chemical periodicity

Chemical Periodicity

Principles of General Chemistry, 2nd ed. By M. Silberberg

Chemistry, 8th ed. by W. Whitten, R. Davis, R., M. L. Peck, and G. Stanley.

2

Lecture Goals

1. More About the Periodic Table

Periodic Properties of the Elements

2. Atomic Radii

3. Ionic Radii

4. Ionization Energy

5. Electron Affinity

6. Electronegativity

3

Periodic Table

• Arranged by Dmitri Mendeliv, a Russian Chemist

• Establish a classification scheme of the elements based on their electron configurations.

Column: FAMILY or GROUP

Row: PERIOD

• Certain characteristics are shared by each families & periods.

4

More About the Periodic Table

Representative Elements

– are the elements in A groups

on periodic chart.

• will have their “last”

electron in an outer s or p

orbital

• have fairly regular variations

in their properties

5


d-Transition Elements

– on periodic chart in B groups

• each metal has d electrons.

– ns (n-1)d configurations

• make the transition from

metals to nonmetals

• exhibit smaller variations

from row-to-row than the

representative elements

6


f - transition metals

– sometimes called inner transition metals

• have electrons in f orbitals & are being added two shells below the valence shell

• very slight variations of properties from one element to another.

The Periodic Table:

Metals, Nonmetals, and Metalloids

• 1869 - Mendeleev & Meyer

– The properties of the elements are periodic

functions of their atomic numbers.

The Periodic Table:


• Groups or families

– Vertical group of elements on

periodic table

– Similar chemical and physical

properties

The Periodic Table:

Metals, Nonmetals,

and Metalloids • Period

– Horizontal group of elements on periodic table

– Transition from metals to nonmetals

The Periodic Table:


Some chemical properties

1. Outer shells contain few electrons

2. Form cations by losing electrons

3. Form ionic compounds with nonmetals

4. Solid state characterized by metallic bonding

Metals • Group IA metals: Li, Na, K, Rb, Cs, Fr

• example of a trend - their reactions with water

• are all shiny, soft, silvery, highly reactive metals

at standard T and P, and readily lose their outermost

electron to form cations (+1)

The Alkali Metals

• Group IIA metals

- Be, Mg, Ca, Sr, Ba, Ra

• they are all shiny, silvery-

white, somewhat reactive

metals at standard

temperature and

pressure and readily lose

their two outermost

electrons to form cations

with charge (+2)

The Alkaline Earth Metals Non Metals

Some chemical properties

1. Outer shells contain four or more electrons

2. Form anions by gaining electrons

3. Form ionic compounds with metals and covalent

compounds with other nonmetals

4. Form covalently bonded molecules; noble gases are

monatomic

Chalcogens

• Group VIA nonmetals - O, S, Se, Te

• electronegative elements are strongly associated with

metal-bearing minerals, where they have formed

water-insoluble compounds with the metals in the ores

Halogens

• Group VIIA nonmetals - F, Cl, Br, I, At

• the only periodic table group which contains

elements in all three familiar states of matter

at standard temperature and pressure

Noble, Inert or Rare Gases

• Group VIIIA nonmetals - He, Ne, Ar, Kr, Xe, Rn

• have completely filled electron shells

The Periodic Table:


• Metals are to the left

of stair step.

~ 80% of the elements

• Best metals are on the

far left of the table.

The Periodic Table:


• Non metals are to the

right of stair step.

~ 20% of the elements

• Best metals are on the

far right of the table.

The Periodic Table:


• Metalloids have one side of the box on the

stair step.

The Periodic Table:


22

Periodic Properties of the Elements

By definition, periodic is appearing or recurring at regular intervals .

to be discussed:

• Metallic Character

• Atomic Radii

• Ionic Radii

• Ionization Energy

• Electron Affinity

• Electronegativity

Metallic Character

decreases

incre

ase

s

24

Atomic Radii • describes the relative sizes of atoms.

decreases

incre

ase

s

25

Atomic Radii

26

Atomic Radii

• increases (down a group) – noted by the increase the number of ‘shells’ occupied by e-s

• decrease (across a period) – due to nuclear charge effect (moving across a period, each element has an increased nuclear charge and the electrons are going into the same shell (2s and 2p or 3s and 3p…)

• shielding or screening effect

- effective nuclear charge, Zeff, experienced by an electron is less than the actual nuclear charge, Z.

- the inner electrons block the nuclear charge’s effect on the outer electrons.

27

Ionic Radii decreases

incre

ase

s

28

Ionic Radii

• Cations (positive ions) are always smaller than their respective neutral atoms.

Element Na Mg Al

Atomic

Radius (Å) 1.86 1.60 1.43

Ion Na+ Mg2+ Al3+

Ionic

Radius (Å) 1.16 0.85 0.68

29

Ionic Radii

• Anions (negative ions) are always larger than their neutral atoms.

Element N O F

Atomic

Radius(Å) 0.75 0.73 0.72

Ion N3- O2- F1-

Ionic

Radius(Å) 1.71 1.26 1.19

30

Ionic Radii

• cation (positive ions) radii decrease from left to right across a period.

– increasing nuclear charge attracts the electrons and decreases the radius.

Ion Rb+ Sr2+ In3+

Ionic

Radii(Å) 1.66 1.32 0.94

31

Ionic Radii

• anion (negative ions) radii decrease from left to right across a period.

– increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius.

Ion N3- O2- F-

Ionic

Radii(Å) 1.71 1.26 1.19

32

Atomic Radii and Ionic Radii

• Arrange these

elements based on

their atomic radii.

–Se, S, O, Te

–P, Cl, S, Si

• Arrange these

elements based on

their ionic radii.

–Ga, K, Ca

–Cl, Se, Br, S

33

Atomic Radii and Ionic Radii

• atomic radii

– Se, S, O, Te

O < S < Se < Te

– P, Cl, S, Si

Cl < S < P < Si

• ionic radii

– Ga, K, Ca

K+ > Ca2+ > Ga3+

– Cl, Se, Br, S

Cl- < S2- < Br1- < Se2-

34

Ionization Energy

1st ionization energy (IE1)

– the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a 1+ ion.

Mg(g) + 738kJ/mol Mg+ + e-

Symbolically: Atom(g) + energy ion+

(g) + e-

35

Ionization Energy

2nd ionization energy (IE2)

– the amount of energy required to remove the second electron

from a gaseous 1+ ion

• Symbolically:

ion+ + energy ion2+ + e-

Mg+ + 1451 kJ/mol Mg2+ + e-

Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies.

36

Ionization Energy

1. IE2 > IE1

It always takes more energy to remove a second electron from an ion than from a neutral atom.

2. IE1 generally increases moving across a period.

Important exceptions at Be & Mg, N & P, etc. due to filled and half-filled subshells.

3. IE1 generally decreases moving down a family.

IE1 for Li > IE1 for Na, etc.

increases

decre

ase

s Ionization Energy

38

Ionization Energy

39

Ionization Energy

Group and

element

(IA)

Na

(IIA)

Mg

(IIIA)

Al

(IVA)

Si

IE1

(kJ/mol)

496 738 578 786

IE2

(kJ/mol) 4562 1451 1817 1577

IE3

(kJ/mol) 6912 7733 2745 3232

IE4

(kJ/mol) 9540 10,550 11,580 4356

40

Ionization Energy

• The reason Na forms Na+ and not Na2+ is that

the energy difference between IE1 and IE2 is so

large.

– Requires more than 9 times more energy to remove

the second electron than the first one.

• The same trend is persistent throughout the

series.

– Thus Mg forms Mg2+ and not Mg3+.

41

Ionization Energy • What charge ion would be expected for an element

that has these ionization energies?

IE1 (kJ/mol) 1680

IE2 (kJ/mol) 3370

IE3 (kJ/mol) 6050

IE4 (kJ/mol) 8410

IE5 (kJ/mol) 11020

IE6 (kJ/mol) 15160

IE7 (kJ/mol) 17870

IE8 (kJ/mol) 92040

42

Ionization Energy

• Arrange these elements based on

their first ionization energies.

– Sr, Be, Ca, Mg

– Al, Cl, Na, P

– B, O, Be, N

43

Ionization Energy

• Arrange these elements based on

their first ionization energies.

– Sr, Be, Ca, Mg

Sr < Ca < Mg < Be

– Al, Cl, Na, P

Na < Al < P < Cl

– B, O, Be, N

B < Be < O < N

44

Electron Affinity

• is the amount of energy absorbed when an

electron is added to an isolated gaseous atom

to form an ion with a 1- charge.

• is a measure of an atom’s ability to form

negative ions.

• Symbolically: atom(g) + e- + EA ion-(g)

increases

decre

ase

s

Electron Affinity

46

Electron Affinity

Mg(g) + e- + EA Mg-(g)

EA = +231 kJ/mol

Br(g) + e- Br-(g) + EA

EA = -323 kJ/mol

Two examples of electron affinity values:

If electron affinity > 0, energy is absorbed; if < 0, energy is released.

47

Electron Affinity

48

Electronegativity

• is a measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another element

• is measured on the Pauling scale

– Fluorine is the most electronegative element.

– Cesium and francium are the least electronegative elements.

increases

decre

ase

s

Electronegativity

* for representative elements

50

Electronegativity

51

Electronegativity

• Arrange these elements based

on their electronegativity.

–Se, Ge, Br, As

–Be, Mg, Ca, Ba

52

Electronegativity

• Arrange these elements based

on their electronegativity.

–Se, Ge, Br, As

Ge < As < Se < Br

–Be, Mg, Ca, Ba

Ba < Ca < Mg < Be

53

In summary…

6 Chemical

Periodicity

“Value has a value only if

its value is valued.”

–Brian G. Dyson

President and CEO, Coca-Cola

Enterprises

questions?

chemical periodicity

Documents