GOB Chemistry Exam 1 Study Guide: Key Concepts and Applications

Significant Figures (Sig Figs)

Rules for Significant Figures

Significant figures are the digits in a measurement that are known with certainty plus one digit that is estimated. They are crucial for expressing the precision of measurements in chemistry.

• Zero Rules: Pay special attention to zeros. Leading zeros are not significant, captive zeros (between nonzero digits) are significant, and trailing zeros are significant only if there is a decimal point.

• Mathematical Functions: When performing calculations, the result should be rounded to the correct number of significant figures based on the operation (addition/subtraction: least decimal places; multiplication/division: least sig figs).

Example: 0.00450 has three significant figures.

Unit Conversions

Temperature and Metric Conversions

Unit conversions are essential for expressing measurements in different units and for solving chemical problems.

• Celsius to Kelvin: $K = ^\circ C + 273.15$

• Metric Prefixes: Convert base units (grams, liters, meters) to kilo (103), milli (10-3), centi (10-2), nano (10-9), etc.

Example: 1.0 g = 1000 mg

Chemical and Physical Properties

Classification of Matter

Matter can be classified based on its composition and properties.

• Pure Element vs Compound: An element contains only one type of atom; a compound contains two or more types of atoms chemically bonded.

• Homogeneous vs Heterogeneous Mixture: Homogeneous mixtures have uniform composition; heterogeneous mixtures have visibly different parts.

• Physical Change vs Chemical Change/Reaction: Physical changes do not alter the chemical identity; chemical changes result in new substances.

Example: Dissolving salt in water is a physical change; burning wood is a chemical change.

Density

Calculating and Applying Density

Density is a physical property that relates mass and volume.

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

• Identifying Substances: Substances can be identified by comparing their densities to known values.

• Specific Gravity: Specific gravity is the ratio of the density of a substance to the density of water.

• Effect of Size/Weight: The size or weight does not affect density or specific gravity, as these are intensive properties.

Example: Water has a density of 1.00 g/mL.

Property	Definition
Density	Mass per unit volume
Specific Gravity	Density of substance / Density of water

Ionic Compounds

Symbols, Naming, and Molar Mass

Ionic compounds are formed from the electrostatic attraction between cations and anions.

• Atomic Symbols: Each element is represented by a unique symbol (e.g., Na for sodium).

• Ionic Symbols: Common ions include Na+, Cl-, Ca2+, etc.

• Writing and Naming: Name the cation first, then the anion (e.g., NaCl is sodium chloride).

• Molar Mass Calculation: Add the atomic masses of all atoms in the formula.

Example: Molar mass of NaCl = 22.99 g/mol (Na) + 35.45 g/mol (Cl) = 58.44 g/mol

Identifying Ionic and Covalent Bonds

Bond Types and Valence Electrons

Chemical bonds can be classified as ionic or covalent based on the nature of the atoms involved.

• Ionic Compounds: Formed between metals and nonmetals; electrons are transferred.

• Covalent Compounds: Formed between nonmetals; electrons are shared.

• Valence Electrons: The number of electrons in the outermost shell; determines chemical reactivity.

Example: H2O is covalent; NaCl is ionic.

Nuclear Decay

Particles and Balancing Equations

Nuclear decay involves the transformation of unstable nuclei into more stable forms.

• Beta Particle: $\beta$ or $e^{-}$ (electron emission)

• Alpha Particle: $\alpha$ or $^4_2He$

• Positron: $e^{+}$

• Balancing Nuclear Equations: Ensure the sum of atomic and mass numbers is equal on both sides.

Example: $^{14}_6C \rightarrow ^{14}_7N + e^{-}$

Covalent Bonding

Lewis Structures, Molecular Shape, and Polarity

Covalent bonding involves the sharing of electrons between atoms.

• Lewis Structures: Diagrams showing the arrangement of electrons in a molecule.

• Molecular Shape: Determined by the VSEPR theory (e.g., linear, bent, tetrahedral).

• Bond Polarity: Determined by the difference in electronegativity between atoms.

• Molecule Polarity: Depends on both bond polarity and molecular shape.

• Naming Covalent Compounds: Use prefixes (mono-, di-, tri-) to indicate the number of atoms.

• Molar Mass: Sum of atomic masses in the molecular formula.

Example: CO2 is linear and nonpolar; H2O is bent and polar.

Nuclear Chemistry

Subatomic Particles, Dosage, and Half-Life

Nuclear chemistry focuses on the properties and reactions of atomic nuclei.

• Protons, Neutrons, Mass Number: Mass number = number of protons + number of neutrons.

• Dosage Unit Conversions: Convert between units such as becquerel, curie, gray, and sievert.

• Half-Life Calculations: Amount remaining after n half-lives: $\text{Remaining} = \text{Initial} \times (\frac{1}{2})^n$

• Balancing Nuclear Equations: As above, ensure atomic and mass numbers are conserved.

Example: If 100 g of a substance has a half-life of 3 years, after 6 years, 25 g remain.

Molecular Chemistry

Chemical Equations and Mole Calculations

Molecular chemistry involves the study of molecules and their reactions.

• Balancing Chemical Equations: Ensure the same number of each atom on both sides of the equation.

• Mole-Gram Conversions: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

Example: To find moles in 18 g of H2O: $\text{Moles} = \frac{18}{18} = 1$ mol