1. Acids & Bases Chemistry 6.0

2. Naming Acids Review: A. Binary – H +one anion Prefix “hydro”+ anion name +“ic”acid Ex) HCl hydrochloric acid Ex) H 3 P hydrophosphoric acid B. Tertiary – H + polyatomic anion no Prefix “hydro” (oxo)end “ate” = “ic” acid end “ite” = “ous” acid Ex) H 2 SO 4 sulfuric acid Ex) H 2 SO 3 sulfurous acid

3. Properties of Acids and Bases: TasteTouch Reactions with Metals Electrical Conductivity Acidsour looks like water, burns, stings Yes- produces H 2 gas electrolyte in solution Base (alkali) bitter looks like water, feels slippery No Reaction electrolyte in solution

4. Indicators : Turn 1 color in an acid and another color in a base. A. Litmus Paper : Blue and Red An aci D turns blue litmus paper reD A B ase turns red litmus paper Blue. B. Phenolphthalein : colorless in an acid and pink in a base C. pH paper : range of colors from acidic to basic D. pH meter : measures the concentration of H + in solution

6. Neutralization: A reaction between an acid and base. When an acid and base neutralize, water and a salt (ionic solid) form. Acid + Base → Salt + Water Ex) HCl + NaOH → NaCl + HOH Reactions

7. Reactions Acids and Metals –HA + M  H 2 + MA A = anion M = metal –2HCl + Mg  H 2 + MgCl 2 Acids and Carbonates –HA + MCO 3  MA + CO 2 + H 2 O –HCl + Li 2 CO 3  2LiCl + CO 2 + H 2 O

8. Arrhenius Definition (1884): A. An acid dissociates in water to produce more hydrogen ions, H +. HCl  H +1 + Cl -1 B. A base dissociates in water to produce more hydroxide ions, OH -. NaOH  Na +1 + OH -1 C. Problems with Definition: Restricts acids and bases to water solutions. Oversimplifies what happens when acids dissolve in water. Does not include certain compounds that have characteristic properties of acids & bases. Ex) NH 3 (ammonia) doesn’t fit

9. Bronsted-Lowry Definition (1923): A. An acid is a substance that can donate hydrogen ions. Ex) HCl → H + + Cl - –Hydrogen ion is the equivalent of a proton. –Acids are often called proton donors. –Monoprotic (HCl), diprotic (H 2 SO 4 ), triprotic (H 3 PO 4 ) B. A base is a substance that can accept hydrogen ions. Ex) NH 3 + H + → NH 4 + –Bases are often called proton acceptors. C. Advantages of Bronsted-Lowry Definition Acids and bases are defined independently of how they behave in water. Focuses solely on hydrogen ions.

10. Hydronium Ion: Hydronium Ion – H 3 O + This is a complex ion that forms in water. H +1 + H 2 O  H 3 O +1 To more accurately portray the Bronsted-Lowry, the hydronium ion is used instead of the hydrogen ion.

11. STRONG Acid/Base versus WEAK Acid/Base Strength refers to the % of molecules that form IONS. A strong acid or base will completely ionize (>95% as ions). This is represented by a single (  ) arrow. HNO 3 + H 2 O  H 3 O + + NO 3 - A weak acid or base will partially ionize (<5% as ions). This is represented by a double (↔) arrow. HOCl + H 2 O ↔ H 3 O + + ClO -

12. HF < HCl < HBr < HI increasing strength 7 Strong Acids HNO 3 H 2 SO 4 HClO 3 HClO 4 HClHBr HI 8 Strong Bases LiOHNaOHKOH RbOHCsOHCa(OH) 2 Sr(OH) 2 Ba(OH) 2

13. Strength vs. Concentration Strength refers to the percent of molecules that form ions Concentration refers to the amount of solute dissolved in a solvent. Usually expressed in molarity. See “Acids: Concentration vs. Strength”

14. Conjugate Acid-Base Pairs : A pair of compounds that differ by only one hydrogen ion A.Acid donates a proton to become a conjugate base. B.Base accepts proton to become a conjugate acid. A strong acid will have a weak conjugate base. A strong base will have a weak conjugate acid.

15. Acid (A), Base (B), Conjugate Acid (CA), Conjugate Base (CB) NH 3 +H 2 O ↔NH 4 + +OH - HCl+H 2 O ↔Cl - + H 3 O + Base and Conjugate Acid are a Conjugate Pair. Acid and Conjugate Base are a Conjugate Pair. B B A A CACB CA

16. 1.H 2 O + H 2 O ↔ H 3 O + + OH − BA CA CB 2.H 2 SO 4 + OH − ↔ HSO 4 − + H 2 O A B CB CA 3.HSO 4 − + H 2 O ↔ SO 4 −2 + H 3 O + A B CB CA 4.OH − + H 3 O + ↔ H 2 O + H 2 O B A CA CB AciDonates & Bases accept

17. Ionization of Acids & Bases H 2 SO 4  2 H + + SO 4 -2 –Sulfuric acid H 3 PO 3  –Phosphorous acid Ca(OH) 2  –Calcium hydroxide 3 H + + PO 3 -3 Ca +2 + 2 OH -1

18. The Self-ionization of Water & pH 1. Water is amphoteric, it acts as both an acid and a base in the same reaction. Ex) H 2 O (l) + H 2 O (l) ↔ H 3 O + (aq) + OH - (aq) K eq = equilibrium constant = [H 3 O + ] [OH - ] Because reactants and products are at equilibrium, liquid water is not included in the equilibrium expression @ 25  C, [H 3 O + ] = 1 x 10 -7 M and [OH - ] = 1 x 10 -7 M K w = ion product constant or equilibrium constant for water K w = [H 3 O + ] [OH - ] = 1 x 10 -14 M 2 1.0 x 10 -14 M 2 = [1.0 x 10 -7 M] [1.0x10 -7 M] 1.0 x 10 -14 = [H 3 O + ] [OH - ]

19. Acids: [H 3 O + ] > 1 x 10 -7 M Bases: [OH - ] > 1 x 10 -7 M Using K w in calculations: If the concentration of H 3 O + in the blood is 4.0 x 10 -8 M, what is the concentration of OH ­ ions in the blood? Is blood acidic, basic or neutral? K w = [H 3 O + ] [OH - ] 1.0 x 10 -14 M 2 = [ 4.0 x 10 -8 M ] [OH - ] 2.5 x 10 -7 M = [OH - ] slightly basic

20. The pH scale (1909): the power of Hydrogen A.Measure of H 3 O + in solution. B.pH = -log[H 3 O + ] C.Range of pH: 0-14 pH < 7: acid pH > 7: base pH = 7: neutral D.pOH = -log[OH - ] E.pH + pOH = 14

21. D. pH = -log [H 3 O + ] E.[H 3 O + ] [OH - ] = 1.0 x 10 -14 F.pH + pOH = 14 H+H+ OH - pH 1 7 14

22. H+H+ OH - pH[H 3 O + ][OH - ] 141x10 -14 1x10 0 131x10 -13 1x10 -1 121x10 -12 1x10 -2 111x10 -11 1x10 -3 101x10 -10 1x10 -4 91x10 -9 1x10 -5 81x10 -8 1x10 -6 71x10 -7 61x10 -6 1x10 -8 51x10 -5 1x10 -9 41x10 -4 1x10 -10 31x10 -3 1x10 -11 21x10 -2 1x10 -12 11x10 -1 1x10 -13 1 14

23. Significant Digits Rule The number of digits AFTER THE DECIMAL POINT in your answer should be equal to the number of significant digits in your original number The number of digits AFTER THE DECIMAL POINT in your answer should be equal to the number of significant digits in your original number Ex -log[8.7x10 -4 M ] Ex -log[8.7x10 -4 M ] Calc Answer = 3.0604807474 Calc Answer = 3.0604807474 Sig Fig pH = 3.06 Sig Fig pH = 3.06

24. Acid-Base Titration 1. An acid-base titration is a carefully controlled neutralization reaction or redox which can determine concentration of an unknown solution. 2. To determine the concentration of an unknown substance, a standard solution is needed. This solution has a known concentration. 3. Titration curve: graph that shows how pH changes during a titration.

25. 4.An indicator, usually phenolphthalein, is used in a titration. Colorless in an acid, pink in a base. 5. The point at which enough standard solution is added to neutralize the unknown solution is called the equivalence point.

26. 6. The point at which the indicator changes color is called the endpoint. 7. Therefore: [H + ] = [OH - ] at the equivalence point

27. Ex) A volume of 50.0 mL of 0.150M sodium hydroxide neutralizes 30.0mL of a sulfuric acid solution. What is the concentration of the sulfuric acid solution? H 2 SO 4 + NaOH  Na 2 SO 4 + HOH Note: Mole ratio between acid an base in not 1:1 22 M b = 0.125M