CHMY141 Lecture Notes: General Chemistry Concepts (Vocabulary Flashcards)
Course Overview and Resources
Fall 2025 General Chemistry – CHMY141 overview and contact points.
Instructors: Dr. Candace Goodman (Office Gaines 215; Office Hours WF 2–4 pm, Tu/Th 9–10:30 am or by appointment; candace.goodman@montana.edu) and Dr. Joan Broderick (Office CBB 219; Office Hours WF 9:30–11:30 am or by appointment; jbroderick@montana.edu).
Online resources provide primary class information and materials:
Canvas / eCat portal: Announcements, Course Handouts (syllabus, schedule), Link to Homework, Textbook, Recommended Reading & Practice Problems, Lecture Slides/Notes, Video Tutorials, Exam Objectives, Discussion Board.
Course materials needed:
Textbook: OER (via Canvas link)
ALEKS access (link in Canvas)
Lab Manual (link in Canvas)
Lab Notebook (Provided)
Calculator (scientific or graphing)
Scratch paper and writing implements for in-class problems
Important course components:
In-class quizzes (Canvas)
ALEKS homework
Unit Assessments (4) and a Final Exam; one dropped, no makeup exams
Unit 1: Sept 11–12; Unit 2: Oct 9–10; Unit 3: Oct 30–31; Unit 4: Dec 4–5
Grading overview (total 1000 pts):
Hour Exams & Final: 160 pts each → 640 pts
In-Class Questions: 15 pts each → 150 pts
Homework: 15 or 26 pts each → 210 pts
Lab (CHMY142) important details:
Cannot miss more than 2 labs (excused or unexcused)
Reports/Worksheets due the day of the next lab
Proper attire required (long pants, closed-toed shoes, goggles & face mask)
Contact your TA ASAP if you anticipate missing a lab or are sick
Chemistry: What is Chemistry and Why Study It?
Chemistry studies matter and its transformations; matter is anything that has mass and occupies space; transformations are chemical changes.
A chemical formula shows the number and type of atoms in the smallest unit of matter, e.g., table salt composed of sodium (Na) and chlorine (Cl) with formula
ext{NaCl}Chemistry acts as the central science linking physics and biology; it enables applications across many fields and real-world problems.
Motivations for taking chemistry in curricula include science core requirements, major prerequisites, and practical scheduling considerations.
Matter, Physical vs Chemical Changes, and Classification
Types of changes:
Physical change: no change in chemical identity; still the same substance (e.g., phase changes: solid ⇄ liquid ⇄ gas)
Chemical change: alters chemical makeup and identity of substances; new substances formed (e.g., Na(s) + Cl₂(g) → NaCl(s))
Matter is composed of particles; basic particles include:
Atoms: basic unit of ordinary matter
Molecules: atoms bound together in specific arrangements
Structure of matter determines properties; structure also links to biological/real-world effects (e.g., development, health)
Pure Substances vs Mixtures
Matter classifications:
Pure substance: single type of matter; constant composition
Element: contains one type of atom (e.g., Au for gold, O₂ for oxygen gas)
Symbols may be one to three letters (C, Al, Fe, W)
Compound: composed of two or more different elements (e.g., H₂O, NaCl)
Mixture: two or more substances physically intermingled; variable composition
Homogeneous mixture (solution): uniform composition throughout
Heterogeneous mixture: nonuniform composition; samples differ in component ratios
Examples:
Pure substances: copper (Cu), table salt (NaCl), sucrose (C₁₂H₂₂O₁₁)
Mixtures: salt+sugar, rocks (minerals)
Matter and Measurement: Numbers, Units, and Uncertainty
Numbers arise from counting (exact) and from measurement (uncertain):
Chemical formulas indicate elements and counts (e.g., NaCl tells 1 Na and 1 Cl, subscript counts atoms per molecule)
Uncertainty is inherent in any measurement due to instrument limitations and variability.
Significant figures (sig figs) convey precision and uncertainty; the last reported digit is the uncertain one.
Exact numbers have unlimited sig figs (e.g., counting discrete objects, defined quantities, some conversion factors).
Unit discussion:
Derived units: e.g., density d = rac{m}{V} with units g/mL (or kg/m³)
Common conversions: 1 L = 1000 mL = 1000 cm³; 1 mL = 1 cm³
Temperature scales: Kelvin (K) is absolute; K = °C + 273.15 ext{ and } °C = K - 273.15
Uncertainty and sig figs together guide reporting of measurements and calculated results.
Significant Figures, Rounding, and Precision vs Accuracy
Key rules:
Leading zeros are not significant; trailing zeros after a decimal are significant; trailing zeros without a decimal can be ambiguous.
Exact numbers have infinite sig figs.
For multiplication/division, report to the least number of sig figs among the factors.
For addition/subtraction, report to the least precise decimal place among the numbers.
Examples from practice:
2.5 cm has 2 sig figs with uncertainty in the last digit (±0.1 cm)
2.52 cm has uncertainty ±0.01 cm, depending on instrument precision
Uncertainty and scientific notation:
Scientific notation helps manage very large/small numbers and preserves sig figs.
Example: 5.41000 × 10^7 preserves five significant figures.
Exact Numbers, Atomic Mass, and Avogadro’s Number
Exact numbers and conversion factors have unlimited sig figs:
1 mol = 6.022 × 10^23 particles (Avogadro’s number)
1 amu is defined relative to 1/12 of the mass of a ¹²C atom: 1 ext{ amu} = rac{1}{12} m(^{12} ext{C})
1 Cal (food Calorie) = 1000 cal = 4184 J
1 eV = 1.602 × 10^-19 J
1 kWh = 3.60 × 10^6 J
Atomic mass units and molar mass:
Atomic weight (atomic mass) in amu corresponds to molar mass in g/mol
Example: Hydrogen ≈ 1.008 g/mol; Oxygen ≈ 15.999 g/mol
Masses, Moles, and Stoichiometry
Moles are a bridge between the atomic world and macroscopic masses:
1 mol contains 6.022 × 10^23 particles (N_A)
Mass of a mole (molar mass): e.g., H ≈ 1.008 g/mol; NaCl ≈ 58.44 g/mol
Converting between mass, moles, and number of particles:
Mass → moles: n = rac{m}{M} where M is molar mass (g/mol)
Moles → particles: N = nN_A
Particles → moles: n = rac{N}{N_A}
Practice problems (conceptual):
How many copper atoms in 0.496 mol? Use N = nNA with NA = 6.022 × 10^23.
How many moles in 7.49 × 10^22 zinc atoms? Use n = rac{N}{N_A}.
Density and volume as needed to link mass to volume and subsequently to moles.
Energy, Light, and Electromagnetic Radiation
Light behaves both as particles (photons) and waves:
Photon energy: E = h
abla = h
u = rac{hc}{ ilde{\lambda}}
Key constants:
Planck’s constant: h = 6.626\times 10^{-34} \text{ J s}
Speed of light: c = 3.00\times 10^{8} \text{ m s}^{-1}
Wavelength and energy relationship:
E = rac{hc}{\lambda}
Atomic emission/absorption and Bohr’s model (hydrogen-like):
Energy change for transitions: \Delta E = -RH\left(\frac{1}{ni^2} - \frac{1}{n_f^2}\right)
Rydberg constant: R_H = 2.179\times 10^{-18}\text{ J}
Emitted photon wavelength: \lambda = \frac{hc}{|\,\Delta E|}
Electromagnetic spectrum and color-energy relationships:
Red light has lower energy and longer wavelength; blue/violet higher energy and shorter wavelength
Photoelectric effect:
For electron emission, photon energy must exceed the binding energy; excess energy becomes kinetic energy: \frac{1}{2}mv^2 = E_{photon} - \phi where φ is the work function
Atomic Theory and the History of the Atom
Key historical milestones:
Democritus proposed indivisible atom (ca. 400 BC)
Dalton proposed modern atomic theory linking to conservation of mass and definite/multiple proportions
Curies discovered radioactivity; atoms not indivisible
Thomson discovered the electron and proposed the plum pudding model
Millikan measured electron charge via oil-drop experiment: q_e = -1.60 \times 10^{-19} \text{ C per drop}
Rutherford discovered nucleus and mostly empty space in atoms via alpha-particle scattering
Quantum mechanics and Schrödinger developed models of electrons in atoms; classical trajectories fail at quantum scales
Atomic structure notation:
Protons (Z), neutrons (N), electrons (e⁻) define the atom
Mass number: A = Z + N
Neutrons: N = A - Z
In neutral atoms, #protons = #electrons; ions have imbalance leading to net charge
Isotopes and ions:
Isotopes have same Z but different N (hence different A)
Ions differ from neutral atoms by electron count; cations have fewer electrons, anions have more
Atomic Structure: Mass, Isotopes, and Atomic Mass Unit
Atomic mass and isotopes:
Atomic mass is the weighted average of isotopic masses (in amu) reflecting natural abundance
Atomic weight on the periodic table is this weighted average
Mass spectrometry (FYI):
Ionize atoms, measure mass-to-charge ratio to determine mass and abundance of isotopes
Electron Configuration and Orbital Theory
Orbitals and quantum numbers define electron location probabilities:
Principal quantum number: n = 1,2,3,4,\ldots (energy and size of orbital)
Angular momentum quantum number: l = 0,1,2,…,n-1 (orbital shape: s, p, d, f)
Magnetic quantum number: m_l = -l, -l+1, …, +l (orientation of orbital)
Spin quantum number: m_s = +\tfrac{1}{2}, -\tfrac{1}{2}
Aufbau principle (The Aufbau “building up” rule): electrons fill the lowest-energy subshells first; e.g., 4s is lower in energy than 3d for many elements so 4s² fills before 3d^x.
Hund’s rule: electrons occupy degenerate orbitals singly before pairing to maximize total spin; no two electrons in an orbital share identical quantum numbers.
Noble gas notation (short form): replace inner shells with [Noble Gas] and continue with the remaining electrons. Example: Cl → [Ne] 3s² 3p⁵
Common orbitals and shapes:
s orbitals (l = 0): spherical
p orbitals (l = 1): two-lobed, orientation along axes
d orbitals (l = 2): four lobes with various orientations; some lobes between axes, others along axes with donut shapes
f orbitals (l = 3): complex shapes with multiple lobes
Example to illustrate electron configuration:
Zn (Z = 30): ground-state configuration is
1s^2\;2s^2\;2p^6\;3s^2\;3p^6\;3d^{10}\;4s^2
or in noble-gas shorthand: [\text{Ar}]\;3d^{10}\;4s^2
Periodic Table and Periodicity
Periodic table organization: horizontal periods and vertical groups/families; elements arranged by atomic number (Henry Moseley’s work).
Groups have similar chemical properties; blocks (s, p, d, f) reflect the subshell being filled.
Notable groups and trends: noble gases in group 18; alkali metals in group 1; halogens in group 17; transition metals occupy d-blocks.
The periodic law: properties of elements show periodic recurrence when elements are ordered by atomic number.
Blocks and electron filling order influence periodic trends:
Atomic radius generally increases down a group and decreases across a period
Ionization energy generally increases across a period and decreases down a group
Electron affinity trends vary but are generally exoergic (negative) for many halogens
Practice of Electron Configuration and Orbitals
Electron configurations are built by filling subshells in order of increasing energy; common pattern: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p …
Example questions:
Ground-state configuration for Zn: [\text{Ar}] 3d^{10} 4s^2
Identify valid quantum number sets: valid sets must satisfy n > 0, 0 ≤ l < n, ml ∈ [-l, l], ms ∈ {+1/2, -1/2}
Masses, Moles, and Density in Practice
Density concept: mass per unit volume; units typical: g/mL or kg/m³
Practice: density calculations from mass and volume; e.g., cube with given dimensions and mass; calculating mass of a displaced volume
Intensive vs extensive properties:
Intensive: do not depend on amount of substance (e.g., density, color, temperature)
Extensive: depend on amount (e.g., mass, volume, length)
Chemical Formulas: Empirical, Molecular, and Structural
Chemical formulas convey composition:
Empirical formula gives the simplest whole-number ratio of atoms (e.g., H₂O₂ → HO)
Molecular formula gives the actual number of atoms in a molecule (e.g., H₂O₂)
Structural formula shows actual bonding and connectivity (e.g., H–O–O–H for hydrogen peroxide)
Examples:
Ethanol: C₂H₆O (molecular formula); could be written as empirical formula C₂H₆O simplifies to C₂H₆O (already simplest)
Dimethyl ether: C₂H₆O (molecular) with a structural representation showing O-linked methyl groups
Nomenclature, Stoichiometry, and Masses
The mole concept is essential in stoichiometry: relate grams to moles and moles to molecules/atoms
Stoichiometric calculations rely on balanced chemical equations and proper significant figures
Dimensional analysis using conversion factors: ratio with a value of 1 to convert units
Chemical Energies and Energy Units
Energy is conserved but can be expressed in various units:
Joules (J) and kilojoules (kJ)
Calories (cal) and kilocalories (kcal; often written Cal for food Calories): 1 Cal = 1000 cal; 1 cal = 4.184 J; 1 Cal = 4184 J
Electron volt (eV): 1 eV = 1.602 × 10^-19 J
Kilowatt-hour (kWh): 1 kWh = 3.60 × 10^6 J
Energy in chemical context:
Kinetic energy: K = \tfrac{1}{2} mv^2
Potential energy and chemical bonds: energy stored in chemical bonds changes during reactions
Gas, Solutions, and Temperature Conversions (Practical Examples)
Temperature scales and conversions:
Kelvin is the SI unit for absolute temperature
Boiling/freezing points and body temperature can be converted to Kelvin or Celsius as needed
Temperature estimation trick: Rough conversion between Celsius and Fahrenheit: °C ≈ 0.5(°F − 32) or rough mnemonic (doubling °F and adding 30) to estimate rough values.
Spectroscopy, Atoms, and Photons
Atomic spectroscopy describes how atoms absorb or emit photons when electrons transition between energy levels:
Absorption: electron promoted to a higher energy level
Emission: electron returns to a lower energy level and emits a photon
Each element has a unique emission/absorption spectrum
Bohr’s model limitations lead to modern quantum mechanical treatment; focus on energy levels and electron orbitals rather than fixed paths
Photon emission/absorption energy relationships allow calculation of emitted/absorbed wavelengths via the same energy equations above
Orbitals, Probability, and Wavefunctions
Orbitals are regions in space where electrons are likely to be found; described by wavefunctions and probability densities
The Schrödinger equation yields allowed energies and wavefunctions; we use quantum numbers to designate orbitals:
n (principal): energy and size
l (angular momentum): orbital shape
m_l (magnetic): orientation
m_s (spin): electron spin
Orbital shapes by l:
s: spherical (l = 0)
p: two lobes (l = 1)
d: four lobes with various orientations (l = 2)
f: more complex, eight-lobed patterns (l = 3)
Electron density vs radial distribution:
Probability density is highest where electrons are likely to be found; but maxima are not always at the nucleus
Radial distribution function gives the probability of finding an electron within a spherical shell at radius r
Practical Quantum Chemistry: Intersection with the Periodic Table
Periodicity arises from how electrons fill orbitals; elemental properties correlate with electron configuration
Orbital filling and valence electrons determine chemical behavior and reactivity
Notable concepts:
The noble gas core notation simplifies electron configurations for heavier elements
The d- and f-blocks involve exceptions to simple filling patterns; beware of irregularities in real elements
Worked Examples and Key Formulas (Summarized)
Energy of light and photons: E = h u = rac{hc}{\lambda}
Planck’s constant: h = 6.626\times 10^{-34}\ \text{J s}
Speed of light: c = 3.00\times 10^{8}\ \text{m s}^{-1}
Hydrogen-like energy change (Bohr/Rydberg):
\Delta E = -RH\left(\frac{1}{ni^2} - \frac{1}{nf^2}\right),\quad RH = 2.179\times 10^{-18}\ \text{J}Wavelength of emitted/absorbed photon:
\lambda = \frac{hc}{|\Delta E|}Atomic mass unit and Avogadro’s number:
1\text{ amu} = \frac{1}{12}\,m(^{12}\text{C})
N_A = 6.022\times 10^{23}
Molar mass, moles, and number of particles:
n = \frac{m}{M}\quad (\text{m in g, M in g/mol})
N = nN_A
Density:
d = \frac{m}{V}
Temperature scales:
K = °C + 273.15\quad \text{and} \quad °C = K - 273.15
Energy units relationships:
1\ \text{cal} = 4.184\ \text{J},\quad 1\ \text{Cal} = 1000\ \text{cal} = 4184\ \text{J}
1\ \text{eV} = 1.602\times 10^{-19}\ \text{J}
1\ \text{kWh} = 3.60\times 10^{6}\ \text{J}
Electron configuration and notation:
Ground-state example (Zn): 1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^6\ 3d^{10}\ 4s^2 = [\text{Ar}]\ 3d^{10}\ 4s^2
Quantum numbers for orbitals:
n\in{1,2,3,\ldots}
l\in{0,1,2,\ldots,n-1}
m_l\in{-l, -l+1,\ldots, l}
m_s\in{+\tfrac{1}{2}, -\tfrac{1}{2}}
Mass from isotopes and average atomic mass:
Average atomic mass is the weighted average of isotopic masses based on abundance
Quick Reference: Key Concepts to Memorize
Matter classifications: pure substances (elements and compounds) vs mixtures (homogeneous vs heterogeneous)
Physical vs chemical changes and their indicators
Sign figs and uncertainty rules for multiplication/division vs addition/subtraction
Fundamental constants and units: h, c, RH, NA, 1\text{ amu}, \text{J}, \text{eV}, \text{kWh}, \text{Cal}
Atomic theory milestones and what each contributed to modern chemistry
Electron configuration rules (Aufbau, Hund’s, Pauli exclusion) and noble gas shorthand
Bohr’s model limitations and the shift to quantum mechanical description
Spectroscopy and the link between energy levels, photons, and atomic spectra
Practice Prompts (Study Aids)
Compute energy and wavelength for a given transition using \Delta E and \lambda = \frac{hc}{|\Delta E|}
Determine the number of neutrons in a given isotope: N = A - Z
Determine the electron count in ions: neutral atom has Z electrons; cations have fewer, anions have more
Convert between mass, moles, and number of particles using: n = \frac{m}{M},\quad N = nN_A
Determine orbital filling order and predict possible sets of quantum numbers; check for validity: e.g., not all triplets of (n,l,ml) are allowed; ml must be in the range for given l
Note on Sources and Relevance
The content aligns with CHMY141 lectures covering foundational chemistry concepts: matter, measurement, atomic theory, periodicity, chemical formulas, moles, energy, light, and quantum orbitals.
Real-world connections include how spectroscopy identifies elements, how isotopic composition informs atomic mass, and how quantum numbers underpin chemical bonding and material properties.
Ethical/practical implications: accurate measurement, uncertainty reporting, and interpretation of data are essential for scientific integrity and informed decision-making in technology, medicine, and environmental science.