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Slide 1 / 123 Slide 2 / 123
Chemistry
Atomic Origins
2015-08-14
www.njctl.org
Slide 3 / 123
Acids and Bases
lactic acid lactate
Lactic acid is one of many metabolities produced when
we exercise. It generally loses an H+ ion to from the lactate ion (one of the chemicals that causes burning sensations in
our muscles.)
-
Slide 4 / 123
Auto-ionization of WaterIn any sample of water, a small number of water molecules will
dissociate into H+ and OH- ions.
H
H
O OH
H
+-
H2O(l) -------> OH-(aq) + H+(aq)
- +The H+ ion then typically binds to a lone pair of electrons on another water molecule to make the hydronium ion: H3O+
2H2O(l) -------> OH-(aq) + H3O+(aq)
H
H
O OH H
- +H
HOH
H
O
Slide 5 / 123
In 1909, a device was invented that could
measure the H+ or H3O+ concentration in an
aqueous solution.
H3O+(aq) = 1.0 x 10 -7 M
@ 25 C
Using this data, the equilibrium constant for the auto-ionization of water can be calculated.
2H2O(l) --> H3O+(aq) + OH-(aq)
Recalling our equilibrium concepts...... Kw = [H3O+][OH-]
Since equal amounts of H3O+ and OH- are created...
[H3O+] = [OH-] = 1.0 x 10 -7 M
Kw = (1.0 x 10 -7)(1.0 x 10 -7) = 1.0 x 10 -14 M
Auto-ionization of Water
Clearly, the equilibrium lies far to the left!
Water does NOT like to dissociate.
Slide 6 / 123
1 What is the concentration of hydronium ions (H3O+) in pure water?
A 1.0 x 10 -2 M
B 1.0 x 10-5 M
C 1.0 x 10-7 M
D 1.0 x 10 -10 M
E 1.0 x 10-14 M
answ
er
Slide 7 / 123
2 Which of the following is the value of Kw for water?A 1.0 x 10-2
B 1.0 x 10-4
C 1.0 x 10-7
D 1.0 x 10-9
E 1.0 x 10-14
answ
er
Slide 8 / 123
3 Which of the following would be true in pure water?A [H3O+] = [OH-]
B [H3O+] < [OH-]
C [OH-] = 1 x10-7 M
D A and C
E B and C
answ
er
Slide 9 / 123
4 The magnitude of K w indicates that _________
A water ionizes to a very small extentB the autoionization of water is exothermicC water ionizes very quicklyD water ionizes very slowly
answ
er
Slide 10 / 123
5 The molar concentration of hydronium ion, [H3O+ ], in pure water at 25 °C is ___________.
A 0B 1
C 7
D 10-7
E 10-14
answ
er
Slide 11 / 123
Calculating H3O+ or OH- In the natural world, we do not find pure water. There are always
things dissolved in it that influence the concentrations of hydronium and hydroxide ions.
The hydronium or hydroxide concentration in a solution can be determined easily if one knows one or the other.
Kw = [H3O+][OH-] = 1.0 x 10-14
Rearranged for [H3O+] Rearranged for [OH-]
[H3O+] = 1.0 x 10-14/[OH-] [OH-] = 1.0 x 10-14/[H3O+]
Slide 12 / 123
Kw = [H3O+][OH-] rearranged to find [OH-] = Kw/[H3O+]
= 1.0 x 10-14/ 3.4 x 10-5 = 2.9 x 10-10 = [OH-]
Kw = [H3O+][OH-] rearranged to find [H3O+] = Kw/[OH-]
= 1.0 x 10-14/ 1.2 x 10-12 = 8.3 x 10-3 = [H3O+]
#1 What is the [OH-] in a solution with [H3O+] = 3.4 x 10-5 M?
#2 What is the [H3O+] in a solution with [OH-] = 1.2 x 10-12
Calculating H3O+ or OH- Let's do some examples!
move for answer
move for answer
Slide 13 / 123
Calculating H3O+ or OH- Application:
Tap water is NOT pure water. There are many things dissolved in it that affect the amount of [H3O+] and [OH-] in the water sample. Can you think of some things that might chloride (Cl-), carbonate (CO3
2-)
be dissolved in tap water?
Flouride (F-), calcium ions (Ca2+), c
The average concentration of H3O+ in New York City tap water is
5.01 x 10-8 M. What is the average [OH-]?
Kw = [H3O+][OH-] rearranged to find [OH-]
= Kw/[H3O+]
= 1.0 x 10-14/5.01x 10-8
= 1.99x 10-7 M
move for answer
move for answer
Slide 14 / 123
6 What is the [H3O+] in an aqueous sample with an [OH-] equal to 3.4 x 10-3 M?
A 3.4 x 10-3 MB 2.9 x 10-12 M
C 1.0 x 10-7 M
D 9.4 x 10-7 M
E 3.4 x 1011 M
answ
er
Slide 15 / 123
7 Which of the following would have the smallest [OH-]?
A solution with [H3O+] = 2.4 x 10-1
B solution with [H3O+] = 2.4 x 10-11
C solution with [H3O+] = 2.4 x 10-6
D solution with [OH-] = 2.4 x 10-3
E solution with [OH-] = 2.4 x 10-12
answ
er
Slide 16 / 123
8 The pacific ocean off the coast of Hawaii has
a [OH-] = 8.32 x 10-9 M.
What is the [H3O+]?
answ
er
Slide 17 / 123
Arrhenius Definition of Acids and Bases
As we have learned, when certain substances are added to water, the H3O+ concentration changes.
Furthermore, if the [H3O+] changes, it would influence the [OH-].
Kw = [H3O+] [OH-] = 1.0x 10-14
Slide 18 / 123
In 1884, Swedish scientist Svante Arrhenius decided to
create definitions for substances that changed the [H3O+] in an
aqueous solution.
Arrhenius Definition of Acids and Bases
Arrhenius labeled anything that increased the [H3O+] an acid
Arrhenius labeled anything that increased the [OH-] a base
By measuring the [H3O+] of a water solution after a substance had been added, he could see if the substance was acidic or basic!
Slide 19 / 123
H3O+(aq) = 2.3x 10-6 M @ 25 C
HCN(aq)
Example 1: Let's add some HCN(aq)
Remember that pure water has an [H3O+] = 1.0 X 10-7M.
Since the [H3O+] is higher than 1.0 X 10-7M, Arrhenius would have described HCN as an acid!
Arrhenius Definition of Acids and Bases
Slide 20 / 123
By measuring the [H3O+] of a water solution after a substance had been added, he could see if the substance was acidic or basic!
H3O+(aq) = 4.1x10-11 M @ 25 C
NaOH(s)
Example 2: Let's add some NaOH(s)
Since the [H3O+] is lower than 1.0 x 10-7 M thereby making the [OH-] higher than 1.0 x 10-7M, Arrhenius would have described
NaOH as a base!
Arrhenius Definition of Acids and Bases
Slide 21 / 123
9 Which of the following solutions would be considered by Arrhenius to be the most basic?
A 0.1 M NH3 [H3O+] = 3.4x10-10 M
B 0.1 M NaOH [H3O+]= 1x10-13M
C 0.1 M HCl [H3O+] =1x10-1 M
D 0.1 M HCN [H3O+]= 2.3x10-6M
E Pure water [H3O+]=1x10-7 M answ
er
Slide 22 / 123
10 Vinegar has a [H3O+] of around 3.4 x10-3 M. Which of the following solutions would be considered by Arrhenius to be MORE acidic than vinegar?
A 0.1 M NaOH [H3O+] = 1x10-13 M
B 0.1 M HCl [OH-] =1.0 x10-13 M
C 0.1 M NaCN [OH-] = 2.6x10-4 M
D 0.1 M NH3 [H3O+] = 7.6x10-9 M
E pure water [OH-] = 1.0 x10-7 M
answ
er
Slide 23 / 123
Bronsted Lowry Definition of an Acid
At this time, most scientists explained Arrhenius acids as possessing H+ ions that could be added to water to produce [H3O+]
Arrhenius acids in action
HF(aq) + H2O(l) --> F-(aq) + H3O+(aq)
Here, the hydroflouric acid (HF) donates one of it's H+ ions to a water molecule increasing the [H3O+](aq)
Two scientists - Bronsted and Lowry, working independently, decided a more appropriate definition of an acid would be that of an H+ donor.
Slide 24 / 123
Bronsted/Lowry Definition of Base
At this time, most scientists explained Arrhenius bases as possessing OH- ions that would increase the [OH-] and decrease the [H3O+].
NaOH(aq) --> Na+(aq) + OH-(aq)
[OH-] causes [H3O+]
Arrhenius base in action
Unfortunately, this view required that all bases had to possess the hydroxide ion. This was clearly not the case. Many substances, like ammonia (NH3) or sodium phosphate (Na3PO4), were known to be
basic but did NOT have any hydroxide ions!
Slide 25 / 123
Bronsted Lowry Definition of Base
Bronsted and Lowry proposed that, insteading of possessing hydroxide ions, a base was a substance that accepted an H+ from
water to produce OH- ions!
NH3(g) + H2O(l) --> NH4+(aq) + OH-(aq)
Bronsted base in action
When ammonia, NH3, accepts the H+ from the water, the water turns into OH- making the solution basic.
Slide 26 / 123
Bronsted Lowry Definition of Acid and Bases
Summary
Acids are defined as H+ (proton) donors.
HC3H6O3(aq) + H2O(l) --> C3H6O3-(aq) + H3O+(aq)
lactic acid
cyanide base
Bases are defined as H+ (proton) acceptors.
CN-(aq) + H2O(l) --> HCN(aq) + OH-(aq)
Slide 27 / 123
11 A Bronsted acid is a substance that...
A accepts H+ ions
B donates OH- ions
C increases the concentration of OH- ions
D donates H+ ions
E accepts OH- ions
answ
er
Slide 28 / 123
12 Which of the following could NOT act as a Bronsted acid?
A HCN
B H2 SO4
C NH4 +
D H3 O+
E BF3
answ
er
Slide 29 / 123
13 A Bronsted-Lowry base is defined as a substance that __________.
A increases [H+ ] when placed in H2 O
B decreases [H+ ] when placed in H2 O
C increases [OH-] when placed in H2 O
D acts as a proton acceptor
E acts as a proton donor
answ
er
Slide 30 / 123
14 Which of the following compounds could never act as a Bronsted acid? A SO4
2-
B HSO4-
C H2 SO4
D NH3
E CH3 COOH
answ
er
Slide 31 / 123
Bronsted Acids and Bases (In Depth)
Acids and Bases go togetherIt should be noted that if an acid donates an H+, that H+ will be
accepted by another substance.So, where there is an acid, there will be a base
N
H
H
H
OH
H
N
H
H
H
H O
H+ +
NH3 acts as a base and accepts an H+ to become NH4+
water acts an acid and donates it's H+ to become OH-
+ -
Slide 32 / 123
Bronsted Acids and Bases (In Depth)Identifying an acid or a base
By examining the products and reactants of a chemical reaction, one can identify if a substance is behaving as an acid or as a
base.
ExampleHSO4 -(aq) + CN-(aq) --> SO4
2- (aq) + HCN(aq)
HSO4 -(aq) donated an H+ to become SO42- = It's an acid!
CN-(aq) accepted an H+ to become HCN = It's a base!
Slide 33 / 123
Bronsted Acids and Bases (In Depth)
Identifying an acid or a base
Identify which reactant behaves as an acid and which behaves as a base in the following reaction!
H2 O(l) + CH3 NH3+(aq) --> CH3 NH2 (aq) + H3 O+(aq)
CH3 NH3+(aq) donated an H+ to become CH3 NH2 = It's an acid!
H2O(aq) accepted an H+ to become H3O+ = It's a base!move for answer
Slide 34 / 123
15 According to the following reaction, which reactant molecule is acting as an acid?
A H2 SO4
B H2 O
C H3 O+
D HSO4-
E None of the above
H2 O + H2 SO4 → H3 O+ + HSO4 -
answ
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Slide 35 / 123
16 According to the following reaction, which reactant molecule is acting as a base?
H2 O + H2 SO4 → H3 O+ + HSO4 -
A H2 SO4
B H2 O
C H3 O+
D HSO4-
E None of the above
answ
er
Slide 36 / 12317 According to the following reaction, which reactant
molecule is acting as a base?
H3 O+ + HSO4 - → H2 O + H2 SO4
A H2 SO4
B H2 O
C H3 O+
D HSO4-
E None of the above
answ
er
Slide 37 / 123
18 For the following reaction, identify whether the circled compound is behaving as an acid or a base.
A Acid
B Base
C Neither
D Both
E None of the above
H3 PO4 + H2 O ⇌ H2 PO4 - + H3 O+
answ
er
Slide 38 / 123
19 For the following reaction, identify whether the circled compound is behaving as an acid or a base.
H3 PO4 + H2 O ⇌ H2 PO4 - + H3 O+
A Acid
B Base
C Neither
D Both
E None of the above
answ
er
Slide 39 / 123
Bronsted Acids and Bases (In Depth)
Identifying an acid or a base in reversible reactions
Reactions are reversible so we must be able to identify acids and bases based on the reverse reaction.
ExampleF-(aq) + H2 O(l) <--> HF(aq) + OH-(aq)
HF(aq) donates an H+ ion to become F-(aq) = It's an acidOH-(aq) accepts an H+ to become H2 O(l) = It's a base
Slide 40 / 123
20 For the following reaction, identify whether the circled compound is behaving as an acid or a base.
H3 PO4 + H2 O ⇌ H2 PO4 - + H3 O+
A Acid
B Base
C Neither
D Both
E None of the above
answ
er
Slide 41 / 123
21 For the following reaction, identify whether the circled compound is behaving as an acid or a base.
H3 PO4 + H2 O ⇌ H2 PO4 - + H3 O+
A Acid
B Base
C Neither
D Both
E None of the above
answ
er
Slide 42 / 123
Conjugate Acids and Bases
The term conjugate comes from the Latin word “conjugare,” meaning “to join together.”
Reactions between acids and bases always yield their conjugate bases and acids.
HNO2(aq) + H2O(l) NO2 - (aq) + H3O+(aq)
donates H+
accepts H+
Acid Base Conjugate base
conjugate acid
Slide 43 / 123
Conjugate Acids and Bases
To find an acid or bases conjugate in a reaction, simply write the formula for the substance left after the H+ has been donated or
accepted.
Example: What is the conjugate base of HSO4 -(aq)?Since we are looking for a conjugate base, HSO4
- must be an acid so let's have it donate an H+
HSO4-(aq) --> SO4
2- (aq) + H+(aq)
conjugate base
Slide 44 / 123
Conjugate Acids and Bases
Example: What is the conjugate acid of CO3 2- (aq)?Since we are looking for a conjugate acid, CO3
2- must be a base so let's have it accept an H+CO3
2- (aq) + H+ --> HCO3-(aq)
conjugate acid
Dealing with chargesIf you accept an H+, you become more positiveIf you donate an H+, you become more negative
Slide 45 / 123
22 Which of the following would be the conjugate base of HNO2 ?
A NO2-
B H2 NO2
C NO2
D NO22-
E HNO2
answ
er
Slide 46 / 123
23 Which would be the conjugate acid of HCO3 -(aq)?
A CO32-
B HCO3
C CO3
D H2 CO3-
E H2 CO3
answ
er
Slide 47 / 123
24 What would be the an acid/conjugate pair in the following reaction?
A NH2 -/H2 O
B NH2 -/NH3
C H2 O/OH-
D H2 O/NH3
E None of these
NH2 - + H2 O --> NH3 + OH-
answ
er
Slide 48 / 123
Lewis Acids and Bases
DefinitionScientists noticed that some substances could create acidic
solutions despite not having any H+ ions to donate. An example of this was the Ca2+ ion.
G.N. Lewis proposed a mechanism for this
Ca2+ + ---> Ca (OH)+ + +
O
H
H
H
The metal ion accepted a pair of electrons from the water molecule, resulting in the donation of one of the
water's H+ ions.
Slide 49 / 123
Lewis Acids and Bases
A Lewis acid is an electron pair acceptor. Metal ions or molecules with incomplete octets (BF3 ) are good examples.
A Lewis base is an electron pair donor. Molecules with unbonded electrons (NH3 , CN-, OH-, H2 O) are good examples.
Lewis Acid (e- pair
acceptor)
Lewis Base (e- pair donor)
Slide 50 / 123
25 A lewis base is a substance that...
A Accepts H+ ions
B Donates H+ ions
C Accepts e- pairs
D Donates e- pairs
E Decreases the concentration of [OH-]
answ
er
Slide 51 / 123
26 Which of the following would likely act as a lewis acid?
A NH3
B OH-C CN-
D H2OE Fe3+
answ
er
Slide 52 / 123
What are Acids and Bases?
Definition Type Acid Base
Arrhenius (traditional)substance that produces
H3O+ ions in aqueous solution
substance that decreases H3O+ ions in aqueous
solution
Bronsted -Lowry substance that donates H+ ions in reaction
substance that accepts H+ ions in reaction
Lewis substance that accepts an electron pair in reaction
substance that donates an electron pair in reaction
Slide 53 / 123
Question 1: Can you think why the Arrhenius definition was considered insuffienct?
It could not explain how a substance without hydroxide could make a solution basic
Question 2: Can you explain why Lewis felt that the Bronsted definition was insufficient?
It required an acid to be in possession of a hydrogen atom.move for answer
move for answer
Class Discussion - Evolution of a definition
Slide 54 / 123
What are Acids and Bases?
All acids are Lewis acids, most are also Bronsted acids, and many are Arrhenius acids
Lewis
Bronsted
Arrhenius
The lewis definition is generally considered the most broad.
Slide 55 / 123
If a substance can act both as an acid and base, it is known as amphoteric. For example, water can act as a base or acid depending on the situation.
Amphoteric Substances
HCl + H2O Cl- + H3O+
Above, water accepts a proton, thus acting as a base.
NH3 +H2O NH4 + + OH-
Above, water donates a proton, thus acting as an acid
Because of water's amphoteric nature, it makes the perfect solvent for most acid base reactions. Its nature allows for easier
exchange of protons between acids and bases.
Slide 56 / 123
Acid and Base Strength
Strong acids are completely ionized in water (They all donate their H+ ions).
Their conjugate bases are very weak. StrongWeakNegligible StrongWeakNegligible
Acid Base
HCl Cl -
H2SO4 HSO 4-
HNO3 NO 3-
H3O+ H 2OHSO4
- SO 42-
H3PO4 H2PO4-
HF F-HC2H3O2 C2H3O2
-
H2CO3 HCO 3-
H2S HS -
H2PO4- HPO 4
2-
NH4+ NH 3
HCO3- CO 3
2-
HPO42- PO 4
3-
H2O OH -
OH- O 2-
H2 H -
CH4 CH 3-
100% protonatedin H2O
Bas
e st
reng
th in
crea
ses
Aci
d st
reng
th in
crea
ses
100% ionized in H2O
Slide 57 / 123
Acid and Base Strength
Weak acids only ionize partially in water.
Their conjugate bases are weak bases.
StrongWeakNegligible StrongWeakNegligible
Acid Base
HCl Cl -
H2SO4 HSO 4-
HNO3 NO 3-
H3O+ H 2OHSO4
- SO 42-
H3PO4 H2PO4-
HF F-HC2H3O2 C2H3O2
-
H2CO3 HCO 3-
H2S HS -
H2PO4- HPO 4
2-
NH4+ NH 3
HCO3- CO 3
2-
HPO42- PO 4
3-
H2O OH -
OH- O 2-
H2 H -
CH4 CH 3-
100% protonatedin H2O
Bas
e st
reng
th in
crea
ses
Aci
d st
reng
th in
crea
ses
100% ionized in H2O
Slide 58 / 123
Substances with negligible acidity do not ionize in water. They will not readily give up protons.
Their conjugate bases are exceedingly strong.
Acid and Base Strength
StrongWeakNegligibleStrongWeakNegligible
Acid Base
HCl Cl -
H2SO4 HSO 4-
HNO3 NO 3-
H3O+ H 2OHSO4
- SO 42-
H3PO4 H2PO4-
HF F-HC2H3O2 C2H3O2
-
H2CO3 HCO 3-
H2S HS -
H2PO4- HPO 4
2-
NH4+ NH 3
HCO3- CO 3
2-
HPO42- PO 4
3-
H2O OH -
OH- O 2-
H2 H -
CH4 CH 3-
100% protonatedin H2O
Bas
e st
reng
th in
crea
ses
Aci
d st
reng
th in
crea
ses
100% ionized in H2O
Slide 59 / 123
Strong Acids
There are seven strong acids:3 contain a H bound to the very electronegative halogens:
HCl hydrochloric acidHBr hydrobromic acidHI hydroiodic acid
HF, or hydrofloric acid, is a weak acid. Although flourine is very electronegative, the bond strength between flourine and hydrogen is too strong for HF to easily give up H+ .
Slide 60 / 123
27 Which of the following is NOT a strong acid?
A HBr
B HFC HI
D HCl
E A and C
answ
er
Slide 61 / 123
Strong AcidsThere are seven strong acids:4 are from the very electron drawing oxyanions:
HNO3 nitric acidH2SO4 sulfuric acidHClO3 chloric acid HClO4 perchloric acid
Each of these anions has a central atom that is highly electronegative compared to hydrogen. The oxygens that are bonded to that central atom draw more electrons from it making it even more electronegative and likely to take electrons from hydrogen forming H+ .
Slide 62 / 123
Strong Acids
The seven strong acids are:
HCl hydrochloric acidHBr hydrobromic acidHI hydroiodic acid
HNO3 nitric acidH2SO4 sulfuric acidHClO3 chloric acid HClO4 perchloric acid
Slide 63 / 123
Monoprotic Acids
The seven strong acids are strong electrolytes because they are 100% ionized. In other words, these compounds exist totally as ions in aqueous solution.
For the monoprotic strong acids (acids that donates only one proton per molecule of the acid), the hydronium ion concentration equals the acid concentration.
[H3O+] = [acid]
So, if you have a solution of 0.5 M HCl, then [H3O+] = 0.5 M
Slide 64 / 123
All alkali metals in Group I form hydroxides that are strong bases: LiOH, NaOH, KOH, etc.
Only the heavier alkaline earth metals in Group II form strong bases: Ca(OH)2 , Sr(OH)2 , and Ba(OH)2 .
Again, these substances dissociate completely in aqueous solution. In other words, NaOH exists entirely as Na+ ions and OH- ions in water.
Strong Bases
All strong bases are group of compounds called "metal hydroxides."
Slide 65 / 123
28 What would be the [H3O+] in a 0.005 M HBr solution?
answ
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Slide 66 / 123
In any acid-base reaction, the proton moves toward the stronger base. In other words, a stronger base will "hold onto" its proton whereas a strong acid easily releases its proton(s).
HCl(aq) + H2O(l) --> H3O+(aq) + Cl-(aq)
acid base conj. acid conj. base
H2O is a much stronger base than Cl-, so the proton
moves from HCl to H2O.
Acid and Base Strength
WeakNegligibleStrongWeakNegligible
Acid Base
HCl Cl -
H2SO4 HSO 4-
HNO3 NO 3-
H3O+ H 2OHSO4
- SO 42-
H3PO4 H2PO4-
HF F-HC2H3O2 C2H3O2
-
H2CO3 HCO 3-
H2S HS -
H2PO4- HPO 4
2-
NH4+ NH 3
HCO3- CO 3
2-
HPO42- PO 4
3-
H2O OH -
OH- O 2-
H2 H -
CH4 CH 3-
100% protonatedin H2O
Bas
e st
reng
th in
crea
ses
Aci
d st
reng
th in
crea
ses
100% ionized in H2O
Slide 67 / 123
29 What would be the [OH-] in a 0.034 M NaOH solution?
answ
er
Slide 68 / 123
CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq)
Acetic acid is a weak acid. This means that only a small percent of the acid will dissociate.
The double headed arrow is used only in weak acid or weak base dissociation equations since the reaction can proceed with both the forward and reverse reactions.
Acid and Base Strength
A single arrow is used for strong acid or strong bases which dissociate completely since the forward reaction is much more favorable than the reverse reaction.
NaOH Na+ (aq) + OH- (aq)
Slide 69 / 123
30 Strong acids have ___________ conjugate bases.
A strongB weakC neutralD negative
answ
er
Slide 70 / 123
31 HBr, hydrobromic acid is a strong acid. This means that _______________.
A aqueous solutions of HBr contain equal concentrations of H+ and OH-
B does not dissociate at all when it is dissolved in water
C cannot be neutralized by a base
D dissociates completely to H+ and Br- when it dissolves in water
answ
er
Slide 71 / 123
pH pH is defined as the negative
base-10 logarithm of the concentration of hydronium ion.
pH = -log [H3O+]
It is a measure of hydrogen ion concentration, [H+ ] in a solution,
where the concentration is measured in moles H+ per liter, or molarity.
The pH scale ranges from 0-14.
Generally when calculating pH we round to two decimal places.
Slide 72 / 123
Slide 73 / 123
What is the pH of the solution with hydrogen ion concentration of 5.67x10-8 M (molar)?
pH = -log [H+]
First, take the log of 5.67x10-8 = -7.25 Now, change the sign from - to +
Answer: pH = 7.25
Calculating pH
Note: If you take the log of
-5.67x10-8 M, you will end up
with an incorrect answer.
The order of operations:
1. Take the log
2. Switch the sign
Slide 74 / 123
32 What is the pH of a solution with hydrogen ion concentration of 1.0 x 10-5 M?
A 1.0 x 10-5
B -5.00C 5.00
D 9.00
E -9.00
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33 What is the pH of a solution with hydrogen ion concentration of 1.0 x 10-12 M?
A 1.0 x 10-12
B 12.00
C 2.00
D -12.00
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34 What is the pH of a solution whose hydronium ion concentration is 7.14 x 10-3 M?
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35 What is the pH of a solution whose hydronium ion concentration is 1.92 x 10-9 M?
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36 What is the pH of a 0.34 M solution of the strong acid HI? (Remember that strong acids ionize completely)
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pHApplication
In order for proteins to be digested in the stomach, the pH must be lower than 2.7. If the pH is too high, proteins will not be broken
down and may cause a food allergy or indigestion.
A patient complains of indigestion and a sample of stomach fluid is taken and the [H3 O+] is found to
be 3.4 X10-3 M. Is there a problem with the pH?
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What is the relationship between [H3O+] and the pH value?
Below are three different [H3 O+]. Find the pH of each. pH = -log [H3O+]
pH
Hydrogen ion concentration, [H3O+]in moles/Liter pH
1.0 x 10-1
1.0 x 10-2
1.0 x 10-10
Clearly, the lower the [H3O+], the _____ the pH.
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What is the relationship between [H3 O+], the pH value, and the acidity and basicity of a solution?
pH
low H3 O+
High H3 O+
high OH-
low OH-
acidic acid
icba
sicbasic
neutral
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pH
These are the pH values for several common substances.
Mor
e ac
idic
Mor
e ba
sic
Battery acid
lemon juice
pure rain or water
distilled water
sea waterbaking soda
household ammonia
household bleach
household lye
gastric fluid
carbonated beveragesvinegarorange juice
beercoffeeegg yolksmilk
blood
milk of magnesia
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For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.
How Do We Measure pH?
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How Do We Measure pH?For less accurate measurements, one can use Litmus paper
“Red” paper turns blue above ~pH = 8“Blue” paper turns red below ~pH = 5
Or an indicator (usually an organic dye)
0 2 4 6 8 10 12 14
pH range for color change
Methyl violet
Thymol blue
Methyl orange
Bromothymol blue
Phenolphthalein
Alizarin yellow R
Methyl red
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pH
Solution type [H +](M) [OH-] (M) pH value
Acidic > 1.0x10-7 <1.0x10-7 <7.00Neutral =1.0x10-7 =1.0x10-7 =7.00Basic <1.0x10-7 > 1.0x10-7 >7.00
[H+] > [OH-]There are excess hydrogen ions in
solution.
[H+] < [OH-]There are excess hydroxide ions in
solution.
BASEACID
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37 Which of the following solutions would be most acidic?A pH = 3
B pH = 2
C pH = 11
D pH = 14
E pH = 1
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38 Which of the following (M) solutions would be LEAST acidic?
A [H3 O+] = 2.3x10-7
B [H3O+] = 9.1x10-3
C [H3O+] = 1.3 x10-2
D [H3O+] = 7.8x10-9
E [H3O+] = 4.5x10-4
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39 Which of the following solutions would have the highest pH?
A [OH-] =3.4x10-3
B [H3O+] = 5.4x10-11
C [OH-] = 3.4x10-12
D [H3O+] =5.4x10-2
E [OH-] =3.4x10-1
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40 Which solution below has the highest concentration of hydroxide ions?
A pH = 3.21
B pH = 7.00
C pH = 8.93
D pH = 12.60
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41 Which solution below has the lowest concentration of hydrogen ions?
A pH = 11.40
B pH = 8.53
C pH = 5.91
D pH =1.98
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42 For a basic solution, the hydrogen ion concentration is ______________ than the hydroxide ion concentration.
A greater than
B less than
C equal to
D Not enough information.
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43 For an acidic solution, the hydroxide ion concentration is ______________ than the hydrogen ion concentration.
A greater than
B less than
C equal to
D Not enough information.
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44 Which of the following would turn blue litmus paper red?
A Solution with [OH-] = 2.3 E-7 M
B Solution with pH = 4
C Solution with pOH = 2
D A and C
E B and C
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Understanding a Log Based Scale
Because of the base-10 logarithm, each 1.0-point value on the pH scale differs by a value of ten.
A solution with pH = 9 has a hydrogen ion concentration, [H+],
that is 10 times more than a pH = 10 solution.
A solution with pH = 8 has a hydrogen ion concentration, [H+],
that is 102 or 100 times more than a pH = 10 solution.
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45 A solution with pH = 3 has a hydrogen ion concentration that is __________than a solution with pH = 5.
A 2x more
B 2x less
C 100x more
D 100x less
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46 A solution with pH = 14 has a hydrogen ion concentration that is __________than a solution with pH = 11.
A 3x more
B 3x less
C 1000x more
D 1000x less
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pOHJust as the pH of a solution can be calculated by:
pH = -log[H3O+]
The pOH of a solution can be calculated by:
pOH = - log[OH-]
Recall that the [OH-] and [H3O+] are inversly related so pH and pOH are as well.
0 7 14low pH
14 7 0
high pOH low pOH
high pH
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Calculating pOH
What is the pOH of a solution that has a [OH-] = 2.3 E-5 M?
pOH = - log[OH-]
pOH = - log(2.3 E-5)
= 4.63
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47 What is the pOH of a solution with a [OH-] = 2.7 x10-2 M?
A 2.7
B 12.43
C 1.57
D -1.57
E -2.7 answ
er
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pOHOnce we have calculated pOH, it is very easy to calculate pH.
Remember that our solvent for all of our reactions is Water. We also know that we have a Kw value for water of 1 x 10-14. This is ALWAYS true for water. We can also determine the following equations:
Kw=[H+][OH-]
Throwing in our logarithms for pH, pOH and pKw we end up with this:
pKw = pH + pOH
Remember that Kw is a constant and if we that the negative log of that constant we get 14 so.....
14 = pH + pOH
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pOH to pH and vice versa
Therefore, if we have a pOH and we want to convert it to pH, so long as we are using water for our solvent, we can use the below equation to determine the pH of the solution.
14 = pH + pOH
In other words, to find the pH of a basic compound, you first must need to determine the pOH of that compound and
then use that to determine the pH. Remember that pOH is calculated using [OH-] and pH is calculated using [H+].
Other then that, there is no difference in the steps used to calculate pOH and pH.
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48 What is the pOH of a solution with a pH =5?
A 5B 15C 7D 8E 9
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49 What is the pH of an aqueous ammonia solution with a [OH-] = 1 x 10-4 M?
A 4B 1 x10-4
C 10D 1 x10-10
E 3 answ
er
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50 What is the pOH of an aqueous HCl solution with a [H3O+] = 2.7 x10-1 M?
A 13.43
B 0.57
C 2.7 x10-1
D 2.7 x10+1
E 12.43
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51 What would be the pH of a 0.045 M NaOH solution? (Recall that NaOH is a strong base and will ionize completely)
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52 Which of the following would be LEAST acidic?
A pOH = 2
B pOH = 4
C pH = 10
D pH = 2
E pH = 11
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Calculating [H3O+] and [OH-] from pH or pOH
If given a pH, one can determine the [H3O+] by:
10-pH = [H3O+]
If given a pOH, one can determine the [OH-] by:
10-pOH = [OH-]
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Calculating [H3O+] and [OH-] from pH or pOH
What is the [H3O+] in a lemon juice solution with a pH = 3.5?
10-3.5 = 3.2x10-4 M
What is the [H3O+] in a bottle of soda with a pOH = 11.4?
14 = pOH + pH 14 = 11.4 + pH
pH = 2.6
10-2.6 = 2.5x10-3 M
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53 What is the OH- ion concentration if the pH of a solution is 6?
A 1 x10-6
B 1 x10-8
C 1 x106
D 1 x1012 answ
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54 What is the OH- concentration if the pH of a solution is 11?
A 1 x 10-4
B 1 x10-3
C 1 x 10-11
D 1 x1011
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55 What is the hydrogen ion concentration (M) in a solution of Milk of Magnesia whose pH = 9.8?
A 9.8 M
B 9.8x10-10 M
C 4.2 M
D 1.6x10-10 M
E 4.2x10-10 M
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56 What is the hydronium ion concentration in a solution whose pH = 4.29?
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57 For a 1.0-M solution of a strong acid, a reasonable pH would be_____.
A 0
B 6
C 7
D 9
E 13
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58 For a 1.0-M solution of a weak base, a reasonable pH would be_____.
A 2
B 6
C 7
D 9
E 14
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Buffers
A buffer is a solution that can maintain a nearly constant pH when diluted or when strong acids or strong bases are added to it.
A buffer solution is made up of a weak acid, HA, and its conjugate base, A-, or a weak base and its conjugate acid.
http://chemcollective .org/activities /tutoria ls /buffe rs /buffe rs3
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Buffers
When a strong base is added to a buffer the hydroxide OH- from the strong base reacts with the weak acid, which gives up its H+ to form water. The weak acid neutralizes the strong base.
http://chemcollective .org/activities /tutoria ls /buffe rs /buffe rs3
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BuffersIf a strong acid is added to a buffer it will react with the weak conjugate base to form a weak acid that does not readily dissociate, and, therefore, does not significantly alter the pH.
http://chemcollective .org/activities /tutoria ls /buffe rs /buffe rs3
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59 Buffers are composed of
A Strong acids to neutralize strong bases
B Strong bases to neutralize strong acids
C A weak acid and its conjugate base
D A strong acid and its conjugate base
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60 A buffer solution contains carbonic acid (H2CO3) and bicarbonate (HCO3-). When a small amount of HCl is added to the buffer
A The HCl dissociates and the H+ significantly lowers the pH of the solution.
B The HCl dissociates and the H+ reacts with the bicarbonate to form a neutral compound.
C The pH of the solution remains stable.
D Both b and c
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61 A buffer solution contains formic acid (HCO2H) and sodium formate (HCO2Na). When a small amount of NaOH is added to the buffer
A The NaOH dissociates and the OH- significantly raises the pH of the solution.
B The formic acid neutralizes the hydroxide to form water.
C The sodium formate neutralizes the hydroxide.
D None of the above
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Buffer systems maintain a constant pH in blood
The body maintains the pH of blood at around 7.4. If the pH level changes just a few tenths of a pH unit, serious health consequences can result. A decrease in blood pH is called acidosis, an increase is called alkalosis.
There are 3 systems that regulate the pH of blood. The bicarbonate system is the most important and is controlled by the rate of respiration. In the bicarbonate system, carbon dioxide combines with water to form carbonic acid, which dissociates to form bicarbonate and hydrogen ions. CO2 + H2O H2CO3 HCO3- + H+
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62 Based on the figure below, holding one's breath can lead to which condition?
A Alkalosis
B Acidosis
C Hemolysis
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63 How does the body's response to the condition in the previous question help restore the pH of the blood?
A Breathing out reduces the amount of CO2 present, thereby reducing the production of carbonic acid.
B Breathing in increases the amount of oxygen in the blood.
C Breathing has no effect on the pH of blood.