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This chapter explores the arrangement of electrons in atoms, starting from Rutherford's atomic model and evolving into Bohr's model and the quantum mechanical model, highlighting key concepts like wave-particle duality, the photoelectric effect, and the Heisenberg Uncertainty Principle. It explains the behavior of electromagnetic radiation, including its spectrum, wavelength, and frequency calculations, and how they relate to the energy of photons. Flame tests and their application in identifying metal ions are also discussed, emphasizing the connection between light and atomic structure.
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1st- Light & Electrons Chapter 4Arrangement of Electrons in Atoms
The Development of A New Atomic Model • Rutherford’s model was an improvement over previous models, but still incomplete. • Where exactly are electrons located? • What prevented the electrons from being drawn into the nucleus?
Wave Description Of Light • Electromagnetic Radiation: • form of energy that exhibits wavelike behavior as it travels through space. • EX: visible light, X-ray, Ultraviolet and inferred light, microwaves, and radio waves. • Travels at a constant speed of 3.0 x 108 m/s • Electromagnetic Spectrum: All the electromagnetic radiation form the ES. (fig 4-1, p. 92)
Wave Calculations • Wavelength (λ) - distance between two peaks . Measured in meters • Frequency (v) - number of peaks that pass a point each second. • Hz = Hertz = s-1 • c = λ v • where c = 3.0 x 108 m/s
Is light really a wave? • Max Planck – did experiments with light-matter interactions where light did not act like a wave • Photoelectric Effect - emission of electrons from a metal when light shines on the metal. • Only emitted at certain energies; wave theory said any energy should do it. • Led to the particle theory of light
Planck suggested that objects emit energy in specific amounts called QUANTA • Quantum - minimum quantity of energy that can be lost or gained by an atom. • led Planck to relate the energy of an electron with the frequency of EMR • E = hv • E= Energy (J, of a quantum of radiation) • v= frequency of radiation emitted • h= Planck’s constant (6.626 x 10-34 J∙s)
leads to Einstein’s dual nature of light (EMR behaves as both a wave and a particle) • Photon - particle of EMR having zero mass and carrying a quantum of energy.
Hydrogen Emission Spectrum • Ground State - Lowest energy state of electron. • Excited State - higher energy than ground state. • Bright-line Spectrum (emission spectrum) • Series of specific light frequencies emitted by elements "spectra are the fingerprints of the elements"
Bohr Model Of H Atom • Bohr explained how the electrons stay in the cloud instead of slamming into the nucleus • Definite orbits; paths • The greater the distance from the nucleus, the greater the energy of an electron in that shell.
Electrons start in lowest possible level - ground state. • Absorb energy - become excited and shift upward. • Dropping back down - emits photons (packets of energies equal to the previously absorbed energy). • Hydrogen Emission Spectrum
Quantum Model of the Atom • Bohr’s model was great, but it didn’t answer the question “why?” • Why did electrons have to stay in specific orbits? • Why couldn’t the electrons exist anywhere within the electron cloud? • Louis de Broglie pointed out that electrons act like waves • Using Planck’s equation (E=hv), Louis proved that electrons can have specific energies and that Bohr’s quantized orbits were actually correct
Heisenberg Uncertainty Principle • Impossible to determine both the exact location and velocity of an electron
Schrodinger Wave Equation • He gave more support to Bohr’s quantized energy levels • Quantum theory – describes the wave properties of electrons using mathematical equations
Equation Practice • What is the energy of yellow light with a wavelength of 548 nm? • Convert nm m • Use wave equation to calculate frequency • Use Planck’s equation to calculate energy
Equation Practice • What is the energy of blue light with a wavelength of 460 nm? • Convert nm m • Use wave equation to calculate frequency • Use Planck’s equation to calculate energy
Equation Practice • What is the energy of magenta light with a wavelength of 691 nm? • Convert nm m • Use wave equation to calculate frequency • Use Planck’s equation to calculate energy
2nd- Flame Tests How colors of Fireworks are made
Background Information • electromagnetic radiation: • form of energy that acts as a wave as it travels • includes: visible light, X rays, ultraviolet and infrared light, microwaves, and radio waves • All forms are combined to form electromagnetic spectrum
Background Information • all forms of radiation travel at a speed of 3.0 x 108 m/s (c: speed of light) • wavelength: () • distance between points on adjacent waves • in nm (109 nm = 1 m) • frequency: () • number of waves that passes a point in a second • in waves/second Inversely proportional!
Background Information • we also have an equation to relate Energy (E) of the radiation and frequency () • where h is Planck’s constant: 6.626 x 10-34 J*s
Flame Tests • used to identify metal ions in unknown compounds • usually for Group 1 and 2 metals • when the metal ions in compounds are heated, they release certain wavelengths of visible radiation • colors are dependent on energy of electrons in ion
Step C2: Convert in nm to m • begin with wavelength in nm • use the statement of equality: 1m = 109 nm • Example: 360 nm
Step C3: Find • Example:
Step C4: Find E • Example:
