Fig. 7.14

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Fig. 7.14
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Chapter 8Electron Configurations and Periodicity
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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 Hunds Rule

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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 Hunds
Rule
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• 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.
• Mendeleevs Predictions from the Periodic Table
• Some Periodic Properties
• Periodicity in the Main-Group Elements

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• The figure shows relative energies for the
  hydrogen atom shells and subshells each orbital
  is indicated by a dashed-line.

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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 4d 4f
• 5s 5p 5d 5f
• 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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Quantum Numbers Noble Gases
Electron Orbitals
Number of Electrons Element
1s2
2
He
1s2 2s22p6
10
Ne
1s2 2s22p6 3s23p6
18 Ar
1s2 2s22p6 3s23p6 4s23d104p6
36 Kr
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6
54 Xe
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6
6s24f14 5d106p6 86 Rn
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6
6s24f145d106p6 7s25f146d10?
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• There are a few exceptions to the building-up
  order prediction for the ground state.
• Chromium (Z24) and copper (Z29) 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.

• 4p5

• 3s2 3p6

• 4s2

• Br 1s2

• 2s2 2p6

• 3d10

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• The noble-gas configuration substitutes the
  preceding noble gas for the core configuration
  and explicitly shows subshells beyond that.
• Br Ar3d104s24p5

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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
• Ar3d10

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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.

For arsenic, As, Z 33.
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p3
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• 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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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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• In 1927, Friedrich Hund discovered, by
  experiment, a rule for determining the
  lowest-energy configuration of electrons in
  orbitals of a subshell.
• Hunds 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

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

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?

• 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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• 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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• Mendeleevs 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.

Te is larger than Se. Se is larger than Br.
Br lt Se lt 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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• 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 lt As lt Br
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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 electron affinity is gt 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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• Broadly speaking, the trend is toward more
  negative electron affinities going from left to
  right in a period.
• Lets 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 Valence
• IA ns1 ns2 stable
• IIIA ns2np1 ns2np2 stable
• IVA ns2np2 ns2np3 stable
• VA ns2np3 ns2np4 not so stable
• VIA ns2np4 ns2np5 very stable
• VIIA 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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• Some oxides of Group IVA

• PbO
• (yellow)

• PbO2
• (dark brown)

• SnO2 (white)

• SiO2 (crystalline solid quartz)

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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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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.