Acids & Bases

Properties ACIDS: Sour taste Conduct electricity React with active metals to form H gas React with carbonate compounds to form CO2 gas React with bases to neutralize forming salt and water

BASES: Bitter taste Conduct electricity Feel slippery Do not react with active metals Do not react with carbonate compounds React with acids to neutralize forming salt and water

Examples ACIDS: citric acid (lemon juice) ascorbic acid (vitamin C) lactic acid (in fatigued muscles, yogurt) acetic acid (vinegar) 3-methyl-2-hexanoic acid (underarm odor)

BASES: Sodium hydroxide (oven cleaner) Sodium hydrogen carbonate (baking soda) Sodium carbonate (washing soda) Ammonia (glass cleaner)

Arrhenius Acids & Bases According to the Arrhenius definition, an acid is any substance, which when dissolved in water, tends to increase the amount of H +. An example is HCl: HCl (g)  H + (aq) + Cl - (aq)

An Arrhenius base is any substance, which when dissolved in water, tends to increase the amount of OH -. An example is NaOH: NaOH (s)  Na + (aq) + OH - (aq)

These definitions are correct but not general enough to include the wide range of acid and base substances which are known to exist. In addition, they rely on the use of water as a solvent, which is also too narrow.

Bronsted-Lowry Acids and Bases The Bronsted-Lowry definition is named for Johannes Bronsted and Thomas Lowry, who independently proposed it in 1923.

A Bronsted-Lowry (BL) acid is defined as any substance that can donate a hydrogen ion (proton). A Bronsted-Lowry (BL) base is any substance that can accept a hydrogen ion (proton).

Thus, according to the BL definition, acids and bases must come in what is called conjugate pairs. For example, consider acetic acid dissolved in water: CH 3 COOH (aq) + H 2 O (l)  H 3 O + (aq) + CH 3 COO - (aq) Acid (H donor) Base (H acceptor) Conjugate acid Conjugate base CONJUGATE PAIR

What Makes a Strong Acid or Strong Base? Strong electrolytes are completely dissociated into ions in water. The acid or base molecule does not exist in aqueous solution, only ions. Weak electrolytes are incompletely dissociated.

Strong Acids Strong acids completely dissociate in water, forming H + and an anion. There are six strong acids. The others are considered to be weak acids. HCl - hydrochloric acid HNO 3 - nitric acid H 2 SO 4 - sulfuric acid HBr - hydrobromic acid HI - hydroiodic acid HClO 4 - perchloric acid

100% dissociation isn't true as solutions become more concentrated. If the acid is 100% dissociated in solutions of 1.0 M or less, it is called strong. Sulfuric acid is considered strong only in its first dissociation step. H 2 SO 4  H + + HSO 4 -

Weak Acids A weak acid only partially dissociates in water to give H + and the anion. Examples of weak acids include hydrofluoric acid, HF, and acetic acid, CH 3 COOH.

Strong Bases Strong bases dissociate 100% into the cation and OH- (hydroxide ion). The hydroxides of the Group I and Group II metals usually are considered to be strong bases. LiOH - lithium hydroxide NaOH - sodium hydroxide KOH - potassium hydroxide RbOH - rubidium hydroxide CsOH - cesium hydroxide

*Ca(OH) 2 - calcium hydroxide *Sr(OH) 2 - strontium hydroxide *Ba(OH) 2 - barium hydroxide * These bases completely dissociate in solutions of 0.01 M or less. The other bases make solutions of 1.0 M and are 100% dissociated at that concentration. There are other strong bases than those listed, but they are not often encountered.

Weak Bases Examples of weak bases include ammonia, NH 3, and diethylamine, (CH 3 CH 2 ) 2 NH. Most weak bases are anions of weak acids. Weak bases do not furnish OH - ions by dissociation. Instead, they react with water to generate OH - ions.

To determine whether a substance acts a an acid or a base, consult the acid-base list.

Substances that are amphiprotic are those having characteristics of both an acid and a base and capable of reacting as either (AKA amphoteric)amphoteric

Lewis Base Gilbert Lewis broadened the definition of acid and base. A Lewis base possesses a pair of electrons that H+ can bind to.

Here we see ammonia (NH 3 ), methoxide (CH 3 O-), fluoride ion (F-), and hydroxide (OH-). They all have a pair of unshared electrons available for H+ to bind to. These unshared pair of electrons are often referred to as a lone pair.

Lewis Acid: On the top, the acid (H+) is being attracted to a Lewis base (a substance with an unpaired electrons (a lone pair) on nitrogen in ammonia. On the bottom, the compound, boron trifluoride, is also attracted to the same lone pair of electrons. So it behaves like the H+ acid and is called a Lewis Acid. Boron has 3 outer electrons that it shares with 3 fluorines.

If boron shares the lone pair electrons of nitrogen, then boron will have a stable octet (eight) outer electrons. That's why boron trifluoride binds with ammonia (bottom right). All atoms have their outer shell full of electrons (hydrogen with two; nitrogen, fluorine, and boron with eight).