CHEM 1150 Exam 1 Study Guide: GOB Chemistry Fundamentals

Chemistry Basics – Matter and Measurements

Classification of Matter

Understanding the classification of matter is fundamental in chemistry. Matter can be categorized based on its composition and properties.

• Pure Substance: A material with a constant composition (element or compound).

• Mixture: A combination of two or more substances that are not chemically bonded. Mixtures can be homogeneous (uniform composition) or heterogeneous (non-uniform composition).

• Example: Water (H2O) is a compound; air is a homogeneous mixture.

Periodic Table and Atomic Structure

The periodic table organizes elements by increasing atomic number and groups elements with similar properties.

• Element Names & Symbols: Each element has a unique symbol (e.g., Na for sodium).

• Periodic Table Features: Includes periods (rows), groups (columns), and the staircase dividing metals and nonmetals.

• Key Elements: Know the symbols and names for elements 1-38, 47, 57-78, 89, 92.

• Example: Carbon (C), Iron (Fe), Gold (Au).

Chemical and Physical Properties

Chemical properties describe a substance's ability to undergo chemical changes, while physical properties can be observed without changing the substance's identity.

• Chemical Properties: Reactivity, flammability.

• Physical Properties: Melting point, density, color.

Measurement and Significant Figures

Accurate measurement is essential in chemistry. Significant figures reflect the precision of a measurement.

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

• Example: 2.50 has three significant figures.

Unit Conversions and Calculations

Unit conversions are necessary for solving problems in chemistry. Common conversions include length, mass, and temperature.

• Length: 1 inch = 2.54 cm (exact), 2.205 lb = 1 kg, 1 L = 1.057 qt

• Temperature: $^ ext{o}C = rac{5}{9} (^ ext{o}F - 32)$

• Example: Convert 100°F to °C: $^ ext{o}C = rac{5}{9}(100 - 32) = 37.8^ ext{o}C$

Density and Specific Gravity

Density is a physical property defined as mass per unit volume. Specific gravity compares the density of a substance to that of water.

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

• Specific Gravity: $\text{Specific Gravity} = \frac{\text{Density of substance}}{\text{Density of water}}$

• Example: If a liquid has a density of 1.2 g/mL, its specific gravity is 1.2.

Heat Capacity

Heat capacity is the amount of heat required to raise the temperature of a substance by one degree Celsius.

• Formula: $q = m \cdot c \cdot \Delta T$

• Where: q = heat (J), m = mass (g), c = specific heat (J/g·°C), ΔT = change in temperature (°C).

Atoms and Radioactivity

Atomic Structure

Atoms consist of protons, neutrons, and electrons. The atomic number is the number of protons, and the mass number is the sum of protons and neutrons.

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

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Radioactivity and Types of Radiation

Radioactive decay involves the emission of particles or energy from unstable nuclei.

• Types of Radiation: Alpha (α), beta (β), gamma (γ), positron, neutron.

• Effects: Radiation can damage biological tissues.

• Example: Alpha decay: $^{238}_{92}U \rightarrow ^{234}_{90}Th + ^{4}_{2}He$

Balancing Nuclear Equations

Nuclear equations must be balanced for both mass and atomic numbers.

• Example: $^{14}_{6}C \rightarrow ^{14}_{7}N + ^{0}_{-1}e$ (beta decay)

Energy and Bonding in Nuclear Reactions

Nuclear reactions involve large energy changes, including emission or decay, bombardment, fusion, and fission.

• Fusion: Combining light nuclei to form heavier nuclei.

• Fission: Splitting heavy nuclei into lighter nuclei.

• Example: $^{235}_{92}U + ^{1}_{0}n \rightarrow ^{141}_{56}Ba + ^{92}_{36}Kr + 3 ^{1}_{0}n$

Units of Radiation

Radiation is measured in units such as the curie (Ci), becquerel (Bq), rad, rem, and gray (Gy).

• Example: 1 Ci = 3.7 × 1010 disintegrations per second.

Electron Structure and Ions

Valence Electrons and the Octet Rule

Valence electrons are the electrons in the outermost shell of an atom. The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Valence Electrons: Determine chemical reactivity.

• Octet Rule: Atoms achieve stability by having eight electrons in their valence shell.

• Example: Sodium (Na) loses one electron to form Na+.

Ions and Polyatomic Ions

Ions are charged particles formed when atoms gain or lose electrons. Polyatomic ions are groups of atoms with a charge.

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

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

• Polyatomic Ions: Examples include NO3-, SO42-, CO32-, PO43-, CH3CO2-, OH-, CN-, NH4+.

Electron Configuration and Isoelectronic Species

Electron configuration describes the arrangement of electrons in an atom. Isoelectronic species have the same number of electrons.

• Example: Na+ and Ne are isoelectronic (both have 10 electrons).

Determining Electron Gain or Loss

Main group elements gain or lose electrons to achieve a stable octet or to match their charge.

• Example: Cl gains one electron to form Cl-; Mg loses two electrons to form Mg2+.

Summary Table: Key Terms and Definitions

Term	Definition	Example
Element	Pure substance consisting of one type of atom	Oxygen (O)
Compound	Pure substance composed of two or more elements chemically bonded	Water (H2O)
Mixture	Physical blend of two or more substances	Air
Isotope	Atoms of the same element with different numbers of neutrons	Carbon-12, Carbon-14
Cation	Positively charged ion	Na+
Anion	Negatively charged ion	Cl-
Polyatomic Ion	Ion composed of two or more atoms	SO42-
Specific Gravity	Ratio of substance's density to water's density	1.2

Additional info: Academic context and examples have been expanded for clarity and completeness. All equations are provided in LaTeX format as required.