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Chapter Seven: Atomic Structure and Periodicity. In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate) 1s 2s3s4s2p3p3d.

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Presentation on theme: "Chapter Seven: Atomic Structure and Periodicity. In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate) 1s 2s3s4s2p3p3d."— Presentation transcript:

1 Chapter Seven: Atomic Structure and Periodicity

2 In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate) 1s 2s3s4s2p3p3d Energy Single-Electron Atom Energy Levels

3 1s 2s 3s 4s 2p 3p 3d E = n + l Screening results in the 4s- orbital having a lower energy that that of the 3d-orbital. 4s: n + l = = 4 vs. 3d: n + l = = 5 Multi-Electron Atom Energy Levels

4 In the H-atom, all subshells of same n have same energy. In a multi-electron atom: 1.subshells increase in energy as value of n + l increases. 2.for subshells of same n + I, the subshell with lower n is lower in energy. Assigning Electrons to Subshells

5 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 Rules:

6 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 “Hund’s Rule”

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

8 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 1s2s2p      1s 2 2s 2 2p 3 Orbitals Electron Configuration (spdf) notation: Electron Configuration: Orbital Box Notation

9 Atomic Electron Configurations

10 Electron Configurations & the Periodic Table

11 Electron Filling Order

12 Group 4A Atomic number = 6 6 total electrons 1s 2 2s 2 2p 2 1s2s2p      Carbon

13 Orbital box notation: 1s2s2p3s3p This corresponds to the energy level diagram: 1s 2s 2p 3s 3p Electron Configuration in the 3 rd Period

14 Orbital box notation: 1s2s2p3s3p Aluminum: Al (13 electrons) 1s 2s 2p 3s 3p                    1s 2 2s 2 2p 6 3s 2 spdf Electron Configuration 3p 1 Electron Configuration in the 3 rd Period

15 3s3p [Ne] 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. Orbital Box Notation   Noble gas Notation [Ne] 3s 2 3p 2 Full electron configuration spdf notation 1s 2 2s 2 2p 6 3s 2 3p 2 Noble Gas Notation

16 The innermost electrons (core) can be represented by the full shell of noble gas electron configuration: 1s 2 2s 2 = [He], 1s 2 2s 2 2p 6 = [Ne], 1s 2 2s 2 2p 6 3s 2 3p 6 = [Ar]... The outermost electrons are referred to as the “Valence” electrons”. Element Full Electron Config. Noble Gas Notation Mg1s 2 2s 2 2p 6 3s 2 [Ne] 3s 2 Core or Noble Gas Notation

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

18 Electron Configurations are written by shell even though the electrons fill by the periodic table: Ni: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 last electron to fill:3d 8 electron configuration by filling: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 electron configuration by shell: (write this way) Transition Elements

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20 Ions with UNPAIRED ELECTRONS are PARAMAGNETIC (attracted to a magnetic field). Ions without UNPAIRED ELECTRONS are DIAMAGNETIC (not attracted to a magnetic field). Fe 3+ ions in Fe 2 O 3 have 5 unpaired electrons. This makes the sample paramagnetic. Electron Configurations of Ions

21 f-block elements: These elements have the configuration [core] ns x (n - 1)d y (n - 2)f z Where n is the period and x, y & z are particular to the element. Cerium: [Xe] 6s 2 5d 1 4f 1 Uranium: [Rn] 7s 2 6d 1 5f 3 Lanthanides & Actinides

22 © 2009, Prentice-Hall, Inc. Some Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

23 © 2009, Prentice-Hall, Inc. Some Anomalies For instance, the electron configuration for copper is [Ar] 4s 1 3d 10 rather than the expected [Ar] 4s 2 3d 9.

24 © 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.

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

26 Formation of ions & Valence electrons Ions are formed when atoms either: 1. Cation (+): Give up (lose) electrons 2.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

27 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]

28 Isoelectronic series:

29 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] 3d 10 Cu 2+ : [Ar] 3d 9

30 © 2009, Prentice-Hall, Inc. Development of Periodic Table Elements in the same group generally have similar chemical properties. Physical properties are not necessarily similar, however.

31 © 2009, Prentice-Hall, Inc. Periodic Trends In this chapter, we will rationalize observed trends in – Sizes of atoms and ions. – Ionization energy. – Electron affinity.

32 Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends

33 © 2009, Prentice-Hall, Inc. 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.

34 © 2009, Prentice-Hall, Inc. Effective Nuclear Charge The effective nuclear charge, Z eff, is found this way: Z eff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

35 © 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.

36 © 2009, Prentice-Hall, Inc. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row (due to increasing Z eff ). …increase from top to bottom of a column (due to increasing value of n).

37 © 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.

38 © 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.

39 © 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, Z eff is essentially the same, but the valence electrons are farther from the nucleus.

40 © 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, Z eff increases.

41 © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

42 © 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.

43 © 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 −

44 © 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.

45 © 2009, Prentice-Hall, Inc. Trends in Electron Affinity There are again, however, two discontinuities in this trend.

46 © 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.

47 © 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.

48 © 2009, Prentice-Hall, Inc. Sizes of Ions Ionic size depends upon: – The nuclear charge. – The number of electrons. – The orbitals in which electrons reside.

49 © 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.

50 © 2009, Prentice-Hall, Inc. Sizes of Ions Anions are larger than their parent atoms. – Electrons are added and repulsions between electrons are increased.

51 © 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.

52 © 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.

53 Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2–

54 Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2–

55 Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2–

56 Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2–

57 Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2–

58 Na + N 3- Mg 2+ F–F– O 2– Ion e/p ratio Since N 3- has the highest ratio of electrons to protons, it must have the largest radius.

59 Na + N 3- Mg 2+ F–F– O 2– Ion e/p ratio Since N 3- has the highest ratio of electrons to protons, it must have the largest radius. Since Mg 2+ has the lowest, it must have the smallest radius. The rest can be ranked by ratio.

60 Na + N 3- Mg 2+ F–F– O 2– Ion e/p ratio Since N 3- has the highest ratio of electrons to protons, it must have the largest radius. Since Mg 2+ has the lowest, it must have the smallest radius. The rest can be ranked by ratio. N 3- >O 2– >F–F– >Na + >Mg 2+ Decreasing size isoelectronic Notice that they all have 10 electrons: They are isoelectronic (same electron configuration) as Ne.

61 Moving through the periodic table: Atomic radii Ionization Energy Electron Affinity Down a groupIncreaseDecrease Becomes less exothermic Across a PeriodDecreaseIncrease Becomes more exothermic Summary of Periodic Trends


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