Chapter 1-7: Electron Configurations, Periodic Trends, and Quantum Rules
Exam-focused notes: Electron configuration, shielding, and periodic trends
Exam mindset and study timing
The transcript opens with an encouragement about starting to study for the second exam as soon as possible; waiting until today is risky and chances of passing are slim but not zero. A realistic, probabilistic view is offered (e.g., “10,000 to 1 odds,” “100,000 to 1 odds”) to emphasize that late starts reduce success probability, but some chance may remain if you begin now.
Emphasize the practical takeaway: begin studying today; there’s always some chance if you put in effort.
Electron configuration: shorthand vs. full notation
Core electrons can be represented with square brackets; the remaining electrons fill the valence shells after the core.
Example: writing the configuration for elements down the periodic table often uses a core shorthand like [Ar] … to represent all electrons up to the noble gas with a closed shell, followed by the valence electrons (e.g., 3s2 3p5 for chlorine, written as [Ne] 3s2 3p5 in many contexts; the transcript shows a simplification).
The general rule: after the core, fill the valence orbitals in order of increasing energy to get the full configuration.
Short-hand configuration usefulness: helps avoid writing dozens of electrons for large atoms; longhand configuration is still correct but unwieldy.
Examples of typical short-hand configurations in the transcript:
Chromium (Cr): [Ar] 3d5 4s1 (instead of the naive [Ar] 3d4 4s2, due to extra stability from a half-filled d-subshell, discussed below).
Copper (Cu): [Ar] 3d10 4s1 (an exception that yields stability via a filled d-subshell).
Pattern: exceptions become more common as you go down the periodic table, especially in the f-block; d-block has fewer but notable exceptions.
Other common pattern notes:
Silver (Ag) and Gold (Au) follow the trend of shuffling an electron to achieve a more stable configuration (e.g., Ag: [Kr] 4d10 5s1; Au: [Xe] 4f14 5d10 6s1 in the conventional picture).
The exceptions Cr, Mo, Cu, Ag, and Au are the main ones you’re expected to remember in this context.
The logic behind these exceptions: promoting stability via half-filled or fully-filled subshells reduces total energy; this is a quantum-mechanical consideration.
Aufbau, Pauli, and Hund: guiding principles
Aufbau principle (filling order)
Electrons fill from the lowest energy levels to higher ones: 4s is lower in energy than 3d, so 4s is filled before 3d; after that, 3d is filled before 4p, etc. The general observed order around the 3d/4s region is:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p …
The exact ordering can vary with the element, but the principle remains: follow energy levels to fill electrons.
Pauli exclusion principle
No two electrons in the same atom can have the same four quantum numbers (n, l, ml, ms).
Practically: if two electrons occupy the same orbital (same n, l, ml), they must have opposite spins: ms = +1/2 and m_s = -1/2.
In orbital diagrams, this is shown as opposite spins (arrows up and down) in the same box of a given orbital.
Hund’s rule (degenerate orbitals)
For orbitals of the same energy (degenerate orbitals), electrons fill singly first with parallel spins before any pairing occurs.
Example for d orbitals: there are 5 degenerate d orbitals (m_l = -2, -1, 0, 1, 2) with the same energy. When placing electrons, you first maximize unpaired electrons by placing one electron in each of the five d orbitals with the same spin before pairing.
Practical demonstration in the transcript: seven electrons in the 3d set. The process fills singly with aligned spins across the five d orbitals, then adds the sixth electron by pairing in one orbital, and the seventh by pairing in another, illustrating Hund’s rule and Pauli’s principle together.
Orbital diagrams vs. electron configurations
Electron configuration notation (e.g., 1s^2 2s^2 2p^6 …) lists the occupancy of each subshell in order of increasing energy.
Orbital diagrams visualize orbitals as boxes (with five 3d boxes, three 4p boxes, etc.). Each box can contain up to two electrons with opposite spins; degenerate orbitals are shown as separate boxes.
When drawing, you often start from a core configuration and then fill valence orbitals. Depending on the question, you may show core electrons, valence electrons, or both.
Example (fluorine): electron configuration is 1s^2 2s^2 2p^5. In an orbital diagram, you would show three 2p orbitals with two electrons in the first two boxes (paired) and one electron in the remaining 2p orbital, with opposite spins where electrons pair.
Valence vs core electrons in diagrams
Some problems show all electrons (core + valence); others show only valence electrons. Decide based on the prompt, as it can save time but increase the risk of error if miscounted.
For argon (Ar, Z = 18): core electrons include the [Ar] core, but if writing full configuration, you would have 1s^2 2s^2 2p^6 3s^2 3p^6; valence electrons are those in the outermost shell (for noble gases, the valence shell is full, so valence count is typically 8 for Ar, corresponding to 3s^2 3p^6).
Electron configurations in practice and examples
Full notation vs shorthand
Full notation shows all electrons; shorthand collapses inner shells into a bracketed core:
Chromium (Cr): [Ar] 3d^5 4s^1 (instead of [Ar] 3d^4 4s^2).
Copper (Cu): [Ar] 3d^10 4s^1 (instead of [Ar] 3d^9 4s^2).
The half-filled (d^5) and fully-filled (d^10) subshells provide extra stability, lowering total energy and thus making these configurations favorable.
Comparison example: nickel vs zinc
Ni (Z = 28): configuration [Ar] 3d^8 4s^2; two valence electrons in the 4s with a filled-ish outer shell configuration, but the inner 3d shell is filled to 8.
Zn (Z = 30): configuration [Ar] 3d^10 4s^2; though Zn has two more protons than Ni, the added electrons go into the 3d shell (a core-like shell relative to the valence 4s), increasing core shielding.
Result: despite increasing Z, the effective nuclear charge felt by the outer electrons in Ni and Zn is similar, so their radii are nearly the same. This illustrates how d-electron shielding affects size trends in transition metals.
Shielding, effective nuclear charge, and periodic trends
Shielding and Z_eff (effective nuclear charge)
Shielding: inner (core) electrons block some of the nuclear positive charge from the outer (valence) electrons, reducing the net pull of the nucleus on the valence electrons.
Zeff = Z - Ncore, where Z is the atomic number (total protons) and N_core is the number of core electrons (non-valence electrons).
Example: beryllium (Be, Z = 4) has two core electrons (1s^2); thus Zeff = 4 − 2 = +2 for the valence electrons, explaining the stronger pull and smaller atomic size compared to lithium (Li, Zeff = 1).
Size (atomic radius) trends along the periodic table
Diagonal trends: many properties vary along diagonals across the table; for size, moving toward the bottom-left (toward francium) increases size (more shells added). Moving toward the top-right (toward fluorine) decreases size (higher Z_eff for the same valence shell).
Size trend explanation: adding shells increases the overall radius; increasing Z_eff pulls electrons closer to the nucleus, reducing radius.
Specific example: Li to Be to B to C shows decreasing size due to increasing Z_eff, while moving toward Francium increases size because of added shells.
Ionization energy trends
Ionization energy is the energy required to remove an electron from a neutral atom.
First ionization energy generally increases up the diagonal toward the right/up (toward the halogens, e.g., fluorine) and to the top of the table, because it becomes harder to remove electrons when they are held more tightly by a higher Z_eff and smaller radius.
Noble gases have very high first ionization energies; alkali metals have very low first ionization energies.
Transition elements show relatively flat first ionization energies due to stable radii from d-electron shielding and variable valence electron configurations.
Deep look: removing successive electrons and the valence/core distinction
When removing electrons successively (second, third, etc.), the energy required changes nonlinearly, creating a zigzag pattern in ionization energies.
The zigzag arises because removing valence electrons is comparatively easy, but once you remove all valence electrons and approach core electrons, the energy required jumps dramatically.
Example with calcium and magnesium: both have two valence electrons readily removed, but removing a third electron (a core electron) requires a large energy jump.
This pattern reinforces the practical rule that chemistry largely involves valence electrons; core electrons participate poorly in bonding and chemical reactions.
Metallic character across the diagonal
Metallic character increases toward francium (bottom-left direction on the standard table orientation) and decreases toward chlorine (top-right).
Metals conduct electricity and heat well, are malleable, and typically form basic oxides; nonmetals are more likely to form acidic oxides and rarely conduct as well.
Magnetic properties and di- vs paramagnetism
Magnetic behavior basics
Atoms/elements with unpaired electrons exhibit paramagnetism (attraction to a magnetic field).
Atoms with all electrons paired exhibit diamagnetism (weak repulsion or very slight interaction with a magnetic field).
Example observations mentioned in the transcript
Liquid oxygen can be attracted to magnets, demonstrating paramagnetism (presence of unpaired electrons).
Diamagnetic behavior occurs when electrons are all paired, canceling magnetic moments.
Spin and pairing in context
The presence or absence of unpaired electrons is directly tied to the spin configuration in the outer shells (especially p, d, and f subshells).
Pauli and Hund rules determine whether these electrons are paired or unpaired in degenerate orbitals, influencing magnetic properties.
Worked examples and quick checks
Argon (Ar) as a count/check example
Argon has Z = 18; in a neutral atom, it has 18 electrons.
Full configuration (for reference) is 1s^2 2s^2 2p^6 3s^2 3p^6.
If illustrating valence electrons, the outer shell (3rd shell) has 8 electrons (3s^2 3p^6).
Depending on the problem, you may show the core electrons ([Ne] core) and the valence electrons separately.
Fluorine as an orbital-diagram example
Electron configuration: 1s^2 2s^2 2p^5.
In an orbital diagram, the 2p subshell consists of three degenerate orbitals (px, py, pz). Fill singly with parallel spins up to maximize unpaired electrons before pairing (Hund’s rule).
We use Pauli: electrons occupying the same orbital (same n, l, m_l) must have opposite spins.
The “alpha” and “beta” notation (up/down arrows) is a visual shorthand for spin states.
Nickel (Ni) vs zinc (Zn) radii and the shielding explanation
Ni to Zn transition adds two more electrons to the 3d subshell, not to the valence 4s, increasing core shielding.
While Z increases from Ni to Zn, the effective nuclear charge felt by the outermost electrons remains similar due to added shielding from the extra 3d electrons, so radii remain nearly identical.
This demonstrates how d-electron shielding flattens size trends across the transition series.
Summary and practical takeaways for exam preparation
Know the Aufbau order and how to apply it to construct configurations; remember the common counterexamples (Cr, Cu) where the observed ground-state configurations differ from the simple fill order due to stability considerations.
Be able to switch between full electron configurations and shorthand (e.g., [Ar], [Kr], [Xe]) and understand when to use each.
Understand shielding and Zeff, and be able to explain why atomic size generally decreases across a period (increasing Zeff with relatively constant shell structure) and increases down a group (adding shells).
Recognize the diagonal trends for size, ionization energy, and metallic character, and be able to explain them qualitatively in terms of Z_eff and electron shielding.
Distinguish between valence and core electrons; know that most chemistry is governed by valence electrons and that core electrons contribute to shielding.
Explain magnetic behavior (paramagnetism vs diamagnetism) in terms of unpaired vs paired electrons.
Use electron configurations to answer questions about orbital occupancy, degeneracy, and spin; be able to draw orbital diagrams and interpret them for given elements.
Be comfortable with common exceptions (Cr, Mo, Cu, Ag, Au) and the rationale behind their configurations.
If you want, I can convert this into a printable one-page cheatsheet with the key rules and a few worked examples for quick study before the exam.