1. Acids and Bases Chapter 19

2. Acids pH less than 7 Sour taste Conduct electricity Reacts with metals to produce hydrogen gas Higher [H + ] concentration

3. Brønsted-Lowry Acid Some Common Acids NameFormula Hydrochloric acidHCl Nitric acidHNO 3 Sulfuric acidH 2 SO 4 Phosphoric acidH 3 PO 4 Ethanoic acidCH 3 COOH Carbonic acidH 2 CO 3 Compounds that donate protons (H + ) Polyprotic compounds

4. Acids a HYDRONIUM ION is the ion formed when a water molecule gains a hydrogen ion H 3 O +

5. To summarize... An acid is a chemical that dissolves in water to create more H + ions than there are in neutral water

6. Bases pH greater than 7 Bitter taste Slippery feel Higher [OH - ] concentration

7. Brønsted-Lowry Bases compounds that accept protons (H+)

8. Some common bases Some Common Bases NameFormulaSolubility in Water Sodium hydroxideNaOHHigh Potassium hydroxideKOHHigh Calcium hydroxideCa(OH) 2 Very low Magnesium hydroxideMg(OH) 2 Very low

9. to summarize... A base is a chemical that dissolves in water to create fewer H + ions than there are in neutral water, or more OH - ions

10. Bronsted-Lowry Theory Ammonia gains a proton, so it is a base Water donates a proton, so it is an acid

11. Conjugate Acid the ion or molecule formed when a base gains a hydrogen ion.

12. Conjugate Base the ion or molecule that remains after an acid loses a hydrogen ion.

13. Conjugate acids and bases

15. Some Conjugate Acid-Base Pairs AcidBase HClCl – H 2 SO 4 HSO 4 – H3O+H3O+ H2OH2O SO 4 2– CH 3 COOHCH 3 COO – H 2 CO 3 HCO 3 − HCO 3 – CO 3 2– NH 4 + NH 3 H2OH2OOH –

16. Amphoteric Substances A substance is AMPHOTERIC if it can act as either an acid or a base, such as water

17. Identify each reactant as a Bronsted- Lowry acid (hydrogen-ion donor) or base (hydrogen-ion acceptor). 1.HNO 3 + H 2 O  H 3 O + + NO 3 - 2.CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 COO - 3.NH 3 + H 2 O ↔ NH 4 + + OH - 4.H 2 O + CH 3 COO - ↔ CH 3 COOH + OH -

18. Identify the conjugate acid-base pairs of each reaction. 1.HNO 3 + H 2 O  H 3 O + + NO 3 - 2.CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 COO - 3.NH 3 + H 2 O ↔ NH 4 + + OH - 4.H 2 O + CH 3 COO - ↔ CH 3 COOH + OH -

19. Acid/Base Reactions Acids and bases react to produce salt and water in a neutralization reaction. H 2 SO 4 + 2NaOH  Na 2 SO 4 + 2H 2 O acid base salt water

20. Self-ionization of water the reaction in which water molecules produce ions

21. Self-ionization of water In pure water at 25°C, the concentration of hydrogen ions is only 1 × 10 −7 M. The concentration of OH − is also 1 × 10 −7 M because the numbers of H + and OH − ions are equal in pure water. Any aqueous solution in which [H + ] and [OH − ] are equal is a neutral solution.

22. Ion-product constant for water The product of the concentrations of the hydrogen ions and the hydroxide ions in water is called the ion-product constant for water (K w ). K w = [H + ] × [OH − ] = 1.0 × 10 −14

23. in acidic solutions... When acids dissolve in water hydrogen ions are released: HCl(aq) → H + (aq) + Cl − (aq) The H+ concentration is greater than the OH- concentration. A solution in which [H + ] is greater than [OH − ] is an acidic solution. – The [H + ] is greater than 1 × 10 −7 M.

24. in basic solutions... When bases dissolves in water, it forms hydroxide ions in solution. NaOH(aq) → Na + (aq) + OH − (aq) The H+ concentration is less than the OH- concentration. A basic solution is one in which [H+] is less than [OH−].

25. let’s practice... If the [H + ] in a solution is 1.0 × 10 −5 M, is the solution acidic, basic, or neutral? What is the [OH − ] of this solution?

26. The pH concept The negative logarithm of the hydrogen-ion concentration. pH = -log [H + ]

27. Calculating pH from [H + ] IF given in scientific notation in this form: 1.0 x 10 –x (x being any negative number) then the exponent on the 10 is the pH value. Example: 1.0 x 10 -3 M pH= 3

28. Example What are the pH values of the following three solutions, based on their hydrogen ion concentrations? 1.[H+]= 1.0 x 10-5 M 2.[H+]= 1.0 x 10-9 M 3.[H+]=.001 M

29. Example What are the pH values of the following three solutions, based on their hydrogen ion concentrations? 1.[H + ]= 1.0 x 10 -5 M pH =5 2.[H + ]= 1.0 x 10 -9 M pH = 9 3.[H + ]=.001 M pH= 3

30. Calculating pH from [H + ] If the coefficient is not 1.0 then do the math. Example: What is the pH of a solution with a hydrogen-ion concentration of 4.2 x10 -10 M?

31. Calculating [H + ] from pH IF the pH value is an integer (i.e. 1, 2, 3, 4, 5…) then the [H + ]= 1.0 x 10 -1,-2,-3,-4,-5… Example pH = 4 [H + ] = 1.0 x 10 -4

32. If NOT, then math [H + ] = antilog(-pH) In calculator, the antilog function: Second, log Example: The pH of an unknown solution is 6.35. What is hydrogen-ion [H + ] concentration?

33. Calculating pH from [OH - ] Use the ion-product constant for water to solve for [H + ] and then find pH. K w = [H + ] × [OH − ] = 1.0 × 10 −14 Example: What is the pH of a solution if [OH - ]= 4.0 x 10 -11 ?

34. Strong/Weak Acids Strong acids will completely ionize in aqueous solutions. Weak acids will not ionize completely in aqueous solutions CH 3 COOH (aq) + H 2 0 (l) ↔ H 3 O + (aq) + CH 3 OO - (aq)

35. Strong/Weak Acids/Bases Strong bases dissociates completely into metal ions and hydroxide ions in aqueous solutions Weak bases reacts with water to form the conjugate acid of the base and hydroxide ions

36. SubstanceFormula Hydrochloric acid Nitric acid Sulfuric acid HCl HNO 3 H 2 SO 4 Phosphoric acidH 3 PO 4 Ethanoic acidCH 3 COOH Carbonic acidH 2 CO 3 Hypochlorous acidHClO AmmoniaNH 3 Sodium silicateNa 2 SiO 3 Calcium hydroxide Sodium hydroxide Potassium hydroxide Ca(OH) 2 NaOH KOH Strong acids Strong bases

37. Identify each compound as a strong or weak acid or base. 1.NaOH 2.NH 3 3.H 2 SO 4 4.HCl

38. Nomenclature for Acids The naming system depends on the suffix of the anion: -ide, -ite, and –ate. H n X

39. Nomenclature for Acids 1.When the name of the anion ends in –ide, the acid name begins with the prefix hydro-. The stem of the anion has the suffix –ic and is followed by the word acid. Ex. HCl anion= chloride Hydrochloric acid

40. Nomenclature for Acids 2. When the anion name ends in –ite, the acid name is the stem of the anion with the suffix –ous, followed by the word acid. Ex. H 2 SO 3 anion= sulfite Sulfurous acid

41. Nomenclature for Acids 3. When the anion name ends in –ate, the acid name is the stem of the anion with the suffix –ic, followed by the word acid. Ex. HNO 3 anion= nitrate Nitric acid

42. Write the name of the following acids 1.HF 2.HNO 3 3.H 2 SO 3

43. Nomenclature for acids Look at the ions being used and write the formula like an ionic compound Ex. Hydrobromic acid anion= bromide (rule 1) H + Br - = HBr Phosphoric acid anion= phosphite (rule 2) H + PO 3 -3 = H 3 PO 3 Sulfuric acid anion= sulfate (rule 3) H + SO 4 -2 = H 2 SO 4

44. Write the formula of the following acids 1.Perchloric acid 2.Hydroiodic acid 3.Chlorous acid

45. Nomenclature for Bases The formula and names are the same as ionic compounds NaOH = sodium hydroxide Aluminum hydroxide = Al +3 OH - Al(OH) 3

46. Write the name of the following bases 1.Ba(OH) 2 2.Ca(OH) 2 3.RbOH

47. Write the formulas of the following bases 1.Cesium hydroxide 2.Beryllium hydroxide 3.Manganese hydroxide