Chapter 3: 3.3 Weak Acids and Bases Flashcards
How are acids and bases often defined?
- Acids: Substances that increase hydrogen ion concentration [H+] in solution
- Bases: Substances that increase hydroxide ion concentration [OH-] in solution
What is the difference between a strong acid/base vs. a weak acid/base?
Strong acids and bases completely ionize in solution, whereas weak acids and bases do not ionize completely (only proceed until equilibrium)
What is used interchangeably for the definition of strong acids?
Strong acids increase the [H+] or [H3O+] concentration of the solution (H+ combining with H2O gives the H3O+)
True or False:
Acid-base equations use equilibrium arrows
Not necessarily, strong acid/base reactions use unidirectional arrows as they ionize completely
Define:
Acids and Bases, according to the Arrhenius theory
- Acid: Produces H3O+ ions in water
- Base: Produces OH- ions in water
Define:
Acids and Bases,
according to the Brønsted-Lowry theory
- Acid: Proton donor
- Base: Proton Acceptor
True or False:
Weak acid/base reactions use two directional arrows
True, as they proceed until an equilibrium is achieved between reactants and products
State the notation for:
Equilibrium constants of weak acids and weak bases
- Weak acids: Ka
- Weak bases: Kb
Define:
Lewis Acid and Lewis Base
- Lewis Acid: A substance that can accept a pair of electrons from another atom to form a bond
- Lewis Base: A substance that donates a pair of electrons to another atom to form a bond
In terms of Lewis definition, what would amines classify as? Why?
Lewis Bases
* They have a lone pair of electrons that they can donate, making them a Lewis Base
In terms of Lewis definition, what would metal cations classify as? Why?
Lewis Acids
* As they are cations they form complex ions with Lewis Bases known as coordination complexes
In terms of Lewis definition, what would oxides of nonmetals classify as? Explain
Lewis acids
* For example, CO2 accepts electrons from OH- to form HCO3-, making it a Lewis acid
In the case of weak acids and weak bases, what would the K value be?
Less than 1, as they cannot ionize completely
In terms of acids/bases, the larger the K value…
The greater the ionization
How are Ka and Kb values often expressed? Show the notation
In a logarithmic scale:
* pKa = -logKa
* pKb = -logKb
True or False:
In terms of weak acids, there are some acids that are “weaker” and some that are “stronger”
True, “stronger” and “weaker” are relative terms
State the Ka and pKa for:
- Stronger weak acid
- Weaker weak acid
- Larger Ka (smaller pKa)
- Smaller Ka (larger pKa)
State calculation method/result of:
- pH
- pOH
- pH + pOH
- -log[H+]
- -log[OH-]
- 14
Defineand give the notation for:
Percent Ionization
The amount of the weak acid ionized (or dissociated) expressed as a percent of the original concentration of the acid, using the relationship:
* (x/c)(100%) = α
q
How are x and c found in the calculation for percent ionization?
c (initial concentration) and x (amount that ionizes) are calculated using Ka/Kb and ICE tables
True or False:
When the quadriatic formula is involved in calculating the amount that ionizes, we only use the positive value
True, as only a positive concentration is possible
In the process of calculating percent ionization, how can we determine if we can assume c-x=c?
- Calculate the ratio for c/Ka. If ratio is less than 400, quadriatic formula should be used
- If α is less than 5%; this is less helpful as it usually requires knowing x already
What is the main rule in calculating the initial concentration of the weak acid?
If you know x, USE IT
What is the relation between diluteness and percent ionization?
The more dilute the solution of a weak acid or base, the greater the percent ionization
