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Biochemistry Exam 4 > Exam 2 - Chapter 11 > Flashcards

Exam 2 - Chapter 11 Flashcards

1
Q

produce hydrogen ions (H+) when they dissolve in water.

A

Arrhenius acids

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2
Q
  • are also electrolytes, because they produce H+ in water.
  • have a sour taste.
  • turn blue litmus paper red.
  • corrode some metals.
A

Arrhenius acids

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3
Q

• Acids with a hydrogen ion (H+) and a nonmetal (or CN−)
ion are named with the prefix hydro and end with ic acid.

HCl (aq) = ?

A

HCl(aq)

hydrochloric acid

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4
Q

Acids with a hydrogen ion (H+) and a polyatomic ion are

named by changing the end of the name of the polyatomic ion from

A

ate to ic acid or ite to ous acid

ClO3− chlorate ClO2− chlorite
HClO3 chloric acid HClO2 chlorous acid

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5
Q

HBr select the correct name
A. bromic acid
B. bromous acid
C. hydrobromic acid

A

HBr Br−, bromide

C. hydrobromic acid

The name of an acid with a hydrogen ion (H+) and a nonmetal uses the prefix hydro and ends with ic acid.

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6
Q

H2CO3 select the correct name

A. carbonic acid B. hydrocarbonic acid C. carbonous acid

A

An acid with a hydrogen ion (H+) and a polyatomic ion ending in ate is called an ic acid.

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7
Q

HBrO2 select the correct name

A. bromic acid B. hydrobromous acid C. bromous acid

A

An acid with a hydrogen ion (H+) and a polyatomic ion ending in ite is called an ous acid.

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8
Q
  • produce hydroxide ions (OH −) in water.
  • taste bitter or chalky.
  • are also electrolytes, because they produce hydroxide ions (OH−) in water.
  • feel soapy and slippery.
  • turn litmus indicator paper blue and phenolphthalein indicator pink
A

Arrhenius bases

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9
Q

produces cations and OH− anions in an aqueous solution.

A

Arrhenius Bases

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10
Q

According to the Brønsted–Lowry theory,

  • an acid is a substance that ____ H+.
  • a base is a substance that _____ H+.
A

donates

accepts

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11
Q

AciD - Donates hydrogen

Base - Brings in Hydrogren or accepts

A

True Dat Yo

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12
Q

AciD - Donates hydrogen

Base - Brings in Hydrogren or accepts

A

True Dat Yo

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13
Q

In any acid–base reaction, there are two conjugate acid–base pairs.

A
  • Each pair is related by the loss and gain of H+.
  • One pair occurs in the forward direction.
  • One pair occurs in the reverse direction.
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14
Q

Substances that can act as both acids and bases are

A

amphoteric or amphiprotic.

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15
Q

For water, the most common amphoteric substance, the acidic or basic behavior depends on the other reactant.

Water donates H+ when…

Water accepts H+ when it reacts with a…

A

Water donates H+ when it reacts with a stronger base.

Water accepts H+ when it reacts with a stronger acid.

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16
Q

_____ only partially dissociate in water.

A

Weak acids

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17
Q

_____ is the only halogen that forms a weak acid.

A

Hydrofluoric acid, HF,

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18
Q

A _____ completely ionizes (100%) in aqueous solutions.

A

strong acid

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19
Q

A ____ dissociates only slightly in water to form a few ions in aqueous solutions.

A

weak acid

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20
Q

In water, the dissolved molecules of HA, a strong acid:

HI(aq) + H2O(l) —-> H3O+(aq) + I−(aq)

A
  • dissociate into ions 100%.
  • produce large concentrations of H3O+ and the anion (A−).

The strong acid HI dissociates completely into ions.

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21
Q

In weak acids, only a few molecules dissociate.

• Most of the weak acids remain as the undissociated
(molecular) form of the acid.

• The concentrations of H3O+ and the anion (A−) are small.

HF is a weak acid:

A

HF(aq) + H2O(l) —>

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22
Q

Some weak acids, such as carbonic acid, are _____that have two H+, which dissociate one at a time.

A

diprotic acids

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23
Q

Because HCO3− is also a weak acid, a _____
can take place to produce another hydronium ion and the
carbonate ion, CO32−.

A

second dissociation

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24
Q

hydronium ion

A

H3O+(aq)

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25
Q

Strong bases as strong electrolytes:

• are formed from metals of
Groups _______-

• include LiOH, NaOH, KOH,
Ba(OH)2, Sr(OH)2, and Ca(OH)2.

• dissociate _____ in water.
KOH(s) —–> K+(aq) + OH −(aq)

A

1A (1) and 2A (2)

completely

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26
Q
  • that are poor acceptors of H+ ions.
  • produce very few ions in solution.
  • include ammonia.
A

Weak bases are weak electrolytes:

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27
Q

Strong acids have____

A

weak conjugate bases that do not readily accept H+.

28
Q

• As the strength of the acid decreases, the strength of its

conjugate base ____.

A

increases

29
Q

In any acid–base reaction, there are two acids and two bases.

  • However, one acid is stronger than the other acid, and one base is stronger than the other base.
  • By comparing their relative strengths, we can determine the _____
A

direction of the reaction.

30
Q

Because the dissociation of strong acids in water is essentially complete, the reaction is not considered to be an…

A

equilibrium process.

31
Q

As with other dissociation expressions,

• the molar concentration of the products is divided by the
molar concentration of the reactants.

• water is a pure liquid with a constant concentration and is
omitted.

• the expression is called acid dissociation constant, Ka.

A

Ka = [H3O+][CHO2-] / [HCHO2]

HCHO2 (aq) + H2O(l) H3O+(aq) + CHO2−(aq)

32
Q

When the value of the Ka

• is ____, the equilibrium lies to the left, favoring the
reactants.

• is ___, the equilibrium lies to the right, favoring the
products.

A

When the value of the Ka

  • is small, the equilibrium lies to the left, favoring the reactants.
  • is large, the equilibrium lies to the right, favoring the products.
33
Q

small Ka

A

favoring the reactants.

34
Q

large Ka

A

favoring the products.

35
Q

When the value of the Kb

  • is small, the equilibrium lies to the left, favoring the reactants.
  • is large, the equilibrium lies to the right, favoring the products.
A

The stronger the base is, the larger will be the Kb value.

36
Q

When the value of the Kb

  • is small, the equilibrium lies to the left, favoring the reactants.
  • is large, the equilibrium lies to the right, favoring the products.
A

The stronger the base is, the larger will be the Kb value.

37
Q

Water is ____—it can act as an acid or a base.

A

amphoteric

38
Q

The ion product constant for water, Kw, is defined as:

A

• the product of the concentrations of H3O+ and OH −.
• equal to 1.0 × 10−14 at 25 °C (the concentration units are
omitted).

39
Q

When • [H3O+] and [OH−] are equal, the solution is ____.
• [H3O+] is greater than the [OH−], the solution is _____.
• [OH−] is greater than the [H3O+], the solution is ____.

A

neutral
acidic
basic

40
Q

In pure water, the ionization of water molecules produces small but equal quantities of H3O+ and OH − ions.

Expressed mathematically? Pure water is?

A

[H3O+] = 1.0 × 10−7 M [OH−] = 1.0 × 10−7 M

[H3O+] = [OH−] Pure water is neutral.

41
Q

Adding an acid to pure water:

A
  • increases the [H3O+].
  • causes the [H3O+] to exceed 1.0 × 10 −7 M.
  • decreases the [OH−]. The solution is acidic.
42
Q

[H3O+] > [OH−]

A

The solution is acidic.

43
Q

Adding a base to pure water:

A
  • increases the [OH−].
  • causes the [OH−] to exceed 1.0 × 10−7 M.
  • decreases the [H3O+].

[H3O+] < [OH−] The solution is basic.

44
Q

[H3O+] < [OH−]

A

The solution is basic.

45
Q

What is the [H3O+] of a solution if [OH−] is 5.0 × 10−8 M?

A

Because the [H3O+] of 2.0 × 10–7 M is larger than the [OH−] of 5.0 × 10–8 M, the solution is acidic.

46
Q

If lemon juice has [H3O+] of 2.0 × 10–3 M, what is the [OH−] of the solution?

A

Because the [H3O+] concentration of 2.0 × 10−3 M is greater than the [OH−] of 5.0 × 10−12 M, the solution is acidic.

47
Q 
The pH of a solution: 
• is used to indicate the 
• has values that usually range from 
• is acidic when the values are 
• is neutral at a pH of  
• is basic when the values are
A

The pH of a solution:
• is used to indicate the acidity of a solution.
• has values that usually range from 0 to 14.
• is acidic when the values are less than 7.
• is neutral at a pH of 7.
• is basic when the values are greater than 7.

48
Q

[H3O+] > 1.0 × 10 −7 M. so pH is

A

pH less than 7.0 so acidic solution

49
Q

[H3O+] = 1.0 × 10 −7 M. so pH is

A

ph = 7 and neutral solution

50
Q

[H3O+] < 1.0 × 10 −7 M. so pH is

A

ph > 7.0 so basic solution

51
Q

• is a logarithmic scale that corresponds to the [H3O+] of
aqueous solutions.

• is the negative logarithm (base 10) of the [H3O+]. pH =?

A

The pH scale:

pH = −log[H3O+]

52
Q

To calculate the pH, the negative powers of 10 in the molar

concentrations are converted to positive numbers. If [H3O+] is 1.0 × 10−2 M: pH = ?

A

pH = −log[1.0 × 10−2 ] = −(−2.00) = 2.00

53
Q

Because pH is a log scale:

  • a change of one pH unit corresponds to …
  • pH decreases as the [H3O+] ….
A

Because pH is a log scale:

  • a change of one pH unit corresponds to a tenfold change in [H3O+].
  • pH decreases as the [H3O+] increases.
54
Q

pH 2.00 is [H3O+] = ?
pH 3.00 is [H3O+] = ?
pH 4.00 is [H3O+] = ?

A

pH 2.00 is [H3O+] = 1.0 × 10 −2 M
pH 3.00 is [H3O+] = 1.0 × 10−3 M
pH 4.00 is [H3O+] = 1.0 × 10−4 M

55
Q

Find the pH of a solution with a [H3O+] of 4.0 × 10−5.

A

pH = −log[4.0 × 10−5] = 4.40

56
Q

Given the pH of a solution, we can reverse the calculation to obtain the [H3O+].

• For whole number pH values, the negative pH value is the power of 10 in the [H3O+] concentration.

A

[H 3O+] = 10 ^−pH

• For pH values that are not whole numbers, the calculation
requires the use of the 10x key, which is usually a 2nd
function key.

57
Q

When an acid or a base is added to water, the pH changes drastically. In a buffer solution, the pH is maintained; pH does not change when acids or bases are added.

A

True

58
Q

Buffers work because:

A

• they resist changes in pH from the addition of an acid or
a base.

• in the body, they absorb H3O+ or OH − from foods and
cellular processes to maintain pH.

• they are important in the proper functioning of cells
and blood.

• they maintain a pH close to 7.4 in blood. A change in the pH of the blood affects the uptake of oxygen and cellular processes.

59
Q

Buffers work because:

A

• they resist changes in pH from the addition of an acid or
a base.

• in the body, they absorb H3O+ or OH − from foods and
cellular processes to maintain pH.

• they are important in the proper functioning of cells
and blood.

• they maintain a pH close to 7.4 in blood. A change in the pH of the blood affects the uptake of oxygen and cellular processes.

60
Q

A buffer solution: • contains a combination of

A

acid–base conjugate pairs, a weak acid, and a salt of its conjugate base

61
Q

The pH of the solution is maintained as long as the added amounts of acid or base are small compared to the concentrations of the buffer components.

A

TRUE

62
Q

The arterial blood plasma has a normal pH of 7.35–7.45. If changes in H3O+ lower the pH below 6.8 or raise it above 8.0, cells cannot function properly and death may result. In our cells, CO2:

A
  • is continually produced as an end product of cellular metabolism.
  • is carried to the lungs for elimination, and the rest dissolves in body fluids such as plasma and saliva, forming carbonic acid, H2CO3.As a weak acid, carbonic acid dissociates to give bicarbonate, HCO3−, and H3O+.
63
Q

Because Ka is constant at a given temperature:

A
  • the [H3O+] is determined by the [HC2H3O2]/[C2H3O2−] ratio.
  • the addition of small amounts of either acid or base changes the ratio of [HC2H3O2]/[C2H3O2−] only slightly.

• the changes in [H3O+] will be small and the pH will be
maintained.

• the addition of a large amount of acid or base may exceed the buffering capacity of the system.

64
Q

If the CO2 level rises, increasing H2CO3, the equilibrium
shifts to produce more H3O+, which lowers the pH. This
condition is called

A

acidosis.

65
Q

A lowering of the CO2 level leads to a high blood pH, a

condition called

A

alkalosis.

Biochemistry Exam 4 (19 decks)

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