Chapter 11 Study Guide Notes

Nature of Light

• Light is a form of energy that travels through space as a wave.
• It travels at a constant speed, the speed of light, denoted as "c".
  • c = 3.0 \times 10^8 \text{ m/s}
• Two key properties:
  • Wavelength ($\lambda$): distance between crests, measured in meters (m) or prefixed units (e.g., nm).
  • Frequency ($\nu$): number of waves passing a point per second, measured in Hertz (Hz), where 1 \text{ Hz} = 1 \text{ s}^{-1}.
• Relationship between wavelength and frequency: \lambda \nu = c
• Energy: Higher frequency light has more energy; lower wavelength light has more energy.
• Photon: A discrete increment of light, sometimes called a particle of light.
  • Represents the quantity of light emitted or absorbed by a single electron.
• Quantum: Energy of a photon of light, calculated using Planck's equation:
  • E = h\nu
  • Where h is Planck's constant: h = 6.626 \times 10^{-34} \text{ J} \cdot \text{s}
• Electromagnetic Radiation: Light is referred to as electromagnetic radiation.
• Spectrum: Separation of light into its various wavelengths.
  • The electromagnetic spectrum ranges from very short wavelengths (

Electromagnetic Spectrum

• The light we see is a small portion of the entire electromagnetic spectrum.

Bohr’s Model of the Atom

• Niels Bohr studied the atomic spectrum of hydrogen.
• Atomic Spectrum of Hydrogen: A "line spectrum" with only a few specific wavelengths.
• Bohr's Postulates:
  • Electrons orbit the nucleus in specific circular orbits.
  • Light is emitted when an electron transitions from a higher orbit to a lower one.
• Energy Levels: Begin at level one and increase, with increasing distance from the nucleus.
• Equation:
  • Bohr used quantum numbers (1, 2, 3, etc.) to predict the energies of light emitted by hydrogen atoms.
  • He considered transitions from an "excited state" (excess energy) to the "ground state" (lowest energy level).
• Significance:
  • Bohr’s model was the first to specifically predict where electrons actually exist in an atom.
• Limitations:
  • Worked perfectly for hydrogen (one electron) but failed for atoms with multiple electrons.

Current Model of the Atom: Wave Mechanical Model (Quantum Mechanical Model)

• Based on Bohr's idea of specific energy amounts, Erwin Schrödinger proposed a more complex equation.
• Key Concepts:
  • Electrons reside in defined regions called orbitals, not fixed orbits.
  • At every energy level except the first, multiple orbitals exist.
• Energy Levels and Orbitals:
  • Energy levels have specific orbitals:
    • Level 1: 1s (1 orbital, 2 electrons)
    • Level 2: 2s, 2p (1 s orbital, 3 p orbitals, 2 + 6 = 8 electrons)
    • Level 3: 3s, 3p, 3d (1 s, 3 p, 5 d orbitals, 2 + 6 + 10 = 18 electrons)
    • Level 4: 4s, 4p, 4d, 4f (1 s, 3 p, 5 d, 7 f orbitals, 2 + 6 + 10 + 14 = 32 electrons)
• Pauli Exclusion Principle: Each orbital can hold only two electrons.
• Orbital Shapes: Represent regions where electrons can be for a given energy.
  • As energy level increases, orbitals occupy space farther from the nucleus.

Electron Configurations

• Definition: Indicates the orbitals occupied by electrons in an atom.
• Importance: Understanding electron configuration is crucial for predicting atomic behavior, especially of the outermost electrons.
• Ground State: Electrons occupy the lowest energy orbitals possible.
• Aufbau Principle: Electrons are placed in orbitals starting from the lowest energy level upwards.
• Aufbau Diagram: Illustrates orbitals in order of increasing energy.
  • Note: 4s orbital is slightly below 3d, and 5s is slightly below 4d.
• Pauli Exclusion Principle: Only two electrons per orbital with opposite spin (indicated by up and down arrows).
• Hund’s Rule: When filling multiple orbitals with the same energy, each orbital is occupied with a single electron before pairing occurs.
• Example: Fe (Atomic Number 26)
  • Note application of Hund’s Rule in the 3d orbitals.

Writing Electron Configurations

• Direct Notation: Example: Fluorine (9 electrons) - 1s^22s^22p^5
  • Large number: energy level, Letter: orbital type, Superscript: number of electrons.
• Abbreviated Configurations: Use noble gas configurations to shorten notation.
  • Example: Magnesium (Mg)- Complete: 1s^22s^22p^63s^2 , Abbreviated: \text{{Ne}}3s^2
• Procedure:
  • Find the noble gas preceding the element in the periodic table and write its symbol in brackets.
  • Continue writing the electron configuration from that point.
  • Example: Fe (26 electrons) - Nearest noble gas is Ar (18 electrons): \text{{Ar}}4s^23d^6

Valence Electrons

• Definition: Highest level s and p electrons in an atom (furthest from the nucleus).
• Importance: Used in forming chemical bonds.
• Examples:
  • Na (1s^22s^22p^63s^1): Valence electron configuration is 3s^1
  • Se (1s^22s^22p^63s^23p^64s^23d^{10}4p^4): Valence electron configuration is 4s^24p^4

Mendeleev’s Periodic Table

• Arrangement: Organized elements by increasing atomic mass and similar properties.
• Periodicity: Found that properties repeated periodically.
• Prediction: Successfully predicted undiscovered elements due to gaps in the table.

Modern Periodic Table

• Arrangement: Similar to Mendeleev’s but with elements in order of increasing atomic number.
• Groups: Columns with elements of similar properties and same valence electron configuration.
• Periods: Rows with elements having different properties but valence electrons in the same energy level.
• Examples:
  • Si and Ge (Group 14): Similar properties, valence configurations of 3s^23p^2 and 4s^24p^2 respectively.
  • Na and P (Same period): Different properties, valence electrons in same energy level (3s^1 for Na, 3s^23p^3 for P).

Blocks on the Periodic Table

• s-block: Groups 1 & 2 + He (last electron in s-orbital).
• p-block: Groups 13-18 (last electron in p-orbital).
• d-block: Groups 3-12 (last electron in d-orbital).
• f-block: Elements 57-70 and 89-102 (last electron in f-orbital).
• Using the Periodic Table for Electron Configuration: Start at 1s and add electrons up to the necessary number, noting that the d-block is always one energy level lower than the s and p blocks.

Metals, Non-metals, and Metalloids

• Metals: Shiny, malleable, conduct heat and electricity.
• Non-metals: In the p-block (except H), gases, liquids (Br), or brittle solids, poor conductors.
• Metalloids: In the p-block between metals and non-metals, have properties of non-metals but can behave like metals under certain conditions.

Special Names for Certain Sets of Elements

• Noble Gases: Group 18 (8A), unreactive, complete valence electrons (s^2p^6).
• Halogens: Group 17 (7A), very reactive, electron configuration of s^2p^5 (one electron short of complete).
• Alkali Metals: Group 1, very reactive, electron configuration of s^1 (one electron more than complete).
• Alkaline Earth Elements: Group 2, very reactive, electron configuration of s^2 (two electrons more than complete).
• Transition Metals: d-block elements, Groups 3-12.
• Lanthanides: Elements following La in the f-block.
• Actinides: Elements following Ac in the f-block.

Periodic Trends

• Atomic Size (Atomic Radius): Increases down a group, decreases across a period.