General Chemistry Fundamentals
Build a solid foundation in general chemistry. This guide covers essential concepts from atomic structure to chemical equilibrium, perfect for introductory chemistry courses.
Learning Objectives
- ✓Understand atomic structure and the periodic table
- ✓Master chemical bonding and molecular geometry
- ✓Solve stoichiometry and equilibrium problems
- ✓Apply thermodynamics to chemical reactions
1. Atoms and the Periodic Table
Atoms consist of protons, neutrons, and electrons. The atomic number (Z) defines the element, while isotopes have different numbers of neutrons. The periodic table organizes elements by electronic structure, which determines chemical properties.
Key Points
- •Atomic number = number of protons = number of electrons (neutral atom)
- •Mass number = protons + neutrons
- •Periods are horizontal rows (same number of shells)
- •Groups are vertical columns (same valence electrons)
2. Chemical Bonding
Atoms bond to achieve stable electron configurations. Ionic bonds form when electrons transfer from metals to nonmetals. Covalent bonds form when nonmetals share electrons. Metallic bonds occur in metals where electrons are delocalized.
Key Points
- •Electronegativity difference determines bond type
- •Ionic bonds: EN difference > 1.7
- •Polar covalent: EN difference 0.4-1.7
- •Nonpolar covalent: EN difference < 0.4
3. Stoichiometry
The mole concept is central to chemistry. One mole contains Avogadro's number (6.022 x 10^23) of particles. Stoichiometry uses balanced equations to relate quantities of reactants and products.
Key Points
- •Molar mass (g/mol) = grams per mole of substance
- •Balanced equations give mole ratios
- •Limiting reagent is completely consumed first
- •Percent yield = (actual/theoretical) x 100%
4. Gases and Gas Laws
Ideal gas behavior is described by PV = nRT. Real gases deviate from ideal behavior at high pressure and low temperature. The kinetic molecular theory explains gas properties in terms of molecular motion.
Key Points
- •STP: 0°C (273 K) and 1 atm
- •At STP, 1 mole of ideal gas = 22.4 L
- •Dalton's law: total pressure = sum of partial pressures
- •Graham's law: rate ∝ 1/√M (lighter gases diffuse faster)
5. Solutions and Concentration
Concentration is the amount of solute in a given amount of solution. Molarity (M) is moles per liter. Dilution follows M1V1 = M2V2. Colligative properties depend on the number of solute particles.
Key Points
- •Molarity (M) = moles solute / liters solution
- •Molality (m) = moles solute / kg solvent
- •Like dissolves like - polar dissolves polar
- •Electrolytes dissociate in water, non-electrolytes don't
6. Chemical Equilibrium
At equilibrium, forward and reverse reaction rates are equal, but concentrations remain constant. The equilibrium constant K indicates the position of equilibrium. Le Chatelier's principle predicts how equilibrium responds to stress.
Key Points
- •K > 1: products favored; K < 1: reactants favored
- •Q vs K predicts direction of shift
- •Temperature changes affect K; other changes don't
- •Catalysts don't change equilibrium position
High-Yield Facts
- ★Avogadro's number: 6.022 x 10^23 particles per mole
- ★R = 8.314 J/(mol·K) or 0.0821 L·atm/(mol·K)
- ★Water: density = 1.00 g/mL, specific heat = 4.18 J/(g·°C)
- ★Common strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4
- ★STP has been redefined to 0°C and 1 bar (not 1 atm) in some contexts
Practice Questions
1. How many moles of NaCl are in 29.25 g of NaCl? (Na = 23, Cl = 35.5)
2. What volume does 2.0 mol of an ideal gas occupy at STP?
3. How do you prepare 500 mL of 0.10 M NaCl from a 1.0 M stock solution?
FAQs
Common questions about this topic
Algebra is essential - you'll solve equations and work with proportions constantly. Basic logarithms are needed for pH calculations. Scientific notation and significant figures are used throughout.
Use patterns: -ate ions have more oxygen than -ite ions. Per- means one more oxygen than -ate, hypo- means one less than -ite. Practice writing formulas until they become automatic.