Chapter 7 - Atomic Structure and Periodicity (1).ppt PDF

Chapter 7 Atomic Structure and Periodicity Section 7.1 Electromagnetic Radiation Different Colored Fireworks Copyright © Cengage Learning. All rights reserved 2 Section 7.1 Electromagnetic Radiation Questions to Consider  Why do we get colors?  Why do different chemic...

Chapter 7 Atomic Structure and Periodicity Section 7.1 Electromagnetic Radiation Different Colored Fireworks Copyright © Cengage Learning. All rights reserved 2 Section 7.1 Electromagnetic Radiation Questions to Consider  Why do we get colors?  Why do different chemicals give us different colors? Copyright © Cengage Learning. All rights reserved 3 Section 7.1 Electromagnetic Radiation Electromagnetic Radiation  One of the ways that energy travels through space.  Three characteristics:  Wavelength  Frequency  Speed Copyright © Cengage Learning. All rights reserved 4 Section 7.1 Electromagnetic Radiation Characteristics  Wavelength ( ) – distance between two consecutive 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) c =  Copyright © Cengage Learning. All rights reserved 5 Section 7.1 Electromagnetic Radiation The Nature of Waves 6 Section 7.1 Electromagnetic Radiation Classification of Electromagnetic Radiation Copyright © Cengage Learning. All rights reserved 7 Section 7.2 The Nature of Matter Pickle Light Copyright © Cengage Learning. All rights reserved 8 Section 7.2 The Nature of Matter  Energy can be gained or lost only in whole hν number multiples of.  A system can transfer energy only in whole quanta (or “packets”).  Energy seems to have particulate properties too. Copyright © Cengage Learning. All rights reserved 9 Section 7.2 The Nature of Matter  Energy is quantized.  Electromagnetic radiation is a stream of “particles” called photons. hc Ephoton = hν =   Planck’s constant = h = 6.626 × 10-34 Js Copyright © Cengage Learning. All rights reserved 10 Section 7.2 The Nature of Matter  Energy has massE = mc2  Dual nature of light:  Electromagnetic radiation (and all matter) exhibits wave properties and particulate properties. Copyright © Cengage Learning. All rights reserved 11 Section 7.3 The Atomic Spectrum of Hydrogen  Continuous spectrum (results 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 Copyright © Cengage Learning. All rights reserved 12 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. Copyright © Cengage Learning. All rights reserved 13 Section 7.3 The Atomic Spectrum of Hydrogen CONCEPT CHECK! Why is it significant that the color emitted from the hydrogen emission spectrum is not white? How does the emission spectrum support the idea of quantized energy levels? Copyright © Cengage Learning. All rights reserved 14 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) Copyright © Cengage Learning. All rights reserved 15 Section 7.4 The Bohr Model Electronic Transitions in the Bohr Model for the Hydrogen Atom a) An Energy-Level Diagram for Electronic Transitions Copyright © Cengage Learning. All rights reserved 16 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 Copyright © Cengage Learning. All rights reserved 17 Section 7.4 The Bohr Model  For a single electron transition from one energy level to another:  1 1  E =  2.178 10 J  2   18  n  final n 2 initial  ΔE = change in energy of the atom (energy of the emitted photon) nfinal = integer; final distance from the nucleus ninitial = integer; initial distance from the nucleus Copyright © Cengage Learning. All rights reserved 18 Section 7.4 The Bohr Model  The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electron.  As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state (free electron). As the electron is brought closer to the nucleus, energy is released from the system. Copyright © Cengage Learning. All rights reserved 19 Section 7.4 The Bohr Model  Bohr’s model is incorrect. This model only works for hydrogen.  Electrons move around the nucleus in circular orbits. Copyright © Cengage Learning. All rights reserved 20 Section 7.4 The Bohr Model EXERCISE! What color of light is emitted when an excited electron in the hydrogen atom falls from: a) n = 5 to n = 2 blue, λ = 434 nm b) n = 4 to n = 2 green, λ = 486 nm c) n = 3 to n = 2 orange/red, λ = 657 nm Which transition results in the longest wavelength of light? Copyright © Cengage Learning. All rights reserved 21 Section 7.5 The Quantum Mechanical Model of the Atom Probability Distribution for the 1s Wave Function Copyright © Cengage Learning. All rights reserved 22 Section 7.5 The Quantum Mechanical Model of the Atom Radial Probability Distribution Copyright © Cengage Learning. All rights reserved 23 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 1s orbital:  Radius of the sphere that encloses 90% of the total electron probability. Copyright © Cengage Learning. All rights reserved 24 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. 25 Section 7.6 Quantum Numbers Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom Section 7.6 Quantum Numbers EXERCISE! For principal quantum level n = 3, determine the number of allowed subshells (different values of l), and give the designation of each. # of allowed subshells = 3 l = 0, 3s l = 1, 3p l = 2, 3d Copyright © Cengage Learning. All rights reserved 27 Section 7.6 Quantum Numbers EXERCISE! For l = 2, determine the magnetic quantum numbers (ml) and the number of orbitals. magnetic quantum numbers = –2, – 1, 0, 1, 2 number of orbitals = 5 Copyright © Cengage Learning. All rights reserved 28 Section 7.7 Orbital Shapes and Energies Three Representations of the Hydrogen 1s, 2s, and 3s Orbitals Copyright © Cengage Learning. All rights reserved 29 Section 7.7 Orbital Shapes and Energies The Boundary Surface Representations of All Three 2p Orbitals Copyright © Cengage Learning. All rights reserved 30 Section 7.7 Orbital Shapes and Energies The Boundary Surfaces of All of the 3d Orbitals Copyright © Cengage Learning. All rights reserved 31 Section 7.7 Orbital Shapes and Energies Representation of the 4f Orbitals in Terms of Their Boundary Surfaces Copyright © Cengage Learning. All rights reserved 32 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. Copyright © Cengage Learning. All rights reserved 33 Section 7.9 Polyelectronic Atoms  Atoms with more than one electron.  Electron correlation problem:  Since the electron pathways are unknown, the electron repulsions cannot be calculated exactly.  When electrons are placed in a particular quantum level, they “prefer” the orbitals in the order s, p, d, and then f. Copyright © Cengage Learning. All rights reserved 34 Section 7.9 Polyelectronic Atoms Penetration Effect  A 2s electron penetrates to the nucleus more than one in the 2p orbital.  This causes an electron in a 2s orbital to be attracted to the nucleus more strongly than an electron in a 2p orbital.  Thus, the 2s orbital is lower in energy than the 2p orbitals in a polyelectronic atom. Copyright © Cengage Learning. All rights reserved 35 Section 7.9 Polyelectronic Atoms A Comparison of the Radial Probability Distributions of the 2s and 2p Orbitals Copyright © Cengage Learning. All rights reserved 36 Section 7.9 Polyelectronic Atoms The Radial Probability Distribution of the 3s Orbital Copyright © Cengage Learning. All rights reserved 37 Section 7.10 The History of the Periodic Table  Originally constructed to represent the patterns observed in the chemical properties of the elements.  Mendeleev is given the most credit for the current version of the periodic table because he emphasized how useful the periodic table could be in predicting the existence and properties of still unknown elements. Copyright © Cengage Learning. All rights reserved 38 Section 7.11 The Aufbau Principle and the Periodic Table Aufbau Principle  As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals.  An oxygen atom has an electron arrangement of two electrons in the 1s subshell, two electrons in the 2s subshell, and four electrons in the 2p subshell. Oxygen: 1s22s22p4 Copyright © Cengage Learning. All rights reserved 39 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. Copyright © Cengage Learning. All rights reserved 40 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: 1s22s22p4 Oxygen: 1s 2s 2p Copyright © Cengage Learning. All rights reserved 41 Section 7.11 The Aufbau Principle and the Periodic Table Valence Electrons  The electrons in the outermost principal quantum level of an atom. 1s22s22p6 (valence electrons = 8)  The elements in the same group on the periodic table have the same valence electron configuration. Copyright © Cengage Learning. All rights reserved 42 Section 7.11 The Aufbau Principle and the Periodic Table The Orbitals Being Filled for Elements in Various Parts of the Periodic Table Copyright © Cengage Learning. All rights reserved 43 Section 7.11 The Aufbau Principle and the Periodic Table EXERCISE! Determine the expected electron configurations for each of the following. a) S 1s22s22p63s23p4 or [Ne]3s23p4 b) Ba [Xe]6s2 c) Eu [Xe]6s24f7 Copyright © Cengage Learning. All rights reserved 44 Section 7.12 Periodic Trends in Atomic Properties Periodic Trends  Ionization Energy  Electron Affinity  Atomic Radius  Electronegativity  Electron Shielding Section 7.12 Periodic Trends in Atomic Properties Ionization Energy  Energy required to remove an electron from a gaseous atom or ion.  X(g) → X+(g) + e– Mg → Mg+ + e– I1 = 735 kJ/mol (1st IE) Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE) Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE) *Core electrons are bound much more tightly than valence electrons. Section 7.12 Periodic Trends in Atomic Properties Ionization Energy  In general, as we go across a period from left to right, the first ionization energy increases.  Why?  Electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons.  Electrons in the same principal quantum level are generally more strongly bound from left to right on the periodic table. Section 7.12 Periodic Trends in Atomic Properties Ionization Energy  In general, as we go down a group from top to bottom, the first ionization energy decreases.  Why?  The electrons being removed are, on average, farther from the nucleus. Section 7.12 Periodic Trends in Atomic Properties The Values of First Ionization Energy for the Elements in the First Six Periods Section 7.12 Periodic Trends in Atomic Properties CONCEPT CHECK! Explain why the graph of ionization energy versus atomic number (across a row) is not linear. electron repulsions Where are the exceptions? some include from Be to B and N to O Section 7.12 Periodic Trends in Atomic Properties CONCEPT CHECK! Which atom would require more energy to remove an electron? Why? Na Cl Section 7.12 Periodic Trends in Atomic Properties CONCEPT CHECK! Which atom would require more energy to remove an electron? Why? Li Cs Section 7.12 Periodic Trends in Atomic Properties CONCEPT CHECK! Which has the larger second ionization energy? Why? Lithium or Beryllium Section 7.12 Periodic Trends in Atomic Properties Successive Ionization Energies (KJ per Mole) for the Elements in Period 3 Section 7.12 Periodic Trends in Atomic Properties Electron Affinity  Energy change associated with the addition of an electron to a gaseous atom.  X(g) + e– → X–(g)  In general as we go across a period from left to right, the electron affinities become more negative.  In general electron affinity becomes more positive in going down a group. Section 7.12 Periodic Trends in Atomic Properties Atomic Radius  In general as we go across a period from left to right, the atomic radius decreases.  Effective nuclear charge increases, therefore the valence electrons are drawn closer to the nucleus, decreasing the size of the atom.  In general atomic radius increases in going down a group.  Orbital sizes increase in successive principal quantum levels. Section 7.12 Periodic Trends in Atomic Properties Atomic Radii for Selected Atoms Section 7.12 Periodic Trends in Atomic Properties CONCEPT CHECK! Which should be the larger atom? Why? Na Cl Section 7.12 Periodic Trends in Atomic Properties CONCEPT CHECK! Which should be the larger atom? Why? Li Cs Section 7.12 Periodic Trends in Atomic Properties CONCEPT CHECK! Which is larger?  The hydrogen 1s orbital  The lithium 1s orbital Which is lower in energy? The hydrogen 1s orbital The lithium 1s orbital Section 7.12 Periodic Trends in Atomic Properties EXERCISE! Arrange the elements oxygen, fluorine, and sulfur according to increasing:  Ionization energy S, O, F  Atomic size F, O, S Section 7.13 The Properties of a Group: The Alkali Metals The Periodic Table – Final Thoughts 1. It is the number and type of valence electrons that primarily determine an atom’s chemistry. 2. Electron configurations can be determined from the organization of the periodic table. 3. Certain groups in the periodic table have special names. Copyright © Cengage Learning. All rights reserved 62 Section 7.13 The Properties of a Group: The Alkali Metals Special Names for Groups in the Periodic Table Copyright © Cengage Learning. All rights reserved 63 Section 7.13 The Properties of a Group: The Alkali Metals The Periodic Table – Final Thoughts 4. Basic division of the elements in the periodic table is into metals and nonmetals. Copyright © Cengage Learning. All rights reserved 64 Section 7.13 The Properties of a Group: The Alkali Metals Metals Versus Nonmetals Copyright © Cengage Learning. All rights reserved 65 Section 7.13 The Properties of a Group: The Alkali Metals The Alkali Metals  Li, Na, K, Rb, Cs, and Fr  Most chemically reactive of the metals  React with nonmetals to form ionic solids  Going down group:  Ionization energy decreases  Atomic radius increases  Density increases  Melting and boiling points smoothly decrease Copyright © Cengage Learning. All rights reserved 66