Chapter 7 Atomic Energies and Periodicity 5.1 Orbital Energies 5.2 Structure of the Periodic Table...

Chapter 7 Atomic Energies and Periodicity

5.1 Orbital Energies

5.2 Structure of the Periodic Table

5.3 Electron Configurations

5.4 Periodicity of Atomic Properties

5.5 Energetics of Ionic Compounds

5.6 Ions and Chemical Periodicity

Chapter 5 Atomic Energies and Periodicity

Chemistry, 2nd Canadian Edition © 2013 John Wiley & Sons Canada, Ltd.


Learning objective:

Explain the effects of nuclear charge and screening on the energies of electrons



In the last chapter, we discussed atoms and energy levels.

When an atom absorbed energy, an electron within the atom was promoted to a higher energy level.

We also learned about the atomic orbitals which are represented by n, l, ml and ms

What are in the orbitals?


Orbital Energy

When a hydrogen atom absorbs light, the energy of the photon converts it from the ground state to an excited state.

In the process, its electron transfers to an orbital that is larger and less stable.

How do we measure the stability of an orbital?


Ionization Energy

Ionization Energy – energy required to remove one electron from an atom

H → H+ + e- IEH = 2.18 x 10-18 J

He+ → He2+(g) + e- IEHe+ = 8.72 x 10-18 J

Why the large difference?


Important Notes on IE

It is more difficult to remove an electron from a positive species.

The total number of electrons in an atom affects the ionization energy.

An electron in a multielectron atom is attracted to the nucleus, but repelled by the other electrons.


ScreeningAs a free electron approaches a positive cation, it will be attracted to the nucleus, but, it is also repelled by any electrons orbiting the nucleus.

The electron-electron repulsion is canceled by the attraction to the nucleus, and this is referred to as screening.

The higher the value of the l quantum number, the more that orbital is screened by electrons in smaller, more stable orbitals.


Electrons are only screened by inner electrons.

Thus, an electron in the 2s orbital is screened only by the 1s.

An electron in the 3s orbital is screened by the 1s and 2s.

Screening


Example 5 - 1

Make an electron density plot showing the 1s, 2p and 3d orbitals to scale. Label the plot in a way that summarizes the screening properties of these orbitals.



Learning objective:

Understand the relationships between the structure of the periodic table and electronic configurations



What is periodicity? Periodic patterns also provide information about electron

arrangements in atoms.


Pauli Exclusion Principle

In general, most systems in nature tend to form a stable state. So, we might expect all electrons to be in the most stable orbital. This is NOT the case! Why?

Pauli Exclusion Principle: each electron in an atom has a unique set of four quantum numbers (n, l, ml, and ms). Stated another way: no two electrons in an atom can have the same set of 4 quantum numbers (n, l, ml, and ms)


The Aufbau Principle

But, we still want the electrons in the most stable state possible (occupying the lowest energy orbitals possible).

The ground state of an atom is the most stable arrangement of its electrons.

Aufbau principle: we place electrons in orbitals starting with the lowest energy orbital following the Pauli Exclusion Principle.


Rules of Ground-State Configurations

1. Each electron in an atom occupies the most stable available orbital. (Aufbau)

2. No two electrons can have identical quantum numbers. (Pauli)

3. Orbital sub-shell capacities are as follows:s: 2 electrons, p set: 6 electrons

d set: 10 electrons, f set: 14 electrons4. The higher the values of n, the less stable the orbital.5. For equal n, the higher the value of l, the less stable

the orbital


Incr

easi

ng E

nerg

y

1 2 3 4

Principle Energy Level, n

1s

2s

2p3s

3p

3d4s


Which orbitals are filled, and which set of orbitals is partially filled in a germanium atom?

Example 5 - 2


Valence Electrons

The chemical properties of an atom are determined by the electrons that are easily accessible to other atoms.

An electron is spatially accessible when it occupies one of the largest orbitals of an atom.

An electron is energetically accessible when it occupies one of the least stable occupied orbitals of the atom.

Accessible electrons are called valence electrons.


Valence Electrons vs. Core Electrons

Valence electrons participate in chemical reactions; core electrons do not.

Valence electrons are those of highest principle quantum number plus those in partially filled d and f orbitals.

Electrons with lower n values are core electrons.



Learning objective:

Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions



Electron configuration: a complete specification of how electrons are distributed within an atom

Quantum numbers Short-hand configurations (s, p, d, f) Energy level diagrams


n = 1, l = 0, ml = 0, ms = + ½ 1s1

n = 1, l = 0, ml = 0, ms = + ½ 1s2

n = 1, l = 0, ml = 0, ms = - ½

n = 1, l = 0, ml = 0, ms = + ½ 1s22s1

n = 1, l = 0, ml = 0, ms = - ½

n = 2, l = 0, ml = 0, ms = + ½


Example 5 - 3

Construct an energy level diagram and the shorthand representation of the ground-state configuration of aluminum. Provide one set of valid quantum numbers for the highest energy electron.


Noble Gas Configurations

As the atomic number increases, so does the number of electrons, and so the electron configuration can get quite lengthy….how can we make it more compact?

Noble Gas Configuration --- specify the noble gas before the element and then build the remaining portion of the electron configuration according to Aufbau.

e.g. Na: 1s22s22p63s1 or Na: [Ne]3s1


Determine the configuration of indium, first in shorthand form and then in full form.

Example 5 - 4


All three are possible, but are they probable? NO

Hund’s Rule: The most stable configuration is the one with the maximum number of electrons with the same spin orientation


Example 5 - 5

Write the shorthand electron configuration and draw the ground-state orbital energy level diagram for the valence electrons in a sulphur atom.


Exceptions…

Orbitals are filled according to the periodic table. However, some elements are bound to have configurations that don’t match the regular progression.

In the first 40 elements, there are only two exceptions: Cr and Cu - what would you expect them to look like?


What is the cause?

Near-Degenerate Orbitals: orbitals with nearly the same energy.

Because the orbitals are so close in energy, it is possible for an electron to be promoted to the higher energy orbital.

More common above Z = 40 because energy levels are closer together.

See Table 5-2 for more exceptions due to near-degenerate orbitals.


Electron Configurations of Ions

1. Write the correct notation for the atom2. For anions: add electrons until you have equaled the

negative charge, fill according to Hund’s Rule3. For cations: remove electrons from the electron shell

of the highest n If there is a choice of a subshell within the nth subshell, the

electron or electrons are removed from the maximum l


Example 5 - 6

What is the ground-state electron configuration of a Cr3+

cation?


Magnetic Properties of Atoms

Substances areDiamagnetic if they are repelled by a strong magnet.Paramagnetic if they are attracted to a magnet, aligns itself in a

magnetic field. Usually paramagnetic materials are weak. Strong paramagnetic materials are called ferromagnetic.

(e.g. refrigerator magnets)


Magnetic Properties of Atoms

Hydrogen has 1 electron Therefore the electron can have either a +½ or a –½ spin This is evident when hydrogen is placed in a magnetic field Hydrogen is paramagnetic, and the electron aligns itself with

the magnetic field

Helium has 2 electrons The electrons are paired and have opposite spins. Therefore, Helium is diamagnetic (the spins cancel each

other)


So…

Paramagnetism – occurs in substances in which the ions or atoms have unpaired electrons

Diamagnetism – occurs in substances in which the ions or atoms have paired electrons


Example 5 - 7

Which of these species is paramagnetic: F-, Zn2+, and Ti?



Learning objective:

Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electronic configuration



Atomic Radii Distance between the nucleus and outer surface of the largest

orbitalIonization Energy

Energy required to remove an electron from an atom in the gas phase

Electron Affinity The energy change when an electron is added to an atom.


Atomic Radius

As n increases, atomic orbitals become larger and less stable.As Z increases, any given atomic orbital becomes smaller and

more stable.


For each of the following pairs, predict which has a larger radius:

Si or ClS or Se Mo or Ag

Example 5 - 7


Ionization Energy

For the main group elements, the IE increases going up and across to the right of the periodic table. Thus, He has the largest IE, and Fr has the smallest IE Going down a group, n increases, thus the atoms are larger and easier

to remove an electron from, thus IE decreases Going across a period, n is constant, but as electrons, protons and

neutrons are added, the result is an increase in Zeff, thus the nucleus has stronger pull on the outer electrons, making them harder to remove


Electron Affinity

For the main group elements, the EA tends to become more negative from left to right across a row of the periodic table.

NO trend is observed for groups


Ionic Radius

Ionic radius follows a similar trend to atomic radii, but

The radius of a cation is always smaller than the corresponding atom e.g. r(Ca2+) < r(Ca)

The radius of an anion is always larger than the corresponding atom e.g. (Br) > r(Br)



Learning objective:

Understand why ionic compounds exist and the energetics of their formation



1. Vaporization and Ionization of atom to form a cation (Na+)Na (s) → Na (g) Esublimation= 105 kJ/molNa (g) → Na+ (g) + e- E = IE = 495.5 kJ/mol

2. Bond breakage and ionization of atom to form an anion (Cl-)½ Cl2(g) → Cl (g) E = ½ BE = 120 kJ/molCl (g) + e- → Cl- (g) E = EA = -348.5 kJ/mol

3. Condensation of the two to form the solid saltNa+(g) + Cl-(g) → NaCl (s) Elattice = -769 kJ/mol

4. Overall Na (s) + ½ Cl2 (g) → NaCl (s)

E = 105 + 495.5 + 120 – 348.5 – 769 = -397 kJ/mol

Experimentally:Na (s) + ½ Cl2 (g) → NaCl (s) Eexp. = -411 kJ/mol


Energy and Ions

k is a constantr is the distance between the two ionsq1 is the electronic charge of the cationq2 is the electronic charge of the anion

1 2electrical

q qE = k

r


Lattice Energy

Lattice Energy – the energy released when an ionic solid decomposes into gaseous ions

NaCl (s) → Na+(g) + Cl-(g) E = LE = 769 kJ/mol

The magnitude of the lattice energy depends upon the charges of the ions, their size, and the particular lattice arrangement.



Learning objective:

Understand the trends in atomic radius, ionization energy and electron affinity and their relationships to nuclear charge and electronic configuration



Alkali Metalsns1

Alkaline Earth Metalsns2

Transition Metals ndm

Lanthanides/Actinidesnfm


Chapter 5 Visual Summary


Chapter 7 Atomic Energies and Periodicity 5.1 Orbital Energies 5.2 Structure of the Periodic Table...

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Transcript of Chapter 7 Atomic Energies and Periodicity 5.1 Orbital Energies 5.2 Structure of the Periodic Table...