General Chemistry Study Guide: Constants, Units, Atomic Theory, Chemical Reactions, and Problem Solving

Fundamental Constants and Units

Physical Constants

Physical constants are essential values used in chemical calculations and problem solving. They provide the basis for quantitative work in chemistry.

• Avogadro's number (Na): $6.0221367 \times 10^{23}$ mol-1 — the number of particles in one mole of substance.

• Electron charge (e): $1.6022 \times 10^{-19}$ C — the charge carried by a single electron.

• Electron mass: $9.1094 \times 10^{-28}$ g — mass of an electron.

• Faraday constant (F): $96,485$ C/mol e- — charge per mole of electrons.

• Gas constant (R): $0.082057$ L·atm/(mol·K) or $8.314$ J/(mol·K) — used in gas law calculations.

• Planck's constant (h): $6.6262 \times 10^{-34}$ J·s — relates energy and frequency in quantum mechanics.

• Proton mass: $1.6726 \times 10^{-24}$ g — mass of a proton.

• Neutron mass: $1.6749 \times 10^{-24}$ g — mass of a neutron.

• Speed of light (c): $2.99792458 \times 10^8$ m/s — speed at which light travels in a vacuum.

Useful Conversion Factors and Relationships

Conversions between units are frequently required in chemistry. The following relationships are commonly used:

Quantity	Conversion
1 lb	453.6 g
1 cal	4.184 J (exactly)
1 in	2.54 cm (exactly)
1 L	1.0567 qt
1 mi	1.609 km
1 km	0.6215 mi
1 J	1 C × 1 V
1 Pa	1 N/m2
1 pm	$1 \times 10^{-12}$ m

Pressure Units

Pressure is a key concept in gas laws and solution chemistry. Common units and their relationships:

Unit	Equivalent
1 atm	760 mmHg = 760 torr
1 atm	101.325 kPa
1 bar	1 × 105 Pa = 100 kPa
1000 Pa	1 kPa
1 bar	0.9869 atm

Equations

Common equations for temperature conversion, density, and molarity:

• Temperature conversions:

  • $^{\circ}C = (^{\circ}F - 32^{\circ}) \frac{5}{9}$

  • $^{\circ}F = (^{\circ}C \frac{9}{5}) + 32^{\circ}$

• Density: $d = \frac{m}{V}$

• Molarity: $M_1 V_1 = M_2 V_2$

Atoms, Elements, and Atomic Theory

Atomic Structure and Subatomic Particles

Atoms are composed of protons, neutrons, and electrons. Their arrangement determines the properties of elements.

• Protons: Positively charged particles found in the nucleus.

• Neutrons: Neutral particles found in the nucleus.

• Electrons: Negatively charged particles found outside the nucleus.

• Atomic number (Z): Number of protons in the nucleus.

• Mass number (A): Sum of protons and neutrons.

Example: Carbon-12 has 6 protons and 6 neutrons.

Dalton's Atomic Theory and Modifications

Dalton's atomic theory laid the foundation for modern chemistry, but has been modified over time:

• Atoms of a given element are identical (now known to be isotopes).

• All matter is made up of very small particles called atoms.

• Compounds are formed from atoms.

• Chemical reactions involve rearrangement of atoms.

Molecules, Compounds, and Chemical Formulas

Chemical Formulas and Nomenclature

Chemical formulas represent the composition of compounds. Nomenclature rules help name compounds systematically.

• Empirical formula: Simplest whole-number ratio of atoms in a compound.

• Molecular formula: Actual number of atoms of each element in a molecule.

• Example: Glucose: Empirical formula CH2O, Molecular formula C6H12O6.

Writing Chemical Formulas

Formulas for compounds are written using element symbols and subscripts to indicate the number of atoms.

• Phosphorous Trioxide: P2O3

• Gold (III) Bromide: AuBr3

• Lithium Permanganate: LiMnO4

• Aluminium Selenide: Al2Se3

• Ammonium Fluoride: NH4F

• Dichlorine Heptoxide: Cl2O7

Chemical Reactions and Stoichiometry

Types of Chemical Reactions

Chemical reactions involve the transformation of substances. Key types include:

• Redox reactions: Transfer of electrons between species.

• Acid-base reactions: Transfer of protons (H+).

• Precipitation reactions: Formation of an insoluble product.

Balancing Chemical Equations

Balancing equations ensures the conservation of mass and charge. Coefficients are used to equalize the number of atoms on both sides.

• Example: $C_3H_8 + O_2 \rightarrow CO_2 + H_2O$

• Balanced: $C_3H_8 + 5O_2 \rightarrow 3CO_2 + 4H_2O$

Stoichiometry and Molarity

Stoichiometry involves quantitative relationships in chemical reactions. Molarity is a measure of concentration:

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: To dilute 0.450 M K2CO3 to 0.315 M, use $M_1V_1 = M_2V_2$.

Lab Techniques and Problem Solving

Empirical and Percent Composition Calculations

Determining the empirical formula and percent composition is essential for analyzing compounds.

• Percent composition: $\%\text{element} = \frac{\text{mass of element}}{\text{total mass}} \times 100$

• Empirical formula: Find the simplest ratio of moles of each element.

Limiting Reactant and Yield Calculations

Identifying the limiting reactant and calculating theoretical and percent yield are key steps in reaction analysis.

• Limiting reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical yield: Maximum amount of product possible from given reactants.

• Percent yield: $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$

Example Problem: Ostwald Process

The Ostwald process for producing nitric acid involves:

• $4NH_3 (g) + 5O_2 (g) \rightarrow 4NO (g) + 6H_2O (g)$

• Given masses of reactants, determine limiting reactant, theoretical yield, and percent yield.

Additional info:

• Some questions and problems reference solubility rules, oxidation numbers, and organization of subatomic particles, which are foundational topics in general chemistry.

• Tables have been recreated and expanded for clarity and completeness.