Modern Atomic Models & Electron Location
Unit 2, pages 10–19 — Bohr, spectra, quantum model, and electron configurations
Lesson 2.5: Modern Atomic Models
Rutherford’s nuclear model explained only a few properties. It could not explain why metals or metal compounds give off color when heated in a flame, or the chemical properties of elements. A better model of electron behavior was needed.
Niels Bohr (Danish physicist, student of Rutherford) proposed that an electron is in a fixed circular path around the nucleus called an orbit. Orbits of different energy are called energy levels or shells.
Electrons can exist in one energy level or another, but nothing in between. A quantum of energy is required to move an electron from one energy level to another.
The energy level an electron normally occupies. An electron closer to the nucleus (n = 1) has less potential energy than an electron farther from the nucleus.
A higher-energy, less stable level or shell that an electron reaches by absorbing energy. Unstable. Electrons become stable again by releasing that energy.
The packet of energy released as light when an excited electron returns to its original ground state. The energy released has a specific wavelength. E = hc/λ or E = hν (E = energy, h = Planck’s constant, c = speed of light, λ = wavelength, ν or f = frequency).
- K (PEL 1) — 2 electrons
- L (PEL 2) — 8 electrons
- M (PEL 3) — 18 electrons
- N (PEL 4) — 32 electrons
Tubes contain elements in the gaseous phase under pressure. When electricity excites the electrons, the tubes glow a unique color. Viewed through a spectroscope, the light is composed of narrow lines of color: bright-line (emission) spectra.
Produced when an electron releases energy as it returns from a higher energy level to a lower energy level. An emission spectrum is an atom’s fingerprint and can identify an element. The arrangement of electrons in the PELs produces spectra specific to that element.
Excited electron moves back to the first principal energy level. Ultraviolet region.
Excited electron moves back to the second principal energy level. Visible light region.
Excited electron moves back to the third principal energy level. Infrared region. Lower energy, higher wavelength.
The range of wavelengths of energy. Energy released by electrons occupies the portion humans detect as visible light. Slight variations in energy appear as different colors. Increasing energy goes with decreasing wavelength (gamma → X-ray → UV → visible → infrared → radio).
- f = c / λ
- E = hc / λ
- E = hf
- E(light) = E_f − E_i
- h = Planck’s constant ≈ 6.63 × 10−34 J·s
- c = 3.00 × 108 m/s
Red ≈ 656 nm, green ≈ 486 nm. Red has longer wavelength, so lower energy and lower frequency than green. Answer: (4) The red line has lower energy and lower frequency than the green line.
Wavelength and energy are inversely related. UV-A (~350 nm) has the shortest wavelength and highest energy. Blue (~450 nm) is in the middle. Red (~650 nm) has the longest wavelength and lowest energy. Blue light has less energy than UV radiation, which is known to cause biological damage.
- Does not explain spectra of atoms with more than one electron.
- Cannot explain why some spectral lines are double or triple.
- Cannot explain why energy levels are so close together.
- An electron moving in a circular path would accelerate, lose energy, and spiral into the nucleus.
Erwin Schrödinger (1926) treated electrons moving around the nucleus as waves. The nucleus is surrounded by orbitals, not orbits. An orbital is a three-dimensional region about the nucleus where there is a high probability of finding an electron. Orbitals differ in size, shape, and orientation from Bohr’s circular orbits. Electron location is described by 4 quantum numbers.
Lesson 2.6: Electron Location
The main energy level surrounding the nucleus. n values are whole numbers: 1, 2, 3, etc. As n increases, distance from the nucleus increases, energy increases, and the size of the orbitals increases.
Regions within a principal energy level: s, p, d, f. Shapes: s is spherical, p is dumbbell, d has four lobes. The number of sublevels (possible orbital shapes) in each main energy level equals the value of n.
Each orbital holds up to 2 electrons. Orbitals within a sublevel have equal energy. Positions are relative to x, y, z axes.
- s sublevel: 1 orbital, 2 electrons (lowest energy)
- p sublevel: 3 orbitals (px, py, pz), 6 electrons
- d sublevel: 5 orbitals, 10 electrons
- f sublevel: 7 orbitals, 14 electrons (highest energy)
Describes the spin of an electron in either of two possible directions. Explains why electrons can be near each other in energy levels without repelling too much. Paired electrons have opposite spins.
- n = 1: s only — 2 electrons
- n = 2: s, p — 8 electrons
- n = 3: s, p, d — 18 electrons
- n = 4: s, p, d, f — 32 electrons
The distribution of electrons of an atom in atomic orbitals. Example: 1s² 2s² 2p⁶. Three rules must be followed: Aufbau principle, Hund’s rule, and Pauli’s exclusion principle.
An electron occupies the lowest-energy orbital before filling a higher-energy orbital. Exceptions at higher principal quantum numbers: 5f and 6d, and the chromium/copper-type exceptions (4s¹ 3d⁵ and 4s¹ 3d¹⁰).
In multielectron atoms, in the 3rd energy level and above, sublevels begin to overlap. The periodic table can be used to predict which sublevel is being filled by a particular element (s-block, p-block, d-block, f-block).
- Sodium: 1s² 2s² 2p⁶ 3s¹
- Aluminum: 1s² 2s² 2p⁶ 3s² 3p¹
- Scandium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
- Chromium (exception): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵
- Copper (exception): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
- Zinc: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰
- Chlorine: 1s² 2s² 2p⁶ 3s² 3p⁵
- Oxygen ion (2−): 1s² 2s² 2p⁶
- Potassium ion (1+): 1s² 2s² 2p⁶ 3s² 3p⁶
Identifies the total number of electrons in each energy level. Core/kernel = non-valence electrons. Valence shell = the outermost energy level. Valence electrons = the outermost electrons (s and p sublevels). Example: Mg is 2-8-2 (core 2-8, valence 2).
The symbol of the last noble gas before that atom, followed by the configuration of the remaining electrons. Examples: Mg is [Ne] 3s²; Ar is [Ne] 3s² 3p⁶; Rb is [Kr] 5s¹; Mn is [Ar] 4s² 3d⁵.
Orbitals of equal energy are each occupied by one electron before any one orbital is occupied by a second electron. All electrons in singly occupied orbitals must have the same spin.
No two electrons in the same atom can have the same set of four quantum numbers. When electrons pair up, they do so with opposite spins.
Sublevels are divided into orbitals. No more than 2 electrons occupy an orbital. Boxes or lines represent orbitals; arrows represent electrons with opposite spins when paired. Follow Aufbau, Hund’s, and Pauli’s rules when drawing.