General Chemistry Study Guide: Atomic Structure, Periodic Table, and Chemical Compounds

Atomic Structure and Electron Configuration

Electron Configurations and Orbital Diagrams

Understanding electron configurations and orbital diagrams is essential for predicting the chemical behavior of elements. Electron configurations describe the arrangement of electrons in an atom's orbitals, while orbital diagrams visually represent these arrangements.

• Electron Configuration: The notation that shows the distribution of electrons among the orbitals of an atom (e.g., 1s2 2s2 2p6).

• Orbital Diagram: Uses arrows to represent electrons in each orbital, following the Pauli exclusion principle and Hund's rule.

• Example: The electron configuration for Oxygen (O) is 1s2 2s2 2p4.

Periodic Table Trends

The periodic table organizes elements by increasing atomic number and reveals recurring trends in their properties. These trends are crucial for predicting element behavior.

• Atomic Radius: Generally decreases across a period and increases down a group.

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

• Electronegativity: Increases across a period, decreases down a group.

• Electron Affinity: Tends to become more negative across a period.

• Example: Sodium (Na) has a larger atomic radius than Chlorine (Cl).

Periodic Properties and Their Explanations

Periodic properties such as atomic radius, ionization energy, and electronegativity are explained by the arrangement of electrons and the effective nuclear charge.

• Effective Nuclear Charge (Zeff): The net positive charge experienced by valence electrons.

• Shielding Effect: Inner electrons shield outer electrons from the nucleus, affecting periodic trends.

• Example: The first ionization energy of Na is lower than that of Mg due to lower Zeff.

Classification of Elements and Chemical Species

Isoelectronic Species

Isoelectronic species are atoms, ions, or molecules with the same number of electrons.

• Example: Ne, Na+, and F- are isoelectronic.

Diamagnetic and Paramagnetic Species

Magnetic properties depend on electron configurations.

• Diamagnetic: All electrons are paired; not attracted to a magnetic field.

• Paramagnetic: Has unpaired electrons; attracted to a magnetic field.

• Example: O2 is paramagnetic due to unpaired electrons.

Periodic Trends and Their Applications

Trends in Atomic and Ionic Radii

Atomic and ionic radii change predictably across periods and groups.

• Atomic Radius: Decreases left to right, increases top to bottom.

• Ionic Radius: Cations are smaller than their parent atoms; anions are larger.

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

Trends in Ionization Energy

Ionization energy is the energy required to remove an electron from an atom.

• Trend: Increases across a period, decreases down a group.

• Example: Mg has a higher ionization energy than Na.

Trends in Metallic and Nonmetallic Character

Metallic character decreases across a period and increases down a group.

• Metals: Good conductors, malleable, tend to lose electrons.

• Nonmetals: Poor conductors, brittle, tend to gain electrons.

• Example: Na is more metallic than Cl.

Chemical Bonding and Compound Formation

Lewis Dot Structures

Lewis dot structures represent valence electrons in atoms and molecules, helping predict bonding and molecular shape.

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

• Example: H2O has two lone pairs on oxygen and two single bonds to hydrogen.

Naming Compounds

Systematic naming of compounds follows IUPAC rules for binary ionic and molecular compounds.

• Binary Ionic Compounds: Name cation first, then anion (e.g., NaCl: sodium chloride).

• Binary Molecular Compounds: Use prefixes to indicate number of atoms (e.g., CO2: carbon dioxide).

Writing Chemical Formulas

Formulas represent the composition of compounds using element symbols and subscripts.

• Empirical Formula: Simplest whole-number ratio of atoms.

• Molecular Formula: Actual number of atoms in a molecule.

• Example: Empirical formula of glucose is CH2O; molecular formula is C6H12O6.

Quantitative Chemistry: Moles, Mass, and Formulas

Mole Concept and Molar Mass

The mole is a fundamental unit for counting particles in chemistry. Molar mass is the mass of one mole of a substance.

• Mole: $1 \text{ mole} = 6.022 \times 10^{23}$ particles (Avogadro's number).

• Molar Mass: $\text{Molar mass} = \sum \text{atomic masses of elements in the formula}$

• Example: Molar mass of H2O is $2 \times 1.008 + 16.00 = 18.016$ g/mol.

Conversions Between Mass, Moles, and Number of Particles

Conversions are essential for quantitative chemical calculations.

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

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

• Example: 36.0 g of water is $\frac{36.0}{18.016} = 2.0$ moles, which is $2.0 \times 6.022 \times 10^{23} = 1.204 \times 10^{24}$ molecules.

Empirical and Molecular Formulas

Empirical formulas show the simplest ratio of elements; molecular formulas show the actual number of atoms.

• Empirical Formula Calculation: Divide the number of moles of each element by the smallest value to get the ratio.

• Molecular Formula Calculation: $\text{Molecular formula} = \text{Empirical formula} \times n$, where $n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}$

• Example: If empirical formula is CH2O and molar mass is 180 g/mol, $n = \frac{180}{30} = 6$, so molecular formula is C6H12O6.

Calculating Mass of Elements in Compounds

To find the mass of a specific element in a compound, use the percent composition or molar ratios.

• Percent Composition: $\text{Percent} = \frac{\text{Mass of element in 1 mole}}{\text{Molar mass of compound}} \times 100\%$

• Example: In H2O, percent H = $\frac{2.016}{18.016} \times 100\% = 11.2\%$

Combustion Analysis

Combustion analysis is used to determine empirical formulas from the masses of products formed when a compound is burned.

• Steps: Measure mass of CO2 and H2O produced, calculate moles of C and H, determine empirical formula.

• Example: Burning 1.00 g of a compound produces 1.47 g CO2 and 0.60 g H2O; use these to find moles of C and H.

Summary Table: Periodic Trends

The following table summarizes key periodic trends for elements in the periodic table.

Property	Across a Period (Left to Right)	Down a Group (Top to Bottom)
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases
Metallic Character	Decreases	Increases

Additional info: These notes expand on the syllabus outline by providing definitions, examples, and formulas for each topic, ensuring a self-contained study guide for exam preparation.