Chemistry Ch. 9

Chapter 9: Electrons in Atoms

9.2 Light: Electromagnetic Radiation

• The interaction of light with atoms played a crucial role in the development of atomic models.

• Definition of Electromagnetic Radiation (Light):

  • A type of energy exhibiting wavelike behavior as it travels through space.

• Types of Electromagnetic Radiation:

  • X-rays

  • Gamma rays

  • Microwaves

  • Infrared radiation

  • Ultraviolet radiation

  • Visible Light

  • Wavelength Range (in meters and nanometers):

  • High energy: Gamma rays (up to 0.01 nm)

  • Visible light: ranges from about 400 nm (violet) to 750 nm (red)

  • Low energy: Radio waves (1 m and above)

• Speed of Light (c):

  • Value: $c = 3.0 imes 10^8 ext{ m/s}$

  • Light demonstrates both wave and particle properties.

Wave Properties of Light

• Wavelength (λ):

  • The distance between adjacent wave crests.

  • SI Unit: Meter

• Diagram:

  • Describes wave properties: crest, trough, direction of travel.

• Color Determination in Visible Light:

  • Wavelength determines the color seen:

  • Red has the longest wavelength.

  • Violet has the shortest wavelength.

  • White light contains all visible wavelengths, leading to various color perceptions.

  • Example:

  • A red shirt appears red because it reflects red light and absorbs other colors.

Frequency and Energy of Light

• Frequency (ν):

  • Defined as the number of cycles or crests passing a stationary point per second.

  • SI Unit: Hertz (Hz)

• Relationship Between Wavelength and Frequency:

  • Expressed by the equation:

  • $c = <br>u imes ext{λ}$

• Example Questions:

  • What wavelength corresponds to a frequency of $4.61 imes 10^{14} ext{ Hz}$?

  • Find the frequency of light with a wavelength of $5.6 imes 10^{-9} ext{ m}$.

• Photon Definition:

  • A massless packet of light energy.

  • Energy content is wavelength-dependent: shorter wavelengths correspond to higher energy.

• Comparison of Light Types:

  • Yellow light has a longer wavelength than violet light, meaning it has less energy per photon.

• Energy of a Photon Calculation:

  • $E_{ ext{photon}} = h <br>u$ where

  • Planck's Constant (h):

    • Value: $h = 6.626 imes 10^{-34} ext{ J} imes ext{s}$

• Practical Example:

  • Calculate energy for violet light with frequency $7.230 imes 10^{14} ext{ Hz}$.

9.4 The Bohr Model: Atoms with Orbits

• When atoms absorb energy (heat, light, electricity), they often emit light.

• Absorption and Emission of Light:

  • Atoms emit light at specific colors due to the interaction between light and the atom's electrons.

• Prism Viewing:

  • Each element emits light at unique wavelengths, forming distinct emission spectra.

• White-Light Spectrum vs. Element Emission Spectrum:

  • White light gives a continuous spectrum, while individual elements provide discrete wavelengths.

The Bohr Model

• Niels Bohr’s Contribution:

  • Developed a model of the atom to explain non-continuous atomic emission spectra.

• Key Features of the Bohr Model:

  • Electrons orbit the nucleus in defined circular paths.

  • Each orbit corresponds to specific energy levels (quantum numbers: n = 1, 2, 3…).

  • Orbit energy is quantized: Electrons cannot exist in between orbits.

• Atomic States:

  • Ground State: Lowest energy state.

  • Excited State: Higher energy state due to absorbed energy.

• Re-emission occurs as electrons transition from excited to lower energy states, emitting photons of light specific to the energy differences involved.

Transitions and Emission

• The emitted light shows discrete lines corresponding to particular transitions between orbits.

• Energy Relationships:

  • Closer orbits release low energy; farther orbits release high energy.

• Hydrogen Transitions:

  • Example transitions include those between n = 3 to n = 2 and n = 3 to n = 1, with differing emitted radiation properties.

• The Bohr model effectively explains hydrogen’s spectrum but fails for multi-electron atoms, leading to the development of more advanced models.

9.5 & 9.6 Quantum Mechanical Model

• Wave-Particle Duality:

  • Electrons can exhibit properties of both waves and particles.

• Electron Path Prediction Limitation:

  • Tracing an electron's exact path is impossible; rather, electron positions are described probabilistically.

• Orbitals Defined:

  • Represent probability distributions where electrons are likely to be found.

  • Effective Volume: Orbitals encompass about 90% probability of finding an electron.

Quantum Mechanic Orbitals

• Principal Quantum Number (n):

  • Describes the electron shell in terms of energy levels (n = 1, 2, 3…).

• Subshells:

  • Indicate orbital shape (s, p, d, f):

  • Each principal shell has a number of subshells equal to n.

• Orbital Shapes and Configurations:

  • 1s Orbital: Lowest energy, closest to nucleus.

  • 2s Orbital: Similar shape but larger than 1s.

  • p Orbitals: Dumbbell shape, three orientations (Px, Py, Pz).

  • d Orbitals: Seasoned with clover-like shapes, five orientations.

  • f Orbitals: Complex shapes inclusive of seven orientations.

• Maximum Capacity of Orbitals:

  • Each can hold up to 2 electrons, categorized as full (2), half full (1), or empty (0).

Electron Configurations

• Definition:

  • Arrangement of electrons within orbitals.

• Key Principles:

  • Aufbau Principle: Fill lowest energy levels first (1s < 2s < 2p…).

  • Pauli Exclusion Principle: No more than two electrons per orbital, emphasizing opposite spins.

  • Hund's Rule: Electrons occupy degenerate (equal energy) orbitals singly before pairing.

• Examples of Electron Configurations:

  • Configurations vary for elements (Li, Cl, etc.), often illustrated via orbital diagrams.

Isoelectronic Species

• Definition:

  • Entities with identical electron counts; often matching a noble gas configuration.

• Ion Configurations:

  • Example includes Al3+ matching the electron configuration of Neon.

• Noble Gas Configuration:

  • Strategy to condense electron configurations by referencing the previous noble gas in the periodic table.

9.7 Electron Configurations and the Periodic Table

• Valence Electrons:

  • Outer shell electrons involved in chemical bonding, distinguished from core electrons.

• Determining Valence Electrons:

  • Methodology illustrated using specific elements like Se and Fe.

• Periodic Patterns:

  • Group properties reflect consistent patterns in valence electron numbers.

• Example of Group Trends:

  • Elements exhibit periodic valence electron configurations determining their chemical properties.

9.8 The Explanatory Power of the Quantum Mechanical Model

• Stability of Atoms:

  • Atoms with 8 valence electrons are especially stable.

  • Elements close to noble gas configurations display high reactivity due to potential electron gain/loss.

• Group Reactivity Examples:

  • Alkali Metals (Group 1):

  • Typically lose one electron, forming 1+ ions corresponding to noble gas configuration.

  • Alkaline Earth Metals (Group 2):

  • Tend to lose two electrons, forming 2+ ions for stability.

  • Halogens (Group 7):

  • Generally gain one electron to reach noble gas configuration, resulting in 1− ions.