Chapter Eight Electron Configurations and Periodicity Electron Configuration

  • Slides: 52
Download presentation
Chapter Eight Electron Configurations and Periodicity

Chapter Eight Electron Configurations and Periodicity

Electron Configuration • An electron configuration of an atom is a particular distribution of

Electron Configuration • An electron configuration of an atom is a particular distribution of electrons among available subshells. – The notation for a configuration lists the subshell symbols sequentially with a superscript indicating the number of electrons occupying that subshell. – For example, lithium (atomic number 3) has two electrons in the “ 1 s” subshell and one electron in the “ 2 s” sub shell 1 s 2 2 s 1. Copyright © Houghton Mifflin Company. All rights reserved. 2

Electron Configuration • An orbital diagram is used to show the orbitals of a

Electron Configuration • An orbital diagram is used to show the orbitals of a subshell are occupied by electrons. – Each orbital is represented by a circle. – Each group of orbitals is labeled by its subshell notation. 1 s 2 s 2 p – Electrons are represented by arrows: up for ms = +1/2 and down for ms = -1/2 Copyright © Houghton Mifflin Company. All rights reserved. 3

The Pauli Exclusion Principle • The Pauli exclusion principle, which summarizes experimental observations, states

The Pauli Exclusion Principle • The Pauli exclusion principle, which summarizes experimental observations, states that no two electrons can have the same four quantum numbers. – In other words, an orbital can hold at most two electrons, and then only if the electrons have opposite spins. Copyright © Houghton Mifflin Company. All rights reserved. 4

The Pauli Exclusion Principle • The maximum number of electrons and their orbital diagrams

The Pauli Exclusion Principle • The maximum number of electrons and their orbital diagrams are: Sub shell Number of Orbitals Maximum Number of Electrons s (l = 0) 1 2 p (l = 1) 3 6 d (l =2) 5 10 f (l =3) 7 14 Copyright © Houghton Mifflin Company. All rights reserved. 5

Aufbau Principle • Every atom has an infinite number of possible electron configurations. –

Aufbau Principle • Every atom has an infinite number of possible electron configurations. – The configuration associated with the lowest energy level of the atom is called the “ground state. ” – Other configurations correspond to “excited states. ” – Table 8. 1 lists the ground state configurations of atoms up to krypton. Copyright © Houghton Mifflin Company. All rights reserved. 6

Aufbau Principle • The Aufbau principle is a scheme used to reproduce the ground

Aufbau Principle • The Aufbau principle is a scheme used to reproduce the ground state electron configurations of atoms by following the “building up” order. – Listed below is the order in which all the possible subshells fill with electrons. 1 s, 2 p, 3 s, 3 p, 4 s, 3 d, 4 p, 5 s, 4 d, 5 p, 6 s, 4 f, 5 d, 6 p, 7 s, 5 f – You need not memorize this order. As you will see, it can be easily obtained. Copyright © Houghton Mifflin Company. All rights reserved. 7

Order for Filling Atomic Subshells 1 s 2 s 3 s 4 s 5

Order for Filling Atomic Subshells 1 s 2 s 3 s 4 s 5 s 6 s 2 p 3 p 4 p 5 p 6 p Copyright © Houghton Mifflin Company. All rights reserved. 3 d 4 d 4 f 5 d 5 f 6 d 6 f 8

Orbital Energy Levels in Multielectron Systems 3 d Energy 4 s 3 p 3

Orbital Energy Levels in Multielectron Systems 3 d Energy 4 s 3 p 3 s 2 p 2 s 1 s (See Animation: Orbital Energies) Copyright © Houghton Mifflin Company. All rights reserved. 9

Aufbau Principle • The “building up” order corresponds for the most part to increasing

Aufbau Principle • The “building up” order corresponds for the most part to increasing energy of the subshells. – By filling orbitals of the lowest energy first, you usually get the lowest total energy (“ground state”) of the atom. – Now you can see how to reproduce the electron configurations of Table 8. 1 using the Aufbau principle. – Remember, the number of electrons in the neutral atom equals the atomic number, Z. Copyright © Houghton Mifflin Company. All rights reserved. 10

Aufbau Principle (Using the Noble Gas Shortcut) • Here a few examples. – Using

Aufbau Principle (Using the Noble Gas Shortcut) • Here a few examples. – Using the abbreviation [He] for 1 s 2, the configurations are: Z=4 Beryllium 1 s 22 s 2 or [He]2 s 2 Z=3 Lithium 1 s 22 s 1 or [He]2 s 1 Copyright © Houghton Mifflin Company. All rights reserved. 11

Aufbau Principle • With boron (Z=5), the electrons begin filling the 2 p subshell.

Aufbau Principle • With boron (Z=5), the electrons begin filling the 2 p subshell. Z=5 Z=6 Z=7 Z=8 Z=9 Z=10 Boron Carbon Nitrogen Oxygen Fluorine Neon Copyright © Houghton Mifflin Company. All rights reserved. 1 s 22 p 1 1 s 22 p 2 1 s 22 p 3 1 s 22 p 4 1 s 22 p 5 1 s 22 p 6 or or or [He]2 s 22 p 1 [He]2 s 22 p 2 [He]2 s 22 p 3 [He]2 s 22 p 4 [He]2 s 22 p 5 [He]2 s 22 p 6 12

Aufbau Principle • With sodium (Z = 11), the 3 s subshell begins to

Aufbau Principle • With sodium (Z = 11), the 3 s subshell begins to fill. Z=11 Sodium 1 s 22 p 63 s 1 or [Ne]3 s 1 Z=12 Magnesium 1 s 22 p 23 s 2 or [Ne]3 s 2 • Then the 3 p subshell begins to fill. Z=13 : : Z=18 Aluminum 1 s 22 p 63 s 23 p 1 or [Ne]3 s 23 p 1 Argon 1 s 22 p 63 s 23 p 6 or [Ne]3 s 23 p 6 Copyright © Houghton Mifflin Company. All rights reserved. 13

Configurations and the Periodic Table • Note that elements within a given family have

Configurations and the Periodic Table • Note that elements within a given family have similar configurations. – For instance, look at the noble gases. Helium Neon Argon Krypton 1 s 22 s 22 p 63 s 23 p 63 d 104 s 24 p 6 Copyright © Houghton Mifflin Company. All rights reserved. 14

Configurations and the Periodic Table • Note that elements within a given family have

Configurations and the Periodic Table • Note that elements within a given family have similar configurations. – The Group IIA elements are sometimes called the alkaline earth metals. Beryllium 1 s 22 s 2 Magnesium 1 s 22 p 63 s 2 Calcium 1 s 22 p 63 s 23 p 64 s 2 Copyright © Houghton Mifflin Company. All rights reserved. 15

Configurations and the Periodic Table • Electrons that reside in the outermost shell of

Configurations and the Periodic Table • Electrons that reside in the outermost shell of an atom—or in other words, those electrons outside the “noble gas core”—are called valence electrons. – These electrons are primarily involved in chemical reactions. – Elements within a given group have the same “valence shell configuration. ” – This accounts for the similarity of the chemical properties among groups of elements. Copyright © Houghton Mifflin Company. All rights reserved. 16

Configurations and the Periodic Table • The following slide illustrates how the periodic table

Configurations and the Periodic Table • The following slide illustrates how the periodic table provides a sound way to remember the Aufbau sequence. – In many cases you need only the configuration of the outer electrons. – You can determine this from their position on the periodic table. – The total number of valence electrons for an atom equals its group number. Copyright © Houghton Mifflin Company. All rights reserved. 17

Figure 8. 5: Periodic Table Copyright © Houghton Mifflin Company. All rights reserved. 18

Figure 8. 5: Periodic Table Copyright © Houghton Mifflin Company. All rights reserved. 18

Configurations and the Periodic Table Copyright © Houghton Mifflin Company. All rights reserved. 19

Configurations and the Periodic Table Copyright © Houghton Mifflin Company. All rights reserved. 19

Orbital Diagrams • Consider carbon (Z = 6) with the ground state configuration 1

Orbital Diagrams • Consider carbon (Z = 6) with the ground state configuration 1 s 22 p 2. – Three possible arrangements are given in the following orbital diagrams. 1 s 2 s 2 p Diagram 1: Diagram 2: Diagram 3: – Each state has different energy and different magnetic characteristics. Copyright © Houghton Mifflin Company. All rights reserved. 20

Orbital Diagrams • Hund’s rule states that the lowest energy arrangement (the “ground state”)

Orbital Diagrams • Hund’s rule states that the lowest energy arrangement (the “ground state”) of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons. – Looking at carbon again, we see that the ground state configuration corresponds to diagram 1 when following Hund’s rule. 1 s 2 s Copyright © Houghton Mifflin Company. All rights reserved. 2 p 21

Orbital Diagrams • To apply Hund’s rule to oxygen, whose ground state configuration is

Orbital Diagrams • To apply Hund’s rule to oxygen, whose ground state configuration is 1 s 22 p 4, we place the first seven electrons as follows. 1 s 2 s 2 p • The last electron is paired with one of the 2 p electrons to give a doubly occupied 1 s 2 s 2 p orbital. – Table 8. 2 lists more orbital diagrams. Copyright © Houghton Mifflin Company. All rights reserved. 22

Magnetic Properties • Although an electron behaves like a tiny magnet, two electrons that

Magnetic Properties • Although an electron behaves like a tiny magnet, two electrons that are opposite in spin cancel each other. Only atoms with unpaired electrons exhibit magnetic susceptibility. – A paramagnetic substance is one that is weakly attracted by a magnetic field, usually the result of unpaired electrons. – A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons. Copyright © Houghton Mifflin Company. All rights reserved. 23

Periodic Trends • The periodic law states that the properties of elements recur in

Periodic Trends • The periodic law states that the properties of elements recur in a repeating pattern when arranged according to increasing atomic number. • We will look at (4) periodic properties: – – Atomic radius Ionization energy Electron affinity Electronegativity Copyright © Houghton Mifflin Company. All rights reserved. 24

Dimitri Mendeleev Copyright © Houghton Mifflin Company. All rights reserved. 25

Dimitri Mendeleev Copyright © Houghton Mifflin Company. All rights reserved. 25

Periodic Trends • Atomic radius – Within each period (horizontal row), the atomic radius

Periodic Trends • Atomic radius – Within each period (horizontal row), the atomic radius tends to decrease with increasing atomic number (nuclear charge). – Within each group (vertical column), the atomic radius tends to increase with the period number. Copyright © Houghton Mifflin Company. All rights reserved. 26

Periodic Trends • Two factors determine the size of an atom. – One factor

Periodic Trends • Two factors determine the size of an atom. – One factor is the principal quantum number, n. The larger is “n, ” the larger the size of the orbital. – The other factor is the effective nuclear charge, which is the positive charge an electron experiences from the nucleus minus any “shielding effects” from intervening electrons. Copyright © Houghton Mifflin Company. All rights reserved. 27

Figure 8. 10: Representation of atomic radii (covalent radii) of the maingroup elements Copyright

Figure 8. 10: Representation of atomic radii (covalent radii) of the maingroup elements Copyright © Houghton Mifflin Company. All rights reserved. 28

Periodic Trends • Ionic radii – parallels the trends in atomic radii within the

Periodic Trends • Ionic radii – parallels the trends in atomic radii within the same group – Positive ions in the same group increase in radii down a group – Negative ions in the same group increase in radii down a group Copyright © Houghton Mifflin Company. All rights reserved. 29

Periodic Trends • Ionic radii • But CATIONS (positive ions) are smaller in radii

Periodic Trends • Ionic radii • But CATIONS (positive ions) are smaller in radii than the atoms they come from • And ANIONS (negative ions) are larger than the atoms they come from Copyright © Houghton Mifflin Company. All rights reserved. 30

Copyright © Houghton Mifflin Company. All rights reserved. 31

Copyright © Houghton Mifflin Company. All rights reserved. 31

Copyright © Houghton Mifflin Company. All rights reserved. 32

Copyright © Houghton Mifflin Company. All rights reserved. 32

Periodic Trends • Ionization energy (IE) – The first ionization energy of an atom

Periodic Trends • Ionization energy (IE) – The first ionization energy of an atom is the amount of energy needed to remove the outermost (highest energy) electron from the neutral atom in a gaseous state. – For a lithium atom, the first ionization energy is illustrated by: Copyright © Houghton Mifflin Company. All rights reserved. 33

Periodic Trends • Ionization energy – There is a general trend that ionization energies

Periodic Trends • Ionization energy – There is a general trend that ionization energies increase with atomic number within a given period. – This follows the trend in size, as it is more difficult to remove an electron that is closer to the nucleus. – For the same reason, we find that ionization energies, again following the trend in size, decrease as we descend a column of elements. Copyright © Houghton Mifflin Company. All rights reserved. 34

Figure 8. 11: Ionization energy versus atomic number Copyright © Houghton Mifflin Company. All

Figure 8. 11: Ionization energy versus atomic number Copyright © Houghton Mifflin Company. All rights reserved. 35

Periodic Trends • Ionization energy – The first IE is the NRG required to

Periodic Trends • Ionization energy – The first IE is the NRG required to remove one electron – The second IE is the NRG required to remove a second electron, and so on. – Table 8. 3 lists the successive ionization energies of the first ten elements. Copyright © Houghton Mifflin Company. All rights reserved. 36

Table 8. 3 Copyright © Houghton Mifflin Company. All rights reserved. 37

Table 8. 3 Copyright © Houghton Mifflin Company. All rights reserved. 37

Periodic Trends • Electron Affinity – The electron affinity is the energy given off

Periodic Trends • Electron Affinity – The electron affinity is the energy given off or absorbed when an electron is added to a neutral atom in the gaseous state to form a negative ion. – For a chlorine atom, the first electron affinity is illustrated by: Copyright © Houghton Mifflin Company. All rights reserved. 38

Periodic Trends • Electron Affinity – The more negative the electron affinity, the more

Periodic Trends • Electron Affinity – The more negative the electron affinity, the more stable the negative ion that is formed. – The general trend: • Across a period, electron affinity increases. • Down a group, electron affinity decreases. – Table 8. 4 gives the electron affinities of the main-group elements. Copyright © Houghton Mifflin Company. All rights reserved. 39

Copyright © Houghton Mifflin Company. All rights reserved. 40

Copyright © Houghton Mifflin Company. All rights reserved. 40

Periodic Trends • Electronegativity (EN) – Attraction for electrons in a chemical bond for

Periodic Trends • Electronegativity (EN) – Attraction for electrons in a chemical bond for another atom’s electrons – It’s an atomic “tug of war” • Trend: – Across a period, EN increases – Down a group, EN decreases Copyright © Houghton Mifflin Company. All rights reserved. 41

Copyright © Houghton Mifflin Company. All rights reserved. 42

Copyright © Houghton Mifflin Company. All rights reserved. 42

The Main-Group Elements • The physical and chemical properties of the main-group elements clearly

The Main-Group Elements • The physical and chemical properties of the main-group elements clearly display periodic behavior. Copyright © Houghton Mifflin Company. All rights reserved. 43

Group IA, Alkali Metals • • • Largest atomic radii React violently with water

Group IA, Alkali Metals • • • Largest atomic radii React violently with water to form H 2 Readily ionized to 1+ Metallic character, oxidized in air Chemical formula is R 2 O in most cases Copyright © Houghton Mifflin Company. All rights reserved. 44

Group IIA, Alkali Earth Metals • • Readily ionized to 2+ React with water

Group IIA, Alkali Earth Metals • • Readily ionized to 2+ React with water to form H 2 Filled s shell configuration Metallic Copyright © Houghton Mifflin Company. All rights reserved. 45

Transition Metals • May have several oxidation states • Metallic • Reactive with acids

Transition Metals • May have several oxidation states • Metallic • Reactive with acids Copyright © Houghton Mifflin Company. All rights reserved. 46

Group III A • Metals (except for boron) • Several oxidation states (commonly 3+)

Group III A • Metals (except for boron) • Several oxidation states (commonly 3+) Copyright © Houghton Mifflin Company. All rights reserved. 47

Group IV A • Form the most covalent compounds • Oxidation numbers vary between

Group IV A • Form the most covalent compounds • Oxidation numbers vary between 4+ and 4 - Copyright © Houghton Mifflin Company. All rights reserved. 48

Group V A • Form anions generally(1 -, 2 -, 3 -), though positive

Group V A • Form anions generally(1 -, 2 -, 3 -), though positive oxidation states are possible • Metals, metalloids, and nonmetals Copyright © Houghton Mifflin Company. All rights reserved. 49

Group VI A • Form 2 - anions generally, though positive oxidation states are

Group VI A • Form 2 - anions generally, though positive oxidation states are possible • React vigorously with alkali and alkali earth metals • Nonmetals Copyright © Houghton Mifflin Company. All rights reserved. 50

Halogens • Form monoanions • High electronegativity (high electron affinity) • Diatomic gases •

Halogens • Form monoanions • High electronegativity (high electron affinity) • Diatomic gases • Most reactive nonmetals (F) Copyright © Houghton Mifflin Company. All rights reserved. 51

Noble Gases • Minimal reactivity • Monatomic gases • Filled valence shell octet rule

Noble Gases • Minimal reactivity • Monatomic gases • Filled valence shell octet rule Copyright © Houghton Mifflin Company. All rights reserved. 52