Electron Configurations, Ions, and Periodic Trends (Chapter 3)

Overview of ptable.com and Visualizations

• The periodic-table tool shown is ptable.com, which provides interactive information for each element: discovery year, groups, physical state, and color-coding (solids: black; liquids: blue; gases: red).
• Liquid elements in natural state: bromine (Br) and mercury (Hg).
• Isotopes tab shows all isotopes (natural and man-made) and information like electron count around the nucleus.
• Electron configuration visualization: clicking an element reveals an orbital box diagram organized by energy levels; you can click individual boxes to view orbital shapes, rotate the 3D shapes, and zoom for detail. The diagram includes features such as nodes and color-coding for orbitals.
• Going to the wide view allows visibility of f orbitals in real-world locations on the table. Double-clicking on an element shows history and additional information.
• The instructor will use this tool during electron-configuration demonstrations, so keep ptable open as a reference.

Core vs. Valence Electrons and Reactions

• Core electrons do not participate in bonding or reactions for our purposes.
• Valence electrons are the reactive users: removing them forms ions (cation formation), adding them forms anions, and combining valence shells enables covalent bonding.
• Ionic bonding: electrons are transferred to achieve full outer shells; covalent bonding: electrons are shared in valence shells.
• The discussion sets up the plan for next week: focus on electron configuration, then apply to ions and covalent bonding.

Aufbau Principle (OPPO Rules) and Electron Configurations

• Aufbau principle: electrons fill the lowest available energy levels first (lowest to highest energy).
• The energy-level schematic used by the instructor guides electron filling across the periodic table.
• The standard rules work for almost all elements (118 total), but exceptions exist when d- and f-orbitals come into play:
  • For transition metals and inner transition metals, the simple Aufbau filling can fail due to orbital energy changes when d and f orbitals are involved.
  • The general rule applies well to main-group elements (s and p blocks).
• Practical note: there are multiple exceptions; knowing a few key exceptions helps with homework and predicting ion charges, though memorizing every exception is unnecessary for the exam.

Key Exceptions to the Aufbau Rules (Transition and Inner-Transition Metals)

• Chromium (Cr) and Copper (Cu) are classic exceptions:
  • Cr: expected configuration would be $[Ar] 4s^2 3d^4$, but the actual stable configuration favors a half-filled d subshell: $[Ar] 3d^5 4s^1$ (one 4s electron is promoted to 3d).
  • Cu: expected $[Ar] 4s^2 3d^9$, but the actual stable configuration is $[Ar] 3d^{10} 4s^1$ (fills d-orbitals first, then 4s).
• Rationale: half-filled or fully-filled d subshells can be energetically favorable, lowering the overall energy of the atom.
• The book’s emphasized examples for transition metals show that the energy considerations trump simple Aufbau ordering in these cases.
• Importantly, you should not memorize all possible exceptions; instead, know that these exceptions exist and use them as a guide when solving ion-electron configurations.

Using the Periodic Table to Determine Electron Configurations

• For a neutral atom, locate the element and identify the noble-gas core preceding it (no-ball gas reference: neutral, closed-shell core).
• Example workflow:
  • Argon with 4s^2 3d^6 corresponds to iron (Fe) as $[Ar] 3d^6 4s^2$.
  • For manganese (Mn), locate Mn on the table, use the preceding noble gas (Argon) as core, then fill the 4s and 3d orbitals according to the across-row electron count:
  • The formal sequence along the row: 4s^2 (two electrons) then 3d orbitals (count across the five d orbitals as you move left-to-right).
• For scandium through zinc (the 3d block), the d-orbitals are filled in the context of the prior s-block fill. The general rule is to follow the periodic-table order rather than attempting to group all s, p, d, f separately.
• Example: Europium (Eu, Z = 63) is in the f-block region; to build Eu’s configuration, you start from the previous noble gas as a core (xenon, Xe) and fill the 6s^2, then the 4f orbitals, etc., using the standard filling sequence top-to-bottom and left-to-right along the table. The instructor notes this as a practical, though lengthy, process; in class, they skip writing the entire full configuration for a heavy lanthanide.

Practice: Reading and Writing Electron Configurations

• Given a configuration like ext{Ar}
  obreakspace [ ext{Ar}]
  ightarrow 4s^2 3d^6 ext{ (Fe)}, identify the element and confirm the correct position in the periodic table.
• Example walkthrough:
  • Starting at Argon (Ar) core, add two electrons in 4s and six electrons in 3d to reach iron (Fe) configuration.
• Practical note: In the lab, the same process is used to map configurations to positions on the periodic table. The order of filling follows the table’s block arrangement.

Ionic Formation: Cations and Anions

• Definition: Ions form to achieve a full outer electron shell configuration.
• Cations: positively charged ions formed by removing valence electrons; typically metals lose valence electrons. Example: Na (atomic number 11) tends to lose 1 electron to form Na⁺.
• Anions: negatively charged ions formed by gaining electrons; typically nonmetals gain electrons. Example: Cl (atomic number 17) gains 1 electron to form Cl⁻.
• Example explanation (NaCl formation):
  • Na starts with 11 protons and 11 electrons; loses one electron → 11 protons, 10 electrons; charge = $Z - N_e = 11 - 10 = +1.$
  • Cl starts with 17 protons and 17 electrons; gains one electron → 17 protons, 18 electrons; charge = $Z - N_e = 17 - 18 = -1.$
  • The transferred electron from Na to Cl yields an ionic bond with full outer shells for both ions.
• General rule: metals tend to form +1, +2, +3 charges depending on valence; nonmetals tend to form negative charges corresponding to gaining enough electrons to complete the outer shell.
• Important nuance: many transition metals and lanthanides/actinides have multiple common oxidation states; charges are often determined by the chemical context (what other ions are present).
• Practical emphasis: counting valence electrons from the outer shell makes it easier to predict likely charges and electron-transfer behavior.

Valence Electron Focus: Oxygen Example and General Trends

• Oxygen (O) has valence electrons in the second energy level: configuration includes 2s^2 and 2p^4; outer shell is not counting the 1s core. Oxygen tends to gain two electrons to achieve a full outer shell, forming O²⁻.
• The pattern across the periodic table according to valence electrons:
  • Group 1 elements (alkali metals) form +1 ions by losing one valence electron.
  • Group 2 elements (alkaline earth metals) form +2 ions by losing two valence electrons.
  • Aluminium (Al) often forms +3 ions by losing three valence electrons, mirroring its neon-core-like outer configuration.
  • Other p-block elements form various stable oxidation states corresponding to how many valence electrons they can gain or lose to reach a noble-gas configuration.
• Example: Aluminum configuration can be described as a neon core with 3s^2 3p^1 valence; it tends to lose three electrons to become Al³⁺, achieving a noble-gas-like core ([Ne]).
• Electron configurations explain why some elements take on certain charges: the charge corresponds to the number of electrons gained or lost to reach a full outer shell.

Magnetic Properties and Unpaired Electrons

• Electron configurations dictate magnetic behavior:
  • If there are unpaired electrons in any orbital, the atom or molecule can be paramagnetic (attracted to magnetic fields) because the spins align with an external magnetic field.
  • If all electrons are paired in orbitals, the substance is diamagnetic (not attracted, or very weakly repelled).
  • Ferromagnetism occurs in certain metals (e.g., iron, cobalt, nickel, neodymium) where spins align spontaneously and produce a magnetic field even without an external field.
• Mnemonic for diamagnetism: "Di equals two" to recall paired electrons; however, this is a rough memory aid rather than a formal rule.
• Real-world demonstration: liquid nitrogen vs liquid oxygen in a magnet shows paramagnetic behavior for oxygen and stronger ferromagnetic behavior in iron-containing substances.

Trends in the Periodic Table and Their Origins

• Atomic size (atomic radius):
  • Across a period (left to right): atoms get smaller due to increasing nuclear charge (more protons) pulling electrons closer while electrons are added to the same principal energy level.
  • Down a group: atoms get larger due to adding shells (higher principal quantum number n), increasing distance from the nucleus.
• Radius definitions:
  • Van der Waals radius: radius for non-bonded atoms/ions; used to estimate distance between two noble gas atoms (e.g., Kr–Kr in a solid).
  • Covalent radius: half the distance between the nuclei of two covalently bonded atoms.
• Shielding and effective nuclear charge: outer electrons feel Z_eff = Z − S, where S is the number of shielding electrons inner to the outermost electron.
  • Inner electrons shield outer electrons from the full nuclear charge; higher shielding reduces the attraction, making outer electrons easier to remove.
• Ionization energy (IE): energy required to remove an electron from an atom in the gas phase.
  • First ionization energy corresponds to removing the first electron to form a cation.
  • Across a period: IE generally increases (atoms become less willing to lose electrons as they near a full outer shell).
  • Down a group: IE generally decreases (outer electrons are further from the nucleus and more shielded, easier to remove).
  • Bumps in IE occur when electrons are promoted to a new subshell (e.g., after filling 2s, moving to 2p, or crossing from 2p to 3s, etc.), which changes energy costs for removing electrons.
• Electron affinity (EA): energy change when an electron is added to an atom in the gas phase.
  • In general, EA becomes more negative (more favorable to gain an electron) as you move toward nonmetals on the right side of the periodic table, especially near halogens.
  • The trend is less smooth down a group; it is more reliable across a period.
  • A more robust way to discuss electronegativity and bonding tendencies is via electronegativity values rather than raw EA values.
• Metallic vs nonmetallic trends and the metal/nonmetal split:
  • Metals lie on the left side of the zigzag line; nonmetals on the right; metalloids lie along the zigzag boundary.
  • Metals tend to form cations and conduct electricity; nonmetals tend to form anions; metalloids have intermediate, often semicondutor-like properties.
• Group and block naming conventions:
  • Noble gases (Group 18), Halogens (Group 17), Alkali metals (Group 1), Alkaline earth metals (Group 2).
  • Transition metals form the d-block; Lanthanides (4f) and Actinides (5f) are inner transition metals.
  • Main-group elements include the s- and p-block; groups colored in green in the reference table are main-group elements.
• Practical implications: knowledge of trends helps predict chemical behavior, bonding patterns, and reactivity in real-world contexts (electronics, materials, and chemistry education).

Metals, Nonmetals, Metalloids, and Material Properties

• Metals: typically dense, high melting/boiling points, good conductors of electricity and heat, malleable and ductile.
• Nonmetals: brittle, poor conductors, lower densities and melting points relative to many metals; various states at room temperature.
• Metalloids (metalloids line): exhibit properties of both metals and nonmetals; silicon is a classic semiconductor and is central to electronics.
• Practical notes on metalloids:
  • Silicon (Si) works as a semiconductor; behavior depends on treatment/conditions (doping, temperature, etc.).
  • Some elements in the middle region behave as metals in some contexts and nonmetals in others (semiconductor behavior).
• Densities, melting/boiling points, and conductivity trends generally follow across the table but with notable exceptions near transition metals and metalloids.

Real-World Relevance and Examples Mentioned

• Europium (Eu) is discussed in the context of CT contrast agents for imaging; europium’s placement and electron structure inform its chemical behavior in complex compounds.
• The periodic table is used to understand ion formation, electron configurations, and magnetism, all of which have direct applications in imaging, electronics, and material science.
• The instructor hints at practical labs and demonstrations (e.g., a video showing sodium reacting with chlorine gas to form NaCl) to illustrate electron transfer and bond formation.

Exam Scope and Study Strategy (From the Instructor)

• Chapter 3 content is a focus, with Chapter 2 practice problems emphasized as well; Chapter 1 (unit conversions) is reviewed but not the main focus for new problem-solving on the exam.
• Expect questions on:
  • Core vs. valence electrons and their roles in bonding and ion formation
  • Aufbau principle and typical electron configurations, including key exceptions for d-block elements
  • How to read and derive electron configurations from the periodic table (including noble-gas cores)
  • Ion formation and predicting common oxidation states for main-group elements
  • Distinguishing metals, nonmetals, and metalloids and their general properties
  • Understanding trends: atomic size, ionization energy, electron affinity, and shielding/effective nuclear charge
  • Magnetic properties and the concept of unpaired electrons (paramagnetism) versus paired electrons (diamagnetism) and ferromagnetism in certain metals
• Practical advice: don’t memorize every numerical trend; focus on why trends happen and how they relate to nucleus charge, electron shielding, and orbital filling.

Quick Practice Prompts (Conceptual)

• Identify the noble-gas core for a given element and outline its electron configuration up to the valence shell (e.g., Fe: $[Ar] 3d^6 4s^2$).
• Predict the common oxidation state of an alkali metal (Group 1) and explain why it forms a +1 charge.
• Explain why chromium and copper deviate from the simple Aufbau filling and what stabilizes their configurations.
• Explain the difference between van der Waals radius and covalent radius with simple examples (e.g., krypton pair vs two covalently bonded atoms).
• Describe why a material like oxygen (O₂) is paramagnetic due to unpaired electrons, whereas a noble gas like neon (Ne) is diamagnetic.

Summary Takeaways

• Electron configurations and ion formation are grounded in the Aufbau principle but with notable exceptions in transition metals; recognizing half-filled or fully-filled subshells explains some of these exceptions.
• The periodic table is a powerful predictive tool for electronic structure, ion formation, bonding, and properties (sizes, energies, magnetism).
• Understanding core vs. valence electrons is crucial for predicting chemical reactivity and bonding patterns.
• Trends in atomic size, ionization energy, and electron affinity arise from nuclear charge, shielding, and orbital energy ordering; these trends guide expectations for chemical behavior across the table.
• Real-world relevance includes semiconductors (metalloids like silicon), MRI/CT contrast agents (europium), and magnetic materials (ferromagnetic elements).

Note on the Session Plan

• A ten-minute break was announced followed by practice problems focused on Chapters 2 and 3; Chapter 1 review remains important but is treated as background in this review session.
• Answer keys for homework problems will be posted for practice after the break.