Chapter 8Electron

Configurations and Periodicity

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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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The two possible spin orientations of an electron and the conventions for ms are illustrated here.

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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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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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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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The lowest-energy configuration of an atom is called its ground state.

Any other configuration represents an excited state.

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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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1s

2s 2p

3s 3p 3d

4s 4p 4d4f

5s 5p 5d5f

6s 6p 6d

7s 7p

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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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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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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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There are several terms describing electron configurations that are important.

The complete electron configuration shows every subshell explicitly.

Br: 1s2 4p52s2 2p6 3s2 3p6 4s23d10

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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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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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The valence configuration consists of the electrons outside the noble-gas or pseudo-noble-gas core.

Br: 4s24p5

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

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3

For arsenic, As, Z = 33.

?

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

Arsenic is in period 4, Group VA.Its valence configuration is 4s24p3.

Zinc, 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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This would give a total of

four elements for the second period.

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.

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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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For nitrogen, the orbital diagram would be

1s 2s 2p

?

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

3s 3p

1s 2s 2p

4s 3d

For nickel, Z = 28.

?

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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?

a.1s22s12p3

b.1s22s12p8

c. 1s22s22p63s23p63d8 d.1s22s22p63s23p63d11

e.

1s 2s

a.Allowed; excited.b. p8 is not allowed.c. Allowed; excited.d. d11 is not allowed.e. Not allowed;

electrons in one orbital must have opposite spins.

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

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

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A representation of atomic radii is shown below.

?

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

35Br

34Se

52Te

Te is larger than Se.Se is larger than Br.

Br < Se < Te

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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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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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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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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

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

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

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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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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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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

35Br

33As

51Sb

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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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The element is in Group IIA.

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.

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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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Groups IIA and VIIIA do not form stable anions; their electron affinities are positive.

Group Valence Anion ValenceIA ns1 ns2 stableIIIA ns2np1 ns2np2 stableIVA ns2np2 ns2np3 stableVA ns2np3 ns2np4 not so stable

VIA ns2np4 ns2np5 very stableVIIA ns2np5 ns2np6 very stable

Except for the members of Group VA, these values become increasingly negative with group number.

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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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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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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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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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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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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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PbO (yellow) PbO2

(dark brown)

SiO2 (crystalline solid quartz)

SnO2 (white)

Some oxides of Group IVA

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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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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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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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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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R is arsenic, As.

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