Chapter Seven: Atomic Structure and Periodicity

Single-Electron Atom Energy Levels 3p 3d Energy 2s 2p In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate) 1s

Multi-Electron Atom Energy Levels 2p 3p 3d Screening results in the 4s-orbital having a lower energy that that of the 3d-orbital. E = n + l 4s: n + l = 4 + 0 = 4 vs. 3d: n + l = 3 + 2 = 5

Assigning Electrons to Subshells In the H-atom, all subshells of same n have same energy. In a multi-electron atom: subshells increase in energy as value of n + l increases. for subshells of same n + I, the subshell with lower n is lower in energy. See Chapter 7 Video Presentation Slide 6

Orbital Filling Rules: Aufbau Principle: Lower energy orbitals fill first. Hund’s Rule: Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before pairing up of electrons can occur. Pauli exclusion principle: Individual orbitals only hold two electrons, and each should have different spin. “s” orbitals can hold 2 electrons “p” orbitals hold up to 6 electrons “d” orbitals can hold up to 10 electrons

Orbital Filling: The “Hund’s Rule” Degenerate orbitals are filled with electrons until all are half-filled before pairing up of electrons can occur. Consider a set of 2p orbitals:    2p Electrons fill in this manner

Orbital Filling: The “Pauli Principle” “Pauli exclusion principle” Individual orbitals only hold two electrons, and each should have opposite spin. Consider a set of 2p orbitals:       2p  = spin up  = spin down *the convention is to write up, then down. Electrons fill in this manner

Electron Configuration: Orbital Box Notation Electrons fill the orbitals from lowest to highest energy. The electron configuration of an atom is the total sum of the electrons from lowest to highest shell. Example: Nitrogen: N has an atomic number of 7, therefore 7 electrons Orbitals 1s 2s 2p        1s2 2s2 2p3 Electron Configuration (spdf) notation:

Atomic Electron Configurations

Electron Configurations & the Periodic Table

Electron Filling Order

Carbon Group 4A Atomic number = 6 6 total electrons 1s2 2s2 2p2 1s 2s      

Electron Configuration in the 3rd Period Orbital box notation: 1s 2s 2p 3s 3p 1s 2s 2p 3s 3p This corresponds to the energy level diagram: See Chapter 7 Video Presentation Slide 7

Electron Configuration in the 3rd Period Orbital box notation: 1s 2s 2p 3s 3p Aluminum: Al (13 electrons)                       spdf Electron Configuration   1s2 2s2 2p6 3s2 3p1  

Noble Gas Notation [Ne] Full electron configuration spdf notation The electron configuration of an element can be represented as a function of the core electrons in terms of a noble gas and the valence electrons. Full electron configuration spdf notation 1s22s22p63s23p2 Orbital Box Notation 3s 3p [Ne]     Noble gas Notation [Ne] 3s23p2

Core or Noble Gas Notation The innermost electrons (core) can be represented by the full shell of noble gas electron configuration: 1s22s2 = [He], 1s22s22p6 = [Ne], 1s22s22p63s23p6 = [Ar]... The outermost electrons are referred to as the “Valence” electrons”. Element Full Electron Config. Noble Gas Notation Mg 1s2 2s2 2p6 3s2 [Ne] 3s2

Transition Metal All 4th period and beyond d-block elements have the electron configuration [Ar] nsx (n - 1)dy Where n is the period and x, y are particular to the element. Chromium Iron Copper

Transition Elements Ni: Electron Configurations are written by shell even though the electrons fill by the periodic table: Ni: last electron to fill: 3d8 electron configuration by filling: 1s2 2s2 2p6 3s2 3p6 4s2 3d8 electron configuration by shell: (write this way) 1s2 2s2 2p6 3s2 3p6 3d8 4s2

Electron Configurations of Ions Ions with UNPAIRED ELECTRONS are PARAMAGNETIC (attracted to a magnetic field). Ions without UNPAIRED ELECTRONS are DIAMAGNETIC (not attracted to a magnetic field). Fe3+ ions in Fe2O3 have 5 unpaired electrons. This makes the sample paramagnetic.

Lanthanides & Actinides f-block elements: These elements have the configuration [core] nsx (n - 1)dy (n - 2)fz Where n is the period and x, y & z are particular to the element. Uranium: [Rn] 7s2 6d1 5f3 Cerium: [Xe] 6s2 5d1 4f1

Some Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row. © 2009, Prentice-Hall, Inc.

Some Anomalies For instance, the electron configuration for copper is [Ar] 4s1 3d10 rather than the expected [Ar] 4s2 3d9. © 2009, Prentice-Hall, Inc.

Some Anomalies This occurs because the 4s and 3d orbitals are very close in energy. These anomalies occur in f-block atoms, as well. © 2009, Prentice-Hall, Inc.

Valence Electrons: Electrons at the highest energy level. Electrons available to be gained/lost/shared in a chemical reaction Valence electron configuration: nsxpx (total of 8 valence electrons) Representation of valence electrons in Lewis dot notation:

Formation of ions & Valence electrons Ions are formed when atoms either: 1. Cation (+): Give up (lose) electrons Anion (-): Gain electrons Cation= Overall goal: stable noble gas notation Elements with < 4 valence electrons- form cations Elements with > 4 valence electrons for anions. Non-metals with 4 valence electrons- do not form ions Noble gases (8 valence electrons) – are unreactive

Ions & Electron Configuration Atoms or groups of atoms that carry a charge Cations- positive charge Formed when an atom loses electron(s) Na  Na+ + e- Na+ : 1s22s22p6 (10 e-) = [Ne] Anions –negative charge Formed when an atom gain electrons(s) F + e-  F- Cl- : 1s22s22p6 (10 e-) = [Ne]

Isoelectronic series:

Transition Metals & Ions Transition metals: multi-valent ions (more than one possible charge) Electrons are removed from the highest quantum number first: Example: Cu+ & Cu 2+ Cu+: [Ar] 3d10 Cu2+ : [Ar] 3d9

Development of Periodic Table Elements in the same group generally have similar chemical properties. Physical properties are not necessarily similar, however. © 2009, Prentice-Hall, Inc.

Periodic Trends In this chapter, we will rationalize observed trends in Sizes of atoms and ions. Ionization energy. Electron affinity. © 2009, Prentice-Hall, Inc.

General Periodic Trends Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly See Chapter 7 Video Presentation Slides 8, 9, & 10 Larger orbitals. Electrons held less tightly.

Effective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. © 2009, Prentice-Hall, Inc.

Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons. © 2009, Prentice-Hall, Inc.

What Is the Size of an Atom? The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. © 2009, Prentice-Hall, Inc.

Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row (due to increasing Zeff). …increase from top to bottom of a column (due to increasing value of n). © 2009, Prentice-Hall, Inc.

Ionization Energy The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. The first ionization energy is that energy required to remove first electron. The second ionization energy is that energy required to remove second electron, etc. © 2009, Prentice-Hall, Inc.

Ionization Energy It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionization energy takes a quantum leap. © 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron. For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus. © 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies Generally, as one goes across a row, it gets harder to remove an electron. As you go from left to right, Zeff increases. © 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies However, there are two apparent discontinuities in this trend. © 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies The first occurs between Groups IIA and IIIA. In this case the electron is removed from a p-orbital rather than an s-orbital. The electron removed is farther from nucleus. There is also a small amount of repulsion by the s electrons. © 2009, Prentice-Hall, Inc.

Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e−  Cl− © 2009, Prentice-Hall, Inc.

Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. © 2009, Prentice-Hall, Inc.

Trends in Electron Affinity There are again, however, two discontinuities in this trend. © 2009, Prentice-Hall, Inc.

Trends in Electron Affinity The first occurs between Groups IA and IIA. The added electron must go in a p-orbital, not an s-orbital. The electron is farther from nucleus and feels repulsion from the s-electrons. © 2009, Prentice-Hall, Inc.

Trends in Electron Affinity The second occurs between Groups IVA and VA. Group VA has no empty orbitals. The extra electron must go into an already occupied orbital, creating repulsion. © 2009, Prentice-Hall, Inc.

Sizes of Ions Ionic size depends upon: The nuclear charge. The number of electrons. The orbitals in which electrons reside. © 2009, Prentice-Hall, Inc.

Sizes of Ions Cations are smaller than their parent atoms. The outermost electron is removed and repulsions between electrons are reduced. © 2009, Prentice-Hall, Inc.

Sizes of Ions Anions are larger than their parent atoms. Electrons are added and repulsions between electrons are increased. © 2009, Prentice-Hall, Inc.

Sizes of Ions Ions increase in size as you go down a column. This is due to increasing value of n. © 2009, Prentice-Hall, Inc.

Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge. © 2009, Prentice-Hall, Inc.

Problem: Rank the following ions in order of decreasing size? Na+, Mg2+, F– O2–

Problem: Rank the following ions in order of decreasing size? Na+, Mg2+, F– O2– Ion # of protons # of electrons ratio of e/p Na+ N3- Mg2+ F– O2–

Problem: Rank the following ions in order of decreasing size? Na+, Mg2+, F– O2– Ion # of protons # of electrons ratio of e/p Na+ N3- Mg2+ F– O2– 11 7 12 9 8

Problem: Rank the following ions in order of decreasing size? Na+, Mg2+, F– O2– Ion # of protons # of electrons ratio of e/p Na+ N3- Mg2+ F– O2– 11 7 12 9 8 10

Problem: Rank the following ions in order of decreasing size? Na+, Mg2+, F– O2– Ion # of protons # of electrons ratio of e/p Na+ N3- Mg2+ 11 7 12 9 8 10 0.909 1.43 0.833 1.11 1.25

Na+ N3- Mg2+ F– O2– Ion 0.909 1.43 0.833 1.11 1.25 e/p ratio Since N3- has the highest ratio of electrons to protons, it must have the largest radius.

Since Mg2+ has the lowest, it must have the smallest radius. Na+ N3- Mg2+ F– O2– Ion 0.909 1.43 0.833 1.11 1.25 e/p ratio Since N3- has the highest ratio of electrons to protons, it must have the largest radius. Since Mg2+ has the lowest, it must have the smallest radius. The rest can be ranked by ratio.

Na+ N3- Mg2+ F– O2– Ion 0.909 1.43 0.833 1.11 1.25 e/p ratio Since N3- has the highest ratio of electrons to protons, it must have the largest radius. Since Mg2+ has the lowest, it must have the smallest radius. The rest can be ranked by ratio. N3- > O2– > F– > Na+ > Mg2+ Decreasing size Notice that they all have 10 electrons: They are isoelectronic (same electron configuration) as Ne.

Summary of Periodic Trends Moving through the periodic table: Atomic radii Ionization Energy Electron Affinity Down a group Increase Decrease Becomes less exothermic Across a Period Becomes more exothermic