Acids and Bases Chapter 15. End-of-Chapter HW: See Next Page; Not Graded 2) Some Deadlines (OL = Online; Time = 11:50 pm unless noted): OL Practice 4/5/2013 OL Math Review 4/8/2013 OL Chemistry Review 4/9/2013 Lab Safety Quiz 4/11/2013 – Thursday, Noon OL HW 15 4/15/2013
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End-of-Chapter HW: See Next Page; Not Graded
2) Some Deadlines (OL = Online; Time = 11:50 pm unless noted):
OL Practice 4/5/2013
OL Math Review 4/8/2013
OL Chemistry Review 4/9/2013
Lab Safety Quiz 4/11/2013 – Thursday, Noon
OL HW 15 4/15/2013
OL available at: www.saplinglearning.com
3) Exam & Quiz over 15, 16 & Labs 1,2 in ~ 4 weeks
4) Download, print, read, do lab material for 1st lab (meet next wk)
Not turned in; will discuss in lecture after ch 15 & 16
1 3 4 7 10 through 23 29 33 35
39 41 53 57 59 61 65 67 71
75 77 81 87 93 103 105 107 116
1. Definition:All digits known with certainty plus the first uncertain one. Assume last digit is in doubt by ± 1 unless told otherwise.
2. Zeros: Assume significant with one exception - not significant when left of first non-zero digit. Always report all SF including zeros. Can use underline to show last SF.
3. Examples: 0.0010030 = 5 SF 1.00 = 3 SF 1.013456 = 4 SF
4. Why Use? So that folks looking at your data will immediately know which digits are meaningful. One can be misled on your accuracy if significant figures are not used.
5. Significant Figures limited by:
a) Technique Used
b) Measuring Devices
c) Convention:Standard Deviation , Rel. Std. Dev., Error: Use 2 SF
6. Use of SF in calculations:
a) Multiplication / Division: Answer can’t be more accurate than least accurate number. Smallest # of SF limits answer to that # of SF.
b) Addition / Subtraction: Set up problem & cut off answer at first doubtful digit. Answer can gain or lose SF.
7. Examples: 1 . 0 3 4 2 7 6(7 SF)
- 0 . 0 1 1(2 SF)
1 . 0 2 3 2 7 6 = 1.023 (4 SF)
5.692 x 0.00220 x 1.0011
= 1.124946126 = 1.1 (2 SF)
9.06x10-2 x (3.333 - 3.21) (use addn/subtn rule 1st, then use mult/div rule)
8. Logs & Inverse Logs: a log is the addn of exponent plus the log value of the #
Examples: Log 3.62x104 = 4.0000… + 0.558709 = 4.559 (gain SF)
Inverse log of 4.559 = 104.559 = 100.559 x 104.000 = 3.62 x104(lose SF)
Use the rule given in Sapling hints of chemistry review.
9. Grading: Ignoring SF in lab & lecture WILL lower grade.
10. Accuracy: Closeness of average value to true value; use RE(relative error) in ppt:
(true value – average exp value) x 1000 = RE in ppt (parts per thousand)
11. Precision: Grouping of measurements; good precision = small grouping.
Precision usually reported as standard deviation, sd (units for sd are same as x).
sd = ± √∑(X - xi)2 / ( n-1) ∑= sum;X = ave; n = # runs; xi = each value
rsdin ppt = (sd x 1000) / X Convention: Report sd,rsd, & RE to 2 SF.
See burette lab for worked out example. Can use Excel for calculations.
A. Arrhenius Acids and Bases
H+ + H2O H3O+Hydronium Ion
Note: 1) I will use H+ in place of H3O+; 2) [ ] = m/L = M
Acid: A substance which increases [H+] when dissolved in water.
Base: A substance which increases [OH-] when dissolved in water.
-The above concepts were
introduced by Svante Arrhenius
in his doctoral thesis in 1884 in
Sweden. His work was assumed
to be wrong, and he was given
the lowest possible passing grade;
awarded the Nobel Prize in 1903
for his work!
- Strong acids completely ionize in water to release H+ (H3O+) and an anion; know the names & formulas of the strong & weak acids & bases.
Strong Acids: HClO4 H2SO4 HI HBr HCl HNO3
HI + H2O H3O+ + I-or HI H+ + I-(100%)
1) Caution: Strong acids react with proteins, carbohydrates and fats.
2) Caution: HClO4 and HNO3 can react explosively with organics.
3) Only the first ionization is 100% with H2SO4
4) Important: Know the Strong & Weak acids/Bases, Formulas & Names.
5) For oxy acids like HNO3 , H2SO4 & HClO4the acidic H is on an O atom.
- Weak acids only partially ionize in water; < 5% ionization.
Weak Acids: HC2H3O2 HF H3PO4 HCN H2S HNO2 NH4+ HClO
HC2H3O2 H+ + C2H3O2- & NH4+NH3 + H+(~ 1% ionized)
Note: - only the H attached to the O in acetic acid is released; H’s attached to C are NOT acidic; most organic acids are weak acids.
- pH, conductivity & van’t Hoff factor (i) measurements, theoretical calculation allow one to determine the % ionization.
- Know the names & formulas of the above weak acids.
Strong bases ionize 100 % in water and weak bases only partially ionize.
Examples of Strong Bases: LiOH NaOH KOH Ca(OH)2 Sr(OH)2 Ba(OH)2 Na3PO4
Examples of Weak Bases: NH3 CO3-2 HCO3- CN- C2H3O2- F-
NaOH Na+ + OH-(strong, 100 % Ionized)
NH3 + H2O NH4+ + OH- (weak, ≤1 % Ionized)
F- + H2O HF + OH- (weak, ≤1 % Ionized)
Notes: 1) Salts of weak acids (like C2H3O2- & F-) are weak bases. These are usually found as group I salts: NaF, KF, LiF, NaC2H3O2, etc.
2) Bases are much more damaging than acids to proteins & triglyceride fats. Eye damage can occur on contact; even with dilute bases.
3) Important; know the strong & weak bases/acids & their names.
- H+ and OH- are in dynamic equilibrium in water; changing the M of one will change the M of the other.
- The following ALWAYS takes place in water regardless of any other chemicals present:
H2O H+ + OH-
M in pure water = 56 10-7 10-7
- [H+] & [OH-] = 1.00x10-7M for pure, neutral water at 25oC
1) Le Chatelier’s Principle predicts effect of adding additional H+ or OH- to H2O.
2) Acid + Base react completely to yield a salt plus water.
Example: 2HCl + Ca(OH)2 ----) CaCl2 + 2H2O ;
Heart & driving force of this reaction (NIE): H+ + O-H- ----) H-O-H (H2O)
- The Arrhenius definition does not apply when water is not present; so, the Brønsted concept was generated.
Acid: A proton donor Note: Proton = H+
Base: A proton acceptor
Example: HCl (g) + NH3 (g) -----) NH4Cl (s)
H+(aq) + :NH3(aq) -----) NH4+(aq)(Net Ionic Eqn in H2O)
Notes: 1) HCl is the proton donor & NH3 is the proton acceptor
2) The products are called conjugates.
3) The Brønsted definition includes all Arrhenius acids & bases plus additional ones (more general).
Examples: HF + HCO3- H2CO3 + F-
HF = Acid F- = Conjugate BaseHCO3- = Base H2CO3 = Conjugate Acid
CH3NH2 + HI ------) CH3NH3+ + I-
Which is conjugate acid & which is conjugate base?
I- = conjugate base CH3NH3+= conjugate acid
C. Lewis Acids and Bases
- The most general acid-base definition is based on electrons and is called the Lewis definition (Also gave us: Lewis structures, covalent bonds, “photons”).
Acid = Electron pair acceptor
Base = Electron pair donor
Example: AlCl3 + Cl- -----) AlCl4-
AlCl3 accepts a pair of electrons = Acid
Cl- donates a pair of electrons = Base
1)The above rxn is an acid base rxn in the Lewis sense; AlCl3 & Cl- are Lewis acids & bases but not Brønsted or Arrhenius acids & bases.
2) Many organic reactions use non-polar solvents, and non-polar Lewis acids such as AlCl3BF3FeCl3 are more soluble than polar acids in the organic solvents.
3) Lewis definition includes all Arrhenius & Bronsted acids & bases.
- Table 15.2 is a partial list of the strengths. Organics are left off. Organic acids such as: HC2H3O2 (acetic), HCO2H (formic), HC4H7O2 (butyric), and HOC6H5 (phenol) are weak acids. Organic bases like N(C2H5)3 or CHO2- are weak bases.
CH3COH HCOH CH3CH2CH2COH -OH
- Organic acids are generally weaker than inorganic acids. Some organic bases like :CH3- are much stronger than inorganic bases.
- Two related factors influence acid strength: (EN = electronegativity)
1) Bond Polarity - The more polar the bond, the stronger the acid.
- HCl (∆EN = 0.9) > > H2S (∆EN = 0.4)
2) Bond Strength - The weaker the bond, the stronger the acid.
- HI = strong & HF = weak acid. HI bond weaker than HF bond.
- Bond strength is the larger of the two factors.
a) Which is the strongest acid: H-O-ClO3 or H-O-ClO2?
b) Which is the strongest acid: FCH2-CO2H orCH3-CO2H?
Answers: a) HClO4 b) FCH2-CO2H (both polarity & bond strength factor in)
A. Autoionization of Water - The following is always true in water:
1H2O 1H+ + 1OH-M H2O = 55.5 m/L ≈ constant
Keq = [H+]1[OH-]1 = [H+][OH-]
Keq x [55.5] = Kw = 1.00x10-14 = [H+][OH-]
1.00x10-14 = [H+][OH-]
Notes: 1) The above equilibrium expression is always true for water at 25oC; the Kw value does change a little with temperature.
2) H2O drops out of all equilibrium expressions when water is the solvent.
3) You need to memorize this Kw value; assume T = 25oC for all problems.
H2O H+ + OH-(equilibrium reaction)
1.00x10-14 = [H+][OH-] (equilibrium expression)
- This expression is true regardless of what is present in the water.
- If we let X = M of H2O that dissociates, then X = [H+] = [OH-] at equilibrium.
1.00x10-14 = [X] [X] = [X]2 X = √1.00x10-14 = 1.00x10-7
x = [H+] = [OH-] = 1.00x10-7M (this is called a neutral solution)
- If add another source of acid or base to water, then 1.00x10-14 = [H+][OH-]is still true; but, [H+] & [OH-] will no longer be equal; can calculate new values of [H+] & [OH-] with equil. expression.
1) If one adds 0.10 mole of HI to 1.0 L of water, then what is the [H+]?
- Strong acids ionize 100% The M of H+ = M of the strong acid for HI.
- 0.10 M HI yields 0.10 M H+ in water; [H+] = 0.10 M
2) What will be the [OH-] in 0.10 M HI? Note: [H+] = 0.10 M
- Need to use: 1.00x10-14 = [H+] [OH-] from: H2O H+ + OH-
1.00x10-14 = [0.10] [OH-] [OH-] = 1.0x10-13M
3) What are the [H+] & [OH-] in 1.0 x 10-2M NaOH ?
- NaOH ionizes 100%. 1.0 x 10-2M NaOH yields 1.0x10-2 M OH-
1.00x10-14 = [H+] [OH-] 1.0x10-14 = [H+] [1.0x10-2] [H+] = 1.0x10-12M
Note: The following summarizes the acidity scale (based on M of H+) :
[H+] > 1.0x10-7M = Acidic [H+] = 1.0x10-7M = Neutral [H+] < 1.0x10-7M = Basic
- General Rule: The M of H+ or OH- is the M of the strong acid or strong base.
- Exception: Above not true when M of the strong acid (or base) is < 10-6M.
Why? Are two sources of H+ in water; one from H2O (~10-7) & one from acid.
[H+]total = [H+]H2O + [H+]StrongAcid [H+]total = (≈10-7) + [H+]S Acid
Example: What will be the [H+] if one adds 1.0x10-9 moles of HI to 1.0 L of water?
Will the [H+] be 1.0x10-9M; No, [H+] = (1.0x10-7+ 1.0x10-9) = 1.0x10-7
C. pH Note: “p” = mathematical operator = take -log
Need a more convenient scale than [H+] for measuring acidity; the [H+] scale is too large – from about 10o to 10-14 in water; can compact using the pH scale.
pH = - log [H+] - log [10-x] = x
ln is to the base e;log is to the base 10; Note: ln or log of 1.000 = 0.000
1) What is pH of: 1.0 x 10-4M HCl?1.0 x 10-5M HCl? 5.0 x 10-5M HCl?
4.005.00-log (5.0 x 10-5) = 4.30
2) if take -log of both sides of: 1.00x10-14 = [H+] [OH-]
(note: can use either equation)14.00 = pH + pOH
Example: What is pH if pOH = 11.0?pH = 14.00 – pOH = 3.0
3) Will have to calculate [H+] from pH. With your calculator, use:
10x or inverse log or antilog functions
Examples: ( use: [H+] = 10 –pH )
-What is [H+] if the pH is 5.50?[H+] = 10-5.50 = 3.2 x 10-6
-What is [H+] if the pH is 9.38?[H+] = 10-9.38 = 4.2 x 10-10
4) Can check acidity with either a pH meter or with an Indicator
-log 4.5x10-3 = -log 4.5 + -log 10-3 = -0.653 + 3.000 = 2.347 = 2.35
4.6x10-3 2.3372 3SF (gain a SF)
- antilog 2.34 = -antilog 3.00 + antilog 0.66 = 10-3 x 4.57 = 4.6x10-3
- inverse log 2.34 = 10-2.34 = 4.6x10-3
2.34 4.57x10-34.5±0.1 (2 SF) (lose a SF)
- pH meter is a device used to
measure potential developed by a set
of electrodes (reference and glass).
Et = Eglass+ Eref Et converted to pH
- The surface of the electrode has
special glass which attracts H+ & sets
up a potential difference across glass
Membrane (Eglass). Eref = constant
-Needs to be calibrated with standard
buffers. Up to ± 0.01 pH accuracy.
- Typical combination electrode ($100
to $1000). Will discuss in more detail
in chapter 19 (end of May).
10-1 ………… 10-7 ………..... 10-14
pH Scale: 1 ……….. 7 …..… 14
pH < 7 (acidic)pH = 7 (neutral) pH > 7 (basic)
Indicators – organic compounds whose color varies with acidity. Can coat paper with several indicators and specific color will tell pH to within ± 1 pH. Predict result of mixing MR & TB.
This mix is my choice.