MC

orgo rec Atomic Structure, Isotopes, Electron Configuration, and Bonding - Study Notes

Atomic Number, Mass Number, Protons, Electrons, Neutrons

  • Elements are denoted by symbol and have key identifiers: atomic number Z, mass number A, and the actual element name (e.g., chlorine is Cl).
  • Chlorine example:
    • Symbol: Cl
    • Atomic number: Z = 17
    • Mass number (given): A = 37
    • Number of neutrons: N = A - Z = 37 - 17 = 20
    • For a neutral chlorine atom, number of electrons equals number of protons: E = Z = 17
  • Three fundamental types of subatomic particles in atoms:
    • Protons: positive charge (in the nucleus)
    • Neutrons: neutral charge (in the nucleus)
    • Electrons: negative charge (orbiting the nucleus)
  • In a neutral atom, protons = electrons (balance positive and negative charges).
  • Atomic mass is effectively the sum of protons and neutrons in the nucleus:
    • A = Z + N
    • Chlorine example: A = 17 + 20 = 37 (consistently given)

Isotopes and Isotopic Form

  • Isotopes are atoms of the same element (same Z) with different numbers of neutrons (N).
  • Isotopes differ in atomic mass (A) but have the same atomic number (Z).
  • Hydrogen isotopes (examples):
    • Protium: ^1_1H with 0 neutrons, Z = 1, A = 1
    • Deuterium: ^2_1H with 1 neutron, A = 2
    • Tritium: ^3_1H with 2 neutrons, A = 3
  • How isotopes differ:
    • Primary change: number of neutrons (N)
    • Consequence: mass number A changes; the element identity (Z) stays the same
  • Problem-type: Given Z, N, and E (electrons), determine mass number and charge:
    • Mass number: A = Z + N
    • Charge: q = Z - E (positive if fewer electrons than protons, negative if more electrons)
  • Worked isotopic problems from transcript:
    • Example 1: 7 protons, 7 neutrons, 10 electrons
    • Element: Nitrogen (Z = 7)
    • Mass: A = Z + N = 7 + 7 = 14
    • Charge: q = Z - E = 7 - 10 = -3
    • Result: Nitrogen-14 an ion with charge -3
    • Example 2: 15 protons, 16 neutrons, 15 electrons
    • Element: Phosphorus (Z = 15)
    • Mass: A = 15 + 16 = 31
    • Charge: q = 15 - 15 = 0
    • Result: Neutral phosphorus-31
    • Example 3: 25 protons, 30 neutrons, 23 electrons
    • Element: Manganese (Z = 25)
    • Mass: A = 25 + 30 = 55
    • Charge: q = 25 - 23 = +2
    • Result: Mn with a +2 charge
  • Important note: The atomic number Z is unique to each element; changing Z changes the element itself (e.g., 34 would be Se, not Br). The atomic mass and the number of electrons can vary in ions or different chemical states, but Z identifies the element on the periodic table.

The Nucleus and Electron Cloud: Basic Atomic Model

  • Atom structure:
    • Nucleus: contains protons (positive) and neutrons (neutral)
    • Electron cloud: electrons (negative) orbit the nucleus
  • The nucleus holds all the positive charge; the electrons balance the charge in neutral atoms
  • Orbitals and energy levels:
    • Electrons occupy orbitals in energy levels (shells)
    • Each orbital has a capacity for electrons:
    • s orbital: maximum of 2 electrons
    • p orbitals: 3 types (px, py, p_z); each can hold 2 electrons; total for p set: 6 electrons
  • Energy-level filling order (increasing energy):
    • Level 1: 1s
    • Level 2: 2s, 2p
    • Level 3: 3s, 3p
    • Level 4: 4s, 4p, 3d (and so on)
  • Rule to fill: lowest energy first, then move to higher energies; within an energy level, follow Hund’s rule (maximize unpaired electrons before pairing)

Example: Sulfur Electron Configuration and Energy Diagram

  • Sulfur (S) has Z = 16 electrons in neutral state
  • Filling order and configuration:
    • 1s^2 2s^2 2p^6 3s^2 3p^4
  • Energy diagram vs. electronic configuration: two ways to represent the same distribution
    • Energy diagram shows electrons in orbitals with arrows for spin (up/down, opposite spins)
    • Electronic configuration expresses the occupancy in chemical notation
  • Step-by-step filling (as illustrated in the transcript):
    • 1s filled with 2 electrons (opposite spins)
    • 2s filled with 2 electrons
    • 2p filled with 6 electrons (three orbitals: px, py, p_z all have one electron first, then pair as needed)
    • 3s filled with 2 electrons
    • 3p has 4 electrons distributed as: 3px^2 3py^1 3p_z^1 (Hund’s rule: first place one electron in each p orbital before pairing)
  • Resulting configuration: 1s^2 2s^2 2p^6 3s^2 3p^4

How Orbitals Work: Filling Rules and Terminology

  • Orbital capacity summary:
    • s orbital: 2 electrons
    • p orbitals: 3 types (px, py, p_z); each holds 2 electrons; total for p set: 6 electrons
  • Practical energy-level progression (short form):
    • 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p …
  • When listing electrons, remember opposite spins in a filled orbital (one up, one down)
  • Summary rule: “Fill lowest energy orbitals first; distribute electrons with opposite spins in the same orbital; apply Hund’s rule for degenerate orbitals (like the three 2p orbitals)”

Ionic vs Covalent Bonding: Concepts and Examples

  • Bonding types depend on how electrons are arranged between atoms:
    • Ionic bond: transfer of electrons from one atom to another, resulting in oppositely charged ions held together by electrostatic attraction (typical metal + nonmetal)
    • Covalent bond: sharing of electrons between two nonmetals
  • Key rule: one bond corresponds to two electrons shared (or transferred) between the two atoms
  • Why metals tend to form ionic bonds:
    • Metals (e.g., Group 1A and Group 2A) have 1 or 2 valence electrons and can lose them easily
    • Nonmetals tend to gain electrons to complete their valence shells
    • Creating ions that achieve the nearest noble-gas electron configuration provides a driving force for bond formation
  • Example: Sodium chloride (NaCl)
    • Sodium (Na) has a single valence electron; chlorine (Cl) needs one more to complete its octet
    • Na donates 1 electron to Cl -> Na^+ and Cl^−
    • Na achieves the noble-gas configuration of neon; Cl achieves the noble-gas configuration of argon
    • Resulting compound is held together by an ionic bond due to electrostatic attraction between Na^+ and Cl^−
  • Example of bonding rationale: the driving force is achieving a stable noble-gas electron configuration for both atoms involved

Lewis Structures and Organic Molecules: A Practical Example

  • Lewis (electron-dot) structures summarize valence electrons and bonding patterns using dots and lines
  • General rules for organic molecules (typical elements):
    • Carbon (C): valence 4; can form up to four bonds
    • Hydrogen (H): valence 1; forms one bond
    • Other common heteroatoms: Oxygen (O), Nitrogen (N), etc.; each with characteristic valences (O often forms two bonds and may have two lone pairs, N forms three bonds, etc.)
  • Practice problem in transcript: drawing a Lewis structure for an organic molecule with CH groups and an OH group (example inferred from context: CH3-CH2-CH(OH)-CH3, i.e., 2-butanol)
    • Count valence electrons:
    • Carbons: 4 imes 4 = 16
    • Hydrogens: 10 ext{ H atoms} imes 1 = 10
    • Oxygen: 1 ext{ O} imes 6 = 6
    • Total valence electrons: 16 + 10 + 6 = 32
    • Sketch the carbon skeleton: CH3-CH2-CH(OH)-CH3
    • In the skeleton, form C-C and C-O bonds: three C-C bonds along the chain, one C-O bond, and an O-H bond
    • Distribute electrons to satisfy octets and complete the molecule with correct bonding
    • Example final (neutral) structure: CH3-CH2-CH(OH)-CH3 with O bearing two lone pairs and single bonds to C and H (standard alcohol structure)
    • Note: In a Lewis diagram for this molecule, all atoms should satisfy typical valence: C with four bonds, H with one bond, O with two bonds and two lone pairs
  • Takeaway: Lewis structures help visualize bonding patterns, electron counts, and formal charges (if present)

Quick Biological/Real-World Relevance and Connections

  • Atomic numbers (Z) identify elements on the periodic table; mass numbers (A) and neutron counts (N) relate to isotopes and nuclear stability
  • Isotopes have practical relevance in dating, medical tracers, and nuclear research due to differing neutron counts
  • Electron configuration and orbital filling underpin chemical reactivity, periodic trends, and bonding behavior (ionic vs covalent)
  • Noble-gas configurations as a unifying concept for stability and bond formation explains why many atoms form ions to reach closed shells
  • Lewis structures are foundational for predicting molecular geometry, reactivity, and functional groups in organic chemistry

Summary of Key Formulas and Concepts

  • Mass number: A = Z + N
  • Neutron count: N = A - Z
  • Electron count (neutral atom): E = Z
  • Ion charge: q = Z - E
    • Positive charge when electrons are fewer than protons
    • Negative charge when more electrons than protons
  • Isotopes differ in neutrons; mass number changes accordingly
  • Orbital capacities: s o 2 ext{ e}^-; ext{each } p ext{ orbital} o 2 ext{ e}^-, ext{3 orbitals in } p o 6 ext{ e}^-
  • Electron configuration order (typical): 1s o 2s o 2p o 3s o 3p o 4s o 3d o 4p o ext{…}
  • For sulfur (S, Z = 16): 1s^2 2s^2 2p^6 3s^2 3p^4
  • Ionic bonding driving force: achieve nearest noble-gas electron configurations (Na^+ to Ne-like, Cl^- to Ar-like in NaCl)
  • Covalent bonds involve sharing electrons between identical or different nonmetals; one bond = two shared electrons
  • Example Lewis approach for CH3-CH2-CH(OH)-CH3 (2-butanol) involves counting valence electrons, constructing the carbon backbone, adding O-H bond, and ensuring octets/valence are satisfied

If you’d like, I can tailor these notes to a specific section you’re most worried about or add more worked examples (with full step-by-step solutions) for isotope problems, electron configurations, or Lewis structures.