General Chemistry I: Exam 1 Study Guide and Key Concepts

Ch. 1: Atoms, Molecules, and Measurement

1.1 Atoms and Molecules

• Definitions: Chemistry is the study of matter, its properties, and the changes it undergoes. Atoms are the smallest units of elements, and molecules are combinations of atoms bonded together.

1.2 The Scientific Approach to Knowledge

• Scientific Method: Involves observation, hypothesis, experimentation, and theory development.

• Law vs. Theory: A law summarizes observations; a theory explains them.

1.3 Classification of Matter

• Definitions: Matter is anything that has mass and occupies space. It can be classified as a pure substance (element or compound) or a mixture (homogeneous or heterogeneous).

• Physical State: Solid, liquid, gas.

• Physical vs. Chemical Properties: Physical properties can be observed without changing the substance; chemical properties describe how a substance reacts.

• Separation of Mixtures: Techniques include decanting, distillation, and filtration.

1.4 Physical and Chemical Changes and Properties

• Physical Change: Alters appearance, not composition (e.g., melting ice).

• Chemical Change: Alters composition (e.g., rusting iron).

1.5 Energy: Fundamental Part of Physical and Chemical Change

• Energy: The capacity to do work. Kinetic energy is energy of motion; potential energy is stored energy.

• SI Unit: Joule (J). 1 calorie = 4.184 J; 1 Calorie (food) = 1000 calories.

• Law of Conservation of Energy: Energy cannot be created or destroyed.

• Kinetic Energy Equation:

$KE = \frac{1}{2}mv^2$

1.6 The Units of Measurement

• SI Units: Standardized units for scientific measurement. Know the seven fundamental SI units (see Table 1.1).

• Common SI Units: meter (m), kilogram (kg), second (s), mole (mol), ampere (A), kelvin (K), candela (cd).

1.7 Reliability of Measurement

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Precision vs. Accuracy: Precision is reproducibility; accuracy is closeness to the true value.

• Density: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

1.8 Solving Chemical Problems

• Dimensional Analysis: Using conversion factors to solve problems.

• Scientific Notation: Expressing numbers as $a \times 10^n$.

Ch. 2: Atoms and the Atomic Theory

2.1 Brownian Motion

• Definition: Random movement of particles suspended in a fluid, evidence for the existence of atoms and molecules.

2.2 Early Ideas about the Building Blocks of Matter

• Historical Figures: Democritus, Leucippus, alchemists, Bacon, Kepler, Galileo, Boyle, Newton.

2.3 Modern Atomic Theory and Laws that Led to It

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

• Law of Definite Proportion: 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 that combine with a fixed mass of the first element are ratios of small whole numbers.

• Dalton's Atomic Theory: All matter is composed of atoms; atoms of a given element are identical; atoms combine in simple ratios to form compounds; atoms are rearranged in chemical reactions.

2.4 Discovery of the Electron

• Cathode Ray Tube Experiment: J.J. Thomson discovered the electron and its charge-to-mass ratio.

• Millikan Oil Drop Experiment: Determined the charge of the electron.

2.5 Structure of the Atom

• Gold Foil Experiment: Rutherford discovered the nucleus and proposed the nuclear model of the atom.

• Subatomic Particles: Proton, neutron, electron.

2.6 Atomic Number, Mass Number, and Isotopes

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

• Mass Number (A): Number of protons plus neutrons.

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

• Ion: Atom or molecule with a net electric charge due to loss or gain of electrons.

2.7 Periodic Law and the Periodic Table

• Mendeleev: Created the modern periodic table, predicted properties of undiscovered elements.

• Groups and Periods: Columns are groups (families); rows are periods.

• Main Groups: Alkali metals, alkaline earth metals, halogens, noble gases.

2.8 Atomic Mass

• Atomic Mass Unit (amu): Standard unit for atomic mass.

• Average Atomic Mass: Weighted average of all isotopes of an element.

• Percent Abundance: Proportion of each isotope in a natural sample.

• Calculating Average Atomic Mass:

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

2.9 Avogadro's Number and the Mole

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

• Mole: SI unit for amount of substance; links mass, number of particles, and volume (for gases).

Ch. 3: Molecules, Compounds, and Chemical Nomenclature

3.1 Hydrogen, Oxygen, and Water

• Compound: Substance composed of two or more elements chemically combined in fixed proportions.

• Mixture vs. Compound: Mixtures are physical blends; compounds have chemical bonds.

3.2 Chemical Bonds

• Covalent Bond: Sharing of electron pairs between atoms.

• Ionic Bond: Transfer of electrons from one atom to another, forming ions.

3.3 Representing Compounds: Formulas and Models

• Structural Formula: Shows how atoms are bonded.

• Molecular Formula: Shows the number and type of atoms.

• Empirical Formula: Simplest whole-number ratio of atoms.

• Space-Filling Model: 3D representation of molecule.

3.4 Atomic Level View of Elements and Compounds

• Classification: Elements can be atomic or molecular; compounds can be molecular or ionic.

• Types of Elements: Metals, nonmetals, metalloids.

• Common Diatomic Elements: H2, O2, N2, F2, Cl2, Br2, I2.

3.5 Ionic Compounds: Formulas and Names

• Type I Binary Ionic Compounds: Metal forms only one type of cation (e.g., NaCl).

• Type II Binary Ionic Compounds: Metal forms more than one type of cation (e.g., FeCl2, FeCl3).

• Polyatomic Ions: Ions composed of more than one atom (e.g., NO3-, SO42-).

• Naming: Cation first, then anion; use Roman numerals for Type II metals.

3.6 Molecular Compounds: Formulas and Names

• Binary Molecular Compounds: Two nonmetals; use prefixes (mono-, di-, tri-, etc.).

• Common Names: H2O (water), NH3 (ammonia), NO (nitric oxide), PH3 (phosphine).

3.7 Summary of Inorganic Nomenclature Rules

• Acids: Naming depends on the anion; -ide becomes hydro-...-ic acid, -ate becomes ...-ic acid, -ite becomes ...-ous acid.

3.8 Formula Mass and the Mole Concept of Compounds

• Formula Mass: Sum of atomic masses in a formula unit.

• Mole Concept: 1 mole = 6.022 x 1023 units.

• Percent Composition:

$\text{Percent Composition} = \frac{\text{mass of element in 1 mole of substance}}{\text{molar mass of substance}} \times 100\%$

3.9 Composition of a Compound

• Empirical Formula: Simplest ratio of elements.

• Molecular Formula: Actual number of atoms of each element.

• Using Percent Composition: Can determine empirical and molecular formulas from percent composition data.

Tables

SI Base Quantity	Unit Name	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Amount of substance	mole	mol
Temperature	kelvin	K
Electric current	ampere	A
Luminous intensity	candela	cd

Additional Info

• Be familiar with the use of non-programmable calculators for exams.

• Practice using the periodic table and SI units for problem-solving.

• Review all lecture notes and textbook chapters as all material is fair game for the exam.