Chem 101 Exam 1 Study Guide: Chapters 1–5 (GOB Chemistry Foundations)

Chapter 1: Introduction to Chemistry

Classification of Matter

• Heterogeneous Mixtures: Mixtures with visibly different components or phases (e.g., salad, oil and water).

• Homogeneous Mixtures: Mixtures with uniform composition throughout (e.g., saltwater, air).

• Element: A pure substance consisting of only one type of atom (e.g., O2).

• Atom: The smallest unit of an element that retains its chemical properties.

• Compound: A substance formed from two or more elements chemically bonded in fixed proportions (e.g., H2O).

Subatomic Particles

• Proton: Charge = +1; Relative mass ≈ 1 amu; Located in the nucleus.

• Neutron: Charge = 0; Relative mass ≈ 1 amu; Located in the nucleus.

• Electron: Charge = –1; Relative mass ≈ 1/1836 amu; Located in electron cloud around nucleus.

States of Matter

• Solid: Definite shape and volume; particles closely packed; vibrate in place.

• Liquid: Definite volume, no definite shape; particles less tightly packed; flow freely.

• Gas: No definite shape or volume; particles far apart; move rapidly and independently.

Physical vs. Chemical Changes

• Physical Change: Change in state or appearance without altering composition (e.g., melting ice, dissolving sugar).

• Chemical Change: Change that forms new substances (e.g., rusting iron, burning wood).

Elements and the Periodic Table

• Know symbols and names for the first 20 elements, plus Fe, Ag, Au, Cu.

• Groups: Vertical columns; elements in a group have similar properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Elements are arranged by increasing atomic number.

• Families:

  • Group 1A: Alkali metals

  • Group 2A: Alkaline earth metals

  • Group 7A: Halogens

  • Group 8A: Noble gases

• Metals: Left and center of the table; good conductors, malleable.

• Non-metals: Right side; poor conductors, brittle.

Chapter 2: Measurement in Science and Medicine

Units and Prefixes

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

• Metric Prefixes:

  • kilo- (k): $10^3$

  • deca- (da): $10^1$

  • centi- (c): $10^{-2}$

  • milli- (m): $10^{-3}$

  • micro- (\mu): $10^{-6}$

  • nano- (n): $10^{-9}$

Scientific Notation

• Expresses numbers as $a \times 10^n$ (e.g., $3.2 \times 10^4$).

• Convert by moving the decimal point and adjusting the exponent accordingly.

Significant Figures

• Indicate the precision of a measurement.

• Rules for counting significant figures:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

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

• In calculations:

  • Multiplication/Division: Answer has the same number of sig figs as the least precise value.

  • Addition/Subtraction: Answer has the same number of decimal places as the least precise value.

Dimensional Analysis

• Method for converting units using conversion factors.

• Set up conversion so units cancel appropriately.

Volume, Mass, and Density

• Volume: Space occupied by a substance (units: L, mL, cm3).

• Mass: Amount of matter (units: g, kg).

• Density: Mass per unit volume.

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

• Can rearrange to solve for mass or volume.

Temperature Conversions

• Fahrenheit to Celsius: $^\circ C = \frac{5}{9} ( ^\circ F - 32 )$

• Celsius to Fahrenheit: $^\circ F = \frac{9}{5} ^\circ C + 32$

Chapter 3: Atoms: The Building Blocks of Chemistry

Atomic Structure

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

• Mass Number (A): Total number of protons and neutrons.

  • Formula: $A = Z + N$ (where N = number of neutrons)

• Atomic Symbol: Written as $^{A}_{Z}\text{X}$ (e.g., $^{12}_{6}\text{C}$ for carbon-12).

• Electrons in a neutral atom = atomic number.

Isotopes

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

• Isotopes have the same atomic number but different mass numbers.

• Example: $^{12}\text{C}$, $^{13}\text{C}$, $^{14}\text{C}$ are isotopes of carbon.

Electron Structure

• Electrons occupy energy levels (shells), sublevels (s, p, d, f), and orbitals.

• Box (orbital) diagrams show electron arrangement using arrows for electrons.

• Electron configuration: e.g., $1s^2 2s^2 2p^6$.

• Abbreviated configuration uses noble gas core (e.g., [Ne] $3s^2 3p^4$).

• Valence electrons: Electrons in the outermost shell; determine chemical properties.

• Electron dot (Lewis) structures show valence electrons as dots around the symbol.

Periodic Trends

• Atomic Size: Increases down a group, decreases across a period.

• Ionization Energy: Energy required to remove an electron; decreases down a group, increases across a period.

Chapter 4: Nuclear Chemistry

Radioactive Decay Processes

• Alpha Decay (\alpha): Emission of an alpha particle ($^4_2\text{He}$); decreases mass number by 4 and atomic number by 2.

• Beta Decay (\beta^-): Emission of an electron; neutron converts to proton.

• Positron Emission (\beta^+): Emission of a positron; proton converts to neutron.

• Gamma Decay (\gamma): Emission of high-energy photons; no change in mass or atomic number.

Balancing Nuclear Equations

• Sum of mass numbers and atomic numbers must be equal on both sides.

• Identify new isotopes formed after decay.

Radiation Energies and Doses

• Relative energies: $\gamma > \beta > \alpha$

• Curie (Ci): Measures radioactivity (disintegrations per second).

• Absorbed Dose: Measured in rads or grays (Gy).

Half-Life

• Time required for half the nuclei in a sample to decay.

• Half-life formula: $N = N_0 \left( \frac{1}{2} \right)^{t/t_{1/2}}$

Medical Applications

• Most used radioisotope: Technetium-99m ($^{99m}\text{Tc}$).

• PET scans use $^{18}\text{F}$ (fluorine-18).

Nuclear Fission and Fusion

• Fission: Splitting of heavy nuclei (e.g., $^{235}\text{U}$); initiated by neutron absorption; produces chain reactions; used in nuclear reactors.

• Fusion: Combining light nuclei (e.g., hydrogen isotopes); powers the sun; requires high temperature and pressure.

Chapter 5: Ionic Compounds (Sections 5.1–5.3)

Valence Electrons and Bonding

• Valence electrons determine chemical reactivity and bonding.

• Atoms gain or lose electrons to achieve a stable octet (8 valence electrons).

Types of Bonds

• Ionic Bonds: Formed by transfer of electrons from metals (lose electrons, become cations) to non-metals (gain electrons, become anions).

• Covalent Bonds: (Covered in later chapters) Formed by sharing electrons.

Formation of Ions

• Metals lose electrons to form positive ions (cations).

• Non-metals gain electrons to form negative ions (anions).

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

Octet Rule

• Atoms tend to gain, lose, or share electrons to have 8 electrons in their valence shell.

Formulas and Charges of Ionic Compounds

• Ionic compounds are electrically neutral; total positive and negative charges balance.

• Polyvalent elements (mainly transition metals) can have multiple possible charges; specified with Roman numerals (e.g., iron(III) chloride: FeCl3).

• To determine the charge of a metal in a compound, use the formula and known charges of other ions (e.g., in CuCl2, Cu is +2 because each Cl is –1).

Element	Common Ion Charge	Type
Na	+1	Metal (cation)
Cl	–1	Non-metal (anion)
Fe	+2, +3	Transition metal (polyvalent)
Cu	+1, +2	Transition metal (polyvalent)
O	–2	Non-metal (anion)