General Chemistry Exam 1 Study Guide: Matter, Measurements, Atoms, Molecules, and Ions

Exam 1 Study Guide: General Chemistry

Overview

This study guide outlines the key topics and concepts for Exam 1 in a General Chemistry course. The exam covers foundational principles including matter, measurements, atoms, molecules, ions, and their properties. Mastery of these topics is essential for success in introductory chemistry.

Chapter 1 – Matter, Energy, and Measurements

Significant Figures and Arithmetic

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit. Used to reflect the precision of measurements.

• Rules:

  • Nonzero digits are always significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Arithmetic with Significant Figures:

  • For addition/subtraction: result has as many decimal places as the measurement with the fewest decimal places.

  • For multiplication/division: result has as many significant figures as the measurement with the fewest significant figures.

Precision vs. Accuracy

• Precision: How closely repeated measurements agree with each other.

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

• Example: If a scale gives the same mass every time (e.g., 5.00 g, 5.00 g, 5.00 g), it is precise. If the true mass is 5.00 g, it is also accurate.

Extensive vs. Intensive Properties of Matter

• Extensive Properties: Depend on the amount of matter (e.g., mass, volume).

• Intensive Properties: Independent of the amount of matter (e.g., density, boiling point).

Dimensional Analysis

• Definition: A method to convert one unit to another using conversion factors.

• Steps:

  1. Identify the given quantity and units.

  2. Multiply by conversion factors so units cancel appropriately.

  3. Ensure the final answer has the correct units and significant figures.

• Example: 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, Molecules, and Ions

Identifying Metals, Nonmetals, and Metalloids

• Metals: Good conductors of heat and electricity, malleable, ductile, shiny.

• Nonmetals: Poor conductors, brittle, dull.

• Metalloids: Properties intermediate between metals and nonmetals (e.g., silicon, arsenic).

Chemical Symbols for Neutral Atoms and Ions

• Neutral Atom: Number of protons equals number of electrons.

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

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

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

• Example: $\text{Na}^+$ is a sodium cation; $\text{Cl}^-$ is a chloride anion.

Predicting Charge from Protons and Electrons

• Charge Calculation: $\text{Charge} = \text{Number of protons} - \text{Number of electrons}$

• Example: An atom with 11 protons and 10 electrons has a charge of $+1$.

Periodic Properties

• Periodic Table Organization: Elements are arranged by increasing atomic number.

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows.

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

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

• Properties of Different PT Regions:

  • Metals: Left and center of the table.

  • Nonmetals: Upper right.

  • Metalloids: Border between metals and nonmetals.

• Group Properties: Elements in the same group often form ions with the same charge (e.g., Group 1 forms $+1$ cations).

Molecular and Ionic Compounds

• Molecular Compounds: Composed of nonmetals; atoms share electrons (covalent bonds).

• Ionic Compounds: Composed of metals and nonmetals; transfer of electrons forms ions held by electrostatic attraction.

• Formulas:

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

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

• Naming Compounds:

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

  • Molecular Compounds: Use prefixes to indicate number of atoms (e.g., carbon dioxide).

• Formulas from Names and Vice Versa: Use rules for naming and formula writing to interconvert.

Calculating Average Atomic Mass

• Definition: Weighted average of the masses of all naturally occurring isotopes of an element.

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

• Example: If an element has two isotopes: 75% at 10 amu and 25% at 11 amu, average mass = $0.75 \times 10 + 0.25 \times 11 = 10.25$ amu.

Properties of Subatomic Particles

• Proton: Positive charge, mass ≈ 1 amu.

• Neutron: No charge, mass ≈ 1 amu.

• Electron: Negative charge, mass ≈ 0.0005 amu (much lighter than protons/neutrons).

Empirical vs. Molecular Formula

• Empirical Formula: Simplest ratio of elements in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule; may be a multiple of the empirical formula.

• Example: Glucose has molecular formula $C_6H_{12}O_6$ and empirical formula $CH_2O$.

• Reducing Molecular Formula: Divide subscripts by their greatest common factor to obtain the empirical formula.

Table: Classification of Elements

Type	Location on Periodic Table	Properties	Examples
Metals	Left and center	Conductive, malleable, ductile, shiny	Na, Fe, Cu
Nonmetals	Upper right	Poor conductors, brittle, dull	O, N, Cl
Metalloids	Border between metals and nonmetals	Intermediate properties	Si, As

Additional info: Some details, such as specific group names and example elements, have been inferred for completeness and clarity.