General Chemistry I: Key Concepts, Formulas, and Practice Questions

General Chemistry I: Essential Concepts and Practice

Introduction

This study guide summarizes foundational topics in General Chemistry, including matter, measurement, atomic structure, chemical bonding, stoichiometry, and essential formulas. It also provides worked examples and practice questions to reinforce understanding.

Matter, Measurement, and Problem Solving

Classification of Matter

• Solid: Definite shape and definite volume. Particles are closely packed in a fixed arrangement.

• Liquid: Definite volume but no definite shape. Particles are close but can move past one another.

• Gas: Neither definite shape nor volume. Particles are far apart and move freely.

Mixtures vs. Pure Substances:

• Homogeneous mixture: Uniform composition throughout (e.g., saltwater).

• Heterogeneous mixture: Non-uniform composition (e.g., salad).

Significant Figures: The number of meaningful digits in a measurement. For example, the number 2 in "2 atoms of carbon" has 1 significant figure.

Measurement Units and Conversions:

• 1 cm3 = 1 mL

• 1 in = 2.54 cm

• 1 L = 1.057 qt

• 1 lb = 453.6 g

Temperature Conversions:

• $T(K) = T(^\circ C) + 273.15$

Example: Convert 35°C to Kelvin: $T(K) = 35 + 273.15 = 308.15\ K$

Atoms and Elements

Atomic Structure

• Proton: Charge = +1

• Neutron: Charge = 0

• Electron: Charge = -1

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

Mass Number (A): Number of protons + neutrons.

Isotopes: Atoms of the same element with different numbers of neutrons.

Ions: Atoms or molecules that have gained or lost electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Example: The Ca2+ ion has 18 electrons (20 protons - 2 positive charge).

The Periodic Table and Periodic Properties

Classification of Elements

• Alkali metals: Group 1 elements (e.g., Li, Na, K).

• Alkaline earth metals: Group 2 elements (e.g., Be, Mg, Ca).

• Nonmetals: Elements that are not metals (e.g., C, N, O, F, Cl).

Example: Calcium (Ca) is an alkaline earth metal; chlorine (Cl) is a nonmetal.

Molecules and Compounds

Chemical Formulas and Naming

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

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

• Naming Binary Ionic Compounds: Name the cation first, then the anion (e.g., sodium chloride).

Example: The formula for lithium dichromate is Li2Cr2O7.

Chemical Reactions and Stoichiometry

Balancing Chemical Equations

• Ensure the same number of each type of atom on both sides of the equation.

Stoichiometric Coefficients: Numbers placed in front of compounds to balance equations.

Example: For the reaction $\mathrm{CH_4(g) + O_2(g) \rightarrow CO_2(g) + H_2O(l)}$, the balanced equation is $\mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$.

Mole Concept and Calculations

• Avogadro's number: $6.022 \times 10^{23}$ particles/mol

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

Example: To find the number of atoms in 2.50 moles of Cl, multiply by Avogadro's number: $2.50 \times 6.022 \times 10^{23} = 1.51 \times 10^{24}$ atoms.

Limiting Reactant and Theoretical Yield

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

• Theoretical yield is the maximum amount of product that can be formed from the given amounts of reactants.

Mathematical Operations and Functions

Key Equations and Constants

• Gas Laws:

  • Ideal Gas Law: $PV = nRT$

  • R (gas constant): $0.0821\ \mathrm{L\cdot atm\cdot mol^{-1}\cdot K^{-1}}$

• Thermochemistry:

  • Heat: $q = mc\Delta T$

  • Enthalpy: $\Delta H = \sum n\Delta H_f^\circ(\text{products}) - \sum n\Delta H_f^\circ(\text{reactants})$

• Equilibrium:

  • $K_w = [H^+][OH^-] = 1 \times 10^{-14}$ at 25°C

  • $pH = -\log[H^+]$

• Electrochemistry:

  • $E_{cell} = E_{cathode} - E_{anode}$

  • $\Delta G = -nFE_{cell}$

  • Faraday's constant: $F = 96485\ \mathrm{C\cdot mol^{-1}}$

Atomic and Molecular Structure

Electron Configuration and Orbitals

• Electrons occupy orbitals in order of increasing energy (Aufbau principle).

• Pauli exclusion principle: Each orbital holds a maximum of 2 electrons with opposite spins.

• Hund's rule: Electrons fill degenerate orbitals singly before pairing.

Sample Table: Atomic Orbitals

Orbital Type	n	l	Number of Orbitals	Possible Electrons
s	1	0	1	2
p	2	1	3	6
d	3	2	5	10

Practice Questions (Selected)

1. Identify a solid. Answer: Definite volume and definite shape.

2. Two or more substances in variable proportions, where the composition is constant throughout, are a: Answer: Homogeneous mixture.

3. Identify the type of energy that is NOT chemical energy. Answer: Ball rolling down hill (kinetic energy).

4. The variable temperature is 35°C. What is the temperature in K? Answer: 308 K

5. Read the length of the metal bar with the correct number of significant figures. Answer: 15.00 cm

6. An ethane molecule contains 2 atoms of carbon. The number 2 represents how many significant figures? Answer: 1

7. The average distance between nitrogen and oxygen atoms is 115 pm in a compound called nitric oxide. What is this distance in nanometers? Answer: $1.15 \times 10^{-1}$ nm

8. Identify the charges of the proton, neutron, and electron. Answer: Proton: +1, Neutron: 0, Electron: -1

9. How many neutrons are in nickel-59? Answer: 31

10. How many electrons does the Ca2+ ion possess? Answer: 18

Additional info: More practice questions are available in the original file, covering topics such as empirical formulas, balancing equations, and stoichiometry.

Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties into columns (groups or families). Periods are horizontal rows.

Conclusion

This guide covers the essential concepts, formulas, and problem-solving strategies for General Chemistry I. Mastery of these topics is foundational for success in subsequent chemistry courses.