Chapter 15 Acids and Bases

1. Chapter 15Acids and Bases Dr. Peter Warburton peterw@mun.ca http://www.chem.mun.ca/zcourses/1011.php

2. Acid-Base Theories • In defining what is considered to be an acid and what is considered to be a base, three theories have been proposed: • Arrhenius acid-base theory • Brønsted-Lowry acid-base theory • Lewis acid-base theory • We will see that each subsequent theorybuilds upon what was stated in the previous theory.

3. The Arrhenius Theory • In the Arrheniustheory of acids, anaciddissolved in waterincreases the concentration of hydronium ions H3O+ in the solution: • Arrhenius acid HA: • HA (aq) + H2O (l) H3O+ (aq) + A- (aq) • In this reaction we seeall Arrhenius acidscontainprotons (H+)that aredonatedto water

4. Arrhenius Theory – Acid Strength • In the Arrheniustheory of acids, a strong acidCOMPLETELYreacts with water, so there is no HA left at the end of the reaction: • HA (aq) + H2O (l) →H3O+ (aq) + A- (aq)

5. Arrhenius Theory – Acid Strength • In the Arrheniustheory of acids, a weak acid reacts with water until an equilibrium is reached where HA is still present in the equilibrium mixture: • HA (aq) + H2O (l) ⇌H3O+ (aq) + A- (aq)

6. The Arrhenius Theory • In the Arrheniustheory of bases, abasedissolved in waterincreases the concentration of hydroxide ions OH- in the solution: • Arrheniusbase M(OH)x: • M(OH)x(aq)Mx+ (aq) + x OH- (aq) • In this reaction we seeall Arrhenius basescontainhydroxide (OH-)

7. Arrhenius Theory – Base Strength • In the Arrheniustheory of acids, a strong baseCOMPLETELYdissociates in water, so there is noM(OH)x left at the end of the reaction: • M(OH)x(aq)→ Mx+ (aq) + x OH- (aq)

8. Arrhenius Theory – Acid Strength • In the Arrheniustheory of bases, a weak base only partially dissociates in water until an equilibrium is reached where M(OH)xis still present in the equilibrium mixture: • M(OH)x(aq)Mx+ (aq) + x OH- (aq)

9. Common strong acids and bases

10. Why do we need to improve on Arrhenius theory? • The Arrhenius theory has a drawback! • Certain compounds that DO NOT contain hydroxide can still increase the hydroxide concentration when placed in water. • Arrhenius theory does not explain this!

11. The Brønsted-Lowry Theory • The Brønsted-Lowry Theory: an acid Brønsted-Lowry Theory: an acid is any substance thatdonates protons (H+) while a base is any substance that canaccept protons. • This means that Brønsted-Lowryacid-base reactions are proton transferreactions.

12. Proton transfer reactions • Pairs of compounds are related to each other through Brønsted-Lowryacid-base reactions. These are conjugateacid-basepairs.

13. Proton transfer reactions • Generally, an acid HA has aconjugate base A- (an H+ has transferred away from the acid). Conversely, a base B has aconjugate acid BH+ (an H+ has transferred toward the base).

14. Water in BL acid-base reactions • When a Brønsted-Lowry acid is placed in water, it donates a proton to the water (which acts as a base)and establishes an acid-baseequilibrium.

15. Water in BL acid-base reactions • In the reverse reaction of the equilibrium, the acid H3O+donates a proton to the base A-to give back water and HA.

16. Water in BL acid-base reactions • When a Brønsted-Lowry base is placed in water, it accepts a proton from water (which acts as an acid)and establishes an acid-baseequilibrium.

17. Water in BL acid-base reactions • In the reverse reaction of the equilibrium, the acidBH+donates a proton to the base OH- to give back water and B.

18. Brønsted-Lowry Bases • To accept a proton(to act as a B-L base) requires a molecule to have an unshared pair of electronswhich can then beused to create a bondto theH+. • All Brønsted-Lowry bases have at least onelone pair of electrons.

19. Brønsted-Lowry Bases • In the previous reactions we’ve seen NH3 has a lone pair of electrons and can act as a B-L base. • Also, water has two lone pairs of electrons, and can act as a B-L base.

20. Amphiprotic substances • Some substances, like water, have protons that can be donated (BL acid), and lone pairs of electrons that can accept protons (BL base). This is why it can act like an acidORa base DEPENDING on the other species present. • Such substances are said to be • amphiprotic.

21. Problem • Write a balanced equation for the dissociation of each of the following Brønsted-Lowry acids in water: • a) H2SO4 • b) HSO4- • c) H3O+ • d) NH4+

22. Problem • What is the conjugate acid of each of the following Brønsted- Lowry bases? • a) HCO3- • b) CO32- • c) OH- • d) H2PO4-

23. Why do we need to improve on Brønsted-Lowry theory? • There are many reactions that behave VERY MUCH LIKE proton transfer reactions that DO NOT involve protons!

24. Lewis Acids and Bases • A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. • These definitions are more general than the Brønsted-Lowry definitions because protons DO NOT need to be involved in Lewis acid-base reactions.

25. In general LA + :LB  LA-LB • BF3 is a Lewis acid where the B atom can accept an electron pair. The N of the NH3 has a lone pair that can be donated, making NH3 a Lewis base.

26. Metal ions as Lewis acids • Many metal ions have the ability to act as Lewisacids.The ions are willing to accept electron pairs from LIGANDS (which act as Lewis bases)because this often stabilizes the ion in solution. The result is often called a complex ion.

27. Problem

28. Comparing theories • Since Arrheniusacids must contain protons, then ALL Arrhenius acids ARE ALSOBrønsted-Lowryacids. • We’ve already seen that NOT ALLBrønsted-LowrybasesareArrheniusbases.

29. Comparing theories • There areLewis acids (like metal ions)that ARE NOT • Brønsted-Lowry acids. • ALLBrønsted-Lowry bases must all have at least one lone pair of electrons, so • ALLBrønsted-Lowrybases • MUST ALSO BE • Lewisbases

30. BL acid and base strength • Brønsted-Lowry acid-baseequilibriaarecompetitions! • The equilibriumis the result of a tug-of war between the two bases in the system as theyfight for protonsgiven awayby the two acids.

31. Acid Strength and Base Strength • The acid that is “better at donating protons” • OR • the base that is “better at accepting protons” • will be found • in lesser amounts • at equilibrium • compared to theother acid(or base).

32. Strong BL acids in water • A strong acid (HA) is one that almost completely dissociatesin water (which acts as a base). • The conjugate base A-will be avery weak base.

33. Strong BL acids in water • At equilibrium, there will be very little to no HA present in the system, and the concentration of A-will essentially be the same as theinitial concentration of HA.

34. Weak BL acids in water • A weak acid (HA) is one that partially dissociatesin water(which acts asa base). The acid is not as good at donating protons to the water. • The conjugate base (A-) will be aweak base. Overall

35. Weak BL acids in water • At equilibrium, there will be some A-andH3O+present in the system. However, the concentration of HA will still be significantat equilibrium.

37. Notice that the strongest acids have the weakest conjugate bases, and the strongest bases have the weakest conjugate acids!

38. Hydrated Protons and Hydronium Ions • The ultimate proton-donor is a proton itself! • In water there is no such thing as H+. • Often more than one water molecule will crowd around the proton to give hydrates with the formula H(H2O)n+where n is 1 to 4.

39. Hydrated Protons and Hydronium Ions

40. Dissociation of Water • It is possible for one water molecule to act as an acid while another water molecule acts as a base at the same time. This leads to the self-ionization of water equilibrium: • H2O (l) + H2O (l)H3O+ (aq) + OH- (aq) • The equilibrium constant for this reaction is called the • ion-product constant for water, Kw. • Kw = [H3O+][OH-]

41. At 25 °C, Kw = 1.0 x 10-14 • so [H3O+] = [OH-] = 1.0 x 10-7 mol/L • Relatively few water molecules are dissociated at equilibrium at room temperature! • We willalways assume that • [H3O+] [OH-]=1.0 x 10-14 at 25 °C.

42. At 25 C Be Careful! • Acidic • [H3O+] > 1.0 x 10-7M • or [OH-] < 1.0 x 10-7 M • Basic • [OH-] > 1.0 x 10-7 M • or[H3O+] < 1.0 x 10-7M • Neutral • [H3O+] =[OH-] =1.0 x 10-7 M

43. At 25 C Be Careful! • We also find, since • [H3O+] [OH-]= 1.0 x 10-14= Kw • then • [H3O+] =1.0 x 10-14/ [OH-] • and • [OH-] =1.0 x 10-14/ [H3O+]

44. At 25 C

45. Problem • The concentration of OH- in a sample of seawater is • 5.0 x 10-6 mol/L. • Calculate the concentration of H3O+ ions, and classify the solution as acidic, neutral, orbasic.

46. Problem • At 50 °C the value of • Kw is 5.5 x 10-14. • What are the • [H3O+] and [OH-] • in a neutral solution • at 50 °C?

47. The pH Scale • [H3O+] in water can range from • very small (strongly basic) • to very large (strongly acidic) • it is sometimes easier to use a negative logarithmic (power of 10) scale to express [H3O+] with a term we call the pH of a solution. • pH = - log [H3O+]

48. At 25 C Be Careful! pH and acidity • pH>7 is basic • pH< 7 is acidic • pH=7 is neutral

49. At 25 C Be Careful! pOH and acidity • pOH< 7 is basic • pOH> 7 is acidic • pOH= 7 is neutral pOH = - log [OH-] Or [OH-] = 10-pOH

50. pH scale