Chemistry Exam Notes

Every measurement has a degree of uncertainty; some instruments have more uncertainty.
Example of measurement uncertainty:
- 5-gallon bucket vs. syringe for measuring water volume.
Significant figures help account for measurement uncertainty:
- All non-zero digits are significant (e.g., 4.397 kg has 4 significant figures).
- Interior zeroes (between non-zero digits) are significant (e.g., 408 s has 3 SF).
- Leading zeroes are not significant (e.g., 0.000000000126 m has 3 significant digits).
- Trailing zeroes rules:
- After a decimal point: significant (e.g., 4.900 has 4 SF).
- Before a decimal point and after a non-zero digit: significant (e.g., 3900. has 4 SF).
- Before an implied decimal point: ambiguous.
Calculation rules for significant figures:
- Addition/Subtraction: Final answer matches the least precise quantity's place.
- Multiplication/Division: Final answer matches the least precise quantity's significant figures.

Isotopes are atoms of the same element with different neutron counts, leading to different mass numbers.
Example: Chlorine has isotopes Cl-35 (34.97 amu) and Cl-37 (36.97 amu).
Periodic tables show weighted average masses, accounting for isotope abundance.
- Average mass calculation does not simply involve averaging mass numbers.
- Use equation: ar{m} = \sum mi ai.

Mass of every atom except H-1 is smaller than the sum of its constituents’ masses.
Molar mass connects the average element mass with 6.022 \times 10^{23} atoms (Avogadro's number).
A mole is analogous to a dozen, representing a fixed number.
Example: 12.0 g of carbon has 6.022 \times 10^{23} C atoms.

Electrons behave as both particles and waves (wave-particle duality).
Light exhibits wave properties (amplitude, wavelength, frequency).
J. C. Maxwell showed light propagates as an electric and magnetic field.
Energy of photons can be calculated with: E = h
u = \frac{hc}{\lambda}.
Significant aspects include:
- n can be small, giving rise to quantized states of electrons.
- Photons exhibit energy proportional to their frequency.

Quantum mechanics presents limits on our knowledge of subatomic particles, expressed through Heisenberg's uncertainty principle:
- \Delta x \Delta (mv) \geq \frac{h}{4\pi}.
Schrödinger's model involves wave functions to describe electron properties.
Quantum numbers specify electron orbitals crucial for periodic table and chemical reactivity.
Electron configurations denote arrangement in orbitals, governed by:
- Pauli Exclusion Principle: No two electrons can share identical quantum numbers.
- Hund's Rule: Electrons will occupy degenerate orbitals singly before pairing.

Types of Bonds:
- Ionic Bonds: Electrostatic attraction between cations and anions (e.g., NaCl).
- Covalent Bonds: Sharing of electrons between nonmetals (e.g., H2).
Covalent bonds vary in strength and can exhibit characteristics of both ionic and covalent bonds depending on the elements involved.

Stoichiometry involves quantitative relationships between reactants and products in a reaction.
Limiting reactants determine product yield based on reactant quantities.
- Example: In the reaction N2 + 3H2 \rightarrow 2NH_3, know how to identify the limiting reactant.

Solutions are homogenous mixtures expressed in molarity (M = moles/volume in liters).
Example: 1.00 M NaCl = 1.00 mol NaCl in 1.00 L solution.
Dilution formula: M1V1 = M2V2.

Energy conservation: energy can be transferred but not created or destroyed.
- Heat transfer alters temperature: q = m c_s \Delta T.
Enthalpy change during reactions (ΔH) relates to energy changes in constant pressure.

Vaporization and condensation occur at boiling points, influenced by pressure.
Enthalpy of vaporization (ΔHvap): energy required to vaporize a unit mass of substance.
Phase diagrams represent the states of matter depending on temperature and pressure, defining transitions such as melting, freezing, vaporization, and sublimation.
Dynamic equilibrium is where rates of phase transitions equalize each other.