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.