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Electron Configuration and New Atomic Model Chapter 4

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The Key to the New Atomic Model A connection between light and electrons! Before the 20th century, it was believed that light behaved only like a wave. To fully understand the connection between light and electrons, we must review the properties of light.

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Electromagnetic Radiation A form of energy that exhibits wavelike behavior as it travels through space is called electromagnetic radiation. All forms of this radiation make up the electromagnetic spectrum. All forms move at the speed of light (c= 3.0 x10 8 m/s), through a vacuum and slightly slower through matter. We can assume that it moves that fast through air because it is mostly empty space.

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Electromagnetic Spectrum

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Features of Light Wavelength ( )- is the distance between corresponding points on adjacent waves. Units are either in meters, centimeters or most commonly nanometers. 1 nm= 10 -9 m

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Features of Light Frequency ( )- the number of waves that pass a given point in a specific amount of time, usually in one second Unit- Hertz (waves/second) Higher frequency equals shorter wavelength. Lower frequency equals longer wavelength. They are related by the following equation: c=

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The Photoelectric Effect One experiment completed in the early 1900s challenged the wave theory of interaction between light and matter. The photoelectric effect refers to the emission of electrons from a metal when light shines on it. The wave theory of light predicted that any frequency of light would supply enough energy to eject an electron. However, in this experiment, electrons weren’t emitted if the light’s frequency was below a certain level.

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Particle Description of Light Max Planck- a German physicist who studied the emission of light from hot objects Hot objects do not continuously emit electromagnetic radiation, as they would if they were in the form of waves. Planck suggested that hot objects emitted energy in small, specific amounts called quanta.

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Relationship between Quanta and Frequency of Radiation A quantum is the minimum energy lost or gained by an atom. (plural=quanta) Planck’s proposed relationship: E=h E= energy in Joules of a quantum of radiation, is the frequency, h is a constant, known as Planck’s constant: 6.626 x 10 -34 Js.

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Wave-Particle Duality Einstein took Planck’s idea further and proposed that electromagnetic radiation has a dual wave and particle nature. Light has wave-like properties but can also be thought of as a stream of particles. Each particle of light carries a quantum of energy. He called these particles photons. A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.

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Einstein explains the Photoelectric Effect Electromagnetic radiation is absorbed by matter only in whole numbers. In order for an electron to be ejected from a metal surface, it must be struck by a photon that has the minimum required energy. This corresponds to a minimum frequency. Since electrons are bound more or less closely, depending on the metal, the frequency required to remove an electron is different for different metals.

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Energy of Atoms The lowest energy state of an atom is called its ground state. The state in which the atom has a higher potential energy than its ground state is called its excited state. When an excited state atom returns to the ground state, it gives off energy in the form of electromagnetic radiation, or light.

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Emission Spectra Emission spectra are bands of light at specific frequencies (and therefore wavelengths) that result when the light emitted from an element. This spectra can be observed when the element is in gaseous form and has a current running through it or when it is being burned.

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Emission Line Spectra of Hydrogen What was expected to be observed from hydrogen was a continuous spectrum. Since this was not observed, attempts to explain this resulted in a new theory of the atom called quantum theory.

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Fixed Energy When an excited hydrogen atom returns to its ground state, a photon of radiation is emitted. The energy of the photon is equal to the difference between the initial and final states. Because hydrogen emits light at specific frequencies, the energy differences between these states must be fixed. So a new model of hydrogen must be created!

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Bohr Model of Hydrogen Niels Bohr proposed a model of the atom that linked it’s electron with the photon emission. According to his model, the electron could circle the nucleus in allowed paths called orbits. The electron is in lowest orbit when closest to the nucleus, and therefore has the lowest energy. When the energy of the electron becomes higher, it orbits further from the nucleus.

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Bohr’s Results He matched the spectral lines produced for hydrogen to the energies that hydrogen would allow for electrons. The mathematics supported the experimental data! The discovery of the new model for the hydrogen atom was thought to apply to all other atoms. However, it was soon discovered that his model did not work for atoms with more than one electron.

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Quantum Model of the Atom Scientists struggled with the notion that electrons could only exist in certain orbits with definite energies. They did not understand why there couldn’t be a limitless number of orbits with slightly different energies. Louis de Broglie, a French scientist, investigated this very question in 1924. What he found would change our understanding of matter forever.

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Quantum Model of the Atom De Broglie suggested that electrons be considered waves that are confined to the spaces around the nucleus. His theory about electrons behaving like waves was soon proven to be correct by experimentation. Electrons can be bent or diffracted. Diffraction refers to the bending of a wave as it passes by the edge of an object. It was also shown that interference occurs when waves overlap, causing a slight decrease in energy.

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Heisenberg Uncertainty Principle Scientists wondered where electrons were in the atom, because they could not be detected with photons since the energy of a photon would send an electron off its current path. The Heisenberg Uncertainty principle says that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

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Schrödinger Wave Equation Schrödinger developed an equation based on the theory that electrons have a dual wave-particle nature. Together with the Heisenberg uncertainty principle, the new equation led to modern quantum theory. Quantum theory mathematically describes the wave properties of electrons. Quantum theory determined that wave equations give only the probability of finding an electron at a given place.

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Quantum Numbers Numbers used to specify the properties of atomic orbitals and properties of electrons in those orbitals There are 4 quantum numbers. First three are from Schrödinger's equation 1) Principal Quantum number- (n) indicates the main energy level occupied by an electron -all are whole numbers - as number increases the average distance from the nucleus increases

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Quantum Numbers 2) Angular Momentum Quantum Number- (l) indicates the shape of an orbital The number of orbital shapes possible is equal to n-1. Zero counts… So for a value of n=2, there are a total of 2 l values, l=0 and l=1. The value of l is then assigned a letter: L=0 s L=1 p L=2 d L=3 f

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Orbital Shapes An s orbital is spherical. A p orbital is dumbbell shaped. A d orbital can be clover shaped. An f orbital is really complex.

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Quantum Numbers 3) Magnetic Quantum numbers- (m) indicates the orientation of an orbital around the nucleus In each s sublevel, there is only 1 orientation because it’s a sphere. In a p sublevel, there are 3 orientations because the dumbbell can be aligned on the x, y or z axis. In a d sublevel, there are 5 orientations. In an f sublevel, there are 7 orientations.

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Spin Quantum Numbers Electrons can be thought to spin on an internal axis. It can spin in one of two possible directions. The spin quantum number can have only 2 possible values: +1/2 or -1/2, which indicate direction of spin. A single orbital can hold a maximum of 2 electrons which must have opposite spins!!!

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The Quantum Model of the Atom Section 4.2. Bohr’s Problems Why did hydrogen’s electron exist around the nucleus only in certain allowed orbits? Why couldn’t.

The Quantum Model of the Atom Section 4.2. Bohr’s Problems Why did hydrogen’s electron exist around the nucleus only in certain allowed orbits? Why couldn’t.

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