Notes: Acids and Bases

Properties of Acids Properties of Bases Sour/Tart Bitter Electrolytes Feel Slippery Change indicator color Electrolytes

React with metals to produce H2 Change indicator color

Naming Acids

Anion Acid ending -ide hydro___ic acid -ate -ic acid -ite -ous acid

Naming Bases

Easy! - Metal + Hydroxide

Calculating pH

pH = - log[H+] pOH = - log[OH_] pH -pOH [H+] = 10- [OH_] = 10 [H+] x [OH_] = 1 x 10-14 M 2 pH + pOH = 14

significant figures in [concentration] = decimal places in pH

Classifying a solution as acid, base or neutral

Acids Bases Accept e_ (Lewis) Donate e_ (Lewis) Donate H+ (B-L) Accepts H+ (B-L)

Acid Neutral Base pH < 7 pH = 7 pH > 7 [H+] > 1 x 10-7 M [H+] = [OH_] = 1 x 10-7 M [H+] < 1 x 10-7 M [OH_] < 1 x 10-7 M [OH_] > 1 x 10-7 M

Water molecule Water molecule Hydronium ion Hydroxide ion Strong acids dissociate/ionize completely in aqueous solution.

MEMORIZE these strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4

+ - HNO3(aq)  H (aq) + NO3 (aq)

If you know the concentration of the acid, you know the concentration of the H+ because one mole of acid produces one mole of H+. Simply take the –log[H+] to find pH.

Weak acids only partially dissociate/ionize in aqueous solution. Equilibrium is established between the ions and the undissociated acid molecule. Any acid that is not on the list of strong acids is considered weak. Some weak acids (called organic or carboxylic acids) have chemical formulas ending in –COOH. The hydrogen at the end of the molecule is the one that breaks off into hydrogen ion in solution.

- HNO2(aq) H+(aq) + NO2 (aq)

HCOOH(aq) HCOO-(aq) + H+(aq)

We can compare the RELATIVE strengths of weak acids using a constant called the acid dissociation constant (Ka). The larger the value of Ka, the more the acid dissociates in solution, therefore the stronger the acid. We won’t be calculating the pH of weak acid solutions.

Strong bases also

2 dissociate/ionize completely in aqueous solution. Group IA and IIA metals with hydroxide ion are the strong bases.

NaOH(aq)  Na+(aq) + OH-(aq)

Be careful calculating the pH of bases in general and Group IIA bases in particular.

2+ - Ca(OH)2(aq)  Ca (aq) + 2OH (aq)

Notice that there are 2 moles of hydroxide ion for every one mole of calcium hydroxide. Therefore, if you are told that the calcium hydroxide concentration is 0.035 M, the [OH-] is actually 2 x .035 or .070 M.

To find the pH you would first find pOH = –log[.070] = 1.15 and then subtract from 14 to get pH = 14 – 1.15 = 12.85

Weak bases only partially dissociate/ionize in solution. Any metal other than group IA or IIA with hydroxide is considered a weak base. Also, there are some bases that don’t contain hydroxide in the chemical formula at all. These bases contain nitrogen. The nitrogen interacts with the water in the solution and strips off a hydrogen ion, creating excess hydroxide ion in the solution. This is an example of the Bronsted-Lowry definition of acids and bases.

ARRHENIUS DEFINITION acid--donates a proton (H+) in water base--donates an hydroxide ion in water (OH-)

BRONSTED-LOWRY DEFINITION acid--donates a proton in water base--accepts a proton in water

When using the Bronsted-Lowry definition we often recognize conjugate acid-base pairs in the equilibrium solution. Conjugate acid-base pairs are on opposite sides of the equilibrium arrow and are related to one another by the gain/loss of one hydrogen ion (proton). acids--donate H+ can be neutral compound, cation, or anion + HNO3 + H20  H3O + NO3- neutral compound. acid base CA CB 3 + + NH4 + H2O  H3O + NH3 cation acid base CA CB

+ -2 H2PO4- + H20  H3O + HPO4 anion acid base CA CB bases--accept H+ can be neutral compound, cation, or anion + NH3 + H2O  NH4 + OH- neutral compound base acid CA CB

-2 CO3 + H2O  HCO3- + OH- anion base acid CA CB

-3 -2 PO4 + H2O  HPO4 + OH- anion base acid CA CB

A reaction between an acid and a base is called neutralization. It is a special category of double replacement reaction and it always produces a salt (ionic compound) and water.

2H3PO4(aq) + 3Sr(OH)2(aq)  Sr3(PO4)2(aq) + 6H2O(l) ACID BASE SALT WATER

STRONG ACID + STRONG BASE = NEUTRAL SALT STRONG ACID + WEAK BASE = ACIDIC SALT WEAK ACID + STRONG BASE = BASIC SALT

The laboratory technique used to find the concentration of an unknown acid or base through a neutralization reaction is called titration. We learned the calculations when we were working on reaction stoichiometry.

In reaction stoichiometry, we use Molarity as one of the routes to get to moles. If you know volume and Molarity, you can always find moles. ALWAYS REMEMBER AND NEVER EVER FORGET…

. . . Volume (in liters) x Concentration (Molarity) = moles

Example What volume of a 0.100 M HCl solution is needed to neutralize 25.0 mL of 0.350 M NaOH?