Exam 1 Survival Skills Notes

Exam 1 Survival Skills

Electronic Structure of Atoms:

  • Bohr's Hydrogen Atom Model:
    • Fundamental model that describes the quantized energy levels for electrons in hydrogen.
  • Calculate Energy of Electron Transitions:
    • Ability to compute energy changes when electrons move from one orbital to another using the formula:
      E = -\frac{KZ^2}{n^2}
      where
    • E is energy,
    • K is a constant,
    • Z is the atomic number, and
    • n is the principal quantum number.
  • Wavelengths of Electron Waves:
    • Derive and calculate wavelengths associated with electron motion or other matter waves using the de Broglie wavelength formula:
      \lambda = \frac{h}{mv}
      where
    • \lambda is the wavelength,
    • h is Planck's constant,
    • m is mass, and
    • v is velocity.
  • Quantum Numbers and Interrelations:
    • Understanding of the four quantum numbers:
    1. Principal quantum number (n)
    2. Azimuthal quantum number (l)
    3. Magnetic quantum number (m_l)
    4. Spin quantum number (m_s)
    • Interrelations between these quantum numbers to define electron configurations.
  • Identifying Shells, Subshells, and Orbitals:
    • Be able to specify the designations for different combinations of quantum numbers (e.g., 1s, 2p, etc.) and their significance.
  • Electron Density and Nodes:
    • Understand electron density distributions in orbitals and the concept of nodes (regions where the probability of finding an electron is zero).
  • Order Relative Energies of Orbitals in Multielectron Atoms:
    • Recognize how orbital energies increase based on penetration and shielding effects and use Aufbau principle, Hund's rule, and Pauli exclusion principle.
  • Orbital Shapes and Sizes:
    • Recognize shapes (spherical, dumbbell, etc.) and relative sizes of different orbitals based on quantum numbers.
  • Spin Quantum Numbers and Pauli Exclusion Principle:
    • Understand the significance of the spin quantum number (±1/2) and how it relates to the Pauli exclusion principle (no two electrons can have the same set of four quantum numbers).
  • Assigning Electron Configuration:
    • Determine electron configurations for atoms in the periodic table using the ترتیبات قوانین.
    • Consequences of Hund's rule: maximize unpaired electrons to minimize electron-electron repulsions in degenerate orbitals.
  • Quantum Numbers for Electron Configurations:
    • Identify quantum numbers of all electrons in a given atom based on its position in the periodic table.
  • Filled and Half-Filled Shells:
    • Recognize the concepts of stability associated with filled and half-filled subshells in electron configurations.
  • Electron Configuration of Main Group and Transition Metals:
    • Write and interpret electron configurations for main group elements (s and p-blocks) and transition metals (d-block).
  • Half-Filled Shell Exceptions:
    • Know the exceptions in electronic configurations due to stability associated with half-filled subshells such as Cr: [Ar] 4s^1 3d^5 and Cu: [Ar] 4s^1 3d^{10}.

Periodic Properties of Elements:

  • Shells and Noble Gas Configurations:
    • Ability to recognize electron configurations for all rows in the periodic table and how they compare with noble gas configurations.
  • Valence Electron Configurations:
    • Identify the configuration of valence electrons based on an atom's position in the periodic table.
    • Valence electrons significantly influence an element's chemical properties.
  • Concept of Electron Screening and Effective Charge:
    • Understand how inner-shell electrons shield outer-shell electrons from the full nuclear charge, impacting their effective nuclear charge experienced by valence electrons.
    • Effective charge can be calculated using Slater's rules or similar models.
  • Arranging Atoms by Size:
    • Be able to rank atoms according to their atomic radius, considering the effects of shell structure and effective nuclear charge.
  • Arranging Ions by Size:
    • Recognize how cations (positive ions) and anions (negative ions) differ in size and arrange them accordingly, considering the loss or gain of electrons.
    • Cations are generally smaller than their neutral atoms, while anions are larger.
  • Ionization Energies Trends:
    • Explain periodic trends in ionization energies, including how energy required to remove an electron generally increases across a period and decreases down a group.
  • Electron Affinities Trends:
    • Describe periodic trends in electron affinities, noting that elements in Groups 1 and 2 often show lower affinities, while halogens have high affinities for gaining electrons.
  • Consequences of Half-Filled and Filled Shells:
    • Explain the effects of half-filled and filled subshells; for instance, half-filled shells can lead to higher ionization energies because of increased stability.
  • Isoelectronic Species:
    • Apply the concept of isoelectronic species, identifying species that have the same electron configuration and comparing their sizes.
  • Transition Metals and Ion Removal:
    • Understand electron removal sequences in transition metals, noting that electrons are removed first from the outermost energy level before the d subshell (typically follows the order 4s > 3d).
  • Basic Properties of Metals vs. Non-Metals:
    • Recognize the fundamental differences between metals and non-metals, including conductivity, malleability, ductility, and electron affinity.
  • Reactivity Trends:
    • Explain general trends in reactivity:
    • Alkali Metals: Highly reactive, especially with water.
    • Alkaline Earth Metals: Reactive, but less than alkali metals.
    • Halogens: Reactivity varies; trend toward higher reactivity as you move up the group.
    • Elements from the Oxygen Group: Varying reactivity but generally less reactive than halogens.
    • Noble Gases: Generally non-reactive due to complete valence electron shells.