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Chapter 61 Electronic Structure of Atoms Chapter 6
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2 The Wave Nature of Light -Visible light is a small portion of the electromagnetic spectrum
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Chapter 63 The Wave Nature of Light Frequency (v, nu – The number of times per second that one complete wavelength passes a given point. Wavelength ( lambda) – The distance between identical points on successive waves. v = c c = speed of light, 2.997 x 10 8 m/s
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Chapter 64 The Wave Nature of Light -We can also say that light energy is quantized -This is used to explain the light given-off by hot objects. -Max Plank theorized that energy released or absorbed by an atom is in the form of “chunks” of light (quanta). E = h v h = plank’s constant, 6.63 x 10 -34 J/s - Energy must be in packets of (hv), 2(hv), 3(hv), etc.
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Chapter 65 Quantized Energy and Photons The Photoelectric Effect
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Chapter 66 Quantized Energy and Photons -The photoelectric effect provides evidence for the particle nature of light. -It also provides evidence for quantization. -If light shines on the surface of a metal, there is a point at which electrons are ejected from the metal. -The electrons will only be ejected once the threshold frequency is reached. -Below the threshold frequency, no electrons are ejected. -Above the threshold frequency, the number of electrons ejected depend on the intensity of the light. The Photoelectric Effect
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Chapter 67 Quantized Energy and Photons -Einstein assumed that light traveled in energy packets called photons. -The energy of one photon, E = h. -This equation means that the energy of the photon is proportional to its frequency. The Photoelectric Effect
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Chapter 68 Bohr’s Model of the Hydrogen Atom Line Spectra
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Chapter 69 Bohr’s Model of the Hydrogen Atom Line Spectra
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Chapter 610 Bohr’s Model of the Hydrogen Atom Line Spectra Line spectra can be “explained” by the following equation:
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Chapter 611 Bohr’s Model of the Hydrogen Atom Line Spectra Line spectra can be “explained” by the following equation: - this is called the Rydberg equation for hydrogen.
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Chapter 612 Bohr’s Model of the Hydrogen Atom -Assumed that a single electron moves around the nucleus in a circular orbit. -The energy of a given electron is assumed to be restricted to a certain value which corresponds to a given orbit. Bohr’s Model
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Chapter 613 Bohr’s Model of the Hydrogen Atom -Assumed that a single electron moves around the nucleus in a circular orbit. -The energy of a given electron is assumed to be restricted to a certain value which corresponds to a given orbit. k = 2.179 x 10 -18 Jz = atomic number n = integer for the orbit Bohr’s Model
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Chapter 614 Bohr’s Model of the Hydrogen Atom -Assumed that a single electron moves around the nucleus in a circular orbit. -The energy of a given electron is assumed to be restricted to a certain value which corresponds to a given orbit. n = integer for the orbita o = 0.0529 angstroms z = atomic number Bohr’s Model
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Chapter 615 Bohr’s Model of the Hydrogen Atom -Quantitized energy and angular momentum -The first orbit in the Bohr model has n = 1 and is closest to the nucleus. -The furthest orbit in the Bohr model has n close to infinity and corresponds to zero energy. -Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta (h ). Bohr’s Model – Important Features
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Chapter 616 Bohr’s Model of the Hydrogen Atom Ground State – When an electron is in its lowest energy orbit. Excited State – When an electron gains energy from an outside source and moves to a higher energy orbit. Bohr’s Model – Line Spectra
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Chapter 617 Bohr’s Model of the Hydrogen Atom Bohr’s Model – Line Spectra
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Chapter 618 Bohr’s Model of the Hydrogen Atom Bohr’s Model – Line Spectra
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Chapter 619 Bohr’s Model of the Hydrogen Atom Bohr’s Model – Line Spectra
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Chapter 620 Bohr’s Model of the Hydrogen Atom Bohr’s Model – Line Spectra
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Chapter 621 Bohr’s Model of the Hydrogen Atom -Since the energy states are quantized, the light emitted from excited atoms must be quantized and appear as line spectra. Bohr’s Model
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Chapter 622 The Wave Behavior of Matter -DeBroglie proposed that there is a wave/particle duality. -Knowing that light has a particle nature, it seems reasonable to assume that matter has a wave nature. -DeBroglie proposed the following equation to describe the relationship:
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Chapter 623 The Wave Behavior of Matter -DeBroglie proposed that there is a wave/particle duality. -Knowing that light has a particle nature, it seems reasonable to assume that matter has a wave nature. -DeBroglie proposed the following equation to describe the relationship:
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Chapter 624 The Wave Behavior of Matter -DeBroglie proposed that there is a wave/particle duality. -Knowing that light has a particle nature, it seems reasonable to assume that matter has a wave nature. -DeBroglie proposed the following equation to describe the relationship: -The momentum, mv, is a particle property, where as is a wave property.
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Chapter 625 The Wave Behavior of Matter Heisenberg’s Uncertainty Principle - on the mass scale of atomic particles, we cannot determine exactly the position, speed, and direction of motion simultaneously. -For electrons, we cannot determine their momentum and position simultaneously. The Uncertainty Principle
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Chapter 626 Quantum Mechanics -These theories (wave/particle duality and the uncertainty principle) mean that the Bohr model needs to be refined. Quantum Mechanics
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Chapter 627 Quantum Mechanics -The path of an electron can no longer be described exactly, now we use the wavefunction( ). Wavefunction ( ) – A mathematical expression to describe the shape and energy of an electron in an orbit. -The probability of finding an electron at a point in space is determined by taking the square of the wavefunction: Probability density =
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Chapter 628 Quantum Mechanics -The use of wavefunctions generates four quantum numbers. Principal Quantum Number (n) - This is the same as Bohr’s n - Allowed values: 1, 2, 3, 4, … (integers) - The energy of an orbital increases as n increases - A shell contains orbitals with the same value of n Quantum Numbers
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Chapter 629 Quantum Mechanics Secondary (Azimuthal) Quantum Number (l) -Allowed values: 0, 1, 2, 3, 4,., (n – 1) (integers) -Each l represents an orbital type Quantum Numbers
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Chapter 630 Quantum Mechanics Secondary (Azimuthal) Quantum Number (l) -Allowed values: 0, 1, 2, 3, 4,., (n – 1) (integers) -Each l represents an orbital type lorbital 0s 1p 2d 3f Quantum Numbers
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Chapter 631 Quantum Mechanics Secondary (Azimuthal) Quantum Number (l) -Allowed values: 0, 1, 2, 3, 4,., (n – 1) (integers) -Each l represents an orbital type -Within a given value of n, types of orbitals have slightly different energy s < p < d < f Quantum Numbers
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Chapter 632 Quantum Mechanics Magnetic Quantum Number (m l ). -This quantum number depends on l. -Allowed values: -l +l by integers. -Magnetic quantum number describes the orientation of the orbital in space. Quantum Numbers
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Chapter 633 Quantum Mechanics Magnetic Quantum Number (m l ). -This quantum number depends on l. -Allowed values: -l +l by integers. -Magnetic quantum number describes the orientation of the orbital in space. lOrbitalmlml 0s0 1p - 1, 0, + 1 2d - 2, - 1, 0, + 1, + 2 Quantum Numbers
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Chapter 634 Quantum Mechanics Magnetic Quantum Number (m l ). -This quantum number depends on l. -Allowed values: -l +l by integers. -Magnetic quantum number describes the orientation of the orbital in space. -A subshell is a group of orbitals with the same value of n and l. Quantum Numbers
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Chapter 635 Quantum Mechanics Spin Quantum Number (s) -Allowed values: - ½ + ½. -Electrons behave as if they are spinning about their own axis. -This spin can be either clockwise or counter clockwise. Quantum Numbers
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Chapter 636 Quantum Mechanics Orbitals and Quantum Numbers
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Chapter 637 Representation of Orbitals -All s-orbitals are spherical. -As n increases, the s-orbitals get larger. -As n increases, the number of nodes increase. -A node is a region in space where the probability of finding an electron is zero. The s Orbitals
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Chapter 638 Representation of Orbitals The s Orbitals
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Chapter 639 Representation of Orbitals -There are three p-orbitals p x, p y, and p z. (The letters correspond to allowed values of m l of -1, 0, and +1.) -The orbitals are dumbbell shaped. The p Orbitals
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Chapter 640 Representation of Orbitals The p Orbitals
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Chapter 641 Representation of Orbitals -There are 5 d- and 7 f-orbitals. -Four of the d-orbitals have four lobes each. -One d-orbital has two lobes and a collar. The d and f Orbitals
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Chapter 642 Representation of Orbitals The d and f Orbitals
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Chapter 643 Orbitals in Many Electron Atoms Effective nuclear charge - The charge experienced by an electron in a many-electron atom. Effective Nuclear Charge
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Chapter 644 Orbitals in Many Electron Atoms -Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge. -The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons. -As the average number of screening electrons (S) increases, the effective nuclear charge (z eff ) decreases. -As the distance from the nucleus increases, S increases and z eff decreases. z eff = z - S Effective Nuclear Charge
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Chapter 645 Orbitals in Many Electron Atoms -Orbitals can be ranked in terms of energy to yield an Aufbau diagram. -As n increases, note that the spacing between energy levels becomes smaller. Orbitals and Quantum Numbers
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Chapter 646 Orbitals in Many Electron Atoms
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Chapter 647 Orbitals in Many Electron Atoms - Pauli’s Exclusions Principle - no two electrons can have the same set of 4 quantum numbers. -Therefore, two electrons in the same orbital must have opposite spins. Pauli Exclusion Principle
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Chapter 648 Electron Configurations Electron configurations tells us in which orbitals the electrons for an element are located. Three rules: -Electrons fill orbitals starting with lowest n and moving upwards. -no two electrons can fill one orbital with the same spin (Pauli Exclusion Principle). -for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule).
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Chapter 649 Electron Configurations
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Chapter 650 Electron Configurations and the Periodic Table
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Chapter 651 Electron Configurations and the Periodic Table -There is a shorthand way of writing electron configurations -Write the core electrons corresponding to the filled Noble gas in square brackets. -Write the valence electrons explicitly. Example, P: 1s 2 2s 2 2p 6 3s 2 3p 3 but Ne is 1s 2 2s 2 2p 6 Therefore, P: [Ne]3s 2 3p 3
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Chapter 652 Homework problems: 6.24, 6.26, 6.34, 6.40, 6.42, 6.44, 6.52, 6.60, 6.64
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