Chapter 1-8: Introduction to Atomic Structure and Light Emission (Fluorescence & Chemiluminescence)

Atomic structure recap

• Protons: positive charge; Neutrons: neutral; Electrons: negative charge.
• Atomic number (Z): number of protons in the nucleus; determines identity of the element.
• Electrons in a neutral atom: number of electrons equals number of protons (Z).
• Mass number (A): number of protons plus neutrons; A = Z + N.
• Neutrons: N = A − Z.
• Atomic mass vs. mass number: mass number is the whole-number count used for most discussions; rounding rules discussed in class: often round down, but if the decimal part were > 0.5 you would round up. In many contexts the mass is treated as the nearest integer.
• Important caveat: the nucleus is tiny and protons/neutrons are not the same as electrons; electrons exist in regions around the nucleus described by energy levels.
• Model to visualize electrons around the nucleus: Bohr–Rutherford model (simplified; not a literal picture of electron positions).

Electron energy levels and the Bohr–Rutherford model

• Electrons occupy regions around the nucleus called energy levels (also called shells or rings).
• Energy level nearest the nucleus is n = 1; then n = 2, n = 3, etc. (n is the principal quantum number).
• In this course, the practical loading rule is: the very first energy level can hold up to 2 electrons; each subsequent level can hold up to 8 electrons (a simplified rule used in class; not the most general quantum-mechanical rule).
• Although we say levels are rings and we know where electrons are likely to be, we cannot pinpoint an exact electron position on a ring. The electron’s location is probabilistic, not a fixed point.
• The closer an electron is to the nucleus, the stronger the attraction to the positively charged protons; this makes the electron more stable.
• Energy levels are not equal in energy; electrons can be excited to higher levels by gaining energy and then relax back, emitting light.
• Load rule examples (per shell capacity):
  • 1st shell: 2 electrons max
  • 2nd shell: up to 8 electrons
  • 3rd shell: up to 8 electrons (and so on for this course)
• Orbital arrangement principle (visual): place electrons to minimize repulsion, filling opposite sides of a ring before pairing when possible.
• Three common synonyms used interchangeably in class: energy level, energy shell, energy ring.

How to build electron configurations (example exercises from the transcript)

• Oxygen (Z = 8; A ≈ 16; N = A − Z ≈ 8):
  • Nucleus: 8 protons, 8 neutrons (mass number ~16).
  • Electrons to place: 8.
  • Shell loading: 1st shell gets 2 electrons; 2nd shell gets the remaining 6 electrons.
  • Configuration: 2 in n = 1, 6 in n = 2 → (2, 6).
• Carbon (Z = 6; A ≈ 12; N = 6):
  • Nucleus: 6 protons, 6 neutrons; 6 electrons total.
  • Shell loading: 2 in n = 1; 4 in n = 2.
  • Configuration: (2, 4).
• Chlorine (Z = 17; A ≈ 35; N = 18):
  • Nucleus: 17 protons, 18 neutrons; 17 electrons total.
  • Shell loading: 1st shell 2 e, 2nd shell 8 e, 3rd shell 7 e.
  • Configuration: (2, 8, 7).
• Magnesium (Z = 12; A ≈ 24; N = 12):
  • Nucleus: 12 protons, 12 neutrons; 12 electrons total.
  • Shell loading: 2 in n = 1, 8 in n = 2, 2 in n = 3.
  • Configuration: (2, 8, 2).
• Hydrogen (Z = 1):
  • Nucleus: 1 proton; typically 1 electron.
  • Shell loading: first shell holds 1 electron (fits within the 2-electron capacity).
  • Configuration: (1).
• Quick checks when drawing:
  • The total of electrons in all shells must equal Z (the atomic number).
  • Do not exceed shell capacities; e.g., for magnesium, (2, 8, 2) sums to 12, not more.

Excitation and fluorescence: how energy transfer leads to color

• Electrons can gain external energy and move from a lower energy level (ground state) to a higher one (excited state).
• External energy sources to excite electrons include:
  • Electrical current
  • Flame (high temperature)
  • High-energy light (ultraviolet or other intense light)
• When the excited electron returns to a lower energy level, it emits a photon. This emission process is collectively known as fluorescence in this context.
• Key terms:
  • Ground state: the lowest energy level (often n = 1).
  • Excited state: a higher energy level (n > 1).
  • Fluorescence: the emission of light as electrons relax from an excited state back to lower energy levels.
• Color of emitted light is determined by the energy difference (ΔE) between the initial excited level and the final level:
  • Higher energy drop (ΔE large) -> shorter wavelength (higher energy photon; e.g., violet/blue).
  • Lower energy drop (ΔE small) -> longer wavelength (lower energy photon; e.g., red/orange).
• Photon energy relationship (useful for understanding colors):
  • E = h\nu = \frac{hc}{\lambda}
  • Shorter wavelength (\lambda) means higher energy (\Delta E).
• Transitions can skip levels (e.g., from n = 4 to n = 1) or proceed stepwise (e.g., n = 3 to n = 1); the color depends on the actual transition(s).
• Hydrogen example notes from the lecture:
  • Hydrogen has several emission lines; two prominent ones discussed are red and blue lines; combinations of lines can produce a pink/purple appearance to the eye when viewed without spectral resolution.
  • Some specific transitions in hydrogen (e.g., going to n = 1 from other levels) produce characteristic colors (e.g., red, blue) that are part of emission spectra.
• Emission spectra vs. what we see:
  • The actual emission spectrum consists of discrete spectral lines (emission lines) unique to each element.
  • When viewed with the naked eye, we often see a mixture that appears as a single color; a diffraction grating or prism can separate the lines into distinct colors.
• For practical visuals: visible colors in a flame or glow depend on the element and the specific transitions available to its electrons.

Visible spectrum, order of energy, and color shorthand

• The visible spectrum (ROYGBIV) maps to colors from long to short wavelengths roughly along the rainbow:
  • Red (lowest energy) -> Orange -> Yellow -> Green -> Blue -> Indigo -> Violet (highest energy among visible light).
• In the lecture, high-energy colored light is described as toward the purple/blue end; low-energy colors toward the red/orange end.
• Practical mnemonic: ROYGBIV (Red, Orange, Yellow, Green, Blue, Indigo, Violet).
• In fireworks, colors are chosen by selecting elements whose electrons produce photons in the desired part of the spectrum (e.g., barium for green, strontium for red, copper for blue).

Real-world spectral lines and the diffraction observation

• Emission spectra are element-specific signals: each element has its own set of lines (e.g., hydrogen has a known set of lines; neon has many lines).
• Diffraction glasses or a diffraction grating can separate many lines, revealing distinct lines (red, blue, etc.) rather than a single blended color.
• Observations in the classroom: neon shows many lines but often appears reddish-orange to the eye because many emitted lines lie in the red/orange region; diffraction reveals more lines beyond what the eye perceives as one color.

Fluorescence versus chemiluminescence (glow sticks and more)

• Fluorescence (light emission following excitation by external energy):
  • External energy sources excite electrons (electric current, flame, or high-energy light).
  • When electrons fall back to lower levels, light is emitted.
  • The color depends on the electronic transitions; the actual color observed by the eye is a combination of many lines.
• Chemiluminescence (glow sticks and related phenomena):
  • Energy for excitation comes from a chemical reaction, not from an external light or heat source.
  • Glow sticks have an inner tube containing hydrogen peroxide and an outer dye with a reactive compound (e.g., DPO). When cracked, the reaction occurs, producing energy that excites the dye.
  • The resulting light color matches the dye used and lasts until reactants are consumed.
  • Glow sticks can glow for extended periods; cooling slows the reaction (glow lasts longer but dimmer).
• Fireflies and glowing jellyfish as biological examples of chemiluminescence.
• Luminol-based detection at crime scenes is another chemiluminescence example: the iron in blood catalyzes luminol to emit light, revealing traces of blood.

Practical applications and caveats for color prediction in fireworks and displays

• Fireworks colors are achieved by selecting salts and elements that produce specific emission colors when excited:
  • Barium compounds -> Green emission
  • Strontium compounds -> Red emission
  • Copper compounds -> Blue emission
  • Others produce various colors depending on element and environment.
• It is not necessary to memorize the exact colors for every element; the crucial idea is understanding that colors arise from specific electron transitions and that different elements produce distinct emission spectra.
• Real-world signals and safety notes:
  • UV-fluorescent inks, passport features, and currency markers rely on fluorescent dyes that glow under UV light.
  • Neon signs, although called neon, may use different gases to produce visible colors and rely on fluorescence/electroluminescence rather than pure neon gas emission.

Demos and takeaways from the session

• Key demonstration takeaways:
  • External energy sources excite electrons, which then emit light when they relax; this is the core mechanism behind flame colors and fireworks.
  • A diffraction grating reveals multiple spectral lines that compose the observed color; the eye blends lines into a single color unless spectral components are resolved.
  • The colors observed depend on the element and the transition levels involved, not simply on the amount of energy supplied.
• Summary of the central process:
  • Ground state (n = 1, 2, 3, …) → Excited state (via energy input) → Emission of photon as electron returns to a lower level → Color determined by ΔE and the line spectrum of the element.

Quick mathematical recap (key formulas used in class)

• Neutrons in nucleus: N = A - Z
• Mass number and composition: A = Z + N
• Electron capacity per shell (as taught in class):
  • 1st shell: 2 electrons
  • 2nd shell: up to 8 electrons
  • 3rd shell: up to 8 electrons (and so on in this course)
• Photon energy and wavelength relationships:
  • E = h\nu = \frac{hc}{\lambda}
• Conceptual relationships: higher energy photons have shorter wavelengths; lower energy photons have longer wavelengths.

Connections to broader concepts and real-world relevance

• The atom’s electron structure explains chemical properties and reaction behavior—electrons are arranged in shells; the shell occupancy influences stability and reactivity.
• Fluorescence and chemiluminescence illustrate how energy transfer can occur via different pathways (external energy vs chemical reactions) to produce light.
• The same physical principles underlie everyday phenomena: glow sticks, neon/fluorescent inks, forensic luminescence, and even biological glow (fireflies, jellyfish).
• Understanding emission spectra informs color selection in displays, lighting design, and safety/curity applications (UV ink, currency verification).

Ethical and practical implications (brief)

• Spectral analysis informs forensics (luminol) and safety (authenticating currency, passports, and glow-in-the-dark markers).
• Knowledge of fluorescence and chemiluminescence affects how we perceive consumer products (glow sticks, cosmetics, and novelty items) and informs safety considerations when handling UV light or chemical reagents.
• Fireworks and public displays rely on precise chemical knowledge to achieve intended colors; this has regulatory and safety implications.