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Chapter 4 Arrangement of Electrons in Atoms

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1 Chapter 4 Arrangement of Electrons in Atoms
Section 4-1 Development of a New Atomic Model

2 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)

3 Electromagnetic Spectrum

4 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

5 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

6 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 J∙s)

7 Equation Practice What is the frequency of yellow light with a wavelength of 548 nm?

8 Equation Practice What is the wavelength of blue light with a frequency of 4.60 x 1023 Hz?

9 Equation Practice What is the energy of magenta light with a wavelength of 691 nm?

10 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.

11 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"

12 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?

13 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.

14 Hydrogen Emission Spectrum
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

15 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), dB proved that electrons can have specific energies and that Bohr’s quantized orbits were actually correct

16 Heisenberg Uncertainty Principle
Impossible to determine both the exact location and velocity of an electron

17 Schrodinger Wave Equation
He gave more support to Bohr’s quantized energy levels Quantum theory – describes the wave properties of electrons using mathematical equations Disproved Bohr’s “train tracks” within those energy levels

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