1. Acids/Bases/Salts Properties

2. Common Acids Lacticsour milk Aceticvinegar Phosphorictart taste in soda Citriccitrus fruits Malicapples Tartaricgrapes Formicant bites

3. Common Base Ammoniawindow cleaner Sodium hydroxidelye (drain and oven cleaners) Sodium BicarbonateBaking soda antacid Milk of magnesiaantacid

4. 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

5. Acid nomenclature –naming acids  Binary acids – is an acid that contains only two different elements. Acids are composed of hydrogen (H + ) followed by an anion (negative ion). Acids are composed of hydrogen (H + ) followed by an anion (negative ion).  Oxyacids- is an acid that is a compound of hydrogen, oxygen and third element (usually a non- metal)  If the acid formula contains oxygen in the anion, such as in H 2 SO 4, it is known as an oxyacid.

6. 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.

7. 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

8. 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”

9. 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

10. 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

11. 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”

12. 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

13. Strength of Acids SSSStrong Acid - is one that ionizes completely in aqueous solution. - completely dissociates. - a strong acid is a strong electrolyte. - increases with increasing polarity and decreasing bond energy.

14. WWWWeak acids - are weak electrolytes - it contains hydronium ions, anions and dissolved acid molecules - Organic acids (acidic carboxyl group) - COOH

15. Aqueous solutions for base  Most bases are ionic compounds containing metal cations and the hydroxide anion.  When a base completely dissociates in water to yield aqueous OH - ions, the solution is called alkaline. NaOH (s) water Na + (aq) + OH - (aq)

16. Aqueous solutions for base  Not all bases are ionic compounds.  Ammonia is one example because it produces hydroxide ions when it reacts with water molecules. NH 3 (g) + H 2 0 (l) NH 4 + (aq) + OH - (aq)

17. Strength of bases  The strength of the base depends on the extent to which it dissociates to its ions.  Strong bases are strong electrolytes  Bases that are not very soluble do not produce a large number of hydroxide ions when added to water

18. 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

19. pH and pH scale 0 7 INCREASING ACIDITY NEUTRAL INCREASING BASICITY 14 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 - ]

20. pH and pOH Calculations  pH is the negative base 10 logarithm of the hydrogen ion concentration: pH = - log 10 [H + ]  pH is the expression of the acidity or alkalinity of a solution in terms of its hydronium ion concentration. pH = - log [H 3 O + ] The sum of the pH and the pOH always equals 14. pH + pOH = 14 pH + pOH = 14

21. Example  Calculate the pH, if the [H 3 O] = 2.4 X 10 -6 M [H 3 O] = 2.4 X 10 -6 M pH = - log [H 3 O + ] 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  Find the pH, the pOH = 5.3 pH + pOH = 14 pH + pOH = 14 pH = 14 - 5.3 pH = 8.7 pH = 8.7

22. pH calculations  Use the reverse of the equation to calculate the [H 3 O + ] when pH is known. [H 3 O + ] = antilog (-pH) [H 3 O + ] = antilog (-pH) = 10 -pH  Use an identical equation to calculate pOH. pOH = - log [OH-] pOH = - log [OH-] Calculate the [H3O+], if the pH is 4.71. [H3O+] = antilog (-pH) [H3O+] = antilog (-pH) = antilog (-4.71) = antilog (-4.71) = 1.95 X 10-5 M = 1.95 X 10-5 M

23. Acid- Base Theories TypeAcidBase Arrhenius(Traditional) H + or H 3 0 + producer OH – OH –producer Bronsted-Lowry proton donor proton acceptor Lewis electron –pair acceptor electron- pair donor

24. Arrhenius Acids and Bases Arrhenius acid – is a chemical compound that increases the concentration of hydrogen ions H +, in an aqueous solution. Arrhenius base – is a substance that increases the concentration of hydroxide ions OH -, in aqueous solution. Acid-Bases Theories

25. Bronsted-Lowry acid -Is a molecule or an ion that is a proton donor (H + ) - Donates proton to water HCl + H 2 O H 3 O + + Cl - HCl + NH 3 NH 4 + + Cl – In both reaction HCl is a Bronsted-Lowry acid

26. Bronsted-Lowry Acid Water could be a Bronsted-Lowry acid also. H 2 O (l) + NH 3 (g) NH 4 + (aq) + OH - (aq) The water donates a hydrogen ion (proton) to the ammonia molecule.

27. Bronsted-Lowry Base A molecule or ion that is a proton acceptor (H + ) A molecule or ion that is a proton acceptor (H + ) H 2 O (l) + NH 3 (g) NH 4 + (aq) + OH - (aq) H 2 O (l) + NH 3 (g) NH 4 + (aq) + OH - (aq) Ammonium (NH 4 ) is a Bronsted-Lowry Base because it accepts the proton from the water.

28. Bronsted-Lowry Acid-Base Reaction  Protons are transferred from one reactant (the acid) to another (base). H 2 O (l) + NH 3 (g) NH 4 + (aq) + OH - (aq

29. Bronsted-Lowry Acid-Base  Monoprotic Acids – an acid that can donate only one proton (HCl, HNO 3 )  Polyprotic Acids – an acid that can donate more than one proton per molecule. - Diprotic – can donate two protons per molecule (Sulfuric acid, H 2 SO 4 ) - Triprotic – can donate three protons per molecule(Phosphoric acid, H 3 PO 4 )

30. Acid –Base Reaction The Bronsted-Lowry definition s of acids and bases provide the basis for studying proton (H + ) transfer reaction. Suppose a Bronsted-Lowry acid gives up a proton, the remaining ion or molecule can re-accept that proton and can act as a base- a conjugate base.

31. Acid –Base Reaction  Conjugate base- the ion or molecule that remains after a Bronsted-Lowry acid has given up a proton is the conjugate base of that acid.  Conjugate acid- the ion or molecule that is formed when a Bronsted-Lowry base gains a proton is the conjugate acid of that base.

32. Acid –Base Reaction HF + H 2 O F - + H 3 O + Acid Base Conjugate Base Conjugate Acid The species remaining after a Brønsted-Lowry acid gives up its proton is the conjugate base of that acid: Take off one H from the acid. The species remaining after a Brønsted-Lowry base accepts its proton is the conjugate acid of that base: Take off one H. Add an H to the base.

33. Acid –Base Reaction HCO 3 - (aq) + H 2 O (l) H 2 CO 3 (aq) + OH - (aq) base acid conjugate acid conjugate base (proton acceptor) + H 2 O (l) F - (aq) + H 3 O + (aq) HF (aq) + H 2 O (l) F - (aq) + H 3 O + (aq) acid base conjugate base conjugate acid (proton donor)

34. Strength of conjugate acids and bases  The stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid.  The weaker an acid is, the stronger its conjugate base; the weaker a base is, the stronger its conjugate acid.

35. Lewis Acids and Base  A Lewis acid is an atom, ion, or a molecule that accepts an electron pair to form a covalent bond.  A Lewis base is an atom, ion or a molecule that donates an electron pair to form a covalent bond.

36. Lewis Acids and Base H + (aq) + : NH 3 (aq) [H-NH 3 + ] (aq) H + (aq) + : NH 3 (aq) [NH 4 + ] (aq) A bare proton (H + ) is a Lewis acid in reactions in which it forms a covalent bond.

37. Lewis Acids and Base.... BF 3 (aq) + : F : - (aq) BF 4 – (aq).... An anion is a Lewis Base in a reaction in which it forms a covalent bond by donating an electron pair. An anion is a Lewis Base in a reaction in which it forms a covalent bond by donating an electron pair.

38. 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.

39. Acid/Bases/Salts Neutralization/Titrations

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

41. 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   H 2 SO 4 + KOH   HNO 3 + Ca(OH) 2  NaCl + H 2 O K 2 SO 4 + H 2 O Ca(NO 3 ) 2 + H 2 O +1 -2 +1 +2 Hydrocholoric acid Sodium hydroxide Sodium chloride Water Sulfuric acid Potassium hydroxide Potassium sulfate Water Nitric acid Calcium hydroxide Calcium nitrate water 2 2 22

43. 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 Titration dramatic change in pHdramatic change in pH

45. 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. 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