Chemistry Lab/Acids and Bases

Acids and Bases is a topic for Chemistry Lab for the 2019 season.

Acids and Bases
Acids and Bases is basically an acid/base titration lab. Be sure you know what a titration is, because it is not a good thing if you do not. This is a fairly quick and simple lab to complete, and it is more than worthwhile to double check your lab if you have enough materials. More repetitions of the lab can result in a more accurate answer. In a free-response style lab report, this might also get you some extra points for style and accuracy. Acid/base questions can range in difficulty from identifying if a solution was an acid based on its pH to balancing advanced reactions trying to find the acidic constant. In order to excel in this event you must be prepared for all levels.

Solutions
For more info on solutions see Chem Lab/Aqueous Solutions

Definition of acids and bases
There are three major definitions of acids: the Arrhenius acid, the Bronsted-Lowry acid, and the Lewis acid. Each is defined differently, although the definitions are related.

Arrhenius acids/bases
Arrhenius acids are defined to be chemicals that, when put in water, produce hydronium ([math]H_3O^+[/math]) ions. Similarly, Arrhenius bases produce hydroxide ([math]OH^-[/math]) ions in aqueous solution.

Bronsted-Lowry acids/bases
Bronsted-Lowry acids are defined to be chemicals that donate protons ([math]H^+[/math]), while Bronsted bases accept protons. This is a broader definition than the Arrhenius definition because the acid-base reaction does not have to occur in water, although we know through observation that in solution, Bronsted acids donate protons to water molecules and form hydronium ions, as defined by the Arrhenius definition. Similarly, Bronsted bases in aqueous solution deprotonate water into hydroxide ions.

Lewis acids/bases
Lewis acids are defined to be chemicals that accept electron pairs, while Lewis bases donate electron pairs. Many acids and bases under this definition do not even have an acidic proton to be exchanged, although Arrhenius and Bronsted acids all fit the definition of a Lewis acid.

Acid/base equilibria
Chemical reactions only very rarely react to completion. Reactants react forward to form products, but products can also react backwards to form reactants. The forward and reverse reactions will set up a state known as equilibrium, where the forward reaction and the reverse reaction balance each other out. The equilibrium constant Keq describes the ratios of the concentrations of all chemical species at equilibrium. Take this reaction:

[math]aA + bB \rightleftharpoons cC + dD [/math]

The equilibrium constant Keq is equal to:

[math]K_{eq}=\frac{[C]^c[D]^d}{[A]^a[B]^b}[/math]

Where all of the concentrations are the concentrations at equilibrium and where solids are excluded.

For more info on equilibrium, see Chem Lab/Equilibrium.

Acid and Base Dissociation Constants
The reaction of an acid with water to form hydronium and a conjugate base also occurs in equilibrium. A special equilibrium constant, the acid dissociation constant Ka of any given acid, is used to record the equilibrium concentrations of the reactants and products of the dissociation of a proton from an acid in aqueous solution.

For the following reaction:

[math]HA + H_2O \rightleftharpoons H_3O^+ + A^-[/math]

The acid dissociation constant (Ka) is equal to

[math]\frac{[H_3O^+][A^-]}{[HA]}[/math]

Note that this is similar to the expression for the general equilibrium expression, although water, which is a reactant, is not included in this expression. That is because the concentration of water does not change in this reaction, since it is usually the solvent for the dissociation.

We can also define a base dissociation constant Kb. For the reaction

[math]B + H_2O \rightleftharpoons BH^+ + OH^-[/math]

The base dissociation constant (Kb) is equal to

[math]\frac{[BH^+][OH^-]}{[B]}[/math]

You will often see Ka and Kb expressed as pKa and pKb, where:

[math] pK_{a} = -log(K_{a}) [/math]

[math] pK_{b} = -log(K_{b}) [/math]

Comparing the Strength of Acids and Bases
Quantitatively, acids with greater Ka (lower pKa) are stronger acids than those with lower Ka (greater pKa). Similarly, bases with greater Kb (lower pKb) are stronger bases than those with lower Kb (greater pKb). One can also compare the strength of an acid or base with the strength of its conjugate base or acid. For example, a stronger base will have a weaker conjugate acid, i.e. its conjugate acid will have a greater Ka (lower pKa).

Qualitatively, stronger acids should have a more stable conjugate base. When comparing two acids, the one with the more stable conjugate base is the stronger acid. Apply the following steps to compare the stability of these conjugate bases:
 * 1) Which atom is the charge on? Deprotonating the two acids leaves each with a negative charge. Determine which atom is negatively charged in each conjugate base. If the two atoms are in the same period, the more electronegative atom is better able to stabilize the charge. If the two atoms are in the same group, however, the larger atom can better stabilize the charge. If the negative charge is on the same atom in both conjugate bases proceed to step 2.
 * 2) Resonance. If one of the conjugate bases has resonance structures then the negative charge will be resonance stabilized. If both conjugate bases have resonance structures proceed to step 3.
 * 3) Induction. If there are electronegative atoms that induct negative charge away from the negatively charged atom that conjugate base will be more stable.
 * 4) Orbitals. Electrons in orbitals held closer to the nucleus are more stable than those where they are held further away. For example, a negative charge on an sp-hybridized orbital is more stable than one in an sp2-hybridized orbital, which is more stable than one in an sp3-hybridized orbital.

pH and pOH
pH + pOH = 14

pH
pH is equal to the [math]-\log [H^+][/math] or [math]-\log [H_3O^+][/math].

pOH
pOH is equal to the [math]-\log[OH^-][/math].

Strong Acids
Strong acids are acids that almost completely disassociate in water. Some examples of strong acids are: [math]HI, HBr, HClO_4, HCl, HClO_3, H_2SO_4[/math], and [math]HNO_3 [/math]

Weak Acids
Weak acids are acids that only dissocite to a small extent in water. Their Ka value determines how much of the acid will dissociate in aqueous solution. Weak acids consist of all acidic species that are not strong acids. For example: [math]HCOOH, CH_3COOH, HOOCCHOHCHOHCOOH,[/math] and [math]{HSO_4}^-[/math]

Bases
All bases have a pH greater than 7.

Arrhenius Bases
Arrhenius Bases are defined to be chemicals that, when put in water, produce hydroxide ([math]OH^-[/math]) ions.

Bronsted-Lowry Bases
Bronsted Lowry Acids are defined to be chemical that accept protons ([math]H^+[/math]). This is a broader definition than the Arrhenius definition because it does not have to involve water.

Lewis Bases
Lewis Bases are defined to be chemicals that donate electron pairs.

Strong Bases
Strong Bases are bases that pretty much completely disassociate in water. Examples include [math]LiOH, NaOH, KOH, RbOH,[/math] and [math]CsOH[/math].

Weak Bases
Weak Bases are bases that only partially disassociate in water. They have a Kb to define how much. A common example of a weak base is [math]NH_3[/math].

Base Dissociation Constant
The base dissociation constant (Kb) is equal to

[math]\frac{[BH^+][OH^-]}{[B]}[/math]

for the following reaction:

[math]B + H_2O \to BH^+ + OH^-[/math]

Dissociation Constant of Water
The dissociation constant of water (Kw) is equal to

[math][H^+][OH^-] = 1*10^{-14}[/math]

for the following reaction:

[math]H_2O \to H^+ + OH^-[/math]

This is why pH + pOH = 14

Relationship between Ka and Kb
[math]\frac{[H^+][NH_3]}{[NH_4^+]} * \frac{[NH_4^+][OH^-]}{[NH_3]} = [H^+][OH^-][/math]

This method works for all acids and bases. Thus,

Ka [math]*[/math] Kb = Kw

Titrations
For more info on Titrations, see Chem Lab/Titration Race.

Links

 * Acid and base links