Chapter 7 Atomic Structure and Periodicity Chapter 7

Chapter 7 Table of Contents 7. 1 7. 2 7. 3 7. 4 7.

Section 7. 1 Electromagnetic Radiation Different Colored Fireworks Return to TOC Copyright © Cengage

Section 7. 1 Electromagnetic Radiation Characteristics • Wavelength ( ) – distance between two

Section 7. 1 Electromagnetic Radiation The Nature of Waves Return to TOC Copyright ©

Section 7. 1 Electromagnetic Radiation Classification of Electromagnetic Radiation Return to TOC Copyright ©

Section 7. 2 The Nature of Matter • Energy is quantized. • Electromagnetic radiation

Section 7. 2 The Nature of Matter The Photoelectric Effect Return to TOC Copyright

Section 7. 2 The Nature of Matter • Energy has mass E = mc

Section 7. 3 The Atomic Spectrum of Hydrogen • Continuous spectrum (when white light

$Section 7. 3 The Atomic Spectrum of Hydrogen Refraction of White Light Return to$

Section 7. 3 The Atomic Spectrum of Hydrogen The Line Spectrum of Hydrogen Return

Section 7. 3 The Atomic Spectrum of Hydrogen Significance • Only certain energies are

Section 7. 4 The Bohr Model • Electron in a hydrogen atom moves around

Section 7. 4 The Bohr Model Electronic Transitions in the Bohr Model for the

Section 7. 4 The Bohr Model • Bohr’s model is incorrect. This model only

Section 7. 5 The Quantum Mechanical Model of the Atom • We do not

Section 7. 5 The Quantum Mechanical Model of the Atom Physical Meaning of a

Section 7. 5 The Quantum Mechanical Model of the Atom Relative Orbital Size •

Section 7. 6 Quantum Numbers • Principal quantum number (n) – size and energy

Section 7. 6 Quantum Numbers for the First Four Levels of Orbitals in the

Section 7. 7 Orbital Shapes and Energies 1 s Orbital Return to TOC Copyright

Section 7. 7 Orbital Shapes and Energies Two Representations of the Hydrogen 1 s,

Section 7. 7 Orbital Shapes and Energies 2 px Orbital Return to TOC Copyright

Section 7. 7 Orbital Shapes and Energies 2 py Orbital Return to TOC Copyright

Section 7. 7 Orbital Shapes and Energies 2 pz Orbital Return to TOC Copyright

Section 7. 7 Orbital Shapes and Energies The Boundary Surface Representations of All Three

Section 7. 7 Orbital Shapes and Energies 3 dx -y Orbital 2 2 Return

Section 7. 7 Orbital Shapes and Energies 3 dxy Orbital Return to TOC Copyright

Section 7. 7 Orbital Shapes and Energies 3 dxz Orbital Return to TOC Copyright

Section 7. 7 Orbital Shapes and Energies 3 dyz Orbital Return to TOC Copyright

Section 7. 7 Orbital Shapes and Energies Orbital Return to TOC Copyright © Cengage

Section 7. 7 Orbital Shapes and Energies The Boundary Surfaces of All of the

Section 7. 7 Orbital Shapes and Energies Representation of the 4 f Orbitals in

Section 7. 8 Electron Spin and the Pauli Principle Electron Spin • Electron spin

Section 7. 11 The Aufbau Principle and the Periodic Table Hund’s Rule • The

Section 7. 11 The Aufbau Principle and the Periodic Table Orbital Diagram • A

Section 7. 11 The Aufbau Principle and the Periodic Table Valence Electrons • The

Section 7. 11 The Aufbau Principle and the Periodic Table The Orbitals Being Filled

Section 7. 12 Periodic Trends in Atomic Properties Hund’s Rule • The lowest energy

Section 7. 12 Periodic Trends in Atomic Properties Orbital Diagram • A notation that

Section 7. 12 Periodic Trends in Atomic Properties Valence Electrons • The electrons in

Section 7. 12 Periodic Trends in Atomic Properties The Orbitals Being Filled for Elements

Section 7. 12 Periodic Trends in Atomic Properties Return to TOC Copyright © Cengage

Section 7. 12 Periodic Trends in Atomic Properties • The arrangement of the periodic

Section 7. 12 Periodic Trends in Atomic Properties 여러 전자가 있는 원자에서 한 전자에

Section 7. 13 The Properties of a Group: The Alkali Metals Special Names for

Section 7. 13 The Properties of a Group: The Alkali Metals Versus Nonmetals Return

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Chapter 7 Atomic Structure and Periodicity

Chapter 7 Table of Contents 7. 1 7. 2 7. 3 7. 4 7. 5 7. 6 7. 7 7. 8 7. 9 7. 10 7. 11 7. 12 7. 13 Electromagnetic Radiation The Nature of Matter The Atomic Spectrum of Hydrogen The Bohr Model The Quantum Mechanical Model of the Atom Quantum Numbers Orbital Shapes and Energies Electron Spin and the Pauli Principle Polyelectronic Atoms The History of the Periodic Table The Aufbau Principle and the Periodic Table Periodic Trends in Atomic Properties The Properties of a Group: The Alkali Metals Copyright © Cengage Learning. All rights reserved 2

Section 7. 1 Electromagnetic Radiation Characteristics • Wavelength ( ) – distance between two peaks or troughs in a wave. • Frequency ( ) – number of waves (cycles) per second that pass a given point in space • Speed (c) – speed of light (2. 9979× 108 m/s) Return to TOC Copyright © Cengage Learning. All rights reserved 4

Section 7. 2 The Nature of Matter • Energy has mass E = mc 2 • Dual nature of light: § Electromagnetic radiation (and all matter) exhibits wave properties and particulate properties. Return to TOC Copyright © Cengage Learning. All rights reserved 9

Section 7. 3 The Atomic Spectrum of Hydrogen • Continuous spectrum (when white light is passed through a prism) – contains all the wavelengths of visible light • Line spectrum – each line corresponds to a discrete wavelength: § Hydrogen emission spectrum Return to TOC Copyright © Cengage Learning. All rights reserved 10

$Section 7. 3 The Atomic Spectrum of Hydrogen Refraction of White Light Return to$

Section 7. 3 The Atomic Spectrum of Hydrogen Significance • Only certain energies are allowed for the electron in the hydrogen atom. • Energy of the electron in the hydrogen atom is quantized. Return to TOC Copyright © Cengage Learning. All rights reserved 13

Section 7. 4 The Bohr Model • Electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. • Bohr’s model gave hydrogen atom energy levels consistent with the hydrogen emission spectrum. • Ground state – lowest possible energy state (n = 1) Return to TOC Copyright © Cengage Learning. All rights reserved 14

Section 7. 4 The Bohr Model Electronic Transitions in the Bohr Model for the Hydrogen Atom a) An Energy-Level Diagram for Electronic Transitions Return to TOC Copyright © Cengage Learning. All rights reserved 15

Section 7. 4 The Bohr Model Electronic Transitions in the Bohr Model for the Hydrogen Atom b) An Orbit-Transition Diagram, Which Accounts for the Experimental Spectrum Return to TOC Copyright © Cengage Learning. All rights reserved 16

Section 7. 4 The Bohr Model • Bohr’s model is incorrect. This model only works for hydrogen. • Electrons do not move around the nucleus in circular orbits. Return to TOC Copyright © Cengage Learning. All rights reserved 17

Section 7. 5 The Quantum Mechanical Model of the Atom • We do not know the detailed pathway of an electron. • Heisenberg uncertainty principle: § There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. Δx = uncertainty in a particle’s position Δ(mν) = uncertainty in a particle’s momentum h = Planck’s constant Return to TOC Copyright © Cengage Learning. All rights reserved 18

Section 7. 5 The Quantum Mechanical Model of the Atom Physical Meaning of a Wave Function • The square of the function indicates the probability of finding an electron near a particular point in space. § Probability distribution – intensity of color is used to indicate the probability value near a given point in space. Return to TOC Copyright © Cengage Learning. All rights reserved 19

Section 7. 5 The Quantum Mechanical Model of the Atom Relative Orbital Size • Difficult to define precisely. • Orbital is a wave function. • Picture an orbital as a three-dimensional electron density map. • Hydrogen 1 s orbital: § Radius of the sphere that encloses 90% of the total electron probability. Return to TOC Copyright © Cengage Learning. All rights reserved 20

Section 7. 6 Quantum Numbers • Principal quantum number (n) – size and energy of the orbital. • Angular momentum quantum number (l) – shape of atomic orbitals (sometimes called a subshell). • Magnetic quantum number (ml) – orientation of the orbital in space relative to the other orbitals in the atom. Return to TOC Copyright © Cengage Learning. All rights reserved 21

Section 7. 8 Electron Spin and the Pauli Principle Electron Spin • Electron spin quantum number (ms) – can be +½ or -½. • Pauli exclusion principle - in a given atom no two electrons can have the same set of four quantum numbers. • An orbital can hold only two electrons, and they must have opposite spins. Return to TOC Copyright © Cengage Learning. All rights reserved 36

Section 7. 11 The Aufbau Principle and the Periodic Table Hund’s Rule • The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals. Return to TOC Copyright © Cengage Learning. All rights reserved 37

Section 7. 11 The Aufbau Principle and the Periodic Table Orbital Diagram • A notation that shows how many electrons an atom has in each of its occupied electron orbitals. Oxygen: 1 s 22 p 4 Oxygen: 1 s 2 p Return to TOC Copyright © Cengage Learning. All rights reserved 38

Section 7. 11 The Aufbau Principle and the Periodic Table Valence Electrons • The electrons in the outermost principal quantum level of an atom. 1 s 22 p 6 (valence electrons = 8) • The elements in the same group on the periodic table have the same valence electron configuration. Return to TOC Copyright © Cengage Learning. All rights reserved 39

Section 7. 11 The Aufbau Principle and the Periodic Table The Orbitals Being Filled for Elements in Various Parts of the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 40

Section 7. 12 Periodic Trends in Atomic Properties Hund’s Rule • The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals. Return to TOC Copyright © Cengage Learning. All rights reserved 41

Section 7. 12 Periodic Trends in Atomic Properties Orbital Diagram • A notation that shows how many electrons an atom has in each of its occupied electron orbitals. Oxygen: 1 s 22 p 4 Oxygen: 1 s 2 p Return to TOC Copyright © Cengage Learning. All rights reserved 42

Section 7. 12 Periodic Trends in Atomic Properties Valence Electrons • The electrons in the outermost principal quantum level of an atom. 1 s 22 p 6 (valence electrons = 8) • The elements in the same group on the periodic table have the same valence electron configuration. Return to TOC Copyright © Cengage Learning. All rights reserved 43

Section 7. 12 Periodic Trends in Atomic Properties • The arrangement of the periodic table means that the physical properties of the elements follow a regular pattern. • We can look at the size of atoms, or their atomic radius. • There are two trends for atomic radii: – Atomic radius decreases as you go up a group. – Atomic radius decreases as you go left to right across a period. Return to TOC Copyright © Cengage Learning. All rights reserved 46

Section 7. 12 Periodic Trends in Atomic Properties 여러 전자가 있는 원자에서 한 전자에 미치는 유효 핵전하 (Z Z Z eff eff , effective nuclear charge)는 가려막기상수(shielding constant) s 만큼 감소. =Z-s 0<s<Z Z – 최외각 전자를 제외한 내부전자의 수 Z 내부전자수 Z eff 원자반지름 Na 11 10 1 186 Mg 12 10 2 160 Al 13 10 3 143 Si 14 10 4 132 같은 주기에서 Zeff 가 증가하면, 반지름이 줄어든다 Return to TOC Copyright © Cengage Learning. All rights reserved 47