Compounds, Nomenclature, and Molecular Mass – General Chemistry 1A Chapter 3 Study Notes

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Compounds: Basics and Types

Introduction to Compounds

Compounds are substances formed when two or more elements combine chemically. The stability of atoms in compounds is often achieved by attaining a specific number of valence electrons, similar to those of noble gases. This concept is central to understanding why atoms form compounds.

Valence Electrons: Electrons in the outermost shell of an atom, crucial for chemical bonding.
Stable Electron Configuration: Atoms are 'happy' (stable) when they have a full valence shell, often matching the configuration of noble gases.
Types of Compounds: Ionic and Covalent compounds are the two main categories.

Ionic vs. Covalent Compounds

Ionic Compounds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in the formation of ions (cations and anions) held together by electrostatic forces.
Covalent Compounds: Formed by the sharing of electrons between nonmetals, resulting in discrete molecules.
Example: Sodium chloride (NaCl) is ionic; water (H2O) is covalent.

Properties of Compounds

Big Idea 1: Unique Properties

Compounds have properties distinct from their constituent elements. For example, water (H2O) has very different physical and chemical properties compared to hydrogen and oxygen gases.

Selected Properties	Hydrogen	Oxygen	Water
Boiling Point	-253°C	-183°C	100°C
State at Room Temperature	Gas	Gas	Liquid
Flammability	Explosive	Necessary for combustion	Used to extinguish flame

Big Idea 2: Diversity of Compounds

With over 100 elements, countless combinations are possible, resulting in a vast array of molecules and compounds. Chemistry provides systematic ways to describe and represent these compounds.

Ionic Compounds

Formation and Structure

Ionic compounds are formed when metals react with nonmetals, transferring electrons to achieve stable electron configurations.

Metals: Lose electrons to form positively charged cations.
Nonmetals: Gain electrons to form negatively charged anions.
Electrostatic Attraction: The force between oppositely charged ions holds the ionic compound together in a three-dimensional lattice.

Cations and Anions

Anion: An atom that gains electrons and becomes negatively charged. Example: Oxygen atom (Z = 8) gains 2 electrons to become O2–.
Cation: An atom that loses electrons and becomes positively charged. Example: Sodium atom (Z = 11) loses 1 electron to become Na+.

Predicting Ion Charge

The periodic table helps predict the ionic charge of main-group elements:

Group 1 (Alkali Metals): Lose 1 electron, form 1+ cations.
Group 2 (Alkaline Earth Metals): Lose 2 electrons, form 2+ cations.
Group 16 (Chalcogens): Gain 2 electrons, form 2– anions.
Group 17 (Halogens): Gain 1 electron, form 1– anions.

Covalent Compounds

Formation and Structure

Covalent compounds are formed when nonmetals share electrons, resulting in molecules with neutral charge.

Covalent Bond: A chemical bond formed by the sharing of electron pairs between atoms.
Discrete Molecules: Covalent compounds exist as individual molecules, unlike the extended lattice of ionic compounds.
Example: Water (H2O), methane (CH4).

Chemical Formulas and Representations

Molecular and Structural Formulas

Molecular Formula: Indicates the types and numbers of atoms in a molecule (e.g., H2O).
Structural Formula: Shows how atoms are connected (e.g., H–O–H for water).
Other Representations: Ball-and-stick models, space-filling models.

Elements That Exist as Molecules

Diatomic Molecules: H2, N2, O2, F2, Cl2, Br2, I2
Polyatomic Molecules: S8 (sulfur), P4 (phosphorus)

Nomenclature (Naming Compounds)

Ionic Compounds

Monatomic Cations: Named after the element (e.g., Na+ is sodium ion).
Monatomic Anions: Element name ending replaced with “-ide” (e.g., Cl– is chloride).
Polyatomic Ions: Groups of atoms with a charge (e.g., SO42– is sulfate).

Naming Ionic Compounds

Name the cation first, then the anion.
For transition metals with variable charge, specify the charge with Roman numerals (e.g., FeCl2 is iron(II) chloride).

Naming Oxyanions

Oxyanions: Polyatomic ions containing oxygen and another element.
If two oxyanions exist for an element: -ate (more oxygen), -ite (less oxygen).
If more than two exist: per- (most oxygen), hypo- (least oxygen).
Examples: NO3– (nitrate), NO2– (nitrite), ClO4– (perchlorate), ClO– (hypochlorite).

Naming Acids

Binary Acids: Contain hydrogen and one other element. Named as "hydro-" + root + "-ic acid" (e.g., HCl is hydrochloric acid).
Oxyacids: Contain hydrogen, oxygen, and another element. Named based on the oxyanion:
- If the anion ends in "-ate", acid name ends in "-ic" (e.g., H2SO4 is sulfuric acid).
- If the anion ends in "-ite", acid name ends in "-ous" (e.g., H2SO3 is sulfurous acid).

Naming Hydrates

Hydrate: An ionic compound containing water molecules in its crystal structure.
Name the anhydrous compound, then add a Greek prefix for the number of water molecules followed by "hydrate" (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

Nomenclature Prefixes for Molecular Compounds

Prefixes indicate the number of atoms: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.
First element uses prefix only if more than one atom; second element always uses prefix and ends in "-ide" (e.g., CO2 is carbon dioxide).

Molecular Mass and the Mole

The Mole Concept

The mole is a counting unit in chemistry, representing 6.022 × 1023 entities (Avogadro's number). It allows chemists to relate mass to number of atoms, molecules, or ions.

1 mole = 6.022 × 1023 particles
Used for atoms, molecules, ions, or formula units.

Molar Mass

Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).
For elements: Molar mass (g/mol) equals atomic mass (amu) from the periodic table.
For compounds: Sum the atomic masses of all atoms in the formula.
Example: For H2O:

Formula Mass for Ionic Compounds

Ionic compounds do not exist as discrete molecules; instead, use "formula mass" for the mass of one formula unit.
Calculate formula mass by summing the atomic masses of the ions in the empirical formula.

Empirical and Molecular Formulas

Definitions

Empirical Formula: Shows the simplest whole-number ratio of elements in a compound.
Molecular Formula: Shows the actual number of atoms of each element in a molecule.
Example: Hydrogen peroxide: Empirical formula = HO; Molecular formula = H2O2

Determining Empirical Formulas

Convert mass or percent composition to moles for each element.
Divide each by the smallest number of moles.
If necessary, multiply by an integer to obtain whole numbers.

Example: A compound contains 1.71 g C and 0.65 g H. Convert to moles and find the ratio to determine the empirical formula.

Determining Molecular Formulas

Find the empirical formula and its mass.
Divide the compound's molar mass by the empirical formula mass to get a whole number.
Multiply the subscripts in the empirical formula by this number.

Example: Butyric acid has a molar mass of 88.11 g/mol and an empirical formula determined from percent composition. Use the ratio to find the molecular formula.

Summary Table: Key Terms and Examples

Term	Definition	Example
Ionic Compound	Compound formed by transfer of electrons between metal and nonmetal	NaCl
Covalent Compound	Compound formed by sharing electrons between nonmetals	H2O
Empirical Formula	Lowest whole-number ratio of elements	CH
Molecular Formula	Actual number of atoms in a molecule	C2H2
Molar Mass	Mass of one mole of substance	H2O: 18.02 g/mol

Additional info: Some context and examples have been expanded for clarity and completeness, including systematic steps for empirical and molecular formula determination and nomenclature rules for acids and hydrates.