General Chemistry Study Guide: Atomic Structure, Periodic Properties, and Compounds

Chapter 2: Atomic Structure

Shapes of s and p Orbitals

The spatial distribution of electrons in atoms is described by atomic orbitals. The s and p orbitals are fundamental to understanding electron arrangement and chemical bonding.

• s Orbitals: Spherical in shape, centered around the nucleus. The probability of finding an electron is highest near the nucleus and decreases outward.

• p Orbitals: Dumbbell-shaped, oriented along the x, y, and z axes. Each principal energy level (n ≥ 2) contains three p orbitals (px, py, pz).

• Visualization: s orbitals have no angular nodes, while p orbitals have one angular node (a plane where the probability of finding an electron is zero).

• Example: The 2s orbital is spherical, while the 2p orbitals are oriented at right angles to each other.

Chapter 3: Periodic Properties of the Elements

Effective Nuclear Charge

Effective nuclear charge (Zeff) is the net positive charge experienced by an electron in a multi-electron atom. It accounts for both the attraction to the nucleus and the repulsion from other electrons.

• Formula: $Z_{\text{eff}} = Z - S$ Where Z is the atomic number and S is the shielding constant.

• Trend: Zeff increases across a period due to increasing nuclear charge and relatively constant shielding.

• Example: Electrons in the 2p orbital of oxygen experience a higher Zeff than those in nitrogen.

Trends in Atomic and Ionic Radius

The size of atoms and ions changes predictably across the periodic table.

• Atomic Radius: Decreases across a period (left to right) due to increased Zeff; increases down a group due to additional energy levels.

• Ionic Radius: Cations are smaller than their parent atoms; anions are larger. Ionic radius increases down a group.

• Example: Na+ is smaller than Na; Cl- is larger than Cl.

Trends in Ionization Energy and Electron Affinity (s and p Block Elements)

Ionization energy is the energy required to remove an electron from an atom. Electron affinity is the energy change when an atom gains an electron.

• Ionization Energy: Increases across a period; decreases down a group.

• Electron Affinity: Generally becomes more negative across a period (stronger attraction for electrons); less negative down a group.

• Example: Fluorine has a high electron affinity and ionization energy.

Electron Configuration of Atoms and Ions

Electron configuration describes the arrangement of electrons in an atom or ion.

• Notation: Use the format 1s2 2s2 2p6, etc.

• For ions: Remove electrons from the highest energy orbital first (for cations); add to the lowest available orbital (for anions).

• Example: Fe: [Ar] 4s2 3d6; Fe2+: [Ar] 3d6

Hund’s Rule and Pauli Exclusion Principle

These principles govern electron arrangement in orbitals.

• Hund’s Rule: Electrons occupy degenerate orbitals singly before pairing.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Example: In the 2p orbitals, three electrons will occupy each orbital singly before any pairing occurs.

Magnetic Properties: Paramagnetic vs. Diamagnetic

Atoms and ions can be classified based on their magnetic properties.

• Paramagnetic: Contains unpaired electrons; attracted to magnetic fields.

• Diamagnetic: All electrons are paired; repelled by magnetic fields.

• Example: O2 is paramagnetic; N2 is diamagnetic.

Core vs. Valence Electrons

Electrons are classified as core or valence based on their location and involvement in bonding.

• Core Electrons: Located in inner shells; not involved in bonding.

• Valence Electrons: Located in the outermost shell; participate in chemical reactions.

• Example: In sodium (Na), 1s2 2s2 2p6 are core; 3s1 is valence.

Chapter 4: Molecules and Compounds

Ionic and Covalent Compounds

Compounds are classified based on the nature of the bonding between atoms.

• Ionic Compounds: Formed from the transfer of electrons between metals and nonmetals; composed of ions.

• Covalent Compounds: Formed from the sharing of electrons between nonmetals.

• Example: NaCl is ionic; H2O is covalent.

Lewis Dot Structures

Lewis dot structures represent valence electrons as dots around atomic symbols, helping visualize bonding and lone pairs.

• Steps: Count valence electrons, arrange atoms, distribute electrons to satisfy the octet rule.

• Example: For H2O: O has two lone pairs and forms two bonds with H.

Naming Ionic and Binary Molecular Compounds

Systematic naming conventions are used for compounds.

• Ionic Compounds: Name cation first, then anion (with -ide ending).

• Binary Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate number of atoms.

• Example: CO2 is carbon dioxide; NaCl is sodium chloride.

Polyatomic Ions: Formula and Charge

Polyatomic ions are groups of atoms with a net charge. Memorizing their formulas and charges is essential.

Ion Name	Formula	Charge
Carbonate	CO3	2-
Hydroxide	OH	1-
Nitrite	NO2	1-
Nitrate	NO3	1-
Phosphate	PO4	3-
Sulfite	SO3	2-
Sulfate	SO4	2-
Cyanide	CN	1-
Ammonium	NH4	1+

Molar Mass and Percent Composition

Molar mass is the mass of one mole of a compound. Percent composition is the percentage by mass of each element in a compound.

• Formula for Molar Mass: $\text{Molar mass} = \sum \text{(atomic mass of each element)} \times \text{number of atoms}$

• Formula for Percent Composition: $\text{Percent composition} = \frac{\text{mass of element in 1 mole}}{\text{molar mass of compound}} \times 100\%$

• Example: For H2O: Molar mass = 2(1.01) + 16.00 = 18.02 g/mol; % H = (2.02/18.02) × 100% = 11.2%

Conversions: Moles, Grams, and Molecules

Stoichiometric calculations require converting between moles, grams, and molecules.

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

• Moles to Molecules: $\text{Number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

• Example: 36.04 g H2O = 2 moles; 2 moles = $1.204 \times 10^{24}$ molecules.

Number of Moles of Water in Hydrates

Hydrates are compounds containing water molecules in their crystal structure. The number of moles of water can be determined experimentally.

• Procedure: Heat the hydrate to remove water; measure mass loss.

• Formula: $\text{Moles of water} = \frac{\text{mass of water lost}}{18.02 \text{ g/mol}}$

• Example: If 9.0 g of water is lost, moles = 9.0/18.02 = 0.5 mol.

Empirical Formula from Combustion Analysis

Combustion analysis determines the simplest ratio of elements in a compound.

• Steps: Measure mass of CO2 and H2O produced; calculate moles of C and H; determine ratio.

• Formula: $\text{Empirical formula} = \text{smallest whole-number ratio of moles of each element}$

• Example: If 1 mol C and 2 mol H are found, empirical formula is CH2.