Chem 101 Exam 1 Study Guide: Chapters 1–5.3 (Atoms, Matter, Measurement, Nuclear Chemistry, and Ionic Compounds)

Chapter 1: Introduction to Chemistry and Matter

Classification of Matter

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

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

• Elements: Pure substances consisting of only one type of atom (e.g., O2, Fe).

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

• Compounds: Substances composed of two or more elements chemically combined in fixed ratios (e.g., H2O).

Structure of the Atom

• 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 ≈ 0.0005 amu; Located in electron cloud around nucleus.

States of Matter

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

• Liquids: Definite volume, no definite shape; particles less tightly packed; can flow.

• Gases: No definite shape or volume; particles far apart; move freely.

Physical and Chemical Changes

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

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

Elements and the Periodic Table

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

• Groups (Columns): Elements with similar properties; numbered 1A–8A.

• Periods (Rows): Horizontal rows; elements 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 table; conduct electricity, malleable, shiny.

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

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 and Significant Figures

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

• Significant Figures (Sig Figs): Indicate precision of a measurement.

• Rules for Sig Figs:

  • Nonzero digits are always significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

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

• Calculations:

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

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

Dimensional Analysis and Conversions

• Use conversion factors to change units (e.g., $1\ \text{in} = 2.54\ \text{cm}$).

• Set up problems so units cancel appropriately.

Volume, Mass, and Density

• Volume: Amount of space occupied; units: L, mL, cm3.

• Mass: Amount of matter; units: g, kg.

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

Temperature Conversions

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

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

Chapter 3: Atoms: The Building Blocks of Chemistry

Atomic Structure and Notation

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

• Mass Number (A): $A = \text{Number of protons} + \text{Number of neutrons}$

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

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

Electrons and Electron Configuration

• Number of Electrons: In a neutral atom, equals the atomic number.

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

• Box (Orbital) Diagrams: Visual representation of electron arrangement using boxes and arrows.

• Electron Configuration: Notation showing distribution of electrons (e.g., 1s2 2s2 2p6).

• Abbreviated Configuration: Uses noble gas core (e.g., [Ne] 3s2 3p4).

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

• Electron Dot Structures: Dots around element symbol represent valence electrons.

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 atomic number by 2, mass number by 4.

• Beta Decay (\beta^-): Emission of an electron; neutron converts to proton; atomic number increases by 1.

• Positron Emission (\beta^+): Emission of a positron; proton converts to neutron; atomic number decreases by 1.

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

Balancing Nuclear Equations

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

Radiation and Measurement

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

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

• Half-Life (t1/2): Time for half of a radioactive sample to decay.

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

Medical Applications

• Most Used Radioisotope: Technetium-99m in medicine.

• PET Scans: Use Fluorine-18.

Nuclear Fission and Fusion

• Fission: Splitting of heavy nuclei (e.g., U-235); initiated by neutron absorption; produces chain reactions; used in nuclear reactors.

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

Chapter 5 (through 5.3): Ionic Compounds

Valence Electrons and Bonding

• Valence Electrons: Outermost electrons; determine reactivity and bonding.

• Bonding: Atoms gain, lose, or share electrons to achieve stable electron configurations.

• Types of Bonds:

  • Ionic Bonds: Formed by transfer of electrons from metals to nonmetals.

  • Covalent Bonds: Formed by sharing electrons (not covered in detail here).

Ions and the Octet Rule

• Octet Rule: Atoms tend to gain, lose, or share electrons to have 8 valence electrons.

• Metals: Lose electrons to form cations (positive ions).

• Nonmetals: Gain electrons to form anions (negative ions).

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

Formulas and Charges of Ionic Compounds

• Ionic Compounds: Electrically neutral; total positive charge equals total negative charge.

• Determining Formulas: Balance charges to write correct formula (e.g., Na+ and Cl– form NaCl).

• Polyvalent Elements: Some metals (e.g., Fe, Cu) can have multiple charges; specified with Roman numerals (e.g., Fe(II), Fe(III)).

• Determining Ion Charge from Formula: Example: In CuCl2, Cl is –1, so Cu must be +2.

Subatomic Particle	Symbol	Charge	Relative Mass (amu)	Location
Proton	p+	+1	1	Nucleus
Neutron	n0	0	1	Nucleus
Electron	e–	–1	0.0005	Electron cloud

*Additional info: Academic context and examples have been added to expand on the brief points in the original study guide. Table of subatomic particles included for clarity. Equations are provided in LaTeX format as required.*