GOB Chemistry Study Guide: Ionic Compounds, Covalent Compounds, Moles, and Molarity

Ionic and Covalent Bonding

Types of Chemical Bonds

Chemical bonds are formed by the interaction of valence electrons between atoms. The two main types are ionic bonds and covalent bonds.

• Ionic Bond: Formed between a metal and a nonmetal. Electrons are transferred from the metal to the nonmetal, resulting in ions of opposite charge that attract each other. Example: NaCl, MgO.

• Covalent Bond: Formed between two nonmetals or a nonmetal and a metalloid. Electrons are shared as pairs between atoms, forming molecules. Example: CO2, NH3.

Valence electrons drive all bonding, aiming for a stable electron configuration (octet).

Comparison Table: Ionic vs Covalent Bonds

Bond Type	Elements Involved	Electron Behavior	What Forms	Examples
Ionic	Metal + Nonmetal	Transferred	Ions	NaCl, CaF2
Covalent	Nonmetal + Nonmetal	Shared	Molecules	CO2, NH3

Ions and Their Charges

Formation of Ions

Atoms gain or lose valence electrons to reach the electron count of the nearest noble gas. Metals lose electrons to become cations (+), while nonmetals gain electrons to become anions (−).

• Group 1A: Lose 1 electron → +1 charge (e.g., Na+)

• Group 2A: Lose 2 electrons → +2 charge (e.g., Ca2+)

• Group 3A: Lose 3 electrons → +3 charge (e.g., Al3+)

• Group 5A: Gain 3 electrons → −3 charge (e.g., N3−)

• Group 6A: Gain 2 electrons → −2 charge (e.g., O2−)

• Group 7A: Gain 1 electron → −1 charge (e.g., Cl−)

Isoelectronic ions have the same electron count as a noble gas.

Electron Configurations of Ions

• Cations lose their entire valence shell (e.g., Na: 1s22s22p63s1 → Na+: 1s22s22p6).

• Anions fill their valence shell (e.g., Cl: 1s22s22p63s23p5 → Cl−: 1s22s22p63s23p6).

Ionic Formulas, Names, and Properties

Writing Ionic Formulas

The total positive charge must equal the total negative charge. Use subscripts to balance charges.

• Write each ion with its charge (e.g., Ca2+, N3−).

• Find the lowest common multiple of the charges.

• Use subscripts to reach it: Ca3N2.

Polyatomic Ions

A polyatomic ion is a group of covalently bonded atoms with an overall charge. When more than one is needed, use parentheses.

• Examples: NH4+ (ammonium), SO42− (sulfate), NO3− (nitrate).

• Ca2+ + PO43− → Ca3(PO4)2

Naming Ionic Compounds

• Name the cation first, then the anion with an –ide ending (e.g., KCl is potassium chloride).

• Metals with variable charge require a Roman numeral (e.g., FeCl3 is iron(III) chloride).

• Never use prefixes for ionic compounds.

Properties of Ionic Compounds

• High melting point due to strong electrostatic attractions.

• Solid does not conduct electricity; molten or dissolved forms do.

Covalent Bonding Basics

Lewis Structures and Bond Types

A covalent bond is a shared pair of valence electrons. Atoms share electrons to achieve an octet, with exceptions for H, He, and some elements like B, P, S, Cl.

• Single bond: 2 shared electrons (H–H)

• Double bond: 4 shared electrons (O=O)

• Triple bond: 6 shared electrons (N≡N)

• Lone pair: Pair of unshared electrons

Bonding and Lone Pairs by Group

Group	Valence e−	Bonds	Lone Pairs	Examples
1A (H)	1	1	0	H in H2O
4A (C, Si)	4	4	0	CH4
5A (N, P)	5	3	1	NH3
6A (O, S)	6	2	2	H2O
7A (F, Cl)	7	1	3	HF

Lewis Structures

Steps to Draw Lewis Structures

1. Count total valence electrons.

2. Pick the central atom (never H).

3. Connect atoms with single bonds, then add lone pairs to complete octets.

4. Add double or triple bonds if needed, ensuring the total electron count matches.

Examples of Lewis Structures

• CO2: O=C=O, each O has 2 lone pairs.

• HCN: H–C≡N, N has 1 lone pair.

• CH2O: C double-bonded to O, single-bonded to H, O has 2 lone pairs.

Molecular Shapes (VSEPR Theory)

VSEPR and Molecular Geometry

Electron groups around a central atom repel each other, determining the molecule's shape. Each bond (single, double, triple) and lone pair counts as one group.

Electron Groups	Bonded Atoms	Lone Pairs	Geometry	Examples
2	2	0	Linear	CO2, HCN
3	3	0	Trigonal planar	BF3
3	2	1	Bent	SO2
4	4	0	Tetrahedral	CH4
4	3	1	Trigonal pyramidal	NH3
4	2	2	Bent	H2O

Wedge-dash notation is used to represent 3D shapes on paper.

Naming Covalent Compounds and Skeletal Structures

Binary Covalent Compound Naming

• Use Greek prefixes for the number of atoms (mono-, di-, tri-, etc.).

• Name the first element in full, then the second with an –ide ending.

• Drop "mono-" on the first element and vowels before "oxide" (e.g., carbon monoxide).

• Examples: CO2 is carbon dioxide, SiCl4 is silicon tetrachloride.

Skeletal (Line) Structures

Skeletal structures simplify organic molecules by omitting carbon and hydrogen atoms bonded to carbon. Each line end or bend represents a carbon atom.

• Each carbon needs 4 bonds; missing bonds are C–H.

• Atoms other than C and H (O, N, OH) are shown explicitly.

Electronegativity and Polarity

Electronegativity Trends

Electronegativity is the ability of an atom to attract shared electrons. It increases across a period and decreases down a group. Fluorine is the most electronegative element.

Bond Polarity

• Nonpolar covalent: Same element or nearly equal electronegativity (e.g., H–H, Cl–Cl).

• Polar covalent: Different elements with a clear electronegativity difference (e.g., H–F, C=O).

In a polar bond, the more electronegative atom becomes δ−, the other δ+.

Molecular Polarity

A molecule is nonpolar if its shape makes the bond dipoles cancel. Ask: Are there polar bonds? Is the arrangement symmetric?

Molecule	Shape	Polar Bonds?	Dipoles Cancel?	Polarity
CO2	Linear	Yes	Yes	Nonpolar
H2O	Bent	Yes	No	Polar
CH4	Tetrahedral	No	Yes	Nonpolar
NH3	Trigonal pyramidal	Yes	No	Polar

Moles, Molar Mass, and Conversions

The Mole Concept

A mole is a counting unit: 1 mol = 6.022 × 1023 particles (Avogadro's number). Moles connect the number of particles to mass.

Molar Mass

Molar mass (g/mol) is the sum of atomic masses for all atoms in a compound. Use subscripts to multiply atomic masses.

• Glucose (C6H12O6): 6(12.011) + 12(1.008) + 6(15.999) = 180.16 g/mol

• CO2: 12.011 + 2(15.999) = 44.01 g/mol

Conversions

• Particles to moles: $\text{moles} = \frac{\text{number of particles}}{6.022 \times 10^{23}}$

• Moles to particles: $\text{number of particles} = \text{moles} \times 6.022 \times 10^{23}$

• Mass to moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Moles to mass: $\text{mass (g)} = \text{moles} \times \text{molar mass (g/mol)}$

Molarity and Dilution

Molarity (M)

Molarity is the concentration of a solution, defined as moles of solute per liter of solution.

• $\text{Molarity (M)} = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Convert mL to L before using M.

Dilution Calculations

Dilution adds solvent, keeping the moles of solute constant. Use the formula:

• $M_1 V_1 = M_2 V_2$

• Where 1 = stock/concentrated, 2 = diluted.

Small Concentrations

• 1 M = 1000 mM = 106 μM

• To convert: multiply by 1000 for mM, by 106 for μM.

Example: 2.5 × 10−4 M = 0.25 mM = 250 μM

Additional info: These notes cover the essential GOB Chemistry topics for exam preparation, including ionic and covalent compounds, molecular shapes, naming conventions, moles, molar mass, and solution concentration calculations.