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

Chapter 1 – Matter, Energy, and Measurements

Significant Figures and Arithmetic

Significant figures are the digits in a measurement that are known with certainty plus one digit that is estimated. Correct use of significant figures is essential for reporting scientific data accurately.

• Definition: Significant figures reflect the precision of a measured or calculated quantity.

• Rules: When performing arithmetic:

  • Addition/Subtraction: The result should have the same number of decimal places as the measurement with the fewest decimal places.

  • Multiplication/Division: The result should have the same number of significant figures as the measurement with the fewest significant figures.

• Example: $2.34 \times 1.2 = 2.8$ (rounded to two significant figures)

Precision vs. Accuracy

Understanding the difference between precision and accuracy is crucial in scientific measurements.

• Precision: The closeness of repeated measurements to each other.

• Accuracy: The closeness of a measurement to the true or accepted value.

• Example: If a scale consistently reads 1.2 g for a 1.0 g standard, it is precise but not accurate.

Extensive vs. Intensive Properties of Matter

Properties of matter can be classified based on their dependence on the amount of substance present.

• 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).

• Example: Density is intensive; mass is extensive.

Dimensional Analysis

Dimensional analysis is a method for converting units using conversion factors.

• Conversion Factors: Ratios that express how many of one unit are equal to another unit.

• Process: Multiply the original value by conversion factors so that units cancel appropriately.

• Significant Figures: The answer should reflect the correct number of significant figures based on the given data.

• Example: Convert 2.50 inches to centimeters: $2.50 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 6.35 \text{ cm}$

Chapter 2 – Atoms, Molecules, and Ions

Classification of Elements: Metals vs. Non-metals

The periodic table classifies elements as metals, non-metals, or metalloids based on their properties.

• Metals: Good conductors, malleable, ductile, shiny.

• Non-metals: Poor conductors, brittle, dull.

• Metalloids: Exhibit properties intermediate between metals and non-metals.

• Example: Sodium (Na) is a metal; Oxygen (O) is a non-metal; Silicon (Si) is a metalloid.

Chemical Symbols and Ions

Chemical symbols represent elements and their ions. Understanding how to interpret these symbols is fundamental in chemistry.

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

• Ions: Atoms or molecules with a net charge due to loss or gain of electrons.

• Predicting Charge:

  • Cation: Fewer electrons than protons (positive charge).

  • Anion: More electrons than protons (negative charge).

• Example: $\text{Na}^+$ has 11 protons and 10 electrons.

Periodic Properties

The periodic table (PT) is organized to reflect recurring trends in element properties.

• Organization: Elements are arranged by increasing atomic number.

• Regions:

  • Metals (left and center), non-metals (right), metalloids (stair-step line).

  • Group names: Alkali metals (Group 1), Alkaline earth metals (Group 2), Chalcogens (Group 16), Halogens (Group 17), Noble gases (Group 18).

• Group Properties: Elements in the same group have similar chemical properties and predictable ionic charges.

• Predicting Ionic Charge: Main group elements tend to form ions that achieve noble gas electron configurations.

  • Group 1: $+1$ charge

  • Group 2: $+2$ charge

  • Group 16: $-2$ charge

  • Group 17: $-1$ charge

Naming Molecular and Ionic Compounds

Chemical nomenclature is the system for naming compounds and writing their formulas.

• Molecular Compounds: Composed of non-metals; use prefixes (mono-, di-, tri-, etc.) to indicate number of atoms.

• Ionic Compounds: Composed of cations and anions; name cation first, then anion (with -ide ending if monatomic).

• Common Cations: Sodium ion ($\text{Na}^+$), Calcium ion ($\text{Ca}^{2+}$), Ammonium ion ($\text{NH}_4^+$)

• Common Anions: Chloride ($\text{Cl}^-$), Sulfate ($\text{SO}_4^{2-}$), Nitrate ($\text{NO}_3^-$)

• Example: $\text{NaCl}$ is sodium chloride; $\text{CO}_2$ is carbon dioxide.

Average Atomic Mass

The average atomic mass of an element is the weighted average of the masses of its isotopes.

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

• Example: If $^{35}\text{Cl}$ (75.78%, 34.969 u) and $^{37}\text{Cl}$ (24.22%, 36.966 u): $\text{Average atomic mass} = (0.7578 \times 34.969) + (0.2422 \times 36.966) = 35.45$ u

Properties of Subatomic Particles

Atoms are composed of protons, neutrons, and electrons, each with distinct properties.

Particle	Charge	Mass (amu)
Proton	+1	1.007
Neutron	0	1.009
Electron	-1	0.0005

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

• Example: $^{12}\text{C}$ and $^{13}\text{C}$ are isotopes of carbon.

Empirical vs. Molecular Formulas

The empirical formula gives the simplest whole-number ratio of atoms in a compound, while the molecular formula gives the actual number of atoms of each element in a molecule.

• Empirical Formula: Simplest ratio (e.g., $\text{CH}_2\text{O}$ for glucose).

• Molecular Formula: Actual composition (e.g., $\text{C}_6\text{H}_{12}\text{O}_6$ for glucose).

• Reducing to Empirical Formula: Divide subscripts by their greatest common factor.

• Example: $\text{C}_2\text{H}_4$ reduces to $\text{CH}_2$.

Empirical Formulas and Ionic Charges

Empirical formulas are also used to represent the simplest ratio of ions in an ionic compound.

• Predicting Ionic Charges from Formula: Use the known charges of ions to determine the formula that results in a neutral compound.

• Predicting Formula from Ionic Charges: Balance the total positive and negative charges.

• Example: For $\text{Mg}^{2+}$ and $\text{Cl}^-$, the formula is $\text{MgCl}_2$.