Kaplan MCAT General Chemistry Review

Ionic compounds are made up of positively charged cations and negatively charged anions. Ionic compounds are held together by the ionic bond, which is the force of electrostatic attraction between oppositely charged particles. The word cation (and especially the organic chemistry variant, carbocation) has tripped up unsuspecting students for years, who, having never heard the word spoken aloud before, almost invariably pronounce it as the last two syllables of “vacation” (as if a carbocation were some fun carbon-based holiday…we humbly admit to having made the mistake ourselves at a point in the less glorious distant past.)

The nomenclature of ionic compounds is based on the names of the component ions.

1. For elements (usually metals) that can form more than one positive ion, the charge is indicated by a Roman numeral in parentheses following the name of the element.

Fe²⁺ Iron (II)Cu⁺ Copper (I)Fe³⁺ Iron (III)Cu²⁺ Copper (II)

2. An older but still commonly used method is to add the endings -ous or -ic to the root of the Latin name of the element to represent the ions with lesser or greater charge, respectively.

Fe²⁺ FerrousCu⁺ CuprousFe³⁺ FerricCu²⁺ Cupric

3. Monatomic anions are named by dropping the ending of the name of the element and adding -ide.

H^– HydrideS^2– SulfideF^– FluorideN^3– NitrideO^2– OxideP^3– Phosphide

4. Many polyatomic anions contain oxygen and are therefore called oxyanions. When an element forms two oxyanions, the name of the one with less oxygen ends in -ite and the one with more oxygen ends in -ate.

NO₂^– NitriteSO₃^2– SulfiteNO₃^– NitrateSO₄^2– Sulfate

5. When the series of oxyanions contains four oxyanions, prefixes are also used. Hypo- and per- are used to indicate less oxygen and more oxygen, respectively.

ClO^– Hypochlorite

ClO₂^– Chlorite

ClO₃^– Chlorate

ClO₄^– Perchlorate

6. Polyatomic anions often gain one or more H⁺ ions to form anions of lower charge. The resulting ions are named by adding the word hydrogen or dihydrogen to the front of the anion’s name. An older method uses the prefix bi- to indicate the addition of a single hydrogen ion.

HCO₃^– Hydrogen carbonate or bicarbonate

HSO₄^– Hydrogen sulfate or bisulfate

H₂PO₄^– Dihydrogen phosphate

ION CHARGES

Ionic species, by definition, have charge. Cations have positive charge, and anions have negative charge. Some elements are found naturally only in their charged forms, while others may exist naturally in the charged or uncharged state. Furthermore, some elements can have several different charges or oxidation states. Some of the charged atoms or molecules that you might commonly encounter on the MCAT include the active metals, the alkali metals (group IA) and the alkaline earth metals (group IIA), which have charge of +1 and +2, respectively, in the natural state. Many of the transition metals, such as copper, iron, and chromium, can exist in different positively charged states. Nonmetals, which are found on the right side of the periodic table, generally form anions. For example, all the halogens (Group VIIA) form monatomic anions with a charge of -1. All elements in a given group tend to form monatomic ions with the same charge (e.g., Group IA elements have charge of +1). Note that there are anionic species that contain metallic elements (e.g., MnO₄^- [permanganate] and CrO₄^2- [chromate]); even so, the metals have positively charged oxidation states. (Also note that in the oxyanions of the halogens, such as ClO^- and ClO₂^-, the halogen is assigned a positive oxidation state.) The trends of ionicity as we’ve described them here are helpful but are complicated by the fact that many elements have intermediate electronegativity and are consequently less likely to form ionic compounds and by the left-to-right transition from metallic to nonmetallic character.

Key Concept

Oxyanions of transition metals like the MnO₄^– and CrO₄^2– ions shown here have an inordinately high oxidation number on the metal. As such, they tend to gain electrons in order to reduce this oxidation number and thus make good oxidizing agents. (See Chapter 11.)

ELECTROLYTES