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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Light and the Electromagnetic Spectrum**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Light and the Electromagnetic Spectrum Electromagnetic energy (“light”) is characterized by wavelength frequency amplitude Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 This is the electromagnetic spectrum. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Light and the Electromagnetic Spectrum**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Light and the Electromagnetic Spectrum Wavelength x Frequency = Speed x = c m s 1 s m c is defined to be the rate of travel of all electromagnetic energy in a vacuum and is a constant value—speed of light. The speed of light is defined to be x 108 m/s. The units for frequency are also called “hertz.” Wavelength and frequency are inversely proportional to each other. s m c = 3.00 x 108 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Light and the Electromagnetic Spectrum**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Light and the Electromagnetic Spectrum The light blue glow given off by mercury streetlamps has a wavelength of 436 nm. What is the frequency in hertz? Wavelength () x Frequency (f) = Speed of light x = c s m 3.00 x 108 c = = You could compare this numerical value to the visible spectrum to show the students that it matches, color-wise. 1 x 109 nm 1 m 436 nm = 6.88 x 1014 s-1 = 6.88 x 1014 Hz Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electromagnetic Energy and Atomic Line Spectra**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electromagnetic Energy and Atomic Line Spectra Excite an atom and what do you get? Light emission Line Spectrum: A series of discrete lines on an otherwise dark background as a result of light emitted by an excited atom. The Bunsen burner was used for elemental analysis. Gas discharge tubes can be used along with inexpensive diffraction grating glasses (search on “rainbow glasses”) for a fun demonstration. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 From left to right- excited hydrogen and excited neon. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electromagnetic Energy and Atomic Line Spectra**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electromagnetic Energy and Atomic Line Spectra Johann Balmer in 1885 discovered a mathematical relationship for the four visible lines in the atomic line spectra for hydrogen. 1 = R∞ n2 22 - Johannes Rydberg later modified the equation to fit every line in the entire spectrum of hydrogen. (m & n are integers where n>m) The top equation is the Balmer equation while the bottom one is the Balmer-Rydberg equation. 1 = R∞ n2 m2 - R (Rydberg Constant) = x 10-2 nm-1 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Particle-like Properties of Electromagnetic Energy**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Particle-like Properties of Electromagnetic Energy Photoelectric Effect: Irradiation of clean metal surface with light causes electrons to be ejected from the metal. Furthermore, the frequency of the light used for the irradiation must be above some threshold value, which is different for every metal. The book uses a nice analogy of the effect of throwing a thousand ping-pong balls at a window versus 1 baseball. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Particle-like Properties of Electromagnetic Energy**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Particle-like Properties of Electromagnetic Energy Copyright © 2010 Pearson Prentice Hall, Inc.

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**Particle-like Properties of Electromagnetic Energy**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Particle-like Properties of Electromagnetic Energy Einstein explained the effect by assuming that a beam of light behaves as if it were a stream of particles called photons. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Particle-like Properties of Electromagnetic Energy**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Particle-like Properties of Electromagnetic Energy E = h E h (Planck’s constant) = x J s Energy and frequency are directly proportional to each other. Electromagnetic energy (light) is quantized. Quantum: The amount of energy corresponding to one photon of light. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Particle-like Properties of Electromagnetic Energy**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Particle-like Properties of Electromagnetic Energy Niels Bohr proposed in 1914 a model of the hydrogen atom as a nucleus with an electron circling around it. In this model, the energy levels of the orbits are “quantized” that is, only certain specific orbits corresponding to certain specific energies are available for the electron. This is often called the “planetary model.” A major problem which had to be overcome was that bodies in orbit constantly radiate energy in classical mechanics. Think about what eventually happens to satellites as they orbit the Earth. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Particle-like Properties of Electromagnetic Energy**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Particle-like Properties of Electromagnetic Energy Copyright © 2010 Pearson Prentice Hall, Inc.

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**Wave-like Properties of Matter**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Wave-like Properties of Matter Louis de Broglie in 1924 suggested that, if light can behave in some respects like matter, then perhaps matter can behave in some respects like light. In other words, perhaps matter is wave-like as well as particle-like. mv h = The de Broglie equation allows the calculation of a “wavelength” of an electron or of any particle or object of mass m and velocity v. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Quantum Mechanics and the Heisenberg Uncertainty Principle**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Quantum Mechanics and the Heisenberg Uncertainty Principle In 1926 Erwin Schrödinger proposed the quantum mechanical model of the atom which focuses on the wavelike properties of the electron. In 1927 Werner Heisenberg stated that it is impossible to know precisely where an electron is and what path it follows—a statement called the Heisenberg uncertainty principle. What would be needed to “see” an electron? Light would have to interact with the electron. But this interaction would affect the electron and thus changes it. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Wave Functions and Quantum Numbers**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Wave Functions and Quantum Numbers Probability of finding electron in a region of space (2) Wave equation Wave function or orbital () solve A wave function is characterized by three parameters called quantum numbers: n, l, m. Since we can’t ever be certain of the electron’s position, we work with probabilities. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Wave Functions and Quantum Numbers**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Wave Functions and Quantum Numbers Principal Quantum Number (n) Describes the size and energy level of the orbital Commonly called the shell Positive integer (n = 1, 2, 3, 4, …) As the value of n increases: The energy of the electron increases The average distance of the electron from the nucleus increases Copyright © 2010 Pearson Prentice Hall, Inc.

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**Wave Functions and Quantum Numbers**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Wave Functions and Quantum Numbers Angular-Momentum Quantum Number (l) Defines the three-dimensional shape of the orbital Commonly called the sub-shell There are n different shapes for orbitals If n = 1, then l = 0 If n = 2, then l = 0 or 1 If n = 3, then l = 0, 1, or 2 Commonly referred to by letter (sub-shell notation) l = 0 s (sharp) l = 1 p (principal) l = 2 d (diffuse) l = 3 f (fundamental) After f, the series goes alphabetically (g, h, etc.). Copyright © 2010 Pearson Prentice Hall, Inc.

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**Wave Functions and Quantum Numbers**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Wave Functions and Quantum Numbers Magnetic Quantum Number (ml ) Defines the spatial orientation of the orbital There are 2l + 1 values of ml and they can have any integral value from -l to +l If l = 0 then ml = 0 If l = 1 then ml = -1, 0, or 1 If l = 2 then ml = -2, -1, 0, 1, or 2 etc. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Wave Functions and Quantum Numbers**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Wave Functions and Quantum Numbers Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 The Shapes of Orbitals Node: A surface of zero probability for finding the electron. s orbitals are spherical and penetrate closer to the nucleus. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 The Shapes of Orbitals Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 The Shapes of Orbitals p orbitals are dumbbell-shaped with a nodal plane running through the nucleus. Introductory Physical Chemistry courses often explore the mathematical relationships which define the regions of space (or shape). Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Quantum Mechanics and Atomic Line Spectra**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Quantum Mechanics and Atomic Line Spectra Copyright © 2010 Pearson Prentice Hall, Inc.

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**Quantum Mechanics and Atomic Line Spectra**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Quantum Mechanics and Atomic Line Spectra Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Spin and the Pauli Exclusion Principle**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Spin and the Pauli Exclusion Principle Electrons have spin which gives rise to a tiny magnetic field and to a spin quantum number (ms). An electron can have a spin quantum number of + or - 1/2. Another way to think about it is that 2 electrons in an orbital must be spin-paired. Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Orbital Energy Levels in Multi-electron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Orbital Energy Levels in Multi-electron Atoms Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron. Zeff = Zactual - Electron shielding Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multi-electron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multi-electron Atoms Electron Configuration: A description of which orbitals are occupied by electrons. Degenerate Orbitals: Orbitals that have the same energy level. For example, the three p orbitals in a given subshell. Ground-State Electron Configuration: The lowest-energy configuration. Aufbau Principle (“building up”): A guide for determining the filling order of orbitals. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multi-electron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multi-electron Atoms Rules of the aufbau principle: Lower-energy orbitals fill before higher-energy orbitals. An orbital can only hold two electrons, which must have opposite spins (Pauli exclusion principle). If two or more degenerate orbitals are available, follow Hund’s rule. Electrons repel each other. Electrons in different orbitals (and thus different spatial regions) will experience fewer repulsive forces than electrons placed into the same orbital. Hund’s Rule: If two or more orbitals with the same energy are available, one electron goes into each until all are half-full. The electrons in the half-filled orbitals all have the same value of their spin quantum number. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multi-electron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multi-electron Atoms Electron Configuration H: 1s1 1 electron s orbital (l = 0) n = 1 Electron configurations show the distribution of the electrons between the subshells. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multi-electron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multi-electron Atoms Electron Configuration H: 1s1 He: 1s2 2 electrons s orbital (l = 0) An s subshell can hold 2 electrons since there is only 1 orbital. n = 1 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multi-electron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multi-electron Atoms Electron Configuration H: 1s1 He: 1s2 Lowest energy to highest energy Fill up the subshell before moving to the next one. You can refer to the orbital filling diagram, but it’s easier to use the periodic table as described in section 5.14 of the text. Li: 1s2 2s1 1 electrons s orbital (l = 0) n = 2 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multi-electron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multi-electron Atoms Electron Configuration H: 1s1 He: 1s2 Li: 1s2 2s1 N: 1s2 2s2 2p3 3 electrons p orbital (l = 1) n = 2 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Orbital-Filling Diagram H: 1s1 1s He: 1s2 Orbital-filling diagrams show the distribution of the electrons between the orbitals. 1 “line” for each orbital. An s subshell has only 1 orbital. Li: 1s2 2s1 N: 1s2 2s2 2p3 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Orbital-Filling Diagram H: 1s1 1s 1s He: 1s2 Pauli exclusion principle. 2 electrons in the same orbital must be spin-paired. Li: 1s2 2s1 N: 1s2 2s2 2p3 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Orbital-Filling Diagram H: 1s1 1s 1s He: 1s2 1s 2s Li: 1s2 2s1 N: 1s2 2s2 2p3 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Orbital-Filling Diagram H: 1s1 1s 1s He: 1s2 1s 2s Li: 3 “lines” since there are 3 orbitals in a p subshell. They follow Hund’s rule for putting multiple electrons into a degenerate set of orbitals. 1s2 2s1 1s 2s 2p N: 1s2 2s2 2p3 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Shorthand Configuration Na: 1s2 2s2 2p6 3s1 [Ne] 3s1 Ne configuration Think of it as a mathematical substitution. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Shorthand Configuration Na: 1s2 2s2 2p6 3s1 [Ne] 3s1 P: 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Shorthand Configuration Na: 1s2 2s2 2p6 3s1 [Ne] 3s1 P: 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3 K: 1s2 2s2 2p6 3s2 3p6 4s1 [Ar] 4s1 Ar configuration Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations of Multielectron Atoms**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations of Multielectron Atoms Electron Configuration Shorthand Configuration Na: 1s2 2s2 2p6 3s1 [Ne] 3s1 P: 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3 K: 1s2 2s2 2p6 3s2 3p6 4s1 [Ar] 4s1 [Ar] 4s2 3d1 Sc: 1s2 2s2 2p6 3s2 3p6 4s2 3d1 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Some Anomalous Electron Configurations**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Some Anomalous Electron Configurations Expected Configuration Actual Configuration Cr: [Ar] 4s2 3d4 [Ar] 4s1 3d5 Cu: [Ar] 4s2 3d9 [Ar] 4s1 3d10 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations and the Periodic Table**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations and the Periodic Table Valence Shell: Outermost shell. Li: 2s1 Na: 3s1 Note that in the U.S. system for numbering Main Group element columns, the column number equals the number of valence electrons. Note the 8 electrons for 8a (ignoring helium). This is where the “octet rule” comes from. Cl: 3s2 3p5 Br: 4s2 4p5 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations and Periodic Properties: Atomic Radii**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations and Periodic Properties: Atomic Radii Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations and Periodic Properties: Atomic Radii**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations and Periodic Properties: Atomic Radii Atomic radii increases down a column because successively larger valence-shell orbitals are occupied. Atomic radii decreases from left-to-right because the effective nuclear charge increases. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 This shows that size varies as Period (row): decreases from left to right. Column: increases from top to bottom. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Electron Configurations and Periodic Properties: Atomic Radii**

Chapter 3: Periodicity and the Electronic Structure of Atoms 4/22/2017 Electron Configurations and Periodic Properties: Atomic Radii column radius radius row Atomic radii increases down a column because successively larger valence-shell orbitals are occupied. Atomic radii decreases from left-to-right because the effective nuclear charge increases. Copyright © 2010 Pearson Prentice Hall, Inc.

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**Chapter 3: Periodicity and the Electronic Structure of Atoms**

4/22/2017 Copyright © 2010 Pearson Prentice Hall, Inc.

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