Hello hello! We’re coming into the last stretch of the semester, with finals week starting this very Friday for us (yoohoo, hi Wellesley students!) So, that leaves me with one more unit of chemistry that I want to cover, which is…. acids and bases!
Go ahead and check out last time’s crash course on Electrochemistry if you haven’t already! 🙂
Throughout history, we have come to give better and better definitions of things as we understand them better. Acid base theory are one of the things that have undergone this kind of progress.
- In the beginning, we described acids as substances with sour, corrosive properties, and bases as kind of soapy. That’s it.
- Arrhenius was the first one to put a real definition on it. An acid is a substance that releases H+ and a base is a substance that releases OH-.
- The Bronsted Lowry acid base definition states that a base accepts a proton from an acid.
Therefore, an acid is a proton donor and a base is a proton receiver.
All Arrhenius acids and bases are also Bronsted Lowry acids and bases, respectively.
- Our last definition, by G.N. Lewis, states that an acid is a substance that accepts an electron pair to form a covalent bond from a base.
Therefore, an acid is an electron pair receiver and a base is an electron pair donor.
All Bronsted Lowry acids and bases are also Lewis acids and bases, respectively.
Now, if you remember from when we discussed equilibrium, we used a term called the equilibrium constant. We can use K to describe the dissociation of an acid or base.
Ka and Kb provide a quantitative measure of the strengths of acids and bases.
Furthermore, water undergoes autoprotolysis/autoionization, and has its own ion-product constant. It is important to specify the temperature because the constant increases with temperature (it is an endothermic process).
When dealing with acid-base equilibria, because the product of the concentrations of H+and OH- ions in any aqueous solution is always equal to the ion product constant (Kw), it is important to note that the addition of H+ (OH-) to an aqueous solution causes a concomitant decrease in OH- (H+).
- Strong [weak] acid = acid with relatively large [small] Ka (most common acids are weak).
- Strong [weak] base = base with relatively large [small] Kb (most common bases are strong).
Now, let me say something really sweet about the acid/base partnership: an acid is nothing without a base present to act upon, and a base is nothing without an acid present to act upon!
- If you remove a proton from an acid, you create a species that can gain a proton. This is known as the conjugate base of the acid.
- If you add a proton to a base, you create a species that can lose a proton. This is known as the conjugate acid of the base.
An acid-base reaction will contain two conjugate acid/base pairs.
The Lewis theory is more general.
Any species that has a hydrogen atom can potentially act as an acid, and any compound that has a lone pair of electrons can potentially act as a base.
Lewis bases form coordinate covalent bonds (also known as dative bonds) to acids by supplying both electrons for the bond.
The pH and pOH Scale
The pH is defined as the negative of the base-10 logarithm of the H+ ion concentration (expressed in moles per liter).
pH = -log [H+]
Similarly, the pOH is the negative of the base-10 logarithm of the OH- ion concentration.
pOH = -log [OH-]
And the equilibrium constants are expressed in terms of logarithms too.
pK = -log K
Based on the autoprotolysis of water,
-log Kw = -log [H+] – log [OH-]
pKw = pH + pOH = 14
So, at a temperature of 25 C, in a enutral solution,
pH = pOH =7
A larger value of p(“anything”) means a smaller value of the “anything”, and by orders of magnitude.
Okay, so I said earlier that the acid dissociation constant is a quantitative measure of the strength of an acid or base. What does this mean?
A strong acid has a very large acid dissociation constant Ka (so, a very small pKa), which means that the equilirium lies very far to the right (Ka>>1).
We can say that a strong acid dissociates completely in an aqueous solution (fully deprotonated).
A weak acid has a very small Ka (Ka<1) and a very large pKa, so the equilibrium lies very far to the left.
In the same way, the weaker the base, the lower the value of Kb, and the higher the value of pKb.
The weaker the acid, the stronger the conjugate base.
For any conjugate acid-base pair,
Kw = Ka × Kb
pKw = pKa + pKb
What Causes an Acid or Base to be Strong or Weak?
ACIDS. The difference between a strong acid and a weak acid is determined by the stability of the conjugate base: a strong acid has a stable (low Gibbs free energy of formation) conjugate base and a very weak acid has an unstable (high Gibbs free energy of formation) conjugate base.
- Electronegativity. In the same row of the periodic table, acid strengths of binary acids correlate positively with EN.
HF > H2O >> NH3 > CH4
F- > OH- > NH2- > CH3-
HF is the strongest because it has the most polar bond and the most stable conjugate base. F is more electronegative than oxygen, and is more capable of stabilizing the negative charge on the conjugate base. O is more electronegative than nitrogen, so an alcohol is more acidic than an amine.
- Bond Dissociation Energy. In the same column of the periodic table, acid strengths of binary acids show a negative correlation with bond strength.
The size of the anion increases down a group, making it easier to break the H-X bond, so the acid strength increases.
HF < HCl < HBr < HI
Cl- is more stable than F- because its negative charge is spread over a larger volume of space.
- When a highly electronegative atom has a negative charge, the conjugate base is stabilized.
- Resonance Effects. May make one conjugate base more stable than the others.
- Inductive Effects. As the number of oxygen atoms available to share the negative force increases, the stability of the conjugate base increases (stronger acid).
- Hybridization. A negative charge on an sp hybridized carbon is more stable than a negative charge on an sp3 or sp2 hybridized carbon.
BASES. Small, highly charged negative species such as H- are strong bases. So O2- is a stronger base than OH- because it has a higher negative charge. And since ionic metal oxides contain O2-, a basic solution is produced when a metal oxide is dissolved in water.
Alkali metals and hte heavier alkaline earth metals (Ca, Sr, Ba) produce strong bases too.
Two categories of weak bases:
- A neutral substance with a donatable lone pair of electrons, such as ammonia and amines.
- Anions of weak acids, such as CH3COO- and ClO-.
Since bases are electron donors, electron-donating groups strengthen a base while electron-withdrawing groups weaken a base.
Halogens are electron withdrawing and reduce basicity (for example, PF3 is much weaker than PH3).
Alkyl groups are electron donating and increase basicity.
So, to summarize from above:
Notes for Calculations
When you do the calculations for strong acids and bases at very low concentrations, you have to remember that you have H+ and OH- from water on top of the H+ or OH- from the strong acid or base which has completely dissociated.
When we have a “normal” concentration of acid or base, we can ignore the autoprotolysis of water, typically.
With weak monoprotic acids at normal concentrations, use an ICE table, since we can’t assume that all the acid dissociates. To simplify the calculations, you can make assumptions as long as it follows the 5% rule.
(I’m not going to go into the math here but if you would like to see a post with some examples on how to use the ICE table please let me know!)
With weak bases at normal concentrations, we can do calculations with exactly the same reasoning.
Weak Polyprotic Acids
A polyprotic acid dissociates one proton at a time in a stepwise manner.
Kind of like having more than one weak monoprotic acid in solution, except all the acid species originate from the same acid molecule.
The ionization constants will be given in a table. Ka for the first dissociation is ALWAYS significantly greater than the second (and third, etc.) dissociations because the susequent protons are lost from a negative species.
Salts Acting as Acids or Bases
Most salts are strong electrolytes and dissociate completely into ions in water. (Salts are ionic compounds that result from reactions between acids and bases.)
Salts of a strong acid and a strong base produce neutral aqueous solutions because these types of salts don’t react with water.
However, in salts such as NaCN and NH4Cl, the anion, cation, or sometimes both, behave in water as a base or an acid. The base, CN-, can accept a proton from water (the ion is a conjugate base) and the acid, NH4+, can donate a proton to water (the cation is a conjugate acid.
CN- + H2O ↔ HCN + OH-
NH4+ + H2O ↔ H3O+ + NH3
Since water is split apart, we often call this behavior of the anions and cations hydrolysis.
Most anions are basic because the negative charge interacts with the polar water molecule to cause an H+ to be transferred to the anion.
Some anions, such as HCO3- (bicarbonate) can act as either acid or base, but since the base constant is larger than the acid constant, NaHCO3 is basic.
The salt of a weak acid and a strong base is a weak base, and the salt of a weak base and a strong acid is a weak acid.
The salt of a weak acid and a weak base can be acidic (if Ka>Kb), basic (if Kb>Ka), or neutral (if Kb∼Ka).
Solutions of weak acids or weak bases that contain a common ion are called buffer solutions.
The mixture of a weak acid and its conjugate base has a different pH than the acid alone, which is called the common ion effect.
These solutions resist pH change when diluted or when small amounts of strong acid or strong base are added to them.
We can use Le Chatelier’s principle to understand the effect of the common ion on the equilibrium.
HA(aq) + H2O(l) ↔ H+(aq) + A-(aq)
If we add A- (the conjugate base of HA), we suppress the dissociation of the acid (equilibrium position shifts to the left).
A-(aq) + H2O ↔ HA(aq) + OH-(aq)
A- will react with water, but since through the first equation we’ve created more HA in the solution, we suppress this reaction too (so both reactions are suppressed) and there is virtually no pH change.
The Henderson-Hasselbalch Equation
pH = pKa + log( [A-]’ / [HA]’ )
pOH = pKb + log( [cation]’ / [base]’ )
Without an ICE table, we can determine the pH of any buffer solution if we know the pKa of the acid and the starting concentrations of both acid and conjugate base species.
This is the amount of H+ or OH- that a buffer can absorb without a significant change of pH.
While the buffer pH is determined by the ratio of the acid/conjugate base (or base/conjugate acid), the buffer capacity is determined by the amount of acid/conjugate base (or base/conjugate acid) used to make the buffer.
Obviously, the buffer capacity increases with the amount used to make the buffer!
Designing Buffer Solutions
If we look at the Henderson-Hasselbalch equation, the buffering action is most effective when the [HA] is approximately equal to [A-].
pH ≈ pKa + log1
pH ≈ pKa
So then the solution will have equal capacity to resist pH change upon the addition of acid or base.
This means we should choose an acid whose pKa is as close as possible to the desired pH, and then the concentration of the acid and the conjugate base must be chosen to yield the exact pH.
Titrations and pH Curves
A titration is a quantitative chemical analysis method used to determine the concentration of a solution.
A base or acid of known concentration (the titrant) from a burette is added to a carefully determined volume of a solution of acid or base of unknown concentration (the analyte) until the [H+] undergoes a very sharp change that can be detected using a pH meter or an indicator (changes color).
The plot of pH versus the volume of titrant added is known as a titration curve.
The equivalence point (aka the stoichiometric point) is the point atwhich the number of moles of H+ (or OH-) added as a titrant is equal to the number of moles of OH- (or H+) initially present in the analyte.
The endpoint is the color at which an indicator changes color. The indicator must be chosen so that it changes color as close as possible to the equivalence point.
In the case of a strong acid-strong base titration (in the example above), the pH at the equivalence point is 7.0.
When a weak base (analyte) is titrated with a strong acid, there will be four stages:
- No acid added yet: we have a solution of a weak base.
- Not enough acid added to neutralize base: we’re in the buffer region.
- Exactly enough acid added to neutralize base: we’re in the equivalence point, hydrolysis of a salt.
- Excess acid added beyond equivalence point: we have a strong acid.
In a titration of a weak diprotic acid and a strong base: Two moles of base are needed to neutralize one mole of diprotic acid, and the titration curve contains two equivalence points.
An acid-base indicator (typically a weak acid) changes color within a small pH interval because the acid and the conjugate base have different colors.
Some common examples of indicators are: bromophenol red, thymolphthalein, phenolphthalein, and bromocresol green.
HIn(aq) + H2O(l) ↔ H3O+ + In-(aq)
Ka = [H3O+][In-] / [HIn]
At its turning point, [HIn] = [In-], so K(In) = [H3O+].
Therefore, the pH of the solution at its turning point is the pK(In).
Just for fun, I’ve included a chart showing all the pretty colors that indicators turn when they reach their pK(In).
Thanks for reading! Leave your questions and love in the comments below 🙂
Good luck on finals! And meanwhile, I’m all about that base, ‘Bout that base, No treble.