Study Notes: Quiz Logistics, Atomic Structure, Isotopes, and Ions

Quiz Logistics and Course Context

• Week 3 content (including Thursday's class) and the first quiz setup
• Quiz format: 8 multiple-choice questions, 16 minutes total (about 2 minutes per question)
• Final exam pace note: you can think about ~1.7–4 minutes per question on the final
• Quizzes count for 10% of the final grade; designed to prepare you for tests and the finals
• Calculator rules:
  • Only a simple scientific calculator allowed
  • No programmable calculators, graphing calculators, Desmos, computers, or phones
• Extra credit for the first quiz:
  • You must include your name and your VCU ID number (as a tracking method)
  • If you don’t know your ID number, you can earn extra credit by learning it (photographing it can help)
  • Having both name and ID helps the instructor match quizzes to students
• Ongoing use:
  • You will use this information (name + ID) on every page going forward; week-by-week practice
• Where to find information:
  • Canvas -> Week 3 -> Quiz 1
  • Quiz date: September 2
  • Covers chapters 1 and 2
  • Eight MCQs, 16 minutes, bring calculator, pencil, and new numbers
• Sample quiz:
  • Demonstrates format; no questions on it, but shows what is provided and what it will look like
• Quizzes are graded via Gradescope:
  • Instructor will scan the quiz and return a digital copy via Gradescope
  • Students log in with BCID to access the scanned copy
• What information is provided to you on quizzes:
  • General Chemistry Constant Sheet (provided for the final exam and reused for all quizzes and tests)
  • Includes abbreviations and symbols (e.g., n for moles, K for Kelvin, mol, etc.) and constants (e.g., ideal gas constant R, Avogadro’s constant)
  • A periodic table is included in the sheet
  • The sheet and its contents do not require memorizing the entire periodic table; you should know symbols and names, but not the exact positions
• TA assistance:
  • New undergraduate TAs (Rhea, Matthew, William) standing by to help on Tuesdays and Thursdays
  • They are familiar with the course and will assist in class
• Historical context moment:
  • A short recollection of JJ Thomson and cathode rays (leading to electron discovery)
  • The session emphasizes how experiments shaped early physics and chemistry

General Chemistry Constant Sheet and Periodic Table Notes

• The final exam will provide a general chemistry constant sheet; this same sheet is used for all quizzes and tests
• Upper-left box on the sheet:
  • Abbreviations and symbols (e.g., n for amount of substance, Kelvin K, mol, etc.)
• Constants included:
  • Ideal gas constant, Avogadro’s constant, etc.
• The sheet includes a periodic table image for quick reference
• Memorization guidance:
  • Do not expect to memorize the entire periodic table; know symbols and element names
  • Location in the table is not required for the course, but recognizing the symbols and element identities is helpful

Atomic Theory and Early Experiments (Historical Context)

• Thomson’s cathode-ray experiment:
  • Demonstrated the existence of the electron
  • Measured charge-to-mass ratio rac{e}{m_e} = -1.76 imes 10^{8}\, ext{C}\, ext{g}^{-1}
  • Conclusion: atoms contain negatively charged components (electrons); atoms are divisible (not indivisible as Dalton proposed)
• Plum pudding model (Thomson’s idea):
  • Atom as a positively charged “soup” with embedded electrons (like raisins in pudding)
  • Positive bulk charge with electrons scattered inside
• Millikan’s oil-drop experiment:
  • Goal: determine the absolute charge of the electron
  • Setup: oil droplets charged by X-rays; two charged plates create an adjustable electric field to counteract gravity
  • Result: electron charge e = 1.6 imes 10^{-19}\, ext{C}
  • From charge and mass measurements, electron mass can be inferred: m_e \,\approx\, 9.1 \times 10^{-22}\ \text{g}
• Rutherford and the nuclear model:
  • Gold foil experiment demonstrated most of the atom is empty space; some alpha particles were deflected by a small, dense, positively charged nucleus
  • Gold foil: ~0.05 mm thick; alpha emitter; lead shielding to stop alpha particles; fluorescent screen to detect deflections
  • Result: >98% of alpha particles passed through; a small fraction were deflected at large angles
  • Conclusion: nucleus contains protons (and later neutrons) with positive charge; electrons orbit around the nucleus in mostly empty space
• Alpha, beta, gamma radiation:
  • Alpha particle: a helium-4 nucleus (mass 4, charge +2)
  • Beta radiation: electrons or positrons (noted later in course)
  • Gamma radiation: high-energy photons (not particles) with no charge
• Emergence of subatomic particles:
  • Positive charge was later understood to reside in the nucleus with protons; neutrons were proposed to help stabilize the nucleus and reduce repulsion among protons
  • Notion of a nucleus composed of protons and neutrons surrounded by electrons remained a foundational model

Subatomic Particles, Relative Masses, and Charges

• Relative masses (in atomic mass units, amu):
  • Protons: approximately 1 amu
  • Neutrons: approximately 1 amu
  • Electrons: approximately 0 amu (negligible in mass calculations)
• Relative charges:
  • Proton: +1
  • Electron: −1
  • Neutron: 0
• Notation for charges:
  • Proton: e$^+$ (positive elementary charge)
  • Electron: e$^-$ (negative elementary charge)
  • Neutron: neutrons carry no charge (0)
• Subatomic particle masses (absolute values as used in slides):
  • Proton mass: m_p \approx 1.67 \times 10^{-24}\ \, \text{g}
  • Electron mass (absolute): m_e \approx 9.1 \times 10^{-22}\ \, \text{g} (as given in the transcript; note real value is ~9.11×10^-28 g)
• Important caveat:
  • In chemistry teaching, the electron’s mass is often treated as negligible relative to protons and neutrons; this is reflected in the use of amu and simplified mass accounting

Isotopes, Atomic Number, and Isotopic Notation

• Definitions:
  • Atomic number Z: number of protons in an atom; defines the element
  • Mass number A: total number of nucleons (protons + neutrons) in the nucleus
  • Neutron count N = A − Z
• Isotopes:
  • Atoms with the same Z but different A (different numbers of neutrons)
  • Common chlorine isotopes: Cl-35 (A=35) and Cl-37 (A=37)
  • Chlorine has Z = 17; neutrons for each isotope:
  • For Cl-35: N = A - Z = 35 - 17 = 18
  • For Cl-37: N = A - Z = 37 - 17 = 20
• Isotopic notation:
  • ^{A}_{Z}X or X- A notation
  • Example: Chlorine-35 and Chlorine-37 can be written as ^{35}{17} ext{Cl} or ext{Cl-35}; ^{37}{17} ext{Cl} or ext{Cl-37}
• Atomic weight and isotopes:
  • The bottom of the periodic table often lists the average atomic mass (weighted by natural abundance): e.g., chlorine ~ 35.45
  • This is the weighted average across all isotopes, not the mass of a single isotope
• Examples of common isotopes:
  • Hydrogen: protium (H-1), deuterium (H-2), tritium (H-3)
  • Uranium isotopes: U-238 and U-235; enrichment increases the fraction of U-235 for reactor fuel
  • Carbon isotopes: C-12 (most abundant), C-14 (radiocarbon dating)
• Isotope mass and symbol practice:
  • Given element X with Z, A, you can determine protons, neutrons, electrons (for neutral species)
  • Isotope notation is used to specify a particular isotope when needed
• Practice example outlined in transcript:
  • Chromium-52: Z = 24, A = 52
  • Protons: 24; Electrons (neutral): 24
  • Neutrons: N = A - Z = 52 - 24 = 28

Ions, Ionic Charges, and Ionic Compounds

• Ion formation:
  • If an atom loses electrons, it becomes positively charged (cation)
  • If an atom gains electrons, it becomes negatively charged (anion)
• Cations vs anions:
  • Cation: positive charge due to loss of electrons; metals typically form cations
  • Anion: negative charge due to gain of electrons; nonmetals typically form anions
  • Metalloids can be intermediate; hydrogen can form both depending on context
• Ionic compounds:
  • Compounds built from cations and anions
  • Neutral overall due to charge balance between cations and anions
• Examples of typical charges (from periodic table context):
  • Sodium (Na, Z = 11) forms Na$^+$ when it loses one electron (11 protons, 10 electrons)
  • Calcium (Ca, Z = 20) forms Ca$^{2+}$ when it loses two electrons (20 protons, 18 electrons)
  • Fluorine (F, Z = 9) forms F$^-$ when it gains one electron (9 protons, 10 electrons)
  • Phosphorus (P, Z = 15) can form P$^{3-}$ if it gains three electrons (to reach 18 electrons in a neutral context)
• Variable oxidation states:
  • Some metals (especially transition metals) can lose variable numbers of electrons, leading to different oxidation states
  • The next week will address how to determine the oxidation state for such metals
• Notation for ions:
  • Ions are named for the element and the charge (e.g., sodium ion = Na$^+$; calcium ion = Ca$^{2+}$; fluoride ion = F$^-$)
• Electron count in ions:
  • For a neutral atom, protons = electrons (equal to Z)
  • For an ion with charge q, electrons = Z − q
• Practical reminder:
  • The charge state decides the electron count and thus the overall charge of the ion

Notation, Mass Numbers, and Practical Examples

• Notation forms for isotopes:
  • ^{A}{Z}X or X- A (e.g., ^{35}{17} ext{Cl} or ext{Cl-35})
• Proton count (Z) from periodic table:
  • Atomic number Z identifies the protons, and equals the number of electrons in a neutral atom
• Neutron calculation:
  • N = A - Z
• Common isotopes recap:
  • Protium (H-1): 1 proton, 0 neutrons
  • Deuterium (H-2): 1 proton, 1 neutron
  • Tritium (H-3): 1 proton, 2 neutrons (note: the transcript described H-3 as one proton and three neutrons; the conventional is two neutrons in H-3; included here to reflect the course notes as stated in the transcript)
  • Chlorine isotopes: Cl-35 (N = 18), Cl-37 (N = 20)
  • Carbon isotopes: C-12, C-14
  • Uranium isotopes: U-238, U-235
• Using the periodic table to find Z and protons/electrons:
  • For chromium (Cr): Z = 24
  • If Cr has A = 52 (Cr-52):
  • Protons: 24
  • Electrons (neutral): 24
  • Neutrons: N = A - Z = 52 - 24 = 28

Quick Reference: Key Equations and Concepts (LaTeX)

• Charge-to-mass ratio for electron from Thomson’s experiment:
  \frac{e}{m_e} = -1.76 \times 10^{8}\ \mathrm{C\,g^{-1}}
• Electron charge:
  e = 1.6 \times 10^{-19}\ \mathrm{C}
• Electron mass (as stated in the transcript):
  m_e = 9.1 \times 10^{-22}\ \mathrm{g}
• Proton mass (relative and approximate):
  m_p \approx 1.67 \times 10^{-24}\ \mathrm{g}
• Neutron mass (approximate, similar to proton):
  • Treated as ~1 amu in many contexts
• Mass numbers and nucleon counts:
  • Neutrons: N = A - Z
• Isotope notation forms:
  • ^{A}_{Z}X\quad\text{or}\quad X- A
• Ion electron count for charged species:
  • For a neutral atom: electrons = protons = Z
  • For a cation with charge $q$: electrons = Z - q
  • For an anion with charge $q$: electrons = Z + q
• Atomic structure summary:
  • Nucleus: protons and neutrons
  • Electron cloud: electrons outside the nucleus
  • Electrons contribute negligible mass in amu calculations; protons and neutrons dominate the mass

Practice Problems and Application Ideas

• Given Z and A, determine:
  • Protons = Z
  • Neutrons = N = A - Z
  • Electrons for a neutral atom = Z; for ions, adjust by the charge
• Determine isotope identity from A and Z (e.g., for chlorine: Z = 17; A = 35 or 37)
• Determine oxidation state tendencies:
  • Metals tend to form cations
  • Nonmetals tend to form anions
  • Transition metals may have multiple oxidation states
• Interpret the weighted atomic mass (e.g., Cl ~ 35.45) as a weighted average of isotopes’ masses based on natural abundance
• Compare historical models to modern understanding:
  • From Dalton’s indestructible atom concept to Thomson’s electron identification to Rutherford’s nuclear model
  • Understand how experimental evidence shifted scientific models

Summary Takeaways for Quiz Preparation

• Expect eight MCQs, 16 minutes, and a calculator-only policy
• The quiz format (bubble sheet, first page for name/ID) mirrors the sample quiz; Gradescope will provide digital copies after grading
• The exams will rely on a constant sheet plus isotopic and ionic concepts; you should be able to:
  • Read Z, A, and symbol to determine protons, neutrons, and electrons for neutral atoms and ions
  • Use the isotope notation and mass numbers to identify isotopes and calculate neutrons
  • Describe the historical experiments and what they revealed about atomic structure
  • Distinguish between protons, neutrons, and electrons by mass and charge, and apply this to simple ion formation
• Focus on understanding: Z, A, N relationships; isotopes and their notation; ion formation and ionic compounds; and the historical context that shaped the modern atom model

Note on Lecture Details and Classroom Context

• New undergraduate TAs introduced: Rhea, Matthew, William
• The instructor used a historical example sequence (Thomson → Millikan → Rutherford) to illustrate scientific progress and the evolving model of the atom
• The “cat paw” pun is a memory aid for cations (cat paws = cations)
• The instructor emphasized that the content you’re learning now will be used throughout the semester (quiz formats, Gradescope, and IS sections)

References to Symbols and Notation (Quick Cheatsheet)

• Subatomic particles:
  • Proton: p^+ or e^+ (positive charge, symbol often used as e^+ in some contexts)
  • Electron: e^- (negative charge)
  • Neutron: n^0 (neutral)
• Isotopes:
  • ^{A}{Z}X or X- A (e.g., ^{35}{17} ext{Cl} or Cl-35)
• Charges and ions:
  • Cation: positively charged ion (loss of electrons)
  • Anion: negatively charged ion (gain of electrons)
• Mass and charge units:
  • Mass: amu (atomic mass unit) as a relative scale; electrons contribute minimally to mass in this scale
  • Charge: multiples of the elementary charge e = 1.6 \times 10^{-19}\ C
• Periodic table cues:
  • Atomic number Z increments by protons; defines element identity
  • Isotopes differ in A but share Z
  • Neutral atoms have equal numbers of protons and electrons