chapter 20 acids and bases. describing acids and bases 1.properties of acids and bases acids bases...
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Chapter 20
Acids and Bases
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• Describing Acids and Bases1. Properties of Acids and Bases Acids Bases Contains H+ Contains OH-
Turns blue litmus red Turns red litmus blue Taste sour Taste
bitter Can be electrolytes Can be electrolytes Reacts with bases to Reacts with acids to form water form water
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2. Quick review of naming acids
Hydrogen ions and acidity1. Hydrogen Ions from water a. When water molecules lose a hydrogen ion it becomes OH-
Anion ending Example Acid name Example
-ide Cl- chloride Hydro- (stem) –ic acid
-Ite SO3-2 (stem)-ous acid
-ate SO4-2 (stem)-ic acid
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b. When water molecules gain a hydrogen ion it becomes H3O+
(called the hydronium ion)
2. Self-ionization a. When two water molecules produce ions
b. H2O (l) H+ (aq) + OH- (aq)
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c. [H+] = 1.0 x 10 -7 M
d. [OH-] = 1.0 x 10 -7 M
e. When [H+] and [OH-] are equal it is a neutral solution
f. When they are independent (not equal) [H+] increases, [OH-] decreases [H+] decreases, [OH-] increases
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3. Ion-product constant a. kw : product of concentration of H+ and OH-
in water
b. Kw = [H+] [OH-] = 1.0 x 10-14 M2
c. Acidic solution: one where [H+] is greater than [OH-]
[H+] > 1.0 x 10-7 M
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d. Basic solution: one where [H+] is less than [OH-]
[H+] < 1.0 x 10-7 M
e. Basic solution also known as Alkaline solution
4. The pH concept a. Better expressed using the pH scale
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b. pH + pOH = 14 pH = -log[H+] pOH = -log[OH-]
c. In a neutral solution [H+] = 1.0 x 10-7 M pH = -log (1 x 10-7) = -(log 1 + log 10-7) = -(0.0 + (-7.0)) = 7.0
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d.
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e.
5. Example problems: a. What is the pH of a solution with a
hydrogen-ion concentration of 1.0 x 10-10M?
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b. The pH of an unknown solution is 6.00. What is its hydrogen-ion
concentration?
c. What is the pOH of a solution if [OH-] = 4.0 x 10-11 M?
d. What is [H+] of a solution if the pH = 3.70?
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6. Measuring pH a. Acid-base Indicator 1. Indicator (In) is an acid or base that
undergoes dissociation in a known pH range
2. Reaction form: HIn (aq) H+ (aq) + In- (aq) Acid form Base form
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3. Types: pH color Thymol blue 1.2-3.0 red yellow 8.0-9.5 yellow blue Bromphenol blue 3.0-4.6 yellow blue Bromcresol green 3.7-5.3 yellow blue methyl red 4.2-6.2 red yellow Alizarin 4.5-6.0 yellow red Bromthymol blue 6.0-7.5 yellow blue Phenol red 6.9-8.2 yellow orange Phenolphthalein 8.0-10.0 colorless pink
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alizarin yellow R 8.0 – 12.2 yellow red
4. Useful at room temperature (25 °C)
b. pH meter 1. Useful to make rapid, accurate pH
measurements
2. more practical than liquid indicators
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Acid-Base Theories1. Arrhenius Acids and Bases a. Acids are hydrogen containing compounds
that ionize to yield H+ in aq solutions
b. Bases are compounds that ionize to yield OH- in aq solutions
Acids c. Monoprotic acids have one hydrogen HCl
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d. Diprotic acids : have two hydrogens H2SO4
e. Triprotic acids: have three hydrogens H3PO4
f. Only very polar bonds will dissociate Hδ+--Clδ- H+ (aq) + Cl- (aq)
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g. C-H bonds weakly polar will not dissociate
ex. Ethanoic acid (CH3COOH):
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Bases h. NaOH (s) Na+ (aq) + OH- (aq) i. Common bases: KOH, NaOH, Ca(OH)2, Mg(OH)2
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2. Bronsted-Lowery Acids and Bases a. Acid is a hydrogen-ion donor
b. Base is a hydrogen-ion acceptor
c. Conjugate acid – particle formed when a base gains a hydrogen ion
d. Conjugate base- particle that remains when an acid has donated a hydrogen ion
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e. Conjugate acid-base pair: two substances related by the loss or gain of
a single hydrogen bond f. Examples: 1. NH3(aq) + H2O (l) NH4
+ (aq) + OH- (aq)
acceptor donor (base) (acid) (CA) (CB)
2. HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)
(acid) (base) (CA) (CB)
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g. Amphoteric: a substance that can act like both an acid and base
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Sample problems• 1. Classify the following as Brønsted acids, bases or
both. a) H2O b) OH- c) NH3 d) NH4
+
• 2. What is the conjugate base of the following acids?
a) HClO4 b) NH4+ c) H2O d) HCO3
-
• 3. What is the conjugate acid of the following bases?
a) CN- b) SO42- c) H2O d) HCO3
-
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3. Lewis Acids and Bases a. Acid: a substance that accepts a pair
of electrons to form a covalent bond
b. Base: a substance that donate a pair of electrons to form a covalent
bond c. Examples: 1. H+ + acid base 2.
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Strengths of Acids and Bases 1. Strong acids and bases a. Strong acids: completely ionize (dissociate) HCl, HNO3, H2SO4, HBr, HI, HClO4
b. Dissociation constant (Ka): the ratio of the concentration of the
dissociated form of an acid to the concentration of the
undissociated form
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***See page 600 Table 20.7
3. Equilibrium-constant expression K = [products] [reactants]
** Remember to raise the concentrations to the coefficient number.
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4. Ka = [H+][A-] (Gives the ratio of ions
[HA] vs molecules) Weak acid has Ka <1
Leads to small [H+] and pH of 2-7 5. Kb = [BH+][OH-]
[B] Weak bases has Kb < 1
Leads to small [OH-] and pH of 12-7**Do not use water in the [ ]
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6. Examples: a. Calculate the [OH-] of a 0.500 M solution of
aqueous ammonia. The Kb is 1.74 x 10-
5.
NH3 + H2O NH4+ + OH-
Kb = [NH4+][OH-]
[NH3]
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b. You have 1.00 M acetic acid (HOAc). Calculate the equilibrium concentrations
of HOAc, H+, OAc-, and the pH. Ka = 1.8 x 10 -5
Step 1 Define equilibrium concentrations . [HOAc] [H+] [OAc-]
Initial: 1.00 0 0Change: -x +x +xEquilib: 1.00-x x x
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Step 2: Write the Ka expression
HC2H3O + H20 H+ + C2H3O-
(HOAc) (OAc-) Ka = [H+][OAc-]
[HoAc] 1.8 x 10-5 = (x)(x) = x2
(1.00 –x) (1.00 – x)This is a quadratic. Solve using the quadratic formula. OR
you can make an approximation if x is very small. (Rule of thumb: 10-5 or smaller is OK)
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1.8 x 10-5 = x2 1.00
x = [H+] = [OAc -] = 4.2 x 10-3M
pH = -log[4.2 x 10-3] = 2.37
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c. You have 0.010 M NH3. Calculate the pH if the Kb = 1.8 x 10-5.
NH3 + H2O NH4+ + OH-
[NH3] [NH4+] [OH-]
Initial 0.010 0 0Change -x x xEquilibrium 0.010-x x x
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Kb = [NH4+] [ OH-]
[NH3]
1.8 x 10-5 = (x)(x) 0.010 – x
x = 4.2 x 10-4 M
At equilibrium: 0.010 -4.2 x 10-4 = 0.00958≈0.01
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Once you find [OH-], you find the pOH
pH + pOH = 14
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pH indicators
1. indicator (In) is an acid or base that dissociates in a known pH range
HIn (aq) acid form
OH-
H+
H+ (aq) + In- (aq) base form
2.
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• 3. Types of indicators a. Methyl red: dye that turns red in acids 0-4.4 : red 4.5-6.1: orange 6.2-above: yellow
b. Phenolphthalein: colorless in acids, pink in bases
below pH 8.2: colorless above pH 10: pink
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c. Bromothymol blue: used for weak acids/bases
below pH of 6.0 = yellow pH of 7.0 = green above pH of 7.6 = blue
d. Universal indicator: used for acids and bases
0-3 3-6 7 8-11 11-14 red orange/ green blue purple yellow
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• Problems with indicators 1. Only work at room temperature (will
change colors at different temp)
2. Salts in the solution may change the dissociation process
pH meter: equipment used to measure pH (best pH measurement)