Exam 2 - Chapter 11 Flashcards by James Ozorkiewicz

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

Arrhenius acids

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

Arrhenius acids

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• 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) = ?

HCl(aq)

hydrochloric acid

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

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

ate to ic acid or ite to ous acid

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

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

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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H2CO3 select the correct name

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

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

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HBrO2 select the correct name

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

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

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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

Arrhenius bases

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produces cations and OH− anions in an aqueous solution.

Arrhenius Bases

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According to the Brønsted–Lowry theory,

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

donates

accepts

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AciD - Donates hydrogen

Base - Brings in Hydrogren or accepts

True Dat Yo

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AciD - Donates hydrogen

Base - Brings in Hydrogren or accepts

True Dat Yo

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

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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Substances that can act as both acids and bases are

amphoteric or amphiprotic.

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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…

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

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

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_____ only partially dissociate in water.

Weak acids

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_____ is the only halogen that forms a weak acid.

Hydrofluoric acid, HF,

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

strong acid

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

weak acid

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

HI(aq) + H2O(l) —-> H3O+(aq) + I−(aq)

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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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:

HF(aq) + H2O(l) —>

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

diprotic acids

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

second dissociation

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hydronium ion

H3O+(aq)

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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)

1A (1) and 2A (2) completely

* that are poor acceptors of H+ ions. * produce very few ions in solution. * include ammonia.

Weak bases are weak electrolytes:

Strong acids have____

weak conjugate bases that do not readily accept H+.

• As the strength of the acid decreases, the strength of its | conjugate base ____.

increases

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 _____

direction of the reaction.

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

equilibrium process.

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.

Ka = [H3O+][CHO2-] / [HCHO2] HCHO2 (aq) + H2O(l) H3O+(aq) + CHO2−(aq)

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.

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.

small Ka

favoring the reactants.

large Ka

favoring the products.

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.

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

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.

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

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

amphoteric

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

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

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 ____.

neutral acidic basic

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

[H3O+] = 1.0 × 10−7 M [OH−] = 1.0 × 10−7 M | [H3O+] = [OH−] Pure water is neutral.

Adding an acid to pure water:

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

[H3O+] > [OH−]

The solution is acidic.

Adding a base to pure water:

* increases the [OH−]. * causes the [OH−] to exceed 1.0 × 10−7 M. * decreases the [H3O+]. [H3O+] < [OH−] The solution is basic.

[H3O+] < [OH−]

The solution is basic.

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

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

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

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.

``` 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 ```

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.

[H3O+] > 1.0 × 10 −7 M. so pH is

pH less than 7.0 so acidic solution

[H3O+] = 1.0 × 10 −7 M. so pH is

ph = 7 and neutral solution

[H3O+] < 1.0 × 10 −7 M. so pH is

ph > 7.0 so basic solution

• is a logarithmic scale that corresponds to the [H3O+] of aqueous solutions. • is the negative logarithm (base 10) of the [H3O+]. pH =?

The pH scale: pH = −log[H3O+]

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 = ?

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

Because pH is a log scale: * a change of one pH unit corresponds to ... * pH decreases as the [H3O+] ....

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.

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

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

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

pH = −log[4.0 × 10−5] = 4.40

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.

[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.

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.

True

Buffers work because:

• 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.

Buffers work because:

A buffer solution: • contains a combination of

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

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.

TRUE

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:

* 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+.

Because Ka is constant at a given temperature:

* 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.

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

acidosis.

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

alkalosis.

Exam 2 - Chapter 11 Flashcards

(65 cards)