxides are oxides of nonmetals, or of metals that are in very high oxidation states. • An example of an acidic metal oxide is CrO3, which reacts with water to produce chromic acid, H2CrO4

  49. Acid-base properties of oxides • Oxides with metalloids and oxygen are amphoteric • The acidity of oxides increases as you move up and to the right across the periodic table.

  50. Lewis Acids-Base Model • Lewis Definition – includes the greatest number of compounds • Acids are electron pair acceptors • Bases are electron pair donors • An acid and a base come together to form a covalent bond using the electron pair donated by the base • Best example: Boron complexes

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