Ch. 3: Periodic Properties of the...

Ch. 3: Periodic Properties of

the Elements

Dr. Namphol Sinkaset

Chem 200: General Chemistry I

I. Chapter Outline

I. Introduction

II. The Periodic Table

III. Electrons in the Atom

IV. Electron Spin

V. Sublevel Energy Splitting

VI. Using the Periodic Table

VII. Periodic Properties and Trends

I. Organizing Chemical Info

• When information of the elements was

organized, chemistry began to advance

quickly.

• Element “triads” and “octaves”

• Mendeleev’s periodic table in 1869

• Quantum mechanics explains why the

periodic table appears as it does.

II. Periodic Law

• Initially, Mendeleev ordered elements

by increasing atomic mass.

• Later work by Moseley showed that they

should be ordered by atomic number.

II. The Modern Periodic Table

II. Major Divisions of the Table

• Main-group elements have properties

that are largely predictable based on

their location.

• Transition and inner-transition elements

have properties that are less predictable

based on their location.

• Each column within the main group

region is known as a family or group.

III. Electrons Occupying Orbitals

• From Chapter 3, we know how orbitals

are ordered for the hydrogen atom

• Since hydrogen has only one e-, the

ground state can be written as an

electron configuration:

III. Many e- Atoms

• The Schrödinger equation can’t solve

multi-e- atoms; we only get approximate

solutions.

• We use quantum #’s from H atom

solution to describe orbitals of other

atoms.

III. New Considerations

• An atom with more than 1 e- is more

complicated.

• Two more concepts are needed to

understand these larger atoms:

1) Electron spin

2) Sublevel energy splitting

IV. H Atoms in a Magnetic Field

IV. e- Spin

• e- generate a small magnetic field as if they were spinning.

• There are two possible directions e- can spin, so there are two possible states.

• spin quantum number (ms) can be either +1/2 or –1/2.

IV. Representing e- Spin

• Orbital diagrams are used to show

electron occupation and spin.

IV. Pauli Exclusion Principle

• No two e- in the same atom can have the same 4 quantum #’s!!

• H: n=1, l=0, ml=0, ms=1/2

• He has two p+, so it needs two e-:

1st e-: n=1, l=0, ml=0, ms=1/2

2nd e-: n=1, l=0, ml=0, ms=-1/2

• The orbital is filled and the e- have paired spins.

IV. Electrons in Helium

V. H vs. He Energy Levels

• One additional e- complicates the He

spectrum greater than expected. Why?

V. Removal of Degeneracy

• In H atom, energy of an orbital depends

only on n.

e.g. Energies of 3s, 3p, 3d are degenerate.

• In every other atom, this is not true.

E (s orbital) < E (p orbital) < E (d orbital) <

E (f orbital), etc.

• What removes the degeneracy?

V. Sublevel Energy Splitting

• Three factors contribute to differing

sublevel energies:

1) Coulomb’s Law (Z)

2) shielding

3) penetration

V. Coulomb’s Law

• The PE of like charges is positive (unstable),

but decreases as they move apart.

• The PE of unlike charges is negative (stable)

and increases as they get closer.

• The magnitude of the interaction increases as

charges on particles increases.

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V. Nuclear Charge

• p+ in nucleus

constantly pull all e-.

• Higher charges

attract more strongly.

• More p+ lowers

orbital E by

increasing e-/nucleus

attraction.

V. Shielding

• Electrons shield each

other from the full

charge of the nucleus.

• The effective nuclear

charge, Zeff, is the

actual positive charge

an e- feels.

V. Penetration

• The movement of an

outer e- into the

region occupied by

inner e- is called

penetration.

• Penetrating e-

experience higher

nuclear charge,

lowering its PE.

V. 2s and 2p Radial Distribution

V. 3s, 3p, 3d Penetration

• This is the reason why energetically, s < p < d.

V. Order of Sublevels

V. The Aufbau Principle

• Since e- are “lazy,” they want to

“occupy” the lowest energy level

possible.

• Thus, if we know the energy order of

sublevels, then we can “build up” the e-

configurations of each atom.

V. Writing e- “in” Orbitals

• Two ways to represent how e- are

situated in atoms:

1) e- configuration, nl#

2) orbital diagram, which uses arrows

indicating e-’s and their spin

V. Hund’s Rule

• In the orbital diagram of C, there was a

choice as to where to place the 2nd p

orbital.

• We follow Hund’s rule.

When filling degenerate orbitals, electrons

fill singly first with parallel spins.

• Hund’s rule leads to lower energy.

V. Examples

VI. The Periodic Table

• As you go left to right on the periodic table,

you are using the Aufbau principle.

VI. The Periodic Table

• Each region of the periodic table indicates

what orbitals are being “filled.”

VI. Using the Periodic Table

• You can use an element’s location to

write its full or condensed electron

configuration/orbital diagram.

VI. Using the Periodic Table

• Therefore, Cl is: [Ne] 3s2 3p5.

• From the orbital diagram, we can write

specific quantum numbers for each e-.

• Which e-’s are identified with the

following quantum #’s {n, l, ml, ms}?

{3, 0, 0, -1/2}

{3, 1, 1, 1/2}

VI. Some Caveats

• Because energy differences between s

and d are small, some exceptions to

how e-’s fill exist.

Same for d and f.

• Remember that d principal quantum #

lags by one.

• Remember that f principal quantum #

lags by two.

VI. Sample Problem 3.1

• Write condensed electron configurations

and orbital diagrams for the following

elements.

Mn

Sb

Nd

VI. The Periodic Table

VI. Important Parts of the Periodic Table

1) Each element placed in box w/ atomic #,

atomic mass, and atomic symbol.

2) Atomic # increases as go L to R.

3) Each horizontal row is period.

4) Each vertical column is a group or family.

5) Main group elements are in groups 1,2

and 13-18 (s and p blocks).

VI. Important Parts of the Periodic Table

6) Transition elements are in groups 3-12 (d

block).

7) Inner-transition elements at the bottom

(lanthanides and actinides, f block).

8) Staircase line separates metals on L

from nonmetals on R. Metalloids or

semimetals lie adjacent to the line.

9) Some groups have special names: alkali

metals, alkali earth metals, halogens,

noble gases.

VI. Types of Elements

VI. Core vs. Valence e-’s

VI. Valence Electrons

• valence electrons: the outermost e- in an atom

• Valence e- determine an atom’s chemistry; thus, atoms in the same vertical column have similar chemical properties.

• Valence e- can be determined from the Group number.

VI. Formation of Ions

• Metals tend to lose e-’s and nonmetals

tend to gain e-’s.

• Main-group ions can be predicted.

VI. Transition Metal Cations

• When forming transition metal cations,

remove e-’s from highest n-value orbital

first!

V: [Ar] 4s2 3d3

V2+: [Ar] 4s0 3d3

VI. Magnetic Properties

• Some metals exhibit magnetism

paramagnetic: atom or ion that has

unpaired e-’s

diamagnetic: atom or ion in which all e-’s

are paired

VI. Sample Problem 3.2

• Draw condensed orbital diagrams for

the following and determine whether

they are diamagnetic or paramagnetic.

Sc3+

Ir2+

Mn4+

VII. Atomic Radii

VII. Trend in Atomic Radii

Trend in Atomic Radii

VII. Trend in Ion Size

• Why?

VII. Trend in Ionization Energy

• ionization energy:

energy in kJ needed to

remove an e- from

gaseous atoms/ions

• Why?

• What about 1st, 2nd, 3rd,

ionization energies?

VII. Successive IE’s

VII. Electron Affinity

• electron affinity:

energy change in kJ

when e- added to a

gaseous atom/ion

(generally negative)

• Why?

VII. Trend in Metallic Character