II. Bohr Model of the Atom (p. 94 - 97) Ch. 4 - Electrons in Atoms II. Bohr Model of the Atom (p. 94 - 97) C. Johannesson

Ernest Rutherford’s Model Discovered dense positive piece at the center of the atom- “nucleus” Electrons would surround and move around it, like planets around the sun Atom is mostly empty space It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

Niels Bohr’s Model Why don’t the electrons fall into the nucleus? Move like planets around the sun. In specific circular paths, or orbits, at different levels. An amount of fixed energy separates one level from another.

The Bohr Model of the Atom I pictured the electrons orbiting the nucleus much like planets orbiting the sun. However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another. Niels Bohr

Bohr Model C. Johannesson

Bohr’s model Energy level of an electron analogous to the rungs of a ladder The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder A quantum of energy is the amount of energy required to move an electron from one energy level to another

B. Bohr Model e- exist only in orbits with specific amounts of energy called energy levels Therefore… e- can only gain or lose certain amounts of energy only certain photons are produced C. Johannesson

Before After Photon wavelength changes Photon Moving Electron Electron velocity changes Fig. 5.16, p. 145

A. Line-Emission Spectrum Hydrogen Gas excited state ENERGY IN PHOTON OUT ground state Atoms are “excited” with an electric current C. Johannesson

These are called the atomic emission spectrum Unique to each element, like fingerprints! Very useful for identifying elements

Light is a Particle? Energy is quantized. Light is a form of energy. Therefore, light must be quantized These smallest pieces of light are called photons. Photoelectric effect? Albert Einstein Energy & frequency: directly related.

Explanation of atomic spectra When we next write electron configurations, we are writing the lowest energy. The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level.

Changing the energy Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

Changing the energy Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be “excited”

Changing the energy As the electron falls back to the ground state, it gives the energy back as light

Changing the energy They may fall down in specific steps Each step has a different energy

This is a simplified explanation! Ultraviolet Visible Infrared The further they fall, more energy is released and the higher the frequency. This is a simplified explanation! The orbitals also have different energies inside energy levels All the electrons can move around.

 = h/mv (from Louis de Broglie) What is light? Light is a particle - it comes in chunks. Light is a wave - we can measure its wavelength and it behaves as a wave If we combine E=mc2 , c=, E = 1/2 mv2 and E = h, then we can get:  = h/mv (from Louis de Broglie) called de Broglie’s equation Calculates the wavelength of a particle.

A. Wave-properties Wavelength () - length of one complete wave Frequency () - # of waves that pass a point during a certain time period hertz (Hz) = 1/s C. Johannesson

A. Waves C. Johannesson

B. EM Spectrum HIGH ENERGY LOW ENERGY C. Johannesson

B. EM Spectrum HIGH LOW ENERGY ENERGY R O Y G. B I V red orange yellow green blue indigo violet C. Johannesson

c =  B. EM Spectrum c: speed of light (3.00  108 m/s) Frequency & wavelength are inversely proportional c =  c: speed of light (3.00  108 m/s) : wavelength (m, nm, etc.) : frequency (Hz) C. Johannesson

Long Wavelength = Low Frequency Low ENERGY Short Wavelength = Wavelength Table Long Wavelength = Low Frequency Low ENERGY Short Wavelength = High Frequency High ENERGY

“wave-particle duality” C. Quantum Theory Einstein (1905) Concluded - light has properties of both waves and particles “wave-particle duality” Photon - particle of light that carries a quantum of energy C. Johannesson

C. Quantum Theory The energy of a photon is proportional to its frequency. E = h E: energy (J, joules) h: Planck’s constant (6.6262  10-34 J·s) : frequency (Hz) C. Johannesson

Practice Problems What is the wavelength of radiation with a frequency of 1.50 x 1013 Hz? w = What is the frequency of radiation with a wavelength of 5.00 x 10-8 m? In what region of the electromagnetic spectrum is this radiation? f =

Examples What is the wavelength of blue light with a frequency of 8.3 x 1015 hz? What is the frequency of red light with a wavelength of 4.2 x 10-5 m? What is the energy of a photon of each of the above?