Study Guide: Ionic and Covalent Compounds, Hydrocarbons, and Functional Groups

Ions and Ionic Compounds

Formation and Properties of Ions

Ions are atoms or groups of atoms that have gained or lost electrons, resulting in a net charge. Ionic compounds are formed from the electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions).

• Cations are typically formed by metals losing electrons.

• Anions are typically formed by nonmetals gaining electrons.

• Common monatomic ions include Na+, Cl-, Mg2+, and O2-.

• Polyatomic ions are groups of atoms with a net charge, such as SO42- (sulfate) and NH4+ (ammonium).

Example: Sodium chloride (NaCl) forms when Na loses one electron to become Na+ and Cl gains one electron to become Cl-.

Naming and Writing Formulas for Ionic Compounds

• Write the cation first, followed by the anion.

• Balance charges to ensure the compound is neutral.

• Use Roman numerals for transition metals with variable charges (e.g., FeCl2 is iron(II) chloride).

• For polyatomic ions, use parentheses if more than one is needed (e.g., Ca(NO3)2).

Example: The formula for aluminum sulfate is Al2(SO4)3.

Covalent Compounds

Formation and Properties

Covalent compounds are formed when two or more nonmetals share electrons. The shared electrons allow each atom to achieve a stable electron configuration.

• Single, double, or triple bonds may form depending on the number of shared electron pairs.

• Common examples include H2O, CO2, and CH4.

Naming Covalent Compounds

• Use prefixes to indicate the number of each atom (mono-, di-, tri, tetra, penta-, hexa-, hepta-, octa-, nona-, deca).

• The first element keeps its name; the second element ends in "-ide."

• Omit "mono-" for the first element.

Example: CO2 is carbon dioxide; N2O4 is dinitrogen tetroxide.

Empirical and Molecular Formulas

Definitions and Calculations

The empirical formula is the simplest whole-number ratio of atoms in a compound. The molecular formula shows the actual number of each atom in a molecule.

• To determine the empirical formula, convert mass percentages to moles and simplify the ratio.

• The molecular formula is a whole-number multiple of the empirical formula.

Example: A compound with 40% C, 6.7% H, and 53.3% O has an empirical formula of CH2O.

Hydrocarbons and Branched Hydrocarbons

Types and Nomenclature

Hydrocarbons are organic compounds consisting only of carbon and hydrogen. They are classified based on the types of bonds between carbon atoms.

• Alkanes: Only single bonds (C–C). General formula: CnH2n+2.

• Alkenes: At least one double bond (C=C). General formula: CnH2n.

• Alkynes: At least one triple bond (C≡C). General formula: CnH2n-2.

Branched hydrocarbons have alkyl groups (side chains) attached to the main chain. Common alkyl groups include methyl (–CH3), ethyl (–C2H5), propyl (–C3H7), and butyl (–C4H9).

• Number the main chain to give the lowest possible numbers to substituents.

• Use prefixes (di-, tri-, tetra-) for multiple identical substituents.

Example: 2-methylpropane (isobutane) has a methyl group on the second carbon of propane.

Functional Groups

Common Functional Groups in Organic Chemistry

Functional groups are specific groups of atoms within molecules that determine the chemical properties of those molecules.

• Halogen: –F, –Cl, –Br, –I

• Hydroxyl: –OH (alcohols)

• Carbonyl: C=O (aldehydes and ketones)

• Carboxylic acid: –COOH

• Ester: –COOR

• Aldehyde: –CHO

• Ketone: RCOR'

• Amine: –NH2

• Amide: –CONH2

Example: Ethanol (CH3CH2OH) contains a hydroxyl group.

Practice Problems Overview

Types of Questions

• Writing Lewis structures for ionic and covalent compounds.

• Identifying compounds as ionic or covalent.

• Naming compounds from formulas and writing formulas from names.

• Calculating empirical and molecular formulas from percent composition.

• Identifying functional groups in organic molecules.

• Naming hydrocarbons and recognizing structural formulas.

Key Equations and Concepts

• Percent Composition:

$\text{Percent Composition} = \frac{\text{Mass of Element}}{\text{Total Mass of Compound}} \times 100$

• Empirical Formula Calculation:

1. Convert mass percentages to grams (assume 100 g sample). 2. Convert grams to moles for each element. 3. Divide by the smallest number of moles to get the simplest ratio.

• Molecular Formula:

$\text{Molecular Formula} = (\text{Empirical Formula})_n$

where $n = \frac{\text{Molar Mass}}{\text{Empirical Formula Mass}}$

Sample Table: Prefixes for Covalent Compounds

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-

Additional info:

• Practice problems in the file cover naming, formula writing, Lewis structures, empirical/molecular formula calculations, and functional group identification, which are all core skills in General Chemistry.

• Some images depict organic structures for functional group identification and hydrocarbon naming.