Chemistry Lab/Acids and Bases

Acids and Bases was a topic for the event Chemistry Lab in 2009.

Acids and Bases (2009)
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

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].

Acids
All acids have a pH less than 7

Arrhenius Acids
Arrhenius Acids are defined to be chemicals that, when put in water, produce hydronium ([math]H_3O^+[/math]) ions.

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

Lewis Acids
Lewis Acids are defined to be chemicals that accept electron pairs.

Strong Acids
Strong Acids are acids that pretty much 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 partially disassociate in water. They have a Ka to define how much. Weak Acids consist of pretty much everything that is not a strong acid. 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].

Equilibrium Constants
Take this reaction:

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

The equilibrium constant is equal to:

[math]k=\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 Dissociation Constant
The acid equilibrium constant (Ka) is equal to

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

for the following reaction:

[math]HA \to H^+ + A^-[/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