Acid Base Equilibrium Flashcards
For Exam 2
1
Q
What is the concept of Lewis acids and bases?
A
- It’s in terms of electrons. (Like Lewis structures!)
-
Acid = an electron acceptor
- Hint: Acid accepts electrons
- Includes metal cations with open d orbitals (like iron in hemoglobin)
- Base = an electron donor.
- Neutralization = formation of a coordinate covalent bond
2
Q
What is the Bronsted-Lowry concept of acids and bases?
A
- Acid-base reactions are about proton transfers
- Hint: Bro, pro
- Acid = proton donor
- Base = proton acceptor
3
Q
What is a conjugate acid-base pair?
A
- Two substances that differ only by ONE proton
- Ex: H3O+ and H2O
- The conjugate acid has one more H+, it’s the proton donor (Ex: H3O+)
- The conjugate base is the proton acceptor (Ex: H2O)
- The stronger a Brønsted-Lowry acid or base is, the weaker its conjugate acid or base. Vice versa.
- Every rxn between a Brønsted acid and a Brønsted base involves two conjugate acid–base pairs, typically on opposite sides of arrows
- Stronger acids and bases tend to react with each other to produce their weaker conjugates
4
Q
What is an amphoteric substance?
(aka amphiprotic)
A
- It can behave as either an acid or a base, depending on the other substance
5
Q
How are the strengths of acids and bases measured?
A
- The more complete its reaction with water, the stronger the acid or the base
- If Ka or Kb > 1, strong. If Ka or Kb < 1, weak.
- The larger the K, the stronger the species
- The smaller the pKa, the stronger the species
- The strongest base in water is OH-, the strongest acid in water is H3O+
- Pam’s fav weak acid: acetic acid. HC2H3O2
6
Q
Which side does an acid-base equilibrium favor,
with weak acid + weak base?
A
- The weaker side.
- Equilibrium lies on the side of the weaker acid and weaker base
- In this case, the left, so K < 1.
- (We don’t deal with equilibrium when there is a strong acid or base involved. Then it just goes that direction.)
7
Q
What is a coordinate covalent bond?
A
- When both electrons in the pair came from one species instead of one from each.
- This is associated with the Lewis concept of acids and bases. The base gives an electron pair, the acid accepts it.
- Neutralization = formation of a coordinate covalent bond
8
Q
What are binary acids and what is the periodic trend for their strength?
A
- Binary acids - binary compounds between hydrogen and nonmetals. (ex: HCl, H2S)
- Strength of the acid increases as move down same group, and right across the same period.
- Down a group: Think about bond length.
- HF < HCl < HBr < HI. HI is the strongest halide acid. It has the weakest bond because of bigger atomic radius. Since HF has the strongest bond (because F is the most electronegative so HF is the most polar) it’s the weakest acid.
- Across a period: Think about electronegativity. The more electronegative, the higher the partial positive charge on the H, the closer that it already is to being H+.
9
Q
What are oxoacids, and what is the trend on the periodic table for their strength?
A
- Acids composed of hydrogen, oxygen, and some other element
- They have O-H groups bonded to some central atom. Ex: H2SO4 sulfuric acid
- Stronger acid if there are more oxygens held by central atom
- Ex: nitric acid (HNO3) is a stronger acid than nitrous acid (HNO2)
- Stronger acid if the central atom is more electronegative
10
Q
What is the Arrhenius definition of acids and bases?
A
- Acid: dissociates to increase [H+]
- Base: dissociates to increase [OH-]
11
Q
What are the equations to find pH?
A
- pH= -log[H+]
- [H+] = 10-pH
- pOH= -log[OH-]
- [OH-] = 10-pOH
- pH + pOH = 14
- [H+][OH-] = 10-14
12
Q
What are Ka and Kb
and how can you solve for them?
A
- Types of equilibrium constants: ionization constants (or dissociation constants)
- Ka: Acid
- Kb: Base
- Weak acids and bases dissociate to an extent that is quantified by their equilibrium constant
- Ka and Kb allows us to compare the relative strengths of weak acids and bases.
- Larger Ka or Kb indicates a stronger acid or base, because more product favored. Smaller Ka or Kb indicates a weaker acid or base
- Kw = 10-14 = Ka * Kb
13
Q
What are pKa and pKb
and how can you solve for them?
A
- Just the -log
- pKa= -log(Ka)
- pKb= -log(Kb)
- The smaller the pK value, the stronger it is
14
Q
What are the relationships between Ka, Kb, and Kw
for conjugate pairs?
A
- Kw = 10-14at 25oC
- KaKb = KW
- pKa + pKb = 14
15
Q
Equation for concentrations of H+ and OH-?
(shortcut)
A
[H+] = √(Ka x M)
[OH-] = √(Kb x M)
- Can use these to calculate pH
16
Q
What is an amphiprotic substance,
and what are examples?
A
- It can act as an acid or a base, donate OR accept a proton.
- Ex: Water.
- Polyprotics always have an amphiprotic form. The acid form is the fully protonated form. The base form is the fully deprotonated form. The amphiprotic form is in between, a conjugate of the acid form and the base form.
- Ex: Hydrogen phosphate HPO42- and dihydrogen phosphate H2PO4- are amphiprotic. (H3PO4 is phosphoric acid, weak but only an acid.)
- Ex: bicarbonate HCO3-, hydrogen sulfide HS-, hydrogen oxalate HC2O4-
17
Q
With polyprotic acids, which form is stronger?
A
- The fully protonated form is the strongest
- Each successive dissociation produces a weaker acid
- Ex: phosphoric acid is triprotic, and will have a larger K value and smaller pKa value than its children. H3PO4 > H2PO4- > HPO42-
18
Q
What is the labile hydrogen?
A
- The hydrogen that is easily removable
19
Q
What is a carboxylic acid, and what is the trend for strength?
A
- A carboxylic acid is an organic compound. The general formula is R–COOH, with R referring to the rest of the molecule.
- If there is an electronegative element on the carbon (like chlorine or fluorine) then they hog the electrons, leaving just a partial negative charge on the other side, and it’s easier for the H to be removed.
- So trichloroacetic acid and trifluoroacetic acid are stronger than acetic acid, which just has hydrogens around the carbon.
20
Q
How does the ammonium ion behave in water?
NH4+
A
- Ammonium ion is acidic in water
- So NH4Cl would also be weakly acidic. But NH4OH would be basic because hydroxide is stronger.
21
Q
How does bicarbonate behave in water?
HCO3-
A
- Amphiprotic, but it acts as a weak base in water, receives a hydrogen
- But if you were to put it in a solution with a stronger base, it would be forced to work as an acid.
- Memorize this, but can also see that tabulated Kb value > Ka value.
22
Q
What are some ions that make a neutral solution with water?
A
- Don’t need to memorize. But getting familiar…
- Note that these are the beginnings of the strong acids and some strong bases. When an acid or a base is extremely strong, the conjugate is too weak to affect the pH.
- Metal ions with small charges are nonacids
- Application: CaCl2 and KNO3 are neutral solutions in water because neither the cation nor anion is able to affect the pH
23
Q
What is an ion that is acidic in water?
A
- Only one we have to memorize is NH4+
- Hydrated Al3+ , Cr3+, Fe3+, because metal ions with high charge densities are weak acids
- The only time an anion might be acidic is if it is from a partially neutralized polyprotic acid, and has another hydrogen to donate.
24
Q
What is an ion that is basic in water?
A
- HCO3- (bicarbonate) acts as a weak base in water
- Don’t need to memorize, but here are some more ions that are basic in water
25
What is the equivalence point of a neutralization reaction?
* Moles of acid = moles of base. Both the acid and the base in the reaction have been completely consumed and neither of them are in excess
* This doesn't mean the solution is neutral
* strong acid + weak base: pH \< 7
* strong base + weak acid: pH \> 7
* strong acid + strong base: pH = 7
* weak acid: weak base:
pH \<7 if Ka \> Kb
pH \>7 if Ka \< Kb
* pH =7 if Ka = Kb
* On graph: vertical region near EP. EP is at the center of t.
26
What is a buffer,
and how can you increase it's capacity?
* Buffers are used to prevent large changes to pH
* Capacity of a buffer = the amount of strong acid or base it can absorb before changing pH
* Adding enough strong acid (or base) to consume the weak base (or acid) will exceed the **buffer capacity** of the system and a large pH change will result.
* To increase buffer capacity, keep ratio of base/acid constant but increase buffer concentration. Add more of both weak acid and conjugate base.
* Usually made of a weak Bronsted acid and a weak Bronsted base
* Usually a conjugate pair, or the molecular acid with a soluble salt of the acid
27
How do you choose the ideal buffer?
* The pKa of the weak acid of the buffer should be within ±1 units of the pH you’re trying to maintain
* Ex: if you want a pH= 4.3, you should select a buffer system whose weak acid pKa is between 3.3-5.3. (So acetic acid 4.74 would work!)
28
What is the equivalence point?
* The equivalence point is the point at which all of the initial acid has been converted to conjugate base
* It's the point at which equal moles of acid and base have been added
* MaVa = MbVb
* Looks vertical on graph
* For titration of a weak acid with strong base, equivalence point will be at ph\>7. Vice versa. If both are strong, equivalence point around ph=7.
29
What is the half-equivalence point?
* The half-equivalence point is the point at which pH = pKa
* It's where you have equal concentrations of your acid and its **conjugate base** (half of the titrant needed)
* Appears as plateau that we refer to as buffering regions
30
Interpret a titration curve for a monoprotic
31
How would you draw (or interpret) a titration curve for a weak polyprotic acid?
* The number of half-equivalence and equivalence points in your titration will be equal to the number of protons on your acid
* The first half equivalence point will have a pH that’s equal to pka1 of the acid. The second half equivalence point will be pka2, associated with second protonation of this acid. The equivalence points are the average of what they're in between.
* Notice that Kas get successively smaller (pKa larger) with deprotonations. Held less tightly.
* Ex: Titration curve for carbonic acid H2CO3
* A) Half-Equivalence Point 1. [H2CO3] =[HCO3-]. pH= pKa1
* B) Equivalence Point 1. All of the H2CO3 has been converted to HCO3-
* C) Half-Equivalence Point 2. [HCO3-]=[CO32-]. pH=pKa2
* D) Equivalence Point 2. All of the HCO32- has been converted to CO32-
* Flip for a weak base because starts with high pH
32
What is the Henderson Hasselbalch equation, and when do you use it?
* Use this equation with buffers to calculate the pH during titration (before the last equivalence point is reached)
33
**How do you determine pH:**
1) Before titration
2) During titration
3) At equivalence point
4) After equivalence point
1. Before titration: use the shortcut.
1. [H+] = √(Ka x [M])
2. During titration: Use HH equation, adjusting molarity
1. pH = pKa + log([conj. base]/[conj.acid])
3. At last EP: use the shortcut, adjust molarity
1. [OH-] = √(Kb x [M]) (slightly basic after titration with strong base. slightly acidic after titration with strong acid.)
4. After last EP: take leftover moles. Calculate molarity and find pH from that.
34
What can you make a buffer from?
* Weak acid and weak base (or salt of conjugate base)
* Weak acid and strong base
* Strong acid and weak base
* Depends on how much you add.
* In order to be an effective buffer, the number of moles of the weak acid and its conjugate base must be large compared to the number of moles of strong acid or base that may be added.
* The best buffering will occur when the ratio of [HA] to [A-] is about 1:1. Buffers are considered to be effective when the ratio between the conjugates is anywhere between 10:1 and 1:10.
35
What is the autoionization of water?
* Water can act as acid or base
* Kw = 10-14
