Chapter 9 Electrons and the Periodic Table
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1 Chapter 9 Electrons and the Periodic Table
2 Quantum Mechanics The Behavior of the Very Small Electrons are incredibly small. Electron behavior determines much of the behavior of atoms. Directly observing electrons in the atom is impossible; the electron is so small that observing it changes its behavior.
3 The Beginnings of Quantum Mechanics The field of quantum mechanics began with the studies of physicists in the early the 20th century. Max Planck (1918) Albert Einstein (1921) Neils Bohr (1922) Arthur Compton (1927) Louis de Broglie (1929) Werner Heisenberg (1932) P. A. M. Dirac (1933) Erwin Schrödinger (1933)
4 The Nature of Light:Its Wave Nature Light is a form of electromagnetic radiation composed of perpendicular oscillating waves, one for the electric field and one for the magnetic field. All electromagnetic waves move through space at the same, constant speed x 10 8 m/s in a vacuum = the speed of light, c
5 Characterizing Waves The amplitude is the height of the wave. The amplitude is a measure of how intense the light is the larger the amplitude, the brighter the light. The wavelength (λ) is a measure of the distance covered by the wave.
6 Characterizing Waves The frequency (ν) is the number of waves that pass a point in a given period of time. The number of waves = number of cycles units are hertz (Hz) or cycles/s = s 1 1 Hz = 1 s 1 The total energy is proportional to the amplitude of the waves and the frequency. The larger the amplitude, the more force it has. The more frequently the waves strike, the more total force.
7 The Relationship Between Wavelength and Frequency The shorter the wavelength, the more frequently waves pass, and the higher the frequency. Wavelength and frequency of electromagnetic waves are inversely proportional. ν 1 λ ν = c λ The proportionality constant is c, the speed of light. Because the speed of light is constant (3.00 x 10 8 m/sec), if we know wavelength we can find the frequency, and vice versa.
8 The Relationship Between Wavelength and Frequency c = 3 x 10 8 m/s
9 Calculate the wavelength of red light with a frequency of 4.62 x s 1.
10 Calculate the frequency (in MHz) of a radio signal with a wavelength of 2.98 m. ν = 3.00 x 10 8 m/sec 2.98 m = 1.01 x 10 8 sec -1 = 1.01 x 10 8 Hz 1.01 x 10 8 Hz x 1.00 MHz = 101 MHz 10 6 Hz
11 The Electromagnetic Spectrum RedOrangeYellowGreenBlueViolet
12 The Electromagnetic Spectrum RedOrangeYellowGreenBlueViolet Shorter wavelengths of energy have higher energy than longer wavelengths: Energy Radiowaves have the lowest energy. Increases Gamma rays have the highest energy.
13 Order the following types of electromagnetic radiation: microwaves, gamma rays, green light, red light, ultraviolet light. By wavelength (short to long) gamma < UV < green < red < microwaves By frequency (low to high) microwaves < red < green < UV < gamma By energy (least to most) microwaves < red < green < UV < gamma
14 White Light Produces a Continuous Spectrum
15 Atomic Spectra When atoms or molecules absorb energy, that energy is often released as light energy. When that emitted light is passed through a prism, a pattern of particular wavelengths of light is seen that is unique to that type of atom or molecule the pattern is called an emission spectrum. non-continuous can be used to identify the material
16 Emission Spectrum Violet line λ = nm Blue line λ = 434 nm Green line λ = 486 nm Red line λ = nm
17 Identifying Elements with Flame Tests Na K Li Ba
18 Emission vs. Absorption Spectra Spectra of Mercury
19 Rutherford s Nuclear Model The atom contains a tiny dense center called the nucleus. The nucleus is essentially the entire mass of the atom. The nucleus is positively charged. The positive charge balances the negative charge of the electrons. The electrons move around in the empty space of the atom surrounding the nucleus.
20 Problems with Rutherford s Nuclear Model of the Atom Electrons are moving charged particles. Moving charged particles give off energy. Therefore electrons should constantly be giving off energy. The electrons should lose energy, crash into the nucleus, and the atom should collapse!!
21 The Bohr Model of the Atom Neils Bohr ( ) The energy of the atom is quantized. The amount of energy in the atom is related to the electron s position in the atom. Quantized means that the atom could only have very specific amounts of energy. Bohr correlated these allowed energy levels with allowed radii of electron orbits.
22 Bohr Model The energy of each Bohr orbit, specified by a quantum number, n = 1, 2, 3 is fixed, or quantized. Bohr orbits are like steps of a ladder, each at a specific distance from the nucleus. It is impossible for an electron to exist between orbits in the Bohr model.
23 Bohr s Model The electrons travel in orbits that are at a fixed distance from the nucleus (stationary states). The energy of the electron is proportional to the distance the orbit was from the nucleus. Electrons emit and absorb radiation when they move between orbits. Emitted radiation is a photon of light. The distance between the orbits determines the energy of the photon of light produced.
24 Bohr s Model
25 Excitation and Emission When an atom absorbs energy, an electron is excited to a higherenergy orbit. The electron relaxes to a lower energy level, emitting a photon of light. n=1 Energy is absorbed!! n=2 n=3 n=4 n=5 Energy is emitted!!
26 Balmer Series Red line λ = nm Green line λ = 486 nm Blue line λ = 434 nm n=1 Violet line λ = nm n=2 n=3 n=4 n=5 n=6 The Hydrogen Spectrum
27 Bohr Model The Bohr model also showed that each principal energy level could hold a maximum number of electrons. This explains the increasing length of the rows of the periodic table.
28 Behavior of main-group elements can be explained in terms of outer electrons.
29 The Bohr Model: Atoms with Orbits The great success of the Bohr model of the atom was that it predicted the lines of the hydrogen emission spectrum. However, it failed to predict the emission spectra of other elements that contained more than one electron. For this and other reasons, the Bohr model was replaced with a more sophisticated model called the quantummechanical or wave-mechanical model.
30 The Quantum-Mechanical Model: Atoms with Orbitals The quantum-mechanical model of the atom replaced the Bohr model in the early twentieth century. In the quantum-mechanical model Bohr orbits are replaced with quantum-mechanical orbitals. Orbitals are different from orbits in that they represent probability maps that show a statistical distribution of where the electron is likely to be found.
31 The Quantum-Mechanical Model: Atoms with Orbitals In the quantum-mechanical model, electrons do not behave like particles flying through space. We cannot, in general, describe their exact paths. An orbital is a probability map (a mathematical model)that shows where the electron is likely to be found when the atom is probed. It does not represent an exact path that an electron takes. Electrons reside in principal energy levels which are subdivided into energy sublevels. The principal energy levels (1-7) can theoretically contain 2,8,18,32,50,72, and 98 electrons each.
32 Principal Energy Levels are Divided Into Subshells The number of subshells in a given principal shell is equal to the value of n.
33 Subshells Hold Different Numbers of Electrons } Electrons in subshells Principal quantum number Electrons in primary shell
34 n=4 Principal Energy Levels Primary energy shells n=3 n=2 n=1
35 Primary energy shells n=4 n=3 n=2 s d p s f d p Energy subshells n=1 s p s
36 6d 6s 6p 5d 5f 5p 5s 4f Energy 4s 4p 3p 4d 3d 3s 2p Degenerate Orbital Energies 2s 1s
37 Real Orbital Energies 7s 6d 6s 6p 5d 5f 5p 4f 5s 4p 4d Energy 4s 3d 3p 3s 2p 2s 1s
38 n=4 n=3 n=2 Bohr Model (shells) Modified Bohr Model (subshells) 4f 4d 4p 3d 4s 3p 3s 2p 2s ] Increasing Energy 4s lower in energy than 3d 5s lower in energy than 4d 6s lower in energy than 4f n=1 1s
39 Filling the Orbitals with Electrons Energy levels and sublevels fill from lowest energy to high. s p d f Aufbau Principle Orbitals that are in the same sublevel have the same energy No more than two electrons per orbital Pauli Exclusion Principle When filling orbitals that have the same energy, place one electron in each before completing pairs. Hund s Rule
40 Electron Configuration & the Periodic Table s d p f
41 Electron Configuration & the Periodic Table s Shell being filled = period number Shell being filled = (period number-1) d p 118 Shell being filled = (period number-2) f
42 Electron Configuration & the Periodic Table s Shell being filled = period number Shell being filled = (period number-1) d p Shell being filled = (period number-2) f What is the highest energy sublevel being filled for each of the following atoms? He K Pd Be Co Si Pt U 1s 4s 4d 2s 3d 3p 5d 5f
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