CHEM 0110 — General Chemistry 1

CHEM 0110 Study Guide

General Chemistry 1

This guide covers the quantitative and structural foundations of general chemistry: matter, atoms, reactions, thermochemistry, electronic structure, bonding, and states of matter. Pitt's published syllabus identifies these topics and includes laboratory work. Course scope: Pitt CHEM 0110 syllabus. Check the chemistry syllabus directory for your section.

These are original study notes with invented practice examples. They are supplementary to your assigned problems and lab manual. Equations are written in editable text. A subscript written after an underscore is a label; Δ means final minus initial.

Core content in a useful study order

Measurement. Write units on every number. Use dimensional analysis to connect a known measurement to a desired quantity. Carry extra digits during calculations and round at the end.
Atoms and formulas. Understand isotopes, ions, average atomic mass, formula units, and naming. A chemical formula gives composition; a balanced equation gives reacting mole ratios.
Moles and reactions. Convert mass to moles before applying reaction coefficients. Determine which reactant limits product formation.
Aqueous reactions. Recognize precipitation, acid–base, and oxidation–reduction patterns. Separate dissolved ions only when appropriate.
Energy. Define the system and surroundings before assigning signs to heat and work. Connect measured temperature changes to reaction energy.
Electronic structure. Relate light to energy, then use electron configurations to explain periodic trends.
Bonding and shape. Draw Lewis structures, count electron domains, infer geometry, and evaluate the overall molecular dipole.
Gases and condensed phases. Use particle models to explain gas behavior and intermolecular forces to explain boiling, melting, and solubility trends.

Important vocabulary

Element: Substance whose atoms have the same number of protons.
Isotope: Atom differing in neutron number from another isotope of the same element.
Ion: Charged atom or group of atoms.
Mole: Amount containing exactly 6.02214076 × 10^{23} specified entities.
Molar mass: Mass per mole, commonly g/mol.
Empirical formula: Simplest whole-number ratio of elements in a compound.
Molecular formula: Actual numbers of atoms in one molecule.
Stoichiometry: Quantitative relationships in a balanced chemical reaction.
Limiting reactant: Reactant that allows the smallest amount of product from the available quantities.
Theoretical yield: Maximum predicted product amount under the reaction model.
Molarity: Moles of solute per liter of solution.
Electrolyte: Substance that produces mobile ions in solution.
Spectator ion: Ion unchanged by the net reaction.
Oxidation: Loss of electrons or increase in oxidation number.
Reduction: Gain of electrons or decrease in oxidation number.
Enthalpy: State function H = U + pV; its change is heat at constant pressure under usual non-expansion-work restrictions.
Exothermic: Releases heat to surroundings; reaction ΔH is negative.
Endothermic: Absorbs heat from surroundings; reaction ΔH is positive.
Orbital: Quantum state describing an electron's spatial probability distribution.
Effective nuclear charge: Net nuclear attraction experienced by an electron after shielding.
Electronegativity: Tendency of a bonded atom to attract electron density.
Formal charge: Electron-bookkeeping charge assigned within a Lewis structure.
Resonance: Several valid electron arrangements representing one delocalized structure.
VSEPR: Model relating electron-domain repulsion to molecular geometry.
Intermolecular force: Attraction between particles, distinct from bonds within a molecule.
Hydrogen bonding: Particularly important attraction involving H bonded to N, O, or F and a suitable acceptor.

Essential formulas and constants

Measurement and composition

Density ρ = m/V. Keep mass and volume units consistent.
n = m/Mmolar; Nparticles = nNA.
NA = 6.02214076 × 10^{23} mol^{-1}.
Average atomic mass = Σ(isotope fraction × isotope mass); fractions sum to 1.
Mass percent = component mass/total mass × 100%.
Percent yield = actual yield/theoretical yield × 100%; use matching units.
Molarity c = n/Vsolution, with V in liters for mol/L.
Dilution c_{1}V_{1} = c_{2}V_{2} when the amount of the same solute is conserved.

Reactions and heat

For aA → bB, nB = nA(b/a), provided A is the relevant limiting amount.
q = mcspΔT for heating without a phase change and approximately constant specific heat.
q = nΔHtransition for a phase change using a molar enthalpy.
ΔU = q + w using the chemistry sign convention: energy entering the system is positive.
Expansion work w = −PexternalΔV for constant external pressure.
Calorimetry: qreaction + qsolution + qcalorimeter = 0 in an insulated model.
Hess's law: Add reaction enthalpies when adding equations; reverse the sign when reversing a reaction.
ΔH°reaction = ΣνΔHf°products − ΣνΔHf°reactants, where ν are balanced coefficients.
Bond-energy estimate: ΔH ≈ Σenergy of bonds broken − Σenergy of bonds formed. This is an approximation using averaged bond energies.

Light, bonding, and gases

c_light = λν; Ephoton = hν = hc_light/λ.
h = 6.62607015 × 10^{-34} J·s; c_light = 2.99792458 × 10^{8} m/s.
De Broglie wavelength λ = h/p, where p is momentum.
Formal charge = valence electrons − nonbonding electrons − (bonding electrons/2).
Ideal gas: PV = nRT; T must be kelvin and P absolute pressure.
R = 0.082057 L·atm/(mol·K) or 8.314 J/(mol·K); choose the form matching units.
Kelvin T = Celsius temperature + 273.15.
For fixed n: P_{1}V_{1}/T_{1} = P_{2}V_{2}/T_{2}.
Dalton's law: Ptotal = ΣPi; Pi = XiPtotal for an ideal mixture, where Xi = ni/ntotal.
Ideal-gas density ρ = PMmolar/(RT).

Worked examples

Example 1 Limiting reactant

For 2H_{2} + O_{2} → 2H_{2}O, combine 3.0 mol H_{2} and 2.0 mol O_{2}. Hydrogen can form 3.0 mol water; oxygen can form 4.0 mol water. Hydrogen limits the reaction, so the theoretical yield is 3.0 mol H_{2}O. Oxygen consumed is 1.5 mol, leaving 0.5 mol. Comparing reactant masses directly would not identify the limiting reactant correctly.

Example 2 Calorimetry

A reaction warms 100.0 g of water-like solution from 22.0°C to 27.0°C. Assume csp = 4.184 J/(g·°C) and negligible calorimeter heat capacity. qsolution = 100.0 × 4.184 × 5.0 = 2092 J ≈ 2.1 kJ. Therefore qreaction ≈ −2.1 kJ. The solution warms because the reaction releases heat.

Example 3 Dilution

Make 250.0 mL of 0.100 M solution from 1.00 M stock. Vstock = (0.100 × 250.0)/1.00 = 25.0 mL. The target is 250.0 mL final solution volume. “Add 250.0 mL water” would give the wrong concentration.

Practice questions and answers

Moles in 5.85 g NaCl, using 58.44 g/mol? 0.100 mol.
Pressure of 0.500 mol ideal gas in 10.0 L at 300 K? 1.23 atm.
A reaction has theoretical yield 8.00 g and actual yield 6.20 g. Percent yield? 77.5%.
Which has higher photon energy: 400 nm or 700 nm light? 400 nm, because energy varies inversely with wavelength.
Why is CO_{2} nonpolar while H_{2}O is polar? CO_{2} has cancelling bond dipoles in a linear geometry; bent H_{2}O has a net dipole.
What is the formal charge on N with four single bonds and no lone pair? +1.

Mistakes to catch

Treating reaction coefficients as mass ratios.
Changing subscripts instead of coefficients to balance a reaction.
Using Celsius in gas laws.
Confusing electron-domain geometry with the arrangement of atoms.
Breaking bonds without accounting for the energy released when new bonds form.
Rounding intermediate results so heavily that the final answer changes.

Review routine

Use three passes: explain the particle-level idea, solve a numerical example with units, then explain why a plausible wrong answer is wrong. For labs, record raw measurements before calculations and connect the result to a chemical claim. A useful notebook separates observation, calculation, and interpretation rather than combining them into one unsupported conclusion.