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Chapter 8 Electron Configurations and Periodicity. Contents and Concepts. Electronic Structure of Atoms

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contents and concepts
Contents and Concepts

Electronic Structure of Atoms

In the previous chapter, you learned that we characterize an atomic orbital by four quantum numbers: n, l, ml, and ms. In the first section, we look further at electron spin; then we discuss how electrons are distributed among the possible orbitals of an atom.

  • Electron Spin and the Pauli Exclusion Principle
  • Building-Up Principle and the Periodic Table
  • Writing Electron Configurations Using the Periodic Table
  • Orbital Diagrams of Atoms; Hund’s Rule

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slide3
Periodicity of the Elements

You learned how the periodic table can be explained by the periodicity of the ground-state configurations of the elements. Now we will look at various aspects of the periodicity of the elements.

  • Mendeleev’s Predictions from the Periodic Table
  • Some Periodic Properties
  • Periodicity in the Main-Group Elements

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schr dinger wave equation
Schrödinger Wave Equation!

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slide5
In 1921, Otto Stern and Walther Gerlach first observed electron spin magnetism. In the diagram below, a beam of hydrogen atoms divides in two while passing through a magnetic field. This correlates with the two values of ms: +½ and -½.

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slide6
The two possible spin orientations of an electron and the conventions for msare illustrated here.

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slide7
An electron configuration of an atom is a particular distribution of electrons among available subshells.
  • An orbital diagram of an atom shows how the orbitals of a subshell are occupied by electrons. Orbitals are represented with a circle; electrons are represented with arrows up for ms= +½ or down for ms= -½.

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slide8
The Pauli exclusion principle summarizes experimental observations that no two electrons in one atom can have the same four quantum numbers.
  • That means that within one orbital, electrons must have opposite spin. It also means that one orbital can hold a maximum of two electrons (with opposite spin).

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slide9
An s subshell, with one orbital, can hold a maximum of 2 electrons.
  • A p subshell, with three orbitals, can hold a maximum of 6 electrons.
  • A d subshell, with five orbitals, can hold a maximum of 10 electrons.
  • An f subshell, with seven orbitals, can hold a maximum of 14 electrons.

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slide10
The lowest-energy configuration of an atom is called its ground state.
  • Any other configuration represents an excited state.

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slide11
The building-up principle (or aufbau principle) is a scheme used to reproduce the ground-state electron configurations by successively filling subshells with electrons in a specific order (the building-up order).
  • This order generally corresponds to filling the orbitals from lowest to highest energy. Note that these energies are the total energy of the atom rather than the energy of the subshells alone.

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slide12
1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s 7p

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slide13
This results in the following order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

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slide14
Another way to learn the building-up order is to correlate each subshell with a position on the periodic table.
  • The principal quantum number, n, correlates with the period number.
  • Groups IA and IIA correspond to the s subshell; Groups IIIA through VIIIA correspond to the p subshell; the “B” groups correspond to the d subshell; and the bottom two rows correspond to the f subshell. This is shown on the next slide.

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slide16
There are a few exceptions to the building-up order prediction for the ground state.
  • Chromium (Z=24) and copper (Z=29) have been found by experiment to have the following ground-state electron configurations:
  • Cr: 1s2 2s2 2p6 3s2 3p6 3d5 4s1
  • Cu: 1s2 2s2 2p6 3s2 3p6 3d10 4s1
  • In each case, the difference is in the 3d and 4s subshells.

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slide17
There are several terms describing electron configurations that are important.
  • The complete electron configuration shows every subshell explicitly.
  • Br: 1s2
  • 2s2 2p6
  • 3s2 3p6
  • 4s2
  • 4p5
  • 3d10

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slide18
The noble-gas configuration substitutes the preceding noble gas for the core configuration and explicitly shows subshells beyond that.
  • Br: [Ar]3d104s24p5

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slide19
The pseudo-noble-gas core includes the noble-gas subshells and the filled inner, (n – 1), d subshell.
  • For bromine, the pseudo-noble-gas core is
  • [Ar]3d10

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slide20
The valence configuration consists of the electrons outside the noble-gas or pseudo-noble-gas core.
  • Br: 4s24p5

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slide21
For main-group (representative) elements, an s or a p subshell is being filled.
  • For d-block transition elements, a d subshell is being filled.
  • For f-block transition elements, an f subshell is being filled.

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slide22
For main-group elements, the valence configuration is in the form
  • nsAnpB
  • The sum of A and B is equal to the group number.
  • So, for an element in Group VA of the third period, the valence configuration is
  • 3s23p3

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slide23
Write the complete electron configuration of the arsenic atom, As, using the building-up principle.

For arsenic, As, Z = 33.

1s2

2s2

2p6

3s2

3p6

3d10

4s2

4p3

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slide24
What are the electron configurations for the valence electrons of arsenic and cadmium?

Arsenic is in period 4, Group VA.

Its valence configuration is

4s24p3.

Cadmium, Z = 30, is a transition metal in the first transition series. Its noble-gas core is Ar, Z = 18.

Its valence configuration is

4s23d10.

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slide25

When n = 2, there are two subshells.

The s subshell has one orbital, which could hold one electron.

The p subshell has three orbitals, which could hold three electrons.

This would give a total of

four elements for the second period.

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slide26
In 1927, Friedrich Hund discovered, by experiment, a rule for determining the lowest-energy configuration of electrons in orbitals of a subshell.
  • Hund’s rule states that the lowest-energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons.

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slide27

1s

2s

2p

  • For nitrogen, the orbital diagram would be

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slide28

1s

3s

2s

4s

3d

3p

2p

  • Write an orbital diagram for the ground state of the nickel atom.

For nickel, Z = 28.

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slide29

1s

2s

  • Which of the following electron configurations or orbital diagrams are allowed and which are not allowed by the Pauli exclusion principle? If they are not allowed, explain why?
  • Allowed; excited.
  • p8 is not allowed.
  • Allowed.
  • d11 is not allowed.
  • Not allowed; electrons in one orbital must have opposite spins.
  • 1s22s12p3
  • 1s22s12p8
  • 1s22s22p63s23p63d8
  • 1s22s22p63s23p63d11

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slide30
Magnetic Properties of Atoms
  • 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.
  • This allows us to classify atoms based on their behavior in a magnetic field.

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slide31
A paramagnetic substance is one that is weakly attracted by a magnetic field, usually as the result of unpaired electrons.
  • A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons.

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slide32
You learned how the organization of the periodic table can be explained by the periodicity of the ground-state configurations of the elements. Now we will look at various aspects of the periodicity of the elements.

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slide33
Mendeleev’s periodic table generally organized elements by increasing atomic mass and with similar properties in columns. In some places, there were missing elements whose properties he predicted.
  • When gallium, scandium, and germanium were isolated and characterized, their properties were almost identical to those predicted by Mendeleev for eka-aluminum, eka-boron, and eka-silicon, respectively.

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slide34
Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically.
  • We will look in more detail at three periodic properties: atomic radius, ionization energy, and electron affinity.

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slide35
Atomic Radius
  • While an atom does not have a definite size, we can define it in terms of covalent radii (the radius in covalent compounds).

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slide36
Trends
  • Within each group (vertical column), the atomic radius increases with the period number.
  • This trend is explained by the fact that each successive shell is larger than the previous shell.

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slide37
Within each period (horizontal row), the atomic radius tends to decrease with increasing atomic number (nuclear charge).

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slide38
Effective Nuclear Charge
  • Effective nuclear charge is the positive charge that an electron experiences from the nucleus. It is equal to the nuclear charge, but is reduced by shielding or screening from any intervening electron distribution (inner shell electrons).

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slide39
Effective nuclear charge increases across a period. Because the shell number (n) is the same across a period, each successive atom experiences a stronger nuclear charge. As a result, the atomic size decreases across a period.

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slide40
Atomic radius is plotted against atomic number in the graph below. Note the regular (periodic) variation.

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slide41

A representation of atomic radii is shown below.

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slide42

34

Se

35

Br

52

Te

  • Refer to a periodic table and arrange the following elements in order of increasing atomic radius: Br, Se, Te.

Te is larger than Se.

Se is larger than Br.

Br < Se < Te

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slide43
First Ionization Energy (first ionization potential)
  • The minimum energy needed to remove the highest-energy (outermost) electron from a neutral atom in the gaseous state, thereby forming a positive ion

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slide44
Trends
  • Going down a group, first ionization energy decreases.
  • This trend is explained by understanding that the smaller an atom, the harder it is to remove an electron, so the larger the ionization energy.

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slide45
Generally, ionization energy increases with atomic number.
  • Ionization energy is proportional to the effective nuclear charge divided by the average distance between the electron and the nucleus. Because the distance between the electron and the nucleus is inversely proportional to the effective nuclear charge, ionization energy is inversely proportional to the square of the effective nuclear charge.

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slide46
Small deviations occur between Groups IIA and IIIA and between Groups VA and VIA.
  • Examining the valence configurations for these groups helps us to understand these deviations:
  • IIA ns2
  • IIIA ns2np1
  • VA ns2np3
  • VIA ns2np4

Ittakes less energy to remove the np1 electron than the ns2 electron.

It takes less energy to remove the np4 electron than the np3 electron.

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slide47

These trends and reversals are visible in the graph of ionization energy versus atomic number.

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slide48
The size of each sphere indicates the size of the ionization energy in the figure below.

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slide49
Electrons can be successively removed from an atom. Each successive ionization energy increases, because the electron is removed from a positive ion of increasing charge.
  • A dramatic increase occurs when the first electron from the noble-gas core is removed.

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slide50
Left of the line, valence shell electrons are being removed. Right of the line, noble-gas core electrons are being removed.

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slide51

33

As

35

Br

51

Sb

  • Refer to a periodic table and arrange the following elements in order of increasing ionization energy: As, Br, Sb.

Sb is larger than As.

As is larger than Br.

Ionization energies:

Sb < As < Br

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slide52
Electron affinity (E.A.)
  • The energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion
  • A negative energy change (exothermic) indicates a stable anion is formed. The larger the negative number, the more stable the anion. Small negative energies indicate a less stable anion.
  • A positive energy change (endothermic) indicates the anion is unstable.

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slide54

The electron affinity is > 0, so the element must be in Group IIA or VIIIA.

The dramatic difference in ionization energies is at the third ionization.

The element is in Group IIA.

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slide55
Broadly speaking, the trend is toward more negative electron affinities going from left to right in a period.
  • Let’s explore the periodic table by group.

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slide56
Groups IIA and VIIIA do not form stable anions; their electron affinities are positive.
  • Group Valence Anion Valence
  • IA ns1ns2 stable
  • IIIA ns2np1 ns2np2stable
  • IVA ns2np2 ns2np3 stable
  • VA ns2np3 ns2np4 not so stable
  • VIA ns2np4ns2np5 very stable
  • VIIA ns2np5ns2np6 very stable
  • Except for the members of Group VA, these values become increasingly negative with group number.

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slide57
Metallic Character
  • Elements with low ionization energies tend to be metals. Those with high ionization energies tend to be nonmetals. This can vary within a group as well as within a period.

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slide58
Oxides
  • A basic oxide reacts with acids. Most metal oxides are basic. If soluble, their water solutions are basic.
  • An acidic oxide reacts with bases. Most nonmetal oxides are acidic. If soluble, their water solutions are acidic.
  • An amphoteric oxide reacts with both acids and bases.

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slide59
Group IA, Alkali Metals (ns1)
  • These elements are metals; their reactivity increases down the group.
  • The oxides have the formula M2O.
  • Hydrogen is a special case. It usually behaves as a nonmetal, but at very high pressures it can exhibit metallic properties.

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slide60
Group IIA, Alkaline Earth Metals (ns2)
  • These elements are metals; their reactivity increases down the group.
  • The oxides have the formula MO.

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slide61
Group IIIA (ns2np1)
  • Boron is a metalloid; all other members of Group IIIA are metals.
  • The oxide formula is R2O3.
  • B2O3 is acidic; Al2O3 and Ga2O3 are amphoteric; the others are basic.

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slide62
Group IVA (ns2np2)
  • Carbon is a nonmetal; silicon and germanium are metalloids; tin and lead are metals.
  • The oxide formula is RO2 and, for carbon and lead, RO.
  • CO2, SiO2, and GeO2 are acidic (decreasingly so).
  • SnO2 and PbO2 are amphoteric.

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slide63

Some oxides of Group IVA

  • PbO
  • (yellow)
  • PbO2
  • (dark brown)
  • SnO2 (white)
  • SiO2(crystalline solid quartz)

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slide64
Group VA (ns2np3)
  • Nitrogen and phosphorus are nonmetals; arsenic and antimony are metalloids; bismuth is a metal.
  • The oxide formulas are R2O3 and R2O5, with some molecular formulas being double these.
  • Nitrogen, phosphorus, and arsenic oxides are acidic; antimony oxides are amphoteric; bismuth oxide is basic.

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slide65
Group VIA, Chalcogens (ns2np4)
  • Oxygen, sulfur, and selenium are nonmetals; tellurium is a metalloid; polonium is a metal.
  • The oxide formulas are RO2 and RO3.
  • Sulfur, selenium, and tellurium oxides are acidic except for TeO2, which is amphoteric. PoO2 is also amphoteric.

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slide66
Group VIIA, Halogens (ns2np5)
  • These elements are reactive nonmetals, with the general molecular formula being X2. All isotopes of astatine are radioactive with short half-lives. This element might be expected to be a metalloid.
  • Each halogen forms several acidic oxides that are generally unstable.

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slide67
Group VIIIA, Noble Gases (ns2np6)
  • These elements are generally unreactive, with only the heavier elements forming unstable compounds. They exist as gaseous atoms.

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slide68

For R2O5 oxides, R must be in Group VA.

R is a metalloid, so R could be As or Sb.

The oxide is acidic, so

R is arsenic, As.

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other resources
Other Resources
  • Visit the student website at http://www.college.hmco.com/pic/ebbing9e

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