AS

Notes from Transcript on Electron Shells and Bonding

Electron Shell Capacities

  • Basis: Electron shells (K, L, M, N, O, P) have capacities given by the formula for the electron capacity of a shell: 2n^2, where n is the principal quantum number (energy level).

  • Shells and capacities (from Table 2.2):

    • n=1 (K): 2 electrons
    • n=2 (L): 8 electrons
    • n=3 (M): 18 electrons
    • n=4 (N): 32 electrons
    • n=5 (O): 50 electrons
    • n=6 (P): 72 electrons

  • Shell labels correspond to letters: K, L, M, N, O, P.

  • Observations:

    • The first shell can hold up to 2 electrons.
    • The second shell can hold up to 8 electrons, and so on following the 2n^2 rule.
    • When an atom has a full outer shell, it is more chemically stable.
    • Elements with complete outer shells are inert (chemically inactive): helium (He) and neon (Ne).
    • Atoms with incomplete outer shells (e.g., hydrogen, carbon, oxygen) are chemically active and tend to gain, lose, or share electrons to achieve a full outer shell.
    • The forces driving electron behavior lead to bond formation between atoms.

  • Planetary model and shell filling (concepts illustrated in Fig. 2.10):

    • The first electron shell (n=1) can hold up to 2 electrons.
    • Hydrogen (atomic number Z=1) has one electron; thus its first shell is not full.
    • Helium (Z = 2) has two electrons; its first shell is full.
    • Lithium (Z = 3) has three electrons; first shell is full (2), second shell has 1 electron.
    • The second shell can hold up to 8 electrons.
    • Carbon (Z = 6) has first shell full (2) and second shell with 4 electrons.
    • Neon (Z = 10) has both first and second shells full (2 + 8).
    • The stability pattern: atoms are most stable when their outer shell is full.

  • Inert vs active tendencies (outer-shell completeness):

    • Helium and neon: outer shells complete → chemically inactive (inert).
    • Hydrogen, carbon, oxygen: outer shells incomplete → chemically active; they tend to complete their outer shell by bonding.
    • The drive to complete outer shells explains bonding and molecular formation.

  • Electron configurations (illustrative examples from the figures):

    • Hydrogen (H, Z=1): outer shell not complete (1 electron total).
    • Helium (He, Z=2): outer shell complete (2 electrons in the first shell).
    • Lithium (Li, Z=3): distribution 2 in the first shell, 1 in the second.
    • Carbon (C, Z=6): distribution 2 in the first shell, 4 in the second.
    • Neon (Ne, Z=10): distribution 2 in the first shell, 8 in the second (both shells complete).
    • Sodium (Na, Z=11): distribution 2 in first, 8 in second, 1 in third shell.
    • These distributions illustrate why some atoms are reactive and how bonding occurs to reach full outer shells.

  • Visual references (Fig. numbers described):

    • Fig. 2.9: Three-dimensional and planetary model depictions of nitrogen’s electron shells (principal quantum numbers).
    • Fig. 2.10: Diagram of atoms showing electron shells (general illustration of shell filling).
    • Fig. 2.11: A molecule of oxygen (O₂) with a planetary model, a 3D model, chemical formula, and a line drawing; shows a double covalent bond (two shared electrons).
    • Fig. 2.12: Carbon dioxide (CO₂) – multiple representations; carbon shares two electrons with each oxygen; two double covalent bonds.
    • Fig. 2.13: Sodium chloride (NaCl) – ionic bonding; electron transfer from Na to Cl; crystal formation; environment in body fluids.

Molecules, Compounds, and Bonding

  • A molecule forms when atoms are joined together by chemical bonds.

    • If two or more atoms of the same element join, the result is a molecule of that element (e.g., O₂).
    • Example: Oxygen gas exists as a molecule of two oxygen atoms; chemical formula is ext{O}_2 where the subscript 2 denotes the number of atoms in the molecule.
    • Molecules can also be formed from atoms of different elements (compounds):
    • Carbon dioxide: chemical formula ext{CO}_2; one carbon atom bonded to two oxygen atoms.
    • Sodium chloride: chemical formula ext{NaCl}; a molecule consisting of one sodium atom bonded to one chlorine atom.

  • A molecule is the smallest unit of a compound that retains the properties of that compound.

    • Important: the properties of a compound can be very different from the properties of the elements it is made from (e.g., table salt differs from sodium metal and chlorine gas).

  • Species examples from the figures:

    • Oxygen (O₂): two oxygen atoms sharing electrons; double covalent bond represented by two lines between atoms.
    • Carbon dioxide (CO₂): carbon shares two electrons with each oxygen; forms two double covalent bonds (O=C=O).
    • Sodium chloride (NaCl): formed via ionic bonding; electrons transferred from Na to Cl; in aqueous environments, this leads to Na⁺ and Cl⁻ ions; NaCl crystals can form cube-shaped crystals.

  • Bond types:

    • Covalent bonds: atoms share electrons (e.g., O₂ with a double covalent bond; CO₂ with two double covalent bonds).
    • Ionic bonds: electrons transferred between atoms, creating oppositely charged ions that attract (e.g., Na⁺ and Cl⁻ forming NaCl).

  • In aqueous environments (e.g., within the body): ionic bonds are frequently formed as electrons are transferred from sodium ions to chloride ions, contributing to the formation and stability of salts like NaCl in solution.

Mixtures

  • Most matter is found as mixtures of two or more substances.
  • Three types of mixtures are common: solutions, colloids, and suspensions (as illustrated in Fig. 2.14).
  • Key concept: mixtures differ from pure compounds in composition and properties; the components can retain their own properties within a mixture.

Connections to Principles and Real-World Relevance

  • Stability and reactivity:
    • Atoms seek to complete their outer electron shells; this drives bond formation and chemical reactivity.
    • Inert gases (He, Ne) illustrate stability due to full outer shells; this explains their low reactivity.
  • Structure and properties:
    • The properties of a compound can be very different from the properties of its constituent elements (e.g., NaCl is table salt, very different from metallic sodium or chlorine gas).
  • Bonding types and biological relevance:
    • Covalent bonding (e.g., O₂, CO₂) underpins the structure and function of many biological molecules.
    • Ionic bonding (e.g., NaCl) is common in physiological contexts and affects solubility, electrical charge, and interactions in body fluids.
  • Foundational numbers and formulas:
    • Electron shell capacity is governed by 2n^2, linking quantum numbers to chemical behavior.
    • Shell labeling (K, L, M, N, O, P) correlates with principal quantum numbers ( n=1,2,3,4,5,6 ).

Summary and Key Takeaways

  • Electron shell capacities increase with shell number according to 2n^2, yielding capacities: K=2, L=8, M=18, N=32, O=50, P=72.
  • Atoms are most stable when their outermost shell is full; inert gases like He and Ne have full outer shells.
  • Bond formation (covalent or ionic) arises from the drive to achieve full outer electron shells; this explains the variety of substances and their properties.
  • Molecules are the smallest units of compounds that retain compound properties; examples include O₂, CO₂, and NaCl.
  • Mixtures (solutions, colloids, suspensions) are pervasive in matter and differ from pure substances in composition and properties.