Arrangement of Electrons in the Atom

Objectives

• Define and explain energy levels in atoms.
• Describe and explain the emission spectrum of the hydrogen atom using the Balmer series.
• Describe and explain the absorption spectrum.
• Use flame tests to demonstrate energy absorption/release in discrete units during electron transitions.
• Relate energy levels in atoms to applications like sodium street lights and fireworks.
• Discuss uses of atomic absorption spectrometry (AAS) as an analytical technique.
• Illustrate how line spectra provide evidence for energy levels.
• Use a spectroscope or spectrometer to view emission spectra of elements.
• Define and explain energy sub-levels.
• State the Heisenberg uncertainty principle.
• State the dual wave-particle nature of the electron.
• Define and explain atomic orbitals.
• Describe the shapes of s and p orbitals.

Bohr's Study of Spectra

• White light through a prism creates a continuous spectrum.
• Bohr observed an emission spectrum with narrow, colored lines from Hydrogen Gas Discharge Tube.
• Spectrometers measure light frequency; spectroscopes view spectra.
• Unique emission spectrum for each element helps identify elements.

Flame Tests

• Method: Dip damp wooden splint in salt sample, place in blue Bunsen flame, record flame color.
• Metal Presence - Flame Color:
  • Lithium: Crimson
  • Potassium: Lilac
  • Barium: Green
  • Strontium: Red
  • Copper: Blue-Green
  • Sodium: Yellow

Bohr's Theory: Evidence for Energy Levels

• Energy Level: Fixed energy value an electron may have.
• Ground State: Electrons in the lowest available energy levels.
• Excited State: Electrons in higher energy levels than ground state.
• Process:
  1. Electron in ground state.
  2. Absorbs energy, jumps to excited state.
  3. Unstable excited state.
  4. Falls to lower energy level, releases photon.
  5. Photon energy corresponds to energy difference between levels (E2-E1=hf).
  6. Each light frequency appears as a colored line in the spectrum.
  7. Separate lines indicate discrete energy values.
  8. Visible light emission due to electrons falling to n=2 level (Balmer Series).

Atomic Absorption Spectroscopy (AAS)

• Principles:
  1. White light through gaseous sample of element.
  2. Missing light frequencies appear as dark lines in absorption spectrum.
  3. Dark lines match colored lines in emission spectrum.
• Uses:
  • Instrument: Atomic Absorption Spectrometer.
  • Detects presence and concentration of elements (e.g., lead, chlorine).
  • Light absorption proportional to concentration.

Energy Sublevels

• Each energy level (except n=1) has sublevels.
• Sublevel Definition: Subdivision of a main energy level; consists of orbitals of the same energy.
• Sublevels for Energy Levels:
  • n=1: s
  • n=2: s, p
  • n=3: s, p, d
  • n=4: s, p, d, f (up to this level required for course)

Wave Nature of the Electron

• De Broglie: Electrons have wave and particle properties.
• Heisenberg Uncertainty Principle: Impossible to measure both velocity and position of an electron simultaneously.

Limitations of Bohr's Theory

1. Only explains hydrogen emission spectrum.
2. Omits wave nature of electron.
3. Ignores existence of sublevels.

Atomic Orbitals

• Definition: Region in space with high electron probability.
• Sublevel Orbitals and Characteristics:
  • s: 1 orbital, spherical, 2 electrons max.
  • p: 3 orbitals (px, py, pz), dumbbell-shaped, 6 electrons max.
  • d: Not detailed, 5 orbitals, 10 electrons max.