Periodic Trends in Ionization Energy and Electron Affinity, Slides of Chemistry

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Periodic Properties of the

Elements

Development of the Periodic Table

Discovering the Elements

Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped. (^1871) O

In this chapter

• Effective Nuclear Charge
• Sizes of Atoms and Ions
• Ionization Energy
• Electron Affinity
• Metals, Nonmetals, and Metalloids
• Trends for Group 1A and Group 2A
• Trends for Selected Nonmetals P exam 3

Effective Nuclear Charge

• Many properties of atoms depend on electron configurations and on how strongly the outer electrons in the atoms are attracted to the nucleus
• The strength of the attraction between electrical charges depends on the magnitudes of the charges and on the distance between them (Coulomb’s Law)
• electrostatic potential energy. Eel ( interaction between charged particles) κ Q 1 Q 2 Eel =

---

d

• The attractive force between an electron and the nucleus depends on the magnitude of the nuclear charge and on the average distance between the nucleus and the electron É É O charges distance

Effective Nuclear Charge

• Electron-electron repulsion

cancel some of the attraction of

the electron to the nucleus so the

electron experience less

attraction.

• Electron experiences a net

attraction that is the result of the

nuclear attraction decreased by

the electron-electron repulsions.

• The partially screened nuclear

charge is called the effective

nuclear charge Z eff

Na

s

nucleus (^) I Il inner protons electrons sat Ft

Effective Nuclear Charge

• The effective nuclear charge, Z eff, is found this way: Z eff = Z − S
• where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons. Na 7 1135 Nt I

atomicnumbers^

protons Test L^ S 11 10 It Zeff actual 2.5T

• Zeff increases from left to right across any period of the periodic table. o The number of core electrons stays the same across the period. o The number of protons increases
• Zeff changes far less going down a column than it does across the period
• Zeff increases slightly as we go down a column because of the more diffuse the core electron cloud is less able to screen the valence electrons from the nuclear charge.
• Li 1.3+
• Na 2.5+
• K 3.5+ Effective Nuclear Charge stinketeledron

is

me jp

g

core 9

Example Li Na K IYÉ Zeff

increase P

ie iKNas __ greatest for Zeff Ex (^) Be (^) N C Becc n

Sizes of Atoms and Ions

Consider a collection of argon atoms

in the gas phase.

• When they undergo collisions,

they ricochet apart because

electron clouds cannot penetrate

each other to a significant extent.

The apparent radius is

determined by the closest

distances separating the nuclei

during such collisions.

• This radius is the nonbonding

radius.

• Nonbonding atomic radii are also called van der Waals radii.
• These are used in space-filling models to represent the sizes of different elements.

Sizes of Atoms and Ions

• Now consider a simple diatomic molecule. - The distance between the

two nuclei is called the

bonding atomic radius.

• It is shorter than the nonbonding radius.
• If the two atoms that make up the molecule are the same, then half the bond distance is called the covalent radius of the atom. a __ bonding atomic radius I

Fig 7.

Bonding Atomic Radii for periods 1 through 5 bond length^ C Cl C C^ bondingradius (^) t bonding rad Cl

C

0.77 0. 1.76 (^) A O O O

Atomic size varies consistently through the periodic table.

• As we move down a group, the atoms become larger.
• As we move across a period, atoms become smaller. O j

I n 2 t

n 3 fisinfreases

Izesitsignly

nss

in a^ period T right Zeff T core é's^

same

size decreases

Using Figure 7.

• We can estimate bond lengths in molecules
• For example, the bonding atomic radii for C and Cl are 1.02 Å and 0.76 Å, respectively.
• The sum of the bonding atomic radii of Cl and C = 1.02 Å + 0.76 Å =1.78 Å
• In CCl 4 the measured length of C-Cl bond is 1.77 Å