1. Acids/Bases/Salts Properties

2. Properties  electrolytes  turn litmus red  sour taste  react with metals to form H 2 gas  slippery feel  turn litmus blue  bitter taste  vinegar, soda, apples, citrus fruits  ammonia, lye, antacid, baking soda  electrolytes  pH less than 7  pH more than 7

3. Naming  Acids are composed of hydrogen (H + ) followed by an anion (negative ion).  If the acid formula contains oxygen in the anion, such as in H 2 SO 4, it is known as an oxyacid.

4. 3 Rules To Naming Acids  If H + anion ending in –ide: Acid name is “hydro_____ic acid” Take the root from the anion name and fill in the blank. Take the root from the anion name and fill in the blank.

5. Acid Naming Example Example:HCl Example:HCl Cl is the anion, its name is chlorideCl is the anion, its name is chloride Name of acid is: hydrochloric acidName of acid is: hydrochloric acid Example: HF Example: HF F is the anion, its name is fluorideF is the anion, its name is fluoride Name of acid is: hydrofluoric acidName of acid is: hydrofluoric acid

6. 3 Rules To Naming Acids  H + anion ending in –ate: Acid name is “_____ic acid” Take the root from the anion name and fill in the blank. Take the root from the anion name and fill in the blank. “What I ATE was ICky” “What I ATE was ICky”

7. Acid Naming Example  Example:HNO 3 NO 3 1- is the anion, its name is nitrate NO 3 1- is the anion, its name is nitrate Name of acid is: nitric acid Name of acid is: nitric acid  Example: H 2 CO 3 CO 3 2- is the anion, its name is carbonate CO 3 2- is the anion, its name is carbonate Name of acid is: carbonic acid Name of acid is: carbonic acid

8. Exceptions  Sulfate (SO 4 2- ) Root is not sulf, but sulfur Root is not sulf, but sulfur Sulfuric acidSulfuric acid  Phosphate (PO 4 3- ) Root is not phosph, but phosphor Root is not phosph, but phosphor Phosphoric acidPhosphoric acid

9. 3 Rules To Naming Acids  H + anion ending in –ite: Acid name is “_____ous acid” Take the root from the anion name and fill in the blank. Take the root from the anion name and fill in the blank. “Don’t bITE; it’s infectiOUS” “Don’t bITE; it’s infectiOUS”

10. Acid Naming Example  Example: HNO 2 NO 2 1- is the anion, its name is nitrite NO 2 1- is the anion, its name is nitrite Name of acid is: nitrous acid Name of acid is: nitrous acid  Example: HClO 2 ClO 2 1- is the anion, its name is chlorite ClO 2 1- is the anion, its name is chlorite Name of acid is: chlorous acid Name of acid is: chlorous acid

11. Writing Formulas  Just work backwards, using the three rules. You must criss-cross charges, just like with ionic compounds You must criss-cross charges, just like with ionic compounds H + + anion- H + + anion-

12. Acid Writing Example  Hydrobromic acid anion is bromide (Br -) anion is bromide (Br -) formula is HBr formula is HBr  Acetic acid anion is acetate (C 2 H 3 O 2 - ) anion is acetate (C 2 H 3 O 2 - ) formula is HC 2 H 3 O 2 formula is HC 2 H 3 O 2  Phosphorous acid anion is phosphite (PO 3 3- ) anion is phosphite (PO 3 3- ) formula is H 3 PO 3 formula is H 3 PO 3

13. Naming  Hydroxide bases are composed of a cation (positive ion) followed by hydroxide (OH-).  Naming bases is much simpler than naming acids. Name the cation and then add “hydroxide.”

14. Exception!!!!  Memorize: NH 3 = ammonia

15. Base Naming Example  NaOH name of base: sodium hydroxide name of base: sodium hydroxide  Mg(OH) 2 name of base: magnesium hydroxide name of base: magnesium hydroxide  Fe(OH) 2 name of base: iron (II) hydroxide name of base: iron (II) hydroxide

16. Writing Formulas  Writing base formulas is straightforward.  Find the formula for the cation and then add OH -.  Remember to criss-cross the formula so that the charges equal out to zero.

17. Base Writing Example  potassium hydroxide cation is potassium (K + ) cation is potassium (K + ) formula is KOH formula is KOH  calcium hydroxide cation is calcium (Ca 2+ ) cation is calcium (Ca 2+ ) formula is Ca(OH) 2 formula is Ca(OH) 2  aluminum hydroxide cation is aluminum (Al 3+ ) cation is aluminum (Al 3+ ) formula is Al(OH) 3 formula is Al(OH) 3

18. Electrolytes  Solutions that have “free ions” in solution.  Those “free ions” conduct electricity.  Found in sports drinks because body cells need ions to do work Ex: Sodium/Potassium Pump Ex: Sodium/Potassium Pump

19. Strong Acids  Strong because there is a strong attraction between the acid and the water molecule.  Every acid molecule ionizes (dissolves)

20. Strong Acids  Perchloric Acid HClO 4  Sulfuric Acid H 2 SO 4  Nitric Acid HNO 3  Hydriodic Acid HI  Hydrobromic Acid HBr  Hydrochloric Acid HCl

21. Strong Acids  If it is not listed. it is considered a weak acid.  Weak- they remain as molecules, they don’t ionize in water.

22. Strong Bases  Strong because there is a strong attraction between the base and the water molecule.  Every base molecule ionizes (dissolves)

23. Strong Bases  Lithium hydroxide, LiOH  Calcium hydroxide, Ca(OH) 2  Sodium hydroxide, NaOH  Strontium hydroxide, Sr(OH) 2  Potassium hydroxide, KOH Barium hydroxide,Barium hydroxide, Ba(OH) 2 Magnesium hydroxide,Magnesium hydroxide, Mg(OH) 2

25. Strong Bases  If it is not listed. it is considered a weak base.  Weak- they remain as molecules, they don’t ionize in water.

26. Indicators  Chemical dyes that change color as pH changes.  Different indicators change colors at different pH levels choose an indicator that will show a color change at the pH that you are interested in. choose an indicator that will show a color change at the pH that you are interested in.  Indicators can be on a strip of paper called pH or litmus paper called pH or litmus paper  Other indicators can be added to the solution directly. Some indicators change color more than once and can be added to solutions so that we can see what is happening over time. Some indicators change color more than once and can be added to solutions so that we can see what is happening over time.

27. pH scale 0 7 INCREASING ACIDITY NEUTRAL INCREASING BASICITY 14

29. pH  Whether or not a solution is acidic, basic, or neutral depends on the balance of H + and OH - ions:  Neutral: [H + ] = [OH - ]  Acid: [H + ] > [OH - ]  Base:[H + ] < [OH - ]

30. pH  pH is the negative base 10 logarithm of the hydrogen ion concentration:  pH = - log 10 [H + ]

31. pH calculations (#1)  The sum of the pH and the pOH always equals 14.  pH + pOH = 14

32. Example  Find the pH, the pOH = 5.3  pH + pOH = 14 pH = 14 - 5.3 pH = 8.7 pH = 8.7

33. pH calculations (#2)  pH is the expression of the acidity or alkalinity of a solution in terms of its hydronium ion concentration.  pH = - log [H 3 O + ]

34. Example  Calculate the pH, if the [H 3 O] = 2.4 X 10 -6 M  pH = - log [H 3 O + ] = - log( 2.4 X 10 -6 ) = - log( 2.4 X 10 -6 ) = -(-5.6) = -(-5.6) = 5.6 = 5.6

35. pH calculations (#3)  Use an identical equation to calculate pOH.  pOH = - log [OH - ]

36. pH calculations (#4)  Use the reverse of the equation to calculate the [H 3 O + ] when pH is known.  [H 3 O + ] =10 (-pH) *** 2 nd log on the calulator! = 10 -pH

37. Example  Calculate the [H 3 O + ], if the pH is 4.71.  [H 3 O + ] = 10 (-pH) = 10 (- 4.71) = 10 (- 4.71) = 1.95 X 10 -5 M = 1.95 X 10 -5 M

38. pH calculations (#5)  Use identical process for [OH - ] when pOH is known.  [OH - ] = 10 (-pOH)

39. Acid/Bases/Salts Neutralization/Titrations

40. A. Neutralization  Chemical reaction between an acid and a base.  Products are a salt (ionic compound) and water.

41. A. Neutralization ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H 2 O HC 2 H 3 O 2 + NaOH  NaC 2 H 3 O 2 + H 2 O Salts can be neutral, acidic, or basic. Salts can be neutral, acidic, or basic. Neutralization does not mean pH = 7. Neutralization does not mean pH = 7. weak strong neutral basic

42. Neutralization  HCl + NaOH  HNO 3 + Ca(OH) 2  NaCl + H 2 O Ca(NO 3 ) 2 + H 2 O Hydrocholoric acid Sodium hydroxide Sodium chloride Water Nitric acid Calcium hydroxide Calcium nitrate water

43. B. Titration  Titration Analytical method in which a standard solution is used to determine the concentration of an unknown solution. Analytical method in which a standard solution is used to determine the concentration of an unknown solution. standard solution unknown solution

44.  End Point – point at which an indicator changes color during a titration point at which an indicator changes color during a titration  Equivalence point Point at which equal amounts of H 3 O + and OH - have been added. Point at which equal amounts of H 3 O + and OH - have been added. when mole ratio exactly equals mole ratio required by reaction when mole ratio exactly equals mole ratio required by reaction Determined by… Determined by… indicator color changeindicator color change B. Titration dramatic change in pHdramatic change in pH

45. B. Titration moles H 3 O + = moles OH - M  V  n = M  V  n M:Molarity V:volume n:# of H + ions in the acid or OH - ions in the base

46. B. Titration  42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H 2 SO 4. Find the molarity of H 2 SO 4. H3O+H3O+H3O+H3O+ M = ? V = 50.0 mL n = 2 OH - M = 1.3M V = 42.5 mL n = 1 MV# = MV# M(50.0mL)(2)=(1.3M)(42.5mL)(1) M = 0.55M H 2 SO 4