Chapter 16 - Acid-Base Equilibria Flashcards
16.1
Define “Arrhenius acid”.
Substance that increases concentration of H+ ions in H2O
Applies only to solutions
16.1
Define “Arrhenius base”.
Substance that increases concentration of OH- ions in H2O.
Applies only to solutions
16.2
Define “Bronsted-Lowry acid”.
- Donates a proton to another substance.
- Doesn’t only apply to solutions
- Must have a H atom that it can lose as an H+ ion
16.2
Define “Bronsted-Lowry base”.
- Accepts a proton from another substance.
- Doesn’t only apply to solutions.
- Must have a pair of nonbonding electrons that can be used to bind the H+ ion
16.2
In the equation HCl(g) + H2O(ll) —> H3O+ + Cl-, what is the Bronsted-Lowry acid and what is the base?
The base is H2O. The acid is HCl.
16.2
Define “amphiprotic”.
- Substance that can act as an acid or a base.
- Acts as a base when combined w/ something more acidic than itself, and it acts as an acid when something more basic than itself.
16.2
Give It Some Thought - p. 670
In the forward reaction, which substance acts as the Bronsted-Lowry base:
HSO4- + NH3 SO4 2- + NH4 +
NH3 is the Bronsted-Lowry base.
16.2
In the equation HNO2 + H2O NO2- + H3O+, what are the conjugates?
NO2- is the conjugate base of HNO2.
H3O+ is the conjugate acid of H2O.
16.2
Give It Some Thought - p. 672
Using the three categories (strong, weak, negligible), specify the strength of HNO3 and the strength of its conjugate base, NO3-.
Strong, negligible.
16.2
In the equation HCl + H2O —> H3O+ + Cl-, where does the equilibrium lie?
It lies to the right because H2O’s conjugate base, H3O+, is stronger than HCl’s conjugate base, Cl-, so Cl- can’t take back the protons to become HCl again.
16.2
In the equation CH3COOH + H2O H3O+ + CH3COO-, where does the equilibrium lie?
To the left because CH3COO-, CH3COOH’s conjugate base, is stronger than H3O+, H2O’s conjugate base, and will accept the H+ more easily.
16.3
What is the autoionization of H2O?
H2O + H2O H3O+ + OH-
- Extremely rapid in both directions
- Only ~2/10^9 molecules are ionized at any given instant… so pure H2O is a bad conductor of electricity
16.3
What is the H2O equilibrium expression/equation?
Kw = [H3O+][OH-] = [H+][OH-]= 1.0 x 10^-14
16.4
What is the formula to find pH?
-log[H+]
16.4
Why do pHs generally have 2 decimal places?
Only the numbers to the right of the decimal point are the sig. figs. in a logarithm.
Ex: -log (1.0 x 10^-7) = -(-7.00) = 7.00
16.4
What is the formula to find pOH?
-log[OH-]
16.4
How do pH & pOH correlate?
pH + pOH = 14.00
16.5
What are the 8 strong acids?
HCl, HBr, HI, H2SO4, HClO3, HClO4, HNO3, HIO4
16.5
What are the 7 strong bases?
Alkali metals + OH-, Ca(OH)2, Sr(OH)2, Ba(OH)2
16.5
How do scientists use ionic metal oxides to form OH-?
The net ionic equation for these metal oxides + H2O is O2- + H2O —> 2 OH-, making a strong base.
16.6
What is Ka, and what does it indicate?
Acid-dissociation constant. It shows the magnitude of an acid’s tendency to ionize in H2O. The larger the value of Ka, the stronger the acid.
16.6
What is percent ionization?
Another measure of acid strength.The stronger the acid, the greater is the percent ionization.
(Concentration ionized/original concentration) x 100%
or
([H+]equilibrium/[HA]initial) x 100%
16.6 - Give It Some Thought p. 685
When and why can we assume that the equilibrium concentration of a weak acid equals its initial concentration?
If Ka is x 10^-5 or less, this indicates that the equilibrium lies to the left, or to the reactants’ side, so any sort of change will be very small compared to the initial concentration.
If x is found to be >~5% of the intial concentration, however, the quadratic formula must be used.
16.6
When does the percent ionization of a weak acid decrease?
When ionization increases.
So, for instance, doubling the concentration of a weak acid will not double the concentration of H+.
16.6
What is a polyprotic acid?
It has more than one ionizable H atom, and thus two acid-dissociation constants. It is always easier to remove the 1st proton from a polyprotic acid than to remove the 2nd, due to electrostatic attractions.
H2SO3 HSO3- + H+ Ka1 = .017
HSO3 H+ + SO32- = 6.4 x 10^-8
16.6
Where does equilibrium lie for the dissociation of H2SO4?
H2SO4 —-> HSO4- + H+
To the right because it dissociates completely.
HSO4- SO4 2- + H+
Still to the right because Ka2 is .012, and still quite large.
16.6
How do you find the pH of polyprotic acids?
As long as Ka values differ by a factor of 10^3 or more, treat them as though they’re monoprotic. Most of the H+ will come from the 1st ionization reaction anyway.
16.7
What is the base-dissociation constant?
Always refers to the equilibrium in which a base reacts with H2O to form the corresponding conjugate acid and OH-.
Ex.: NH3 + H2O NH4+ + OH-
[NH4+][OH-]/[NH3]
16.7
What are the two general categories of weak bases?
- Neutral substances with an atom that has a nonbonding pair of electrons that can serve as a proton acceptor. Most contain N. These substances include NH3 & a related class of compounds called amines.
- Anions of weak acids. For example, ClO- from HClO.
16.7
What is an amine?
In organic amines, 1+ N-H bonds in NH3 is replaced with a bond between N and C. So, the replacement of one N-H bond in NH3 with an N-CH3 bond produces NH2CH3, or methylamine.
16.7
Give the chemical formula for the reaction of methylamine with water.
CH3NH2 + H2O CH3NH3+ + OH-
16.8
What’s the relationship between Ka and Kb?
(Ka)(Kb) = Kw
Applies only to conjugate acid-base pairs.
16.8
What are pKa and pKb?
-log Ka, -log Kb
pKa + pKb + pKw = 14.00
