1. Acids, Bases and Salts

2. Naming Acids

3. Naming Acids
Names of binary acids (two atoms) have the form: hydro-…ic acid.
Ex: HF is hydrofluoric acid
HCl is hydrochloric acid
Try:
H2S HI
HBr
Hydrosulfuric Acid
Hydroiodic Acid
Hydrobromic Acid

4. Naming Acids
When naming ternary acids (H and a poly-ion). identify polyatomic ion in the formula.
For polyions that end in “ate” change the ending to “ic”. (Something I “ate” made me feel “ic”.)
For polyions that end in “ite” change the ending to “ous”.(A snake b“ite” is poison“ous”) 
EX: HNO3  nitric acid
HNO2  nitrous acid
H2SO4 
H2SO3 
Sulfuric Acid
Sulfurous Acid

5. Name the following acids:
HCl
H2CO3
H3PO4 HF
HNO3
Hydrochloric Acid
Carbonic Acid
Phosphoric Acid
Hydrofluoric Acid
Nitric Acid

6. Properties of Acids
Acid Property #1. The word acid comes from the Latin word acere. which means "sour“. All acids taste sour.
Acid Property #2. In Robert Boyle wrote that acids would make a blue vegetable dye called "litmus" turn red.
Acid Property #3. Acids destroy the chemical properties of bases.

7. Properties of Acids Cont.
Acid Property #4. Acids conduct an electric current. (electrolytes)
Acid Property #5. Upon chemically reacting with an active metal. acids will evolve hydrogen gas (H2).

8. #6. Acids Have a pH less than 7

9. Properties of Bases
Base Property #1. The word "base" has a more complex history and its name is not related to taste. All bases taste bitter.
Base Property #2. Bases are substances which will restore the original blue color of litmus after having been reddened by an acid.
Base Property #3. Bases destroy the chemical properties of acids.

10. Properties of Bases Cont.
Base Property #4. Bases conduct an electric current. (electrolytes)
Base Property #5. Bases feel slippery. sometimes people say soapy. This is because they dissolve the fatty acids and oils from your skin and this cuts down on the friction between your fingers as you rub them together.

11. #6. Bases have a pH greater than 7
VIDEO

12. Properties of A Salt
A salt is the combination of a cation(+ metal ion) and an anion (- nonmetal ion).
Salts are products of the reaction between acids and bases.
Solid salts usually make crystals.
If a salt dissolves in water solution. it usually dissociates into the anions and cations that make up the salt.

13. Theories of Acid/Base Acids – produce H+ 1. Arrehenius
Bases - produce OH-
Acids – donate proton (H+)
Bases – accept protons (H+)
Acids – accept e- pair
Bases – donate e- pair
1. Arrehenius
only in water
2. Bronsted-Lowry
any solvent
3. Lewis
used in organic chemistry.
wider range of substances

14. The Acid Base Theory Brønsted-Lowry
Two chemists. independent of one another. proposed a new definition of an acid and a base.
An acid is a substance which can donate a proton (H+) to a base.
A base is a substance that can accept a proton (H+) from an acid.

15. Reactions Based on Brønsted-Lowry
HCl + H2O   H3O+ + Cl¯
HCl - this is the acid, because it has a proton (H+) available to donate.
H2O - this is the base, since it accepts the proton (H+) that the acid lost.
Now, here comes an interesting idea:
H3O+ - this is an acid, because it can donate a proton.
Cl¯ - this is a base, since it has the capacity to accept a proton.

16. HCl + H2O  H3O+ + Cl¯
Notice that each pair (HCl and Cl¯ as well as H2O and H3O+ differ by one proton (H+). These pairs are called conjugate pairs.
HNO3 + H2O  H3O+ + NO3¯
The acids are HNO3 and H3O+ and the bases are H2O and NO3¯.

17. The Bronsted-Lowry Concept
Acids and bases are identified based on whether they donate or accept H+.
“Conjugate” acids and bases differ only with the addition or elimination of one H+.

18. Bases and Conjugate Acid
Name
Conjugate acid
CH3COO-
Acetate ion
CH3COOH
Acetic acid
NH3
Ammonia
NH4+
Ammonium
H2PO4-
Dihydrogen phosphate ion
H3PO4
Phosphoric
acid
HSO4-
Hydrogen sulfate ion
H2SO4
Sulfuric acid
OH-
Hydroxide ion
H20
water
NO3-
Nitrate ion
HNO3
Nitric acid
H2O
H30+
Hydronium ion

19. When a substance can donate or accept a
proton (H+) it is an amphoteric substance:
Acting like
a base
Acting like an acid
HCO3-
+ H+
- H+
H2CO3
CO3-2
accepts H+
donates H+

20. Water is considered an amphoteric substance because it can act as either an acid or a base.
Last slide

21. Strong Acids and Bases
Strong acids are those that ionize completely in water.
The dissociation of a strong acid looks like the diagram at the right in that it dissociates completely into positive and negative ions.

22. Let’s examine the behavior of an acid. HA. in aqueous solution.
CHM 101
Let’s examine the behavior of an acid. HA. in aqueous solution.
Sinex HA
What happens to the HA molecules in solution?

23. 100% dissociation of HA HA H+ Strong Acid A-
Would the solution conduct electricity?

24. Weak Acids and Bases
Some acids and bases ionize only slightly in water.
These are considered weak.
The most important weak base is ammonia (NH3).

25. Partial dissociation of HA
Weak Acid A-
Would the solution conduct electricity?

26. At any one time. only a fraction of the molecules are dissociated.
HA  H A- HA
Weak Acid H+
At any one time. only a fraction of the molecules are dissociated. A-

27. Strong Acids
*HNO3 - nitric acid *HCl - hydrochloric acid *H2SO4- sulfuric acid *HClO4 - perchloric acid *HBr - hydrobromic acid *HI - hydroiodic acid
*HClO3- chloric acid
*Refer to solubility rules.

28. Strong Bases
*LiOH - lithium hydroxide *NaOH - sodium hydroxide *KOH - potassium hydroxide *RbOH - rubidium hydroxide *CsOH - cesium hydroxide *Ca(OH)2 - calcium hydroxide *Sr(OH)2 - strontium hydroxide *Ba(OH)2 - barium hydroxide
*Refer to solubility rules for soluble hydroxides

29. Neutralization Reactions
The word "neutralization" is used to describe the double replacement reaction of an acid plus a base because the acid and base properties of H+ and OH- are destroyed or neutralized.
In the reaction. H+ and OH- combine to form water and a salt.

30. Bases Neutralize Acids
Milk of Magnesia contains magnesium hydroxide. Mg(OH)2. which neutralizes stomach acid. HCl.
2 HCl + Mg(OH)2
MgCl2 + 2 H2O

31. Bases Neutralize Acids
In general: Acid + Base  Water + Salt
HCl NaOH  NaCl HOH
H2O
HCl Mg(OH)2  H2O + ______
MgCl2
H2SO NaOH  H2O + ______
Na2SO4

32. HCl + NaOH --> H2O + NaCl HF + Mg(OH)2 --> H2O + MgF2
When acids and bases are equal in strength and concentration. a neutral (pH = 7) solution is formed.
EX:
HCl + NaOH --> H2O + NaCl
(Strong acid) (Strong Base)
HF + Mg(OH)2 --> H2O + MgF2
(Weak acid) (Weak Base)

33. Neutral solutions have an equal number of hydroxide & hydronium ions
Acidic Neutral
Solution Solution

34. Calculations of Neutralization Reactions
Utilize equation: MAVA = MBVB
MA and VA= Molarity and Volume of Acid
MB and VB= Molarity and Volume of Base
*Must adjust the Molarity of strong acids or strong bases based on the number of moles of H+ or OH- they contribute.

35. Adjust the molarity (M)
Strong Acid: H2SO4 contributes 2 moles of H+ ions. Multiply the molarity (M) of H2SO4 by 2. Strong Bases: There are three strong soluble bases that contribute 2 moles of OH- ions. Ca(OH)2. Sr(OH)2. and Ba(OH)2 molarities (M) must be multiplied by 2.

36. Sample Problems
1. How many mL of a 1.5 M HCl acid is needed to neutralize a 500. mL 1.5 M NaOH solution?
2. How many mL of a 2.0 M H2SO4 acid is needed to neutralize a 500. mL 1.5 M KOH solution?
3. How many mL of a M HNO3 acid is needed to neutralize a 275 mL 1.5 M Ca(OH)2 solution?
4. Calculate concentration of acid if 300. mL of HBr is used to neutralize a 500. mL 2.50 M NaOH solution.

37. pH and pOH

38. pH and pOH
The scale is measured on a log scale of 0 to 14. with each unit representing a ten-fold change.

39. pH Scale Hydrogen ion concentration indicates:
The pH scale is a measure of hydrogen ion [H+] concentration (acid molarity) as well as the hydroxide [OH-] concentration (base molarity).
Hydrogen ion concentration indicates:
Acids will have a pH of 0-6
Hydroxide ion concentration indicates:
Bases will have a pH of 8-14
The higher the [H+]. the higher the acidity. the higher the [OH-]. the higher the basicity.

40. Some common pH values

41. pH = -log [H+] (acid Molarity)
Calculating pH
The concentration (M or mol/L) of H+ is expressed in powers of 10. from to 100.
Scientists use the following formula to calculate pH using acid molarity.
pH = -log [H+] (acid Molarity)
(Remember that the [ ] mean Molarity)

42. Example pH Calculations pH = -log [H+]
EX: 0.50M HCl is added to water to make a final volume of 1 liter. What is the pH of this solution?
Step 1: Identify that you have an acid (starts with an H)
Step 2: Identify [H+] = 0.50 M
Step 3: Place value in equation and solve.
pH = -log[0.50] = pH (Acidic) (pH less than 7)

43. Practice pH Calculations pH = -log [H+]
Find pH of the following solutions if [H+] is:
1.00 x 10-3 M = 3.00 pH
6.59 x 10-6 M = 5.18 pH
9.47 x M = 9.02 pH
I-phones & Calculators: -log(6.59ee-6)enter = answer (2 decimal places)

44. Calculate pH of Strong Acids pH = -log [H+]
If you have the strong polyprotic acid, H2SO4, you must adjust the molarity by the number of moles of H+ contributed.
EX: 0.250M H2SO4 is added to water to make a final volume of 1 liter. What is the pH of this solution?
Step 1: Identify [H+] = M x 2 = 0.5 M
Step 2: Place values in equation and solve.
pH = -log[0.50] = 0.30 pH (Acidic)

45. Calculate pH of Weak Acids pH = -log [H+]
Weak acids will not dissociate 100%. A dissociation factor will be included in these problems.
EX: Calculate pH of a M HNO2 solution (dissociation is 5.00%).
Step 1: Identify [H+] =
0.150 M x 5% = M
Step 2: pH = -log[0.0075] = 2.13 pH (Acidic)

46. Calculate H+ Concentration from pH
If given pH you can calculate the hydrogen ion concentration by performing the anti-log function. Calculator (10^)… I-Phone (alog)
[H+] = alog (-pH)
Calculator (10^)… I-Phone (alog)

47. Example [H+] Calculations [H+] = alog (-pH)
EX: Find the [H+] if the pH is 2.00.
Alog (-2.00) = 1.00 x 10-2M
10^(-2.00) = 1.00 x 10-2M
Find [H+] if the pH is:
6.678 pH = 2.1 x 10-7 M
2.533 pH = 2.9 x 10-3 M
10.0 pH = 1.0 x M
2.56 pH - = 2.75 x 10-3 M
Remember the unit for concentration is M.

48. pOH Just like pH can be calculated from the [H+] concentration…
pH = -log [H+]
pOH can be calculated from the [OH-] concentration…
pOH = -log [OH-]

49. pOH Just like [H+] concentration can be calculated from pH…
[H+] = alog (-pH)
[OH-] concentration can be calculated from pOH…
[OH-] = alog (-pOH)

50. Connecting pH to pOH
You can calculate the pH of a solution if you know the concentration of hydroxide ion. [OH-]
If we use the ion product constant of water we can derive this equation:
[pH] x [pOH] = 1.00 x 10-14
Working with this equation leads to:
pH + pOH = 14

51. Practice pH Calculations Using pOH
EX: Find the pH of a solution with an [OH-] of 1.0 x 10-8 M.
Step 1: Calculate pOH by using equation:
pOH = -log[OH-] = -log(1.0 x 10-8 )= 8 pH
Step 2: Subtract the pOH from 14 to find pH:
pH = 14 - pOH = = 6 pH

52. Practice pH Calculations Using pOH
Find the pH of the following solutions with [OH-] of:
1.00 x 10-4 M
-log[1.00x10-4] = 4 pOH, 14-4=10 pH
2.64 x M
-log[2.64x10-13] = 12 pOH, 14-12=2 pH
5.67 x 10-2 M
-log[5.67x10-2] = 1.25 pOH, =12.75 pH
3.45 x M
-log[3.45x10-11] = pOH, =3.54 pH

53. Calculate [OH-] from pH or pOH
If given pH:
14 – pH = pOH
[OH-] = alog(-pOH) or 10^(-pOH)
If given pOH:
Remember unit for concentration is M.

54. Summary of pH and pOH pH = -log[H+] (acid concentration)
pOH = -log[OH-] (base concentration)
pH + pOH = 14
[H+] = alog(-ph) or 10^(-pH)
[OH-] = alog(-pOH) or 10^(-pOH)
*Hints:
Identify initial substance as acid or base to determine if you have H+ or OH-.
Label all concentrations with M unit.

55. Titrations

56. Titrations
Titration is a standard laboratory method of quantitative/chemical analysis which can be used to determine the concentration of a unknown reactant (acid or base).
An acid or base of known concentration (a standard solution) and volume is used to react with a measured volume of an unknown concentration of an acid or base.

57. Titrations

58. Titrations
Using a buret to add the [unknown]. it is possible to determine the exact amount (V) that has been consumed when the endpoint is reached.
The endpoint is the point at which the titration is stopped.
This is classically a point at which the number of moles of [unknown acid or base] is equal to the number of moles of [known acid or base].

59. Titrations

60. Titrations
Due to the logarithmic nature of the pH curve. the transitions are generally extremely sharp. and thus a single drop of unknown just before the endpoint can change the pH by several points - leading to an immediate color change in a chosen indicator.

61. Titrations of Strong Acids and Strong Bases
Titration of a strong acid and a strong base will result in an equivalence point of 7 because a neutral salt water solution is formed.

63. Acids & Bases Review MAVA = MBVB pH = -log[H+] (acid concentration)
(Don’t forget to adjust M x 2 if necessary…example H2SO4 or Ba(OH)2)
pH = -log[H+] (acid concentration)
pOH = -log[OH-] (base concentration)
pH + pOH = 14
[H+] = alog(-ph) or 10^(-pH)
[OH-] = alog(-pOH) or 10^(-pOH)
(Remember Acids (begin with H) & Bases (end with OH)