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Chemistry 3202 > Unit 2 - Acids & Bases > Flashcards

Unit 2 - Acids & Bases Flashcards

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0
Q

Observable Properties of Acids

A
  • Turns blue litmus paper red
  • pH < 7
  • Taste sour
  • React with active metals to produce H2(g)
  • React with carbonates to produce CO2 + H2O
  • Neutralize bases
  • Form electrically conductive solutions when dissolved in water (electrolytic)
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1
Q

Operational Definitions

A
  • definitions based on observable properties. (E.g., acids were sour, bases bitter and slippery)
  • developed by alchemists
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2
Q

Observable Properties of Bases

A
  • Turn red litmus paper blue
  • pH > 7
  • Taste bitter
  • Feel slippery
  • Neutralize acids
  • React with fats and oils to produce soaps and detergents
  • Are (themselves) electrolytic
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3
Q

Conceptual/theoretical definitions

A
  • Explanations for observable properties
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4
Q

Antoine Lavoisier

A
  • Suggested that acids are substances that contain oxygen combined with
    nonmetallic elements
  • e.g., SO2 in water produced sulfurous acid
  • This idea did not explain the acidic behaviour of substances like HCl
  • Theory eventually rejected
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5
Q

Sir Humphrey Davy

A
  • Suggested the presence of hydrogen in a compound accounted for the behaviour of acids
  • Did not explain why hydrogen contained compounds like methane,
    sucrose, ethanol and water did not behave like acids
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6
Q

Svante Arrhenius

A
  • Worked with electrolytes
  • Stated that acids produce hydrogen ions (H+) when dissolved in water and
    bases produce hydroxide ions (OH-) when dissolved in water
  • Definitions are still in use today, but they have their limitations
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7
Q

Arrhenius Theory of Acids and Bases

A
  • An acid is a substance that dissociates in water to produce one or more hydrogen ions
  • A base is a substance that dissociates in water to form one or more hydroxide ions
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8
Q

Arrhenius Theory and Polarity

A

Don’t assume that all compounds containing hydrogen ions are acids or that all the hydrogen ions in an acid are released as hydrogen ions.

  • Only the hydrogen ions in very polar bonds are ionizable.
  • These are bonds in which hydrogen is joined to a very electronegative element
  • For example, hydrogen ions are released when molecules that contain such bonds are dissolved in water
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9
Q

Define Monoprotic Acid, Diprotic Acid & Triprotic Acid

A

Monoprotic Acid: Contains one ionizable hydrogen

Diprotic acid: Contains 2 ionizable hydrogens. Sulfuric acid is an example

Triprotic acid: Contains 3 ionizable hydrogens. Phosphoric acid is an example

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10
Q

Limitations of Arrhenius Theory

A

1) H+ can’t exist as an ion in water. The positive H+ ions are attracted to the polar water molecules forming HYDRONIUM ions or H3O+(aq)

2) CO2 dissolves in water to produce
an acid. NH3 dissolves in water to produce a base. Neither of these observations can be explained by Arrhenius theory

3) Some acid-base reactions can occur in solvents other than water. Arrhenius theory can explain only aqueous acids or bases.

4) Arrhenius theory is not able to
predict whether certain species are
acids or bases. eg. NaHSO4

5) Some compounds that form basic solutions (such as bicarbonate) can actually neutralize stronger bases. Arrhenius theory, therefore, cannot accurately predict whether a given aqueous compound is acidic or basic

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11
Q

Modern Arrhenius Theory

A
  • Acknowledges role of water and production of hydronium ions
  • Also, bases are substances that react with water to produce hydroxide ions
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12
Q

Bronsted Lowry Theory

A

•According to Bronsted-Lowry theory there is only one requirement for an acid-base reaction - one substance must provide a proton and another substance must receive the same proton.

  • in other words an acid-base reaction involves the transfer of a proton (the
    nucleus of a hydrogen ion only in the case of acid-base reactions)
  • any substance can act as an acid, but only if another substance behaves as a base at the same time
  • any substance can act as a base, but only if another substance behaves as
    an acid at the same time

• Like an Arrhenius acid, a Bronsted-Lowry acid must contain H in its formula. This means that all Arrhenius acids are also Bronsted-Lowry acids. However, any negative ion, (not just OH) can be a Bronsted-Lowry base.

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13
Q

Conjugate acid (of a base)

A

• The particle that forms when a base gains a hydrogen ion (proton) from an acid

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14
Q

Conjugate base (of an acid)

A

The particle that remains when an acid has donated a hydrogen ion (i.e., when a proton is removed from an acid)

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15
Q

Conjugate acid-base pair

A
  • Two substances that are related by the loss or gain of a single hydrogen ion
  • Conjugate acids and bases are always paired with an acid and a base
  • The conjugate base of an acid-base pair always has one less hydrogen than the acid. It also has one more negative charge than the acid
  • The conjugate acid of an acid-base pair always has one more hydrogen than the base. It also has one less negative charge than the base
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16
Q

Amphoteric species

A

• A substance can only be classified as a Bronsted-Lowry acid or base for a
specific reaction

  • Protons may be gained in a reaction with one substance, but lost in a reaction with another substance
  • Must have at least one ionizable hydrogen atom
  • Example: Water
  • Behaves as a base in acid-base reactions with strong acids
  • Behaves as an acid in acid-base reactions with strong bases
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17
Q

How do you know when an amphoteric species will behave like an acid or a base in an acid-base reaction?

A
  • You have to know whether the substance is a stronger or weaker acid or base
  • This can be determined using a data table: acids are ranked on their percent reaction (percent ionization) with water. The rank of an acid determines the rank of the strength of its conjugate base
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18
Q

What does the strength of an acid or base mean?

A

Strength refers to the extent to which a substance dissociates in its solvent.

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19
Q

What is a strong acid?

A

Strong acid - an acid that dissociates completely (e.g., HCl)

  • strong acids are listed on the table handed out in class
    • the concentration of hydronium ions in a dilute solution of a strong acid is equal to the concentration of the acid
      e. g., a 1.0 mol/L solution of hydrochloric acid contains 1.0 mol/L of hydronium ions and 1.0 mol/L of chloride ions
  • that is, no HCl remains; 100% of the solution is H3O+ and Cl- ions.
20
Q

What is a weak acid?

A

Weak acid - an acid that dissociates only slightly in a water solution.
- only a small percentage of the acid molecules dissociate
- most of the acid molecules remain intact
- the majority of acids are weak acids
- the concentration of hydronium ions in a weak acid is always less than the
concentration of the dissolved acid*
e.g. acetic acid is a weak acid with only about a 1% dissociation in a 0.1 mol/L solution. Note: the number of acid molecules that dissociate depends on the concentration and temperature of the solution.
* so, in the case of acetic acid, 99% remains as the acid (CH3COOH) and 1% dissociates into CH3COO- and H3O+

21
Q

What is a strong base?

A

Strong base
- dissociates completely into ions in water
- all oxides and hydroxides of the alkali metals (Group 1A) are strong bases
as are the oxides and hydroxides of Group 2A below beryllium
- the concentration of hydroxide ions in a dilute solution of a strong base is
equal to the concentration of the base*
e.g., a 1.0 M solution of NaOH contains 1.0 mol/L of hydroxide ions * that is, no NaOH remains. It’s all Na+ and OH_ ions

22
Q

What is a weak base?

A

weak bases
- most bases are weak
- a weak base dissociates only very slightly in water
- the most common weak base is ammonia; in a 0.1 M solution only about 1%
of the ammonia molecules react with water to form hydroxide ions

23
Q

How can the table be used to predict the direction in which an acid-base reaction will proceed i.e., whether products or reactants are favored in the reaction

A
  • The direction of an acid-base reaction usually proceeds from a stronger acid and a stronger base to a weaker acid and a weaker base.
  • if the reaction proceeds to the right (that is, if the stronger acid and stronger base are on the left side of the equation), the products are favored
  • if the reaction goes to the left (if the stronger acid and base are on the right side of the equation), reactants are favored
24
Q

Calculations that involve strong acids or bases

A

You will recall that when a strong acid dissociates in water, the concentration of H3O+ is equal to the concentration of the strong acid.
Also, when a strong base dissociates in water, the concentration of OH- is equal to the concentration of the strong base.
You must therefore first determine that you are working with a strong acid or base
If you know the concentration of a solution, you can calculate the concentrations of the dissociated or ionized species using the mole ratio.

25
Q

Kw

A
  • The self-ionization of water produces a system at equilibrium for which we can write an equilibrium constant for water, Kw.
  • Kw will be a very small value. This means that the left side of the equilibrium is highly favored. However, we can tell that this equilibrium does in fact exist because pure water will conduct electricity to a very small degree.
26
Q

Self ionization of water

A

Water + water hydronium + hydroxide

  • This equation represents the auto-ionization of water - the formation of ions by a molecular substance without the addition of another substance.
  • Since the products of the reaction are hydronium and hydroxide ions, the equation shows that water is an amphoteric substance (i.e., can act as either an acid or a base)
27
Q

What is the concentration of hydronium ions in pure water

A

1.0 x 10-7 mol/l. (Therefore, the concentration of hydroxide ions in the same solution must also be 1.0 x 10-7 mol/l). At 25 degrees celsius

28
Q

What is Kw called?

A

Kw, the product of the concentrations of the hydronium ions and the hydroxide ions in water, is called the ion-product constant for water

29
Q

How to determine equilibrium constant for water?

A

(H3O+)(OH-) = (1.0 x 10-7) = (1.0 x 10-14)

31
Q

The concentration of H3O+ ions in the solution of a strong acid is equal to …

A

the concentration of the dissolved acid.
Consider [H3O+] in a solution of 0.1 mol/L hydrochloric acid
- all the HCl molecules will dissociate in water forming a [H3O+] = 0.1 M
- the increased [H3O+] pushes the dissociation reaction between water molecules to the left (Le Chatlier’s Principle)
- therefore the [H3O+] that results from the dissociation of water is even less than 1 x 10-7 M. This concentration is negligible compared with the 0.1 M concentration of the HCl

32
Q

[H3O+] and [OH-] in aqueous solutions at 25 oC

A

In an acidic solution, [H3O+] is greater than 1.0 x 10-7 M and [OH-] is less than 1.0 x 10-7 M In a neutral solution, both [H3O+] and [OH-] are equal to 1.0 x 10-7 M
In a basic solution, [H3O+] is less than 1.0 x 10-7 M and [OH-] is greater than 1.0 x 10-7 M
You can use the ion product constant for water, Kw, to determine the concentration of H3O+ or OH- when the concentration of the other ion is known.

33
Q

Find the [H3O+] and [OH-] in each of the following solutions:

a. 2.5 M nitric acid
b. 0.16 M barium hydroxide

A

You know that nitric acid is a strong acid and barium hydroxide is a strong base. Since both dissociate completely in aqueous solutions, you can use their molar concentrations to determine [H3O+] and [OH-], respectively. You can find the concentration of the other ion using Kw.
Note: For a solution of a strong acid,as in H3O+ should be greater than 1.0x10-14 and [OH-] should be less than 1.0 x 10-14 . For a solution of a strong base, [OH-] should be greater than 1.0 x 10-14 and [H3O+] should be less than 1.0 x 10-14 .

34
Q

Background of pH and pOH

A

The concentration of hydronium ions ranges from about 10 M for a concentrated strong acid to about 10-15 M for a concentrated strong base.

  • this wide range of concentrations, and the negative powers of 10, are not very convenient to work with
  • in 1909, a Danish biochemist named Soren Sorensen suggested a method for converting concentrations to positive numbers; this involved the use of logarithms
  • values that represent the concentrations of hydronium and hydroxide ions in a solution
  • base 10 logarithms (exponents whose base is 10)
35
Q

Logarithms

A

the logarithm of a number is the power to which you must raise 10 to equal that number
e.g., the logarithm of 10 is 1 because 101 = 10 the logarithm of 100 is 2 because 102 = 100

36
Q

How to calculate pH

A

pH = -log [H3O+]

37
Q

logarithmic scales

A
  • used to compress values that range over several orders of magnitude
  • pH scale and the Richter scale are both logarithmic scales
38
Q

pH scale

A
  • developed by Soren Sorenson in early 1900s to provide a convenient way of showing different levels of acidity and alkalinity.
39
Q

Relationship between hydronium and hydroxide ions

A
  • at low pH, the concentration of hydronium is high and the concentration of hydroxide is low
  • as pH increases, hydronium ion concentration decreases and hydroxide ion concentration increases
  • this relationship is closely related to Kw (more later)
40
Q

Sig Figs in pH

A

Count only the digits to the right of the decimal point

41
Q

Dilution

A
  • the process of decreasing the concentration of a solution
42
Q

Strength vs Concentration

A

The strength of an acid, or base, is a function of the extent of its reaction with water.
The concentration of a solution is a function of the number of moles of solute dissolved per unit of solvent.

43
Q

Concentrated and dilute

A
  • refer to how much of an acid or base is dissolved, to the number of moles of the acid or base in a given volume
44
Q

Strong and weak

A
  • refer to the extent of ionization of the acid or base, indicate how many of the molecules dissociate into ions
45
Q

pH can be determined by using:

A
  • pH meter. Operates on the principle of relative conductivities of solutions
  • pH paper, pH indicators. Rely on chemical reactions involving special pigments that result in distinctive color changes
46
Q

A buffer

A

A mixture if a weak conjugate acid-base pair that maintains a relatively constant pH when moderate amounts of acid or base are added to it.

47
Q

Buffer capacity

A
  • The number of moles of strong acid (H3O+ ) or strong base (OH-) that the buffer can absorb before the ratio of weak acid to conjugate base (or weak
    base to conjugate acid) differs by more than a factor of 10
  • since a buffer may contain different concentrations of its 2 components, its ability to absorb hydroxide will differ from its ability to absorb hydronium
  • when the ratio of the concentration of the buffer components is close to 1,
    the buffer capacity is at its maximum
  • also influenced by the amount of a buffer used
  • 100 mL of a buffer has 2x the buffering capacity of 50 mL
  • higher concentrations of the 2 components of a buffer mixture will increase its buffering capacity (i.e., a more concentrated buffer resists pH changes more than a dilute buffer).
48
Q

An unbuffered solution can’t neutralize even small amounts of acid or base

A
  • a dramatic change in pH is observed in an unbuffered solution immediately upon the addition of a strong acid or a strong base
  • for example, a neutral solution of saline whose initial pH is 7.0 will have a pH of 2.0 after the addition of 1 mL of 10 M HCl.
    S however, blood plasma, whose initial pH is 7.4, will have a pH of 7.2 after the addition of the same amount of HCl
49
Q

Buffers in out blood

A

We rely on the carbonic acid/hydrogen carbonate ion buffer to keep our blood plasma at a stable pH of 7.4. A change of more than 0.4 units of pH can cause death.
Actually, arterial blood has a pH of 7.4; venous blood pH is slightly lower due to the acidifying effect of higher carbon dioxide concentration.

Chemistry 3202 (4 decks)

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