CHE121 Exam 1 Study Guide

Exam 1 Study Guide

Chapter 1 – Matter and Measurements

Physical and Chemical Properties

Understanding the nature of matter and its transformations is fundamental to chemistry. Matter can be described by its physical and chemical properties, which help distinguish between different substances and predict their behavior in various conditions.

• Physical Properties: Characteristics that can be observed or measured without changing the substance’s identity (e.g., melting point, density, color).

• Chemical Properties: Characteristics that describe a substance’s ability to undergo chemical changes (e.g., flammability, reactivity).

• Physical Change: A change that does not alter the chemical composition of a substance (e.g., melting ice).

• Chemical Change: A process in which one or more substances are converted into new substances (e.g., rusting of iron).

• Example: Boiling water is a physical change; burning wood is a chemical change.

Classification of Matter

• Element: A pure substance that cannot be broken down into simpler substances by chemical means (e.g., oxygen, O2).

• Compound: A substance composed of two or more elements chemically combined in fixed proportions (e.g., water, H2O).

• Mixture: A combination of two or more substances that are not chemically bonded (e.g., air, saltwater).

• Homogeneous Mixture: Uniform composition throughout (solution).

• heterogeneous Mixture: Non-uniform composition (e.g., salad).

Measurement and Significant Figures

• Measurements in chemistry must be reported with the correct number of significant figures to reflect the precision of the instrument used.

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

• Use significant figures in calculations and when reporting results.

• Example: 0.00450 has three significant figures.

Units and Conversions

• SI units are the standard units used in chemistry (meter, kilogram, second, mole, etc.).

• Temperature conversions:

  • Fahrenheit to Celsius: $T(^{\circ}F) = 1.8 \times T(^{\circ}C) + 32$

  • Celsius to Kelvin: $T(K) = T(^{\circ}C) + 273.15$

• Mass, volume, and density:

  • Density: $d = \frac{m}{V}$

  • 1 mL = 1 cm3

• Energy:

  • 1 cal = 4.184 J

Problem Solving and Dimensional Analysis

• Use dimensional analysis (factor-label method) to convert between units.

• Set up conversion factors so that units cancel appropriately.

• Example: To convert 257°F to °C: $T(^{\circ}C) = \frac{T(^{\circ}F) - 32}{1.8}$

Density and Specific Gravity

• Density: $d = \frac{m}{V}$ (mass per unit volume)

• Specific Gravity: Ratio of the density of a substance to the density of water (unitless).

• Example: If gasoline has a density of 0.68 g/cm3, the mass of 10.0 gallons can be found by converting gallons to cm3 and multiplying by density.

Chapter 2 – Atoms and the Periodic Table

Atomic Theory and Structure

The atomic theory explains the nature of matter by describing the structure and behavior of atoms, the fundamental building blocks of all substances.

• Modern Atomic Theory: Atoms are composed of subatomic particles: protons, neutrons, and electrons.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

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

• Atomic Mass Unit (amu): Standard unit for atomic and molecular masses.

• Average Atomic Mass: Weighted average of the masses of all naturally occurring isotopes of an element.

The Periodic Table

• Elements are arranged in order of increasing atomic number.

• Groups (columns) and periods (rows) organize elements with similar properties.

• Metals, Nonmetals, Metalloids: Classification based on physical and chemical properties.

• Predict chemical and physical characteristics based on position in the table.

Electron Configuration

• Describes the arrangement of electrons in an atom.

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

• Example: Electron configuration for tin (Sn): [Kr] 5s2 4d10 5p2

• Valence electrons are those in the outermost shell and determine chemical reactivity.

Orbitals and Electron Capacity

• Each orbital can hold a maximum of two electrons (Pauli exclusion principle).

• Subshells (s, p, d, f) have different shapes and capacities:

  • s: 2 electrons

  • p: 6 electrons

  • d: 10 electrons

  • f: 14 electrons

Key Constants and Equations

Quantity	Equation/Value
Temperature (°F to °C)	$T(^{\circ}F) = 1.8 \times T(^{\circ}C) + 32$
Temperature (°C to K)	$T(K) = T(^{\circ}C) + 273.15$
Density	$d = \frac{m}{V}$
Heat (q)	$q = m s \Delta T$
1 cal	$= 4.184 \text{ J}$
1 amu	$= 1.660539 \times 10^{-24} \text{ g}$
1 kg	$= 2.205 \text{ lb}$
1 in.	$= 2.54 \text{ cm}$ (exactly)
1 mL	$= 1 \text{ cm}^3$

Sample Problems and Applications

• Classify changes as physical or chemical.

• Write numbers in scientific notation.

• Convert between temperature scales.

• Calculate mass, volume, or density using the appropriate formula.

• Determine the number of atoms in a given mass using Avogadro’s number.

• Write electron configurations and determine valence electrons.

Additional info: These notes are based on a course outline, review problems, and a constants/equations sheet for a General Chemistry I course. They cover foundational topics in matter, measurement, atomic structure, and the periodic table, suitable for exam preparation.