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Do Now: Take out your vocab 1. What is light? 2. How is it related to x-rays or radio waves?

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**Wave-Particle Duality**

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**What is light? Light consists of electromagnetic waves.**

Electromagnetic radiation includes:

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**Electromagnetic Radiation**

Consist of particles that move as waves of energy Electromagnetic radiation (EM radiation or EMR) is a form of energy emitted and absorbed by charged particles Charge particle has an electric field surrounding it As the particle moves it creates a magnetic field

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**Waves can be described by amplitude, wavelength, and frequency**

Amplitude – height of the wave Wavelength(l) – distance between crests Frequency(n) – number of wave cycles to pass a given point in a certain amount of time Waves moving and 3-d but we’ll model.

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**Draw a Wave in the Box Measure the wavelength and amplitude.**

Then draw waves in box B of the same amplitude but greater frequency. What happened to the wavelength from A to B. How are wavelength and frequency related? How is energy related to frequency?

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**Using the equation: c = ln**

Find the wavelength of a radio wave that is brodcasted at 95.5 x 106Hz Find the frequency of blue light. Blue light has a wavelength of 6.43x 10-7m.

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Which are not visible? Which are more energetic than visible light? How do you know? What do all the rays in the electromagnetic spectrum have in common?

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**Compare the atomic models of Thomson and Rutherford. **

Do Now: Compare the atomic models of Thomson and Rutherford. Explain the Bohr model of the atom What colors make up the continuous spectrum of visible light? p362

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**Neils Bohr Studied hydrogen and its emission spectrum**

Proposed the Planetary Model Electrons orbit the nucleus Electrons travel in successively larger orbits and when an electron jumps from an outer orbit to an inner one, it emits light.

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CONTINUOUS SPECTRUM

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**Each element emits a unique collection of lines.**

The atomic spectrum can be used to determine the composition of a material, since it is different for each element of the periodic table.

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**Why do we so only certain lines of color?**

Each color has a specific wavelength Each wavelength is associated with a specific amount of energy That energy is released when the electron jumps from a higher energy level to a lower energy level. Only specific energy levels are allowed in the atom! Energy of the atom is quantized!

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Chapter 5, Figure 5.4 Electrons absorb a specific amount of energy to move to a higher energy level. When electrons lose energy, photons with specific energies are emitted. Energy of quantum of electromagnetic radiation is proportional to frequency

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**So the electron behaves like a wave, but where is it?**

Werner Heisenberg proposed Heisenberg uncertainty principle There is a limitation to knowing where the electron is (its position) and where its going (its momentum) Erwin Schrodinger developed a mathematical equation to describe the electron’s wave-like behavior

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**Probability of finding electron at different points is calculated **

Some points will have higher probability than others If connect all points of high probability, three dimensional shapes are formed The most probable place to find the electron will be some place in that shape Atomic orbitals

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Zip Code 07080 The fourth and fifth digits representing a group of delivery addresses within that region First digit (0-9) represents a group of states in the US. 0 is northeastern states and 9 is western states Second and third numbers represent a region in that group, perhaps a large city Quantum numbers are used to describe locations of high probability of finding the electrons

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**3. Last to an orbital. 6d __ __ __ __ __ Orbital Diagram**

5f __ __ __ __ __ __ __ 7s __ 1. Assign electrons to an Energy Level 6p __ __ __ 5d __ __ __ __ __ 4f __ __ __ __ __ __ __ 2. Next to a sublevel: s p d f 6s __ 3. Last to an orbital. 5p __ __ __ 4d __ __ __ __ __ 5s __ 4p __ __ __ 3d __ __ __ __ __ 4s __ 3p __ __ __ 3s __ 2p __ __ __ 2s __ 1s __

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Atomic Orbitals Region around nucleus where electrons are likely to be found Each orbital holds two electrons These electrons have opposite spin

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6d __ __ __ __ __ 5f __ __ __ __ __ __ __ 7s __ 6p __ __ __ 5d __ __ __ __ __ 4f __ __ __ __ __ __ __ 6s __ 5p __ __ __ 4d __ __ __ __ __ 5s __ 4p __ __ __ 3d __ __ __ __ __ 4s __ 3p __ __ __ 3s __ 2p __ __ __ 2s __ 1s __

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**Follow 3 rules to configure the electrons**

1. Aufbau Principle - electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states 2. Pauli Exclusion Principle - two electrons cannot share the same set of quantum numbers within the same system. Therefore, there is room for only two electrons in each orbital and the electrons have opposite spin. 3. Hund’s Rule – in equal energy orbitals, arrange the electrons to achieve the maximum number of unpaired electrons.

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