General Chemistry I: Study Guide for Chapters 1–4

Exam Coverage and Study Strategies

This study guide outlines the key concepts and skills required for Exam 1, covering Chapters 1–4 of a General Chemistry I course. It includes major topics, essential definitions, and problem-solving strategies.

Chapter 1: Matter, Measurement, and Problem Solving

Classification and Properties of Matter

• Matter is anything that has mass and occupies space. It can be classified by state (solid, liquid, gas) and composition (element, compound, mixture).

• Physical properties can be observed without changing the substance's identity (e.g., melting point, density).

• Chemical properties describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Physical changes do not alter the chemical composition, while chemical changes result in new substances.

Separation Techniques

• Filtration separates solids from liquids.

• Distillation separates substances based on differences in boiling points.

Measurement and Units

• Use the SI system for standard units: meter (m), kilogram (kg), second (s), mole (mol), etc.

• Prefixes (e.g., milli-, centi-, kilo-) indicate multiples or fractions of units.

• Derived units are combinations of base units (e.g., m/s, g/cm3).

Precision, Accuracy, and Significant Figures

• Precision refers to the consistency of repeated measurements.

• Accuracy is how close a measurement is to the true value.

• Random error affects precision; systematic error affects accuracy.

• Use significant figures to reflect the precision of measurements.

Dimensional Analysis

• Use conversion factors to convert between units.

• Set up calculations so that units cancel appropriately.

• Example: To convert 25.0 cm to meters: $25.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.250\ \text{m}$

Chapter 2: Atoms and Elements

Atomic Theory and Structure

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of Definite Proportions: A compound always contains the same elements in the same proportion by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of the second element combine with a fixed mass of the first element in small whole numbers.

• Key experiments: Thomson (discovery of the electron), Millikan (charge of the electron), Rutherford (nuclear model of the atom).

• Atoms consist of protons (positive), neutrons (neutral), and electrons (negative).

• Atomic number (Z): Number of protons; Mass number (A): Protons + neutrons.

• Isotopes are atoms of the same element with different numbers of neutrons.

Periodic Table and Atomic Mass

• Elements are arranged by increasing atomic number.

• Groups/families have similar chemical properties.

• Average atomic mass is calculated using isotopic abundances:

$\text{Average atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

• Mass spectrometry is used to determine isotopic composition and atomic masses.

Calculations with Atoms

• Mole (mol): $1\ \text{mol} = 6.022 \times 10^{23}$ entities (Avogadro's number).

• Key calculations:

  • Number of moles: $\text{mol} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

  • Number of atoms: $\text{atoms} = \text{mol} \times 6.022 \times 10^{23}$

Chapter 3: Molecules and Compounds

Types of Chemical Bonds and Formulas

• Ionic bonds form between metals and nonmetals; covalent bonds form between nonmetals.

• Empirical formula: Simplest whole-number ratio of elements.

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

• Structural formula: Shows how atoms are connected.

• Be able to write and interpret chemical formulas and names.

Elements, Compounds, and Ions

• Distinguish between elements (single type of atom) and compounds (two or more elements chemically combined).

• Know the names and symbols for common elements and polyatomic ions.

• Be able to write names from formulas and vice versa.

Calculations with Compounds

• Calculate:

  • Number of molecules: $\text{molecules} = \text{mol} \times 6.022 \times 10^{23}$

  • Mass percent: $\text{mass \%} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100\%$

  • Empirical formula from experimental data (e.g., combustion analysis).

• Use conversion factors between mass, moles, and number of particles.

Organic and Inorganic Compounds

• Distinguish between organic (carbon-containing) and inorganic compounds.

• Identify simple hydrocarbons (alkanes, alkenes, alkynes).

Chapter 4: Chemical Reactions and Stoichiometry

Types of Chemical Reactions

• Be able to identify and write balanced chemical equations.

• Common reaction types: combination, decomposition, single replacement, double replacement, combustion.

Stoichiometry

• Use balanced equations to relate moles of reactants and products.

• Determine limiting reagent and calculate amounts in excess.

• Calculate theoretical yield and percent yield:

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

• Perform stoichiometric calculations using molar ratios.

Combustion and Other Reactions

• Write balanced equations for combustion reactions (hydrocarbon + O2 → CO2 + H2O).

• Predict products and write equations for reactions of alkali metals and halogens.

Key Tables

Table: Common SI Prefixes

Prefix	Symbol	Factor
kilo-	k	103
centi-	c	10-2
milli-	m	10-3
micro-	μ	10-6
nano-	n	10-9

Table: Types of Chemical Formulas

Type	Description	Example
Empirical	Simplest ratio of elements	CH2O
Molecular	Actual number of atoms	C6H12O6
Structural	Shows connectivity	H–C–O–H

Additional info: Some explanations and tables have been expanded for clarity and completeness based on standard General Chemistry I curriculum.