The Periodic Table

Chapter Outline..

-To know the periodic classification of the elements.

-Able to construct its electron configuration.

-Identify the periodic variation in physical properties:– Effective nuclear charge (Zeff)

– Atomic & ionic radii– Ionization Energy– Electron Affinity

- Identify the periodic variation in chemical properties:04/22/23

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The Current Periodic Table

• Mendeleev wasn’t too far off.• Now the elements are put in rows by increasing

ATOMIC NUMBER!!• The horizontal rows are called periods and are

labeled from 1 to 7.• The vertical columns are called groups are

labeled from 1 to 18.

Groups…Here’s Where the Periodic Table Gets Useful!!

• Elements in the same group have similar chemical and physical properties!!

• (Mendeleev did that on purpose.)

Why??• They have the same

number of valence electrons.

• They will form the same kinds of ions.

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Classification of the Elements

ns1

ns2

ns2

np1

ns2

np2

ns2

np3

ns2

np4

ns2

np5

ns2

np6

d1

d5 d10

4f

5f

Ground State Electron Configurations of the Elements

PERIODIC CLASSIFICATION OF THE ELEMENTS

n = principal quantum number of the outermost

subshell

Table 8.2 Summary of Quantum Numbers of Electrons in Atoms

Name Symbol Permitted Values Property

principal n positive integers(1,2,3,…)

orbital energy (size)

angular momentum

l integers from 0 to n-1 orbital shape (The l values 0, 1, 2, and 3 correspond to s, p, d, and f orbitals, respectively.)

magnetic mlintegers from -l to 0 to +l

orbital orientation

spin ms+1/2 or -1/2 direction of e- spin

Refer textbook: Martin S. Silberberg, Chemistry, 4th Edition, McGraw-Hill Higher Education, New York (2006).

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SAMPLE PROBLEM 8.1 Determining Quantum Numbers from Orbital Diagrams

PLAN:

SOLUTION:

Use the orbital diagram to find the third and eighth electrons.

PROBLEM: Write a set of quantum numbers for the third electron and a set for the eighth electron of the F atom.

9F

1s 2s 2p

The third electron is in the 2s orbital. Its quantum numbers are

n = l = ml = ms= 2 0 0 + or -1/2

The eighth electron is in a 2p orbital. Its quantum numbers are

2n = l = ml = ms= -1, 0, or +11 + or -1/2

Refer textbook: Martin S. Silberberg, Chemistry, 4th Edition, McGraw-Hill Higher Education, New York (2006).

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In the ground state of the atom, electrons will occupy the lowest energy orbitals first, and only fill the higher energy orbitals when no lower energy orbitals are left.

Hund’s first rule:- electrons occupy all the orbitals of a given subshell singly before pairing begins. These unpaired electrons have parallel spins. Pauli exclusion principle:- no

two electrons in the same atom may have the same set of n, l, ml, ms quantum numbers

Fig. Order for filling energy sublevels with electrons.

Orbitals:Order of Filling

Electron Configurations of Cations and Anions

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that cation has a stable noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Atoms gain electrons so that anion has a stable noble-gas outer electron configuration.

Of Representative Elements

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+1

+2

+3 -1-2-3

Cations and Anions of Representative Elements..

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Isoelectronic ions

• Iso- means “the same”• Isoelectronic ions have the same # of electrons• Al3+ Mg2+ Na1+ Ne F1- O2- and N3-

– all have 10 electrons• all have the same configuration: 1s22s22p6 (which is the noble

gas: neon)

Electron Configurations of Cations of Transition Metals

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When a cation is formed from an atom of a transition metal,electrons are always removed first from the ns orbital and then from the (n –

1)d orbitals.

Fe: [Ar]4s23d6

Fe2+: [Ar]4s03d6 or [Ar]3d6

Fe3+: [Ar]4s03d5 or [Ar]3d5

Examples:-

Mn: [Ar]4s23d5

Mn2+: [Ar]4s03d5 or [Ar]3d5

Mn2+: [Ar]4s23d3

3d orbital is more stable than the 4s orbital in transition metal ions..

……1

….. 2

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ALL Periodic Table Trends • Influenced by three factors:

1. Energy Level– Higher energy levels are further away from the nucleus.

2. Charge on nucleus (# protons)– More charge pulls electrons in closer. (+ and – attract each

other)

3. Shielding effect

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Periodic Variation in Physical PropertiesA. Effective nuclear charge (Zeff) It is the “positive charge” felt by an electron… Given by Zeff = Z – Z= actual nuclear charge,

= shielding/screening constant

(0 < < Z)

Zeff Z – number of inner or core electrons

Na

Mg

Al

Si

11

12

13

14

10

10

10

10

1

2

3

4

186

160

143

132

ZeffCoreZ Radius

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In a many-electron atom,electrons are both attractedto the nucleus and repelledby other electrons.• The nuclear charge that anelectron “feels” depends onboth factors.• It’s called Effective nuclearcharge.• electrons in lower energylevels “shield” outerelectrons from positivecharge of nucleus.

Continued……

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increasing Zeffin

crea

sing

Ze

ff Li Be B C N O F Ne

3 4 5 6 7 8 9 10Core electron = 1s2

1. Core electron > closer to nucleus than valence electron, thus core e shield valence e> than valence e shield each other.2. Moving across the period, core e remains constant , but Z increases.3. The added e will be valence e, and due to valence e does not shield each other, thus, moving across the period, > Zeff will be felt by valence e.4. Moving the group, Zeff . As n increases, large shells increases, thus valence e are added to these large shells. Thus, electrostatic attraction between nuclues & valence e decreases.

Atomic Size

• Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

}Radius

Atomic Size - Group trends

• Increases down a given group. – the number of electrons and

filled electron shells increases– electrons are found farther from

the nucleus – Therefore, the atomic radii

increase.

HLi

Na

K

Rb

#1. Atomic Size - Period Trends• Decreases across a period from left to right • Electrons are in the same energy level. and unshielded towards

attraction by protons. protons are being added to the nucleus thus creates a "higher

effective nuclear charge." – stronger force of attraction pulling the electrons closer to the

nucleus

Na Mg Al Si P S Cl Ar

Ionic Radii

• 1) Anions (negative ions) are larger than their respective atoms. – WHY? – Electron-electron repulsion forces them to spread further apart. – the protons cannot pull the extra electrons as tightly toward the nucleus.

• 2) Cations (positive ions) are smaller than their respective atoms.WHY? – There is less electron-electron repulsion.– Protons outnumber electrons; the protons can pull the fewer electrons

toward the nucleus more tightly.

First Ionization Energy Definition: The energy required to remove the outermost (highest energy) electron

from a neutral atom in its ground state (from a gaseous atom) . • 1) down a group, first ionization energy decreases.

– Electrons are further from the nucleus – more shielding – easier to remove the outermost electron

• 2) across a period, first ionization energy increases. – the atomic radius decreases– The outer electrons are closer to the nucleus and more strongly attracted to

the center– Similar shielding effect– more difficult to remove the outermost electron.

Figure 8.12

First ionization energies of the main-group elements.

What factors determine IE

• The greater the nuclear charge, the greater IE.

• Greater distance from nucleus decreases IE

• Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.

• Shielding effect

Electron Affinity

• Definition: The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.

• 1) down a group, electron affinity decreases.

• 2) across a period, electron affinity increases.

Trends in metallic behavior.

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Summary…..

Trends in three atomic properties.

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