1. Warm Up 4/14 How many protons and neutrons would an H+ ion have?

2. AcidsBases Definitions Neutralization Strong vs Weak Litmus Paper

4. Common Acids

5. Properties of Acids - Corrosive (burns skin and “eats” metals) - Sour taste - Turns blue litmus paper red - Electrolyte -pH less than 7

6. Common Bases

7. Properties of Bases -Caustic (leaves a white residue on metals) -Bitter taste -Slippery feeling -Turns red litmus blue -Electrolyte -pH greater than 7

8. Arrhenius Acid: a solution made of any solute with H + as its cation. Base: a solution of any ionic salt with hydroxide (OH - ) as its anion.

9. Bronsted Acid: molecule or ion that is a proton (H + ) donor Base: a molecule or ion that is a proton acceptor

10. Neutralization Reaction – When an acid and base react to form a neutral solution (water and a salt). Examples: NaOH + HCl → H 2 O + NaCl 2NaOH + H 2 SO 4 → 2H 2 O + Na 2 SO 4

11. Acids and bases can be classified as weak or strong. Strong acid/base- ionizes (dissociates) completely ex.: HCl + H 2 O → H + (aq) + Cl - (aq)

12. Weak acid/base- only ionizes (dissociates) partially ex.: HF + H 2 O → H + (aq) + F - (aq) + HF

15. Warm Up 3/26 Give the salt that will be formed in the following neutralization reaction: 2 HBr + Ba(OH) 4 → 2 H 2 O + ___?___

16. pH stands for “power of hydronium” Hydronium is H 3 O+ (it is the correct way to write the formula of a H+ ion that has been donated in water) pH Scale

17. Acids have a pH less than 7 7 is neutral Bases have a pH higher than 7 Each step on the pH scale is a step of 10 Ex.: Lemons (pH 2) are ten times more acidic than vinegar (pH 3) and 100 times more acidic than tomatoes (pH 4)

18. Calculating pH pH or “power of hydronium” is based on the concentration of H 3 O + ions H 3 O + is another way to write H + Formula: pH = -log [H + ]

19. Calculator Steps 1)Press the (-) button (not the minus sign button!) 2)Press the Log button 3)Enter the concentration and close the parenthesis

20. Example #1 Calculate the pH of a solution whose [H + ] = 1.0 x 10 -4

21. Example #2 Calculate the pH of a solution whose [H + ] = 3.1 x 10 -4

22. Left Side Practice Find the pH of these concentrations of H+, and tell whether it’s an acid or a base. 1) [H + ] = 1.0 ×10 −6 M 2) [H + ] = 2.19 ×10 −4 M 3) [H + ] = 9.18 ×10 −11 M 4) [H + ] = 4.71 ×10 −7 M 5) [H + ] = 1.0 M

23. Lesson: pH Calculations

24. K = [H + ][OH - ] = 1.0 ×10 −14 w Ionization Constant of Water: (We will be using this in our calculations)

25. Finding pH from [OH - ] Divide 1.0 ×10 −14 by the [OH - ] concentration. Then solve like a regular H+ problem: -log [H + ] Example: Calculate the pH of a solution with an [OH - ] of 1.0 ×10 −3.

26. pOH Calculations Formula: pOH = -log [OH - ] If given the [OH - ], then just put that number in the parenthesis. If given the [H + ], then divide it by 1.0 ×10 −14 and put the answer in the parenthesis.

27. Example: Calculate the pOH of a solution with an [OH - ] of 2.7×10 −4. Example: Calculate the pOH of a solution with an [H + ] of 9.18 ×10 −11.

28. To go from pH back to [H + ] (or from pOH back to [OH - ]) just put in: 10^- given Example: Calculate the [H + ] of a solution whose pH = 6.3

29. Calculating pOH pOH or “power of hydroxide ion” Formula: pOH = -log [OH - ]

30. Example #3 What is the pH of a 0.250M solution of KOH?

31. Example Problem #4 Calculate the pH and pOH of a 0.315M H 2 SO 4 solution.