General Chemistry Core Concepts: Structured Study Notes

Density and Conversions

Density

Density is a fundamental property of matter, defined as mass per unit volume. It is commonly used to identify substances and solve problems involving mass and volume.

• Definition: Density ($d$) is given by $d = \frac{m}{V}$, where $m$ is mass and $V$ is volume.

• Units: Common units are g/cm3 or kg/m3.

• Example: Water has a density of 1.00 g/cm3 at 4°C.

Conversions

Conversions are essential for working with different units in chemistry.

• Key Point: Use conversion factors to change units (e.g., 1 cm = 0.01 m).

• Example: To convert 100 mL to L: $100 \text{ mL} \times \frac{1 \text{ L}}{1000 \text{ mL}} = 0.1 \text{ L}$

Elements and Atomic Structure

Element Symbols

Each element is represented by a unique one- or two-letter symbol (e.g., H for hydrogen, O for oxygen).

• Periodic Table: The periodic table organizes elements by increasing atomic number.

Subatomic Particles and Isotopes

Atoms consist of protons, neutrons, and electrons. Isotopes are atoms of the same element with different numbers of neutrons.

• Proton: Positively charged, found in nucleus.

• Neutron: Neutral, found in nucleus.

• Electron: Negatively charged, found in electron cloud.

• Isotope: Same number of protons, different number of neutrons.

• Ion: Atom with a net charge due to loss or gain of electrons.

• Example: $^{12}C$ and $^{13}C$ are isotopes of carbon.

Common/Stable Ions

Many elements form ions with characteristic charges (e.g., Na+, Cl-).

• Cations: Positively charged ions (e.g., Na+).

• Anions: Negatively charged ions (e.g., Cl-).

Atomic Data and Periodic Trends

Basic Data: Mass, Charge, Distances

Atomic mass, charge, and atomic radii are key properties for understanding chemical behavior.

• Atomic Mass: Measured in atomic mass units (amu).

• Charge: Protons (+1), electrons (-1), neutrons (0).

• Atomic Radius: Distance from nucleus to outermost electron.

Periodic Table Trends

The periodic table reveals trends in atomic size, ionization energy, and electronegativity.

• Atomic Size: Decreases across a period, increases down a group.

• Ion Size: Cations are smaller, anions are larger than parent atoms.

• Electronegativity: Tendency to attract electrons; increases across a period, decreases down a group.

• Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

Atomic Mass, Mole, and Molar Mass

Reading Data from the Periodic Table

The periodic table provides atomic mass, mass number, and other essential data.

• Atomic Mass: Weighted average of isotopes.

• Mass Number: Total number of protons and neutrons.

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

• Molar Mass: Mass of one mole of a substance (g/mol).

Conversions: Atom/Object Mass & Mole

Relating mass, moles, and number of particles is fundamental in chemistry.

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

• Example: 18 g of H2O is 1 mole.

Electron Configuration and Valence Electrons

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom.

• Condensed Notation: Uses noble gas core (e.g., [Ne]3s2).

• Rules: Aufbau principle, Pauli exclusion, Hund's rule.

Valence Electrons

Valence electrons are the outermost electrons, crucial for chemical bonding.

• Key Point: Group number often indicates number of valence electrons.

• Example: Oxygen has 6 valence electrons.

Effective Nuclear Charge (Screening)

Effective nuclear charge ($Z_{eff}$) is the net positive charge experienced by valence electrons.

• Equation: $Z_{eff} = Z - S$, where $Z$ is atomic number and $S$ is shielding constant.

Isoelectronic Atoms and Ions

Isoelectronic species have the same number of electrons.

• Example: Na+ and Ne are isoelectronic.

Unpaired Electrons (Spin)

Unpaired electrons contribute to magnetic properties.

• Spin Quantum Number: $+\frac{1}{2}$ or $-\frac{1}{2}$.

Quantum Numbers and Electron Configuration Rules

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and electrons.

• Principal (n): Energy level.

• Angular (l): Shape of orbital.

• Magnetic (ml): Orientation.

• Spin (ms): Electron spin.

Electron Configuration Boxes

Hund's rule, Pauli exclusion principle, and Aufbau principle govern electron filling.

• Hund's Rule: Maximize unpaired electrons in degenerate orbitals.

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

• Aufbau Principle: Fill lowest energy orbitals first.

Chemical Bonding

Ionic and Covalent Bonds

Chemical bonds form between atoms to create compounds.

• Ionic Bond: Transfer of electrons between metals and nonmetals.

• Covalent Bond: Sharing of electrons between nonmetals.

• Example: NaCl (ionic), H2O (covalent).

Energy of Photon

The energy of a photon is related to its frequency and wavelength.

• Equation: $E = h\nu$ and $c = \lambda \nu$

• h: Planck's constant, $\nu$: frequency, $\lambda$: wavelength, $c$: speed of light.

Naming Compounds and Lattice Energy

Naming Compounds

Compounds are named according to systematic rules.

• Ionic: Metal + nonmetal (e.g., NaCl).

• Covalent: Prefixes indicate number of atoms (e.g., CO2 is carbon dioxide).

Lattice Energy

Lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions.

• Equation: $U = \frac{kQ_1Q_2}{r}$

• Depends on: Ion charges and distance between ions.

Lewis Structures and Molecular Geometry

Lewis Rules and Octet Rule

Lewis structures represent valence electrons and predict molecular bonding.

• Octet Rule: Atoms tend to have 8 electrons in their valence shell.

• Exceptions: Less than 8 (e.g., H, B) or more than 8 (expanded octet).

Resonance and Electron Domains

Resonance structures depict delocalized electrons; electron domains determine molecular shape.

• Resonance: Multiple valid Lewis structures for a molecule.

• Electron Domain: Regions of electron density around a central atom.

VSEPR Rule and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes.

• Key Shapes: Linear, trigonal planar, tetrahedral, etc.

• Example: CH4 is tetrahedral.

Polarity and Formal Charges

Bond polarity and molecular polarity affect chemical properties; formal charges help identify the best Lewis structure.

• Polar Bond: Unequal sharing of electrons.

• Nonpolar Bond: Equal sharing of electrons.

• Formal Charge: $FC = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}\text{bonding electrons}$

Hybridization and Molecular Orbitals

Hybridization

Atomic orbitals mix to form hybrid orbitals (sp, sp2, sp3).

• Example: CH4 has sp3 hybridization.

Bond Angles and Sigma/Pi Orbitals

Bond angles depend on molecular geometry; sigma and pi bonds describe electron sharing.

• Sigma (σ) Bond: Direct overlap of orbitals.

• Pi (π) Bond: Side-by-side overlap of p orbitals.

• Example: Double bond = 1 σ + 1 π bond.

Molecular Orbitals (MO)

Molecular orbital theory explains bonding using combined atomic orbitals.

• Bonding MO: Lower energy, stabilizing.

• Antibonding MO: Higher energy, destabilizing.

Summary Table: Key Atomic and Molecular Properties

Property	Definition	Example
Density	Mass per unit volume	Water: 1.00 g/cm3
Atomic Number	Number of protons	Carbon: 6
Isotope	Same element, different neutrons	$^{12}C$, $^{13}C$
Ionization Energy	Energy to remove electron	Na: 496 kJ/mol
Electronegativity	Ability to attract electrons	F: 3.98
Mole	6.022 × 1023 particles	1 mol H2O
Valence Electrons	Outer shell electrons	Oxygen: 6
Hybridization	Mixing of orbitals	sp3 in CH4

Additional info: Some content was inferred and expanded for completeness and clarity, including definitions, examples, and equations.