CHEMISTRY. CHM201 Class #16 CHEMISTRY. Chapter 7 Continued. Chapter 7 Outline for Class #16

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1 CHEMISTRY Fifth Edition Gilbert Kirss Foster Bretz Davies CHM201 Class #16 Chemistry, 5 th Edition Copyright 2017, W. W. Norton & Company CHEMISTRY Fifth Edition Gilbert Kirss Foster Bretz Davies Chapter 7 Continued A Quantum Model of Atoms: Waves, Particles, and Periodic Properties Chemistry, 5 th Edition Copyright 2017, W. W. Norton & Company Chapter 7 Outline for Class # The Periodic Table and Filling the Orbitals of Multielectron Atoms 7.9 Electron Configurations of Ions 7.10 The Sizes of Atoms and Ions 7.11 Ionization Energies 7.12 Electron Affinities 3 1

2 Aufbau Principle Aufbau principle: Method of building electron configurations by adding one electron at a time as atomic number increases When adding electrons to an atom: Electrons always go in lowest-energy orbitals available. A maximum of two electrons per orbital 4 Orbital Diagrams Orbital diagram depiction of the arrangement of electrons in an atom or ion using boxes to represent orbitals Up/down arrows indicate e with +/ spin. Hund s Rule: Degenerate orbitals orbitals of the same energy (e.g., 2p) The lowest-energy configuration maximizes the number of unpaired electrons. 5 Orbital Diagrams for Hydrogen and Helium Atoms 6 2

3 Orbital Diagrams for Multielectron Atoms 7 Electron Configurations of Second Row Elements 8 Orbital Energies in Multielectron Atoms Orbitals are filled from lowest energy to highest energy: Within a given shell, subshells fill in order of s, p, d, f. Due to small energy differences there is some overlap of subshells at higher values of n. This causes electrons to be placed first into the higher n value for larger n values: 4s is filled before 3d. 6s is filled before 4f. 9 3

4 Orbitals Ordered by Energy 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 10 Periodic Table: Filling of Orbitals 11 Periodic Table: s, p, and d Blocks 12 4

5 Anomalies in Configurations Chromium, copper do not follow pattern: Cr = 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 Cu = 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 Anomalies arise from stability associated with half-filled and completely filled d- subshells. 13 Electron Configurations: Ions Formation of ions: The gain or loss of valence electrons to achieve stable electron configuration (filled shell) Cations: Na Na + + e [Ne]3s 1 [Ne] + e Anions: Cl + e Cl [Ne]3s 2 3p 5 + e [Ne]3s 2 3p 6 = [Ar] 14 Isoelectronic Ions/Atoms Main group elements: Form ions by gain/loss of e to obtain noble gas configuration: Mg Mg e = [Ar] O + 2e O 2 = [Ne] Isoelectronic describes atoms/ions having identical electron configurations: Na +, Mg 2+, O 2, F, Ne = 1s 2 2s 2 2p 6 (= [Ne]) K +, Cl, Ca 2+, Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 (= [Ar]) 15 5

6 Cations of Transition Metals Transition metal cations: When electrons are lost to make cations they lose the electrons from the orbitals with the highest n value first: Fe: [Ar]3d 6 4s 2 Fe 2+ : [Ar]3d 6 4s 2 [Ar]3d 6 Fe 3+ : [Ar]3d 6 4s 2 [Ar]3d 5 16 Practice: Electron Configurations Write the electron configuration for the following atoms/ions: a) Pt b) I c) Rh 2+ Collect and Organize: We are writing electron configurations for one transition metal atom (Pt), one main group anion (I ), and one transition metal cation (Rh 2+ ). 17 Practice: Electron Configurations Write the electron configuration for the following atoms/ions: a) Pt b) I c) Rh 2+ Analyze: We can write electron configurations for elemental atoms based on their position in the periodic table by using the Aufbau principle and the orbital-filling diagram. For the iodide anion, we add one extra electron to the 5p orbitals. For the Rh 2+ cation, we will remove two electrons from the outer-shell orbital (5s). 18 6

7 Practice: Electron Configurations Write the electron configuration for the following atoms/ions: a) Pt b) I c) Rh 2+ Solve: Pt = [Xe]5d 8 6s 2 I = [Kr]4d 10 5s 2 5p 6 Rh 2+ = [Kr]4d 7 19 Practice: Electron Configurations Write the electron configuration for the following atoms/ions: a) Pt b) I c) Rh 2+ Think About It: The number of valence electrons is consistent with the locations of these elements in the periodic table and their charge (if ionic). 20 Sizes of Atoms and Ions Atomic radius: Half the distance between identical nuclear centers in a molecule Metallic radius: Half the distance between nuclear centers in the solid metal Ionic radius: Derived from the distance between nuclear centers in solid ionic compound 21 7

Atomic, Metallic, and Ionic Radii 22 Atomic Radius: Periodic Trends Atomic

8 Atomic, Metallic, and Ionic Radii 22 Atomic Radius: Periodic Trends Atomic radii: Increase going down a column As atomic number increases, so does the number of valence electrons. More electrons mean more electron-electron repulsion which increases the size of the atom. Decrease going across a row Increased effective nuclear charge (Z eff ) moving across row Increased attraction for valence electrons decreased atomic size 23 Atomic Radius: Periodic Trends 24 8

Atomic Radius: Periodic Trends Ion radii: Cations Cations of main group elements are much smaller than their parent atom. The electrostatic attraction is greater for the remaining electrons.

9 Atomic Radius: Periodic Trends Ion radii: Cations Cations of main group elements are much smaller than their parent atom. The electrostatic attraction is greater for the remaining electrons. Anions Anions of main group elements are much larger than their parent atom. The increase in electrons increases the electronelectron repulsion. 25 Radii of Atoms and Ions 26 Ionization Energy (IE) Ionization energy: The amount of energy needed to remove 1 mole of e from 1 mole of ground-state atoms or ions in the gas phase (kj/mol): X(g) X + (g) + e 1st Ionization Energy (IE 1 ): Mg(g) Mg + (g) + e IE 1 = 738 kj/mol 2nd Ionization Energy (IE 2 ): Mg + (g) Mg 2+ (g) + e IE 2 = 1451 kj/mol Note: IE 2 > IE

10 Ionization Energy Trends 28 Ionization Energies 29 Successive Ionization Energies 30 10

Electron Affinities (EA) Electron affinity: Energy change that occurs when 1 mole of electrons combine with 1 mole of atoms or ions in the gas phase Cl(g) + e Cl (g) EA 1 = 349 kj/mol Periodic

11 Electron Affinities (EA) Electron affinity: Energy change that occurs when 1 mole of electrons combine with 1 mole of atoms or ions in the gas phase Cl(g) + e Cl (g) EA 1 = 349 kj/mol Periodic trends EA values become more negative moving to the right and up in the periodic table. 31 Periodic Trends in EA 32 Visual Problems 1. Which representations depict the loss of an electron? 33 11

12 Visual Problems 2. Which representations can be discussed using only a quantized model of light? 34 Visual Problems 3. Which representations are consistent with the model of an atom that discusses probability with regard to the location of electrons? 35 12

Gilbert Kirss Foster. Chapter3. Atomic Structure. Explaining the Properties of Elements

Gilbert Kirss Foster. Chapter3. Atomic Structure. Explaining the Properties of Elements Gilbert Kirss Foster Chapter3 Atomic Structure Explaining the Properties of Elements Chapter Outline 3.1 Waves of Light 3.2 Atomic Spectra 3.3 Particles of Light: Quantum Theory 3.4 The Hydrogen Spectrum