Module 5

Page 1: Compound and Flame Test Colors

• Flame Test: Analytical procedure to detect the presence of metal ions through flame color.

  • When heated, electrons in metal ions gain energy and transition to higher energy levels.

  • Electrons fall back to lower levels, releasing energy as light of characteristic colors.

Common Metal Ions and Their Flame Colors:

• Lithium (Li"): Red color

• Sodium (Na"): Yellow color

• Potassium (K+): Lilac color

• Rubidium (Rb+): Red-Violet color

• Caesium (Cs+): Blue color

• Calcium (Ca²+): Orange color

• Strontium (Sr²+): Red color

• Barium (Ba²+): Green color

• Radium (Ra²+): Intense color

• Copper (Cu²+): Green/blue color

• Iron (Fe²+/Fe³+): Yellow to brown color

• Boron (B³+): Green color

• Indium (In³+): Blue color

• Lead (Pb²+): Blue color

• Arsenic (As³+): Blue color

• Antimony (Sb³+/Sb5+): Pale blue color

• Selenium (Se²/Se⭑+): Red color

• Zinc (Zn²+): Blue color

• Note: Some metal ion colors are faint and hard to distinguish.

Page 2: Announcements

• Exam 1: Scheduled for next Friday at 4:00 PM covering Modules 1 – 5 with 30 multiple-choice questions.

  • Calculator Requirement: Only approved syllabus calculators allowed.

• Token Opportunity#1: Extended to Sunday night.

• Token Opportunity#2: Open now – Complete the Chemical Naming video and practice assignment on Blackboard by Wednesday, February 19.

Page 3: Module 5 - Electromagnetic Energy and the Bohr Model of the Atom

Page 4: Wave Properties

• Wave Definition: Oscillation or periodic movement transporting energy.

• Described by:

  • Wavelength (λ): Distance between two consecutive crests/troughs of a wave.

  • Frequency (ν): Number of wave cycles passing a stationary point per second (Hz).

  • Amplitude: Half distance between peaks and troughs.

Page 5: Electromagnetic Radiation

• Electromagnetic Spectrum: All types of electromagnetic radiation.

• Speed of light (c): Constant speed for electromagnetic waves in a vacuum.

• Relationship: 𝑐 = 𝜆ν

  • Where 𝑐 = 2.998 × 10^8 m/s.

Page 6: The Particle Nature of Light

• Photon: Quantum of electromagnetic radiation, representing a "chunk" of energy.

• Energy and frequency relationship: 𝑬 = 𝒉ν

  • h: Plank's constant = 6.626 × 10⁻³⁴ J·s.

Page 7: Wavelength, Frequency, and Energy Relationship Summary

• Energy calculated for one photon:

  • 𝑐 = 2.998 × 10^8 m/s

  • 𝑬 = ℎ𝜈

  • 𝜈 = 𝑐/𝜆

  • h = 6.63 × 10⁻³⁴ J·s.

Page 8: Frequency Calculation Example

• Find frequency of electromagnetic radiation with wavelength 530.0 nm:

  • Use formula: 𝜈 = 𝑐/𝜆

  • Calculation: 𝜈 = 2.998 × 10^8 m/s / 5.30 × 10⁻⁷ m = 5.66 × 10¹⁴ s⁻¹

Page 9: Energy Calculation Example

• Calculate energy of one mole of photons of red light (632.8 nm):

  • Use formulas: 𝑬 = ℎν and 𝐸 = (6.63 × 10⁻³⁴ J·s) (3.0 × 10^8 m/s) / 6.328 × 10⁻⁷ m

  • Result: 1.89 × 10² kJ.

Page 10: Photoelectric Effect

• Light directed at metal surface can emit electrons if it has enough energy.

• Demonstrates wave-particle duality of light.

• Equation: 𝐸 = ℎν = ℎ𝑐/𝜆.

Page 11: Spectra Types

• Continuous Spectrum: Unbroken series of wavelengths.

• Line Spectrum: Narrow lines throughout spectral regions.

Page 12: Line Spectra

• Each emission line corresponds to a single wavelength of light, indicating discrete energies of light emitted by gases.

Page 13: Historical Models of the Atom

• Overview of atomic theory development:

  • John Dalton (1803): Indivisible atoms; different elements have varying atoms.

  • J.J. Thomson (1904): Electrons discovered; "Plum Pudding" model.

  • Ernest Rutherford (1911): Nucleus discovery through gold foil experiment.

  • Niels Bohr (1913): Electrons in quantized orbits; proposed modifications.

Page 14: Bohr Model Features

• Electrons orbit nucleus in specific circular orbits with allowed energies.

• The hydrogen atom does not emit energy while in fixed orbits.

• Transition of electrons occurs via photon emission/absorption equal to energy difference between orbits.

Page 15: Bohr’s Model of Hydrogen Atom

• Ground state (n=1): Lowest energy.

  • Energy level equation: 𝐸𝑛 = -𝑘/n².

Page 16: Electron Energy and Distance

• As n increases, energy and distance from nucleus increase; energy levels get closer while orbits expand.

• Excited state: Higher energy level absorption.

Page 17: Energy Transitions

• Photons emitted during electron transitions from higher to lower energy levels.

• Conditions: Photon absorbed when ni < nf, emitted when ni > nf.

Page 18: Hydrogen Emission Spectrum

• Emission lines correspond to transitions from excited states down to lower states.

Page 19: Energy Change from Level Transitions

• Absorbed energy: ΔE positive; emitted energy: ΔE negative.

• Photon energy = magnitude of ΔE.

• Calculate frequency/wavelength using E = hν and λν = c.

Page 20: Emission Spectra of Elements

• Each element has a unique line spectrum.

Page 21: Bohr’s Model Contributions and Shortcomings

• Contributions:

  • Energy of electrons quantized.

  • Ability to transition between levels.

• Shortcomings:

  • Only explains hydrogen's emission spectrum.

  • Circular orbits don't account for electron wave properties.

Page 23: Photon Energy Calculation

• Find energy of a photon with a wavelength of 361 nm:

  • Calculation resulting in energy of 5.51 × 10⁻¹⁹ J.

Page 24: Photon Emission

• Photons emitted when electron falls to lower energy levels.

Page 25: Absorbing Photon Transition

• Ground state transitions viable when a photon absorbs energy to a higher state.

Page 26: Frequency Calculation

• Calculate frequency of photon with energy 1.93 × 10⁻¹⁷ J:

  • Result: 2.91 × 10¹⁶ s⁻¹.