Acids and Bases In any sample of water, a small number of water - - PDF document

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Acids and Bases

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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

• ur muscles.)
• Slide 4 / 123

Auto-ionization of Water

In any sample of water, a small number of water molecules will dissociate into H+ and OH- ions.

H H O O H 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 O H H

• +

H H O H 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

answer

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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

answer

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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

answer

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4 The magnitude of K w indicates that _________

A water ionizes to a very small extent B the autoionization of water is exothermic C water ionizes very quickly D water ionizes very slowly

answer

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5 The molar concentration of hydronium ion,

[H3O+], in pure water at 25 °C is ___________.

A 0

B

1

C 7 D 10-7

E 10-14

answer

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+]

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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

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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 (CO32-) 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

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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 M B 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

answer

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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

answer

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8 The pacific ocean off the coast of Hawaii has

a [OH-] = 8.32 x 10-9 M. What is the [H3O+]?

answer

Slide 17 / 123 Arrhenius Definition of Acids and Bases

As we have learned, when certain substances are added to water, the H

3O+ concentration changes.

Furthermore, if the [H3O+] changes, it would influence the [OH-]. Kw = [H3O+] [OH-] = 1.0x 10-14

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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!

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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

answer

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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

answer

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)

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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

answer

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12 Which of the following could NOT act as a Bronsted

acid?

A HCN B H2SO4 C NH4+ D H3O+ E BF3

answer

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13 A Bronsted-Lowry base is defined as a

substance that __________.

A increases [H+] when placed in H2O B decreases [H+] when placed in H2O C increases [OH-] when placed in H2O D acts as a proton acceptor E acts as a proton donor

answer

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14 Which of the following compounds could never

act as a Bronsted acid?

A SO4 2- B HSO4 - C H2SO4 D NH3 E CH3COOH

answer

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Bronsted Acids and Bases (In Depth)

Acids and Bases go together It 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 O H 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-

+

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Bronsted Acids and Bases (In Depth)

Identifying an acid or a base By examining the products and reactants of a chemical reaction,

• ne can identify if a substance is behaving as an acid or as a

base. Example HSO4-(aq) + CN-(aq) --> SO4 2- (aq) + HCN(aq) HSO4-(aq) donated an H+ to become SO4 2- = 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! H2O(l) + CH3NH3 +(aq) --> CH3NH2(aq) + H3O+(aq) CH3NH3 +(aq) donated an H+ to become CH3NH2 = It's an acid! H2O(aq) accepted an H+ to become H3O+ = It's a base! move for answer

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15 According to the following reaction, which

reactant molecule is acting as an acid?

A H2SO4 B H2O C H3O+ D HSO4 - E None of the above

H2O + H2SO4 → H3O+ + HSO4 -

answer

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16 According to the following reaction, which

reactant molecule is acting as a base? H2O + H2SO4 → H3O+ + HSO4 -

A H2SO4 B H2O C H3O+ D HSO4 - E None of the above

answer

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17 According to the following reaction, which reactant

molecule is acting as a base? H3O+ + HSO4 - → H2O + H2SO4

A H2SO4 B H2O C H3O+ D HSO4 - E None of the above

answer

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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

H3PO4 + H2O # H2PO4 - + H3O+

answer

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19 For the following reaction, identify whether the

circled compound is behaving as an acid or a base. H3PO4 + H2O # H2PO4 - + H3O+

A Acid B Base C Neither D Both E None of the above

answer

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. Example F-(aq) + H2O(l) <--> HF(aq) + OH-(aq) HF(aq) donates an H+ ion to become F-(aq) = It's an acid OH-(aq) accepts an H+ to become H2O(l) = It's a base

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20 For the following reaction, identify whether the

circled compound is behaving as an acid or a base. H3PO4 + H2O # H2PO4 - + H3O+

A Acid B Base C Neither D Both E None of the above

answer

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21 For the following reaction, identify whether the

circled compound is behaving as an acid or a base. H3PO4 + H2O # H2PO4 - + H3O+

A Acid B Base C Neither D Both E None of the above

answer

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) + H

2O(l) NO 2 - (aq) + H 3O+(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 charges If you accept an H+, you become more positive If you donate an H+, you become more negative

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22 Which of the following would be the conjugate

base of HNO2?

A NO2 - B H2NO2 C NO2 D NO2 2- E HNO2

answer

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23 Which would be the conjugate acid of HCO3 -(aq)?

A CO3 2- B HCO3 C CO3 D H2CO3 - E H2CO3

answer

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24 What would be the an acid/conjugate pair in the

following reaction?

A NH2-/H2O B NH2-/NH3 C H2O/OH- D H2O/NH3 E None of these NH2 - + H2O --> NH3 + OH-

answer

Slide 48 / 123 Lewis Acids and Bases

Definition Scientists 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-, H2O) are good examples. Lewis Acid (e- pair acceptor) Lewis Base (e- pair donor)

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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-]

answer

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26 Which of the following would likely act as a lewis

acid?

A NH3 B OH- C CN- D H2O E Fe3+

answer

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

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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

• f a definition

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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.

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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 + H

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.

Strong W e a k Negligible S t r

• n

g W e a k Negligible

Acid Base HCl Cl - H2SO4 HSO 4- HNO3 NO 3- H3O+ H 2O HSO4- SO 42- H3PO4 H2PO4- HF F- HC2H3O2 C2H3O2- H2CO3 HCO 3- H2S HS - H2PO4- HPO 42- NH4+ NH 3 HCO3- CO 32- HPO42- PO 43- H2O OH - OH- O 2- H2 H - CH4 CH 3- 100% protonated in H

2O

Base strength increases Acid strength increases 100% ionized in H

2O

Slide 57 / 123 Acid and Base Strength

Weak acids only ionize partially in water. Their conjugate bases are weak bases.

S t r

• n

g W e a k Negligible S t r

• n

g W e a k Negligible

Acid Base HCl Cl - H2SO4 HSO 4- HNO3 NO 3- H3O+ H 2O HSO4- SO 42- H3PO4 H2PO4- HF F- HC2H3O2 C2H3O2- H2CO3 HCO 3- H2S HS - H2PO4- HPO 42- NH4+ NH 3 HCO3- CO 32- HPO42- PO 43- H2O OH - OH- O 2- H2 H - CH4 CH 3- 100% protonated in H

2O

Base strength increases Acid strength increases 100% ionized in H

2O

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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

Strong W e a k Negligible S t r

• n

g W e a k Negligible

Acid Base HCl Cl - H2SO4 HSO 4- HNO3 NO 3- H3O+ H 2O HSO4- SO 42- H3PO4 H2PO4- HF F- HC2H3O2 C2H3O2- H2CO3 HCO 3- H2S HS - H2PO4- HPO 42- NH4+ NH 3 HCO3- CO 32- HPO42- PO 43- H2O OH - OH- O 2- H2 H - CH4 CH 3- 100% protonated in H

2O

Base strength increases Acid strength increases 100% ionized in H

2O

Slide 59 / 123 Strong Acids

There are seven strong acids: 3 contain a H bound to the very electronegative halogens: HCl hydrochloric acid HBr hydrobromic acid HI 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+.

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27 Which of the following is NOT a strong acid?

A HBr B HF C HI D HCl E A and C

answer

Slide 61 / 123 Strong Acids

There are seven strong acids: 4 are from the very electron drawing

• xyanions:

HNO3 nitric acid H2SO4 sulfuric acid HClO3 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 acid HBr hydrobromic acid HI hydroiodic acid HNO3 nitric acid H2SO4 sulfuric acid HClO3 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."

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28 What would be the [H3O+] in a 0.005 M HBr solution?

answer

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In any acid-base reaction, the proton moves toward the stronger base. In

• ther words, a stronger base will "hold
• nto" 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

W e a k N e g l i g i b l e Strong Weak Negligible

Acid Base HCl Cl - H2SO4 HSO 4- HNO3 NO 3- H3O+ H 2O HSO4- SO 42- H3PO4 H2PO4- HF F- HC2H3O2 C2H3O2- H2CO3 HCO 3- H2S HS - H2PO4- HPO 42- NH4+ NH 3 HCO3- CO 32- HPO42- PO 43- H2O OH - OH- O 2- H2 H - CH4 CH 3- 100% protonated in H

2O

Base strength increases Acid strength increases 100% ionized in H

2O

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29 What would be the [OH-] in a 0.034 M NaOH solution?

answer

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)

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30 Strong acids have ___________ conjugate

bases.

A strong

B

weak

C

neutral

D negative

answer

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31 HBr, hydrobromic acid is a strong acid. This

means that _______________.

A

aqueous solutions of HBr contain equal concentrations

• f 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

answer

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

pH

Recall that in pure water, the ion-product is Kw = 1.0*10-14 = [H3O+][OH-] In pure water, the hydronium ion concentration and hydroxide ion concentrations are equal: [H3O+] = [OH-] Therefore, in pure water, [H3O+] = [OH-] = 1.0 x 10-14 = 1.0 x 10-7 M pH = -Log[H3O+] = - Log(1.0 x10-7 ) = 7 A pH of 7 is considered neutral on the pH scale

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

• 5.00

C 5.00 D 9.00 E

• 9.00

answer

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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

answer

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34 What is the pH of a solution whose hydronium

ion concentration is 7.14 x 10-3 M?

answer

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35 What is the pH of a solution whose hydronium

ion concentration is 1.92 x 10-9 M?

answer

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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)

answer

Slide 79 / 123 pH

Application 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 [H3O+] is found to be 3.4 X10-3 M. Is there a problem with the pH?

Slide 80 / 123

What is the relationship between [H3O+] and the pH value?

Below are three different [H3O+]. 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.

Slide 81 / 123

What is the relationship between [H3 O+], the pH value, and the acidity and basicity of a solution?

pH

low H3O+ High H3O+ high OH- low OH- acidic acidic basic basic neutral

Slide 82 / 123

pH

These are the pH values for several common substances.

More acidic More basic

Battery acid lemon juice pure rain or water distilled water sea water baking soda household ammonia household bleach household lye gastric fluid carbonated beverages vinegar

• range juice

beer coffee egg yolks milk blood milk of magnesia

Slide 83 / 123

For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

How Do We Measure pH? Slide 84 / 123 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)

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

Slide 85 / 123

pH

Solution type [H +](M) [OH-] (M) pH value Acidic > 1.0x10-7 <1.0x10-7 <7.00 Neutral =1.0x10-7 =1.0x10-7 =7.00 Basic <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. BASE ACID

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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

answer

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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

answer

Slide 88 / 123

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

answer

Slide 89 / 123

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

answer

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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

answer

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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.

answer

Slide 92 / 123

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.

answer

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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

answer

Slide 94 / 123 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.

Slide 95 / 123

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

answer

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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

answer

Slide 97 / 123

pOH

Just 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 14 low 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

answer

Slide 100 / 123

pOH

Once 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

Slide 101 / 123

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 5 B 15 C 7 D 8 E 9

answer

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49 What is the pH of an aqueous ammonia solution with

a [OH-] = 1 x 10-4 M?

A 4 B 1 x10-4 C 10 D 1 x10-10 E 3

answer

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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

answer

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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)

answer

Slide 106 / 123

52 Which of the following would be LEAST acidic?

A pOH = 2 B pOH = 4 C pH = 10 D pH = 2 E pH = 11

answer

Slide 107 / 123 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-]

Slide 108 / 123 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

answer

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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

answer

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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

answer

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56 What is the hydronium ion concentration in a

solution whose pH = 4.29?

answer

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57 For a 1.0-M solution of a strong acid, a

reasonable pH would be_____. A B 6 C 7 D 9 E 13

answer

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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

answer

Slide 115 / 123

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/tutorials/buffers/buffers3

Slide 116 / 123 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/tutorials/buffers/buffers3

Slide 117 / 123 Buffers

If 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/tutorials/buffers/buffers3

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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

Slide 119 / 123

60 A buffer solution contains carbonic acid (H2CO3) and bicarbonate (HCO3-). When a small amount

• f 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

Slide 121 / 123 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 CO

2 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.