Periodic Table

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The Periodic Table Warmup:

Overview of the Periodic Table

Metals Metalloids Nonmetals Noble Gases

Periodic Table

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History of the Periodic Table I. Early attempts

Made the task a little easier: Jöns Jakob Berzelius 1828 Swedish - developed a table of atomic weights - introduced letters to symbolize elements

a) Johann Döbereiner 1829 German - described triads of elements

(e.g. Cl, Br, I; Ca, Ba, Sr; S, Se, Te) – first indication that elements were related to one another – atomic mass is related to chemical properties

Karlsruhe Congress (big Chemistry Conference) 1860 Germany

b) John Newlands 1865 English

- arranged elements in order of relative atomic masses;

- described the Rule of Octaves – every 8th element has similar properties

c) Julius Lothar Meyer 1870 German graph of atomic volume (atomic weight/density) against

atomic weight periodic trends in elements’

properties; established concept of valency

II. Dmitri Mendeleev 1869 Russian a) How:

While writing a book on inorganic chemistry

to get organized, wrote elements on notecards with

some properties and atomic weight/mass: ULTIMATE SOLITAIRE

arranged elements in order of atomic masses

noticed a repetition of properties every 8 or 18

elements

elements with similar properties in horizontal rows

b) The amazing part: he predicted 3 elements not yet

discovered (eka-aluminum, eka-boron, eka-silicon)

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c) Problems : Ar/K, Te/I, Co/Ni

1st element in each pair has greater atomic mass

places reactive K in unreactive noble gases

d) Importance –

1) realized elements yet to be discovered;

2) characteristics of element could be predicted from its atomic weight (and

position on the tables) Properties of Some Elements Predicted by Mendeleev

Predicted Elements

Element and year discovered

Properties

Predicted Properties

Observed Properties

Eka-aluminum Gallium, 1875 Density of metal 6.0 g/mL 5.96 g/mL Melting point

Low 30oC

Oxide formula

Ea2O3 Ga2O3

Eka-boron Scandium, 1877 Density of metal 3.5 g/mL 3.86 g/mL Oxide formula

Eb2O3 Sc2O3

Solubility of oxide Dissolves in acid Dissolves in acid Eka-silicon Germanium, 1886 Melting point High 900oC Density of metal 5.5 g/mL 5.47 g/mL Color of metal Dark gray Grayish white Oxide formula

EsO2 GeO2

Density of oxide 4.7 g/mL 4.70 g/mL Chloride formula EsCl4 GeCl4

Discovery of the Noble Gases 1890s

• Lord Rayleigh (physicist) and Sir William Ramsay (chemist)

• 1894 - Argon “the lazy one”, discovered when Ramsay was trying to isolate nitrogen

• 1895 - Helium – found on earth in uranium minerals (found in the sun in 1868)

• 1898 - Neon “the new one”, Krypton “the hidden one”, Xenon “the alien one”

• 1910 – Radon

Properties: Largely unreactive, 8 electrons in valence shell, low boiling and melting points

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Nucleus discovered – 1910 - Rutherford predicted that the charge of an atom is proportional to its mass III. Henry Moseley 1913 English (worked with Rutherford)

a) λ of emitted X-rays corresponded to # protons

atomic number

“Do other properties match atomic numbers?” Yes!

arranged the periodic table by atomic #’s, not mass

b) Law of Atomic Numbers (Law of Chemical Periodicity)

- the properties of elements are periodic functions of their atomic numbers

- corrected incorrect placement of cobalt and nickel, and iodine and tellurium

IV. Glenn Seaborg 1940s American 1912-1999

a) “transuranium” elements – formation of elements beyond uranium (93-103)

reorganization of periodic table to include both series of radioactive elements

(lanthanides and actinides)

b) note the names of elements 95-103, reflect Seaborg’s academic life – scientists

and institutions (UC-Berkeley)

Trends of the Periodic Table “periodic” = repeating pattern

Overall theme = electrons’ positions relative to each other and the nucleus determine the

following properties:

1. Electron configuration ( reactivity and bonding)

2. Atomic radius

3. Ionic radius

4. Ionization energy

5. Electronegativity

Periodic Table

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1. Electron configuration

1A 1

2A 2

B 3-12

3A 13

4A 14

5A 15

6A 16

7A 17

8A 18

H 1s1

Li 2s1

Na 3s1

K 4s1

Rb 5s1

Cs 6s1

Fr 7s1

1+

He 1s2

Be 2s2

Mg 3s2

Ca 4s2

Sr 5s2

Ba 6s2

Ra 7s2

2+

Sc-Zn 3d

Y-Cd 4d

La-Hg 5d

Ac-Une 6d

(followed

by s)

1+ or 2+

(generally)

B 2s22p1

3s23p1

4s24p1

5s25p1

6s26p1

3+

C 2s22p2

3s23p2

4s24p2

5s25p2

6s26p2

2+

4+

4-

N 2s22p3

3s23p3

4s24p3

5s25p3

6s26p3

3-

O 2s22p4

3s23p4

4s24p4

5s25p4

6s26p4

2-

F 2s22p5

3s23p5

4s24p5

5s25p5

6s26p5

1-

Ne 2s22p6

3s23p6

4s24p6

5s25p6

6s26p6

0

B Lanthanide/

Actinide series

Begin in

period 6

Ce-Lu 4f

Th-Lr 5f

(followed by s)

1+ or 2+

(generally)

The position of a valence electron and the ability to remove it from an atom are related to

• the number of protons in the nucleus

• the extent to which the valence electron is shielded from the positively-charged

nucleus by the negatively-charged core electrons

Noble gas configuration = [core] electrons

“outer” electrons = “valence” electrons Elements of groups 1A-8A have valence

electrons in s and p orbitals.

(Note: USA uses A and B.

Rest of world uses 1-18)

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Isoelectronic Series

= same number of electrons

(iso = equal)

1. Draw the complete electron configuration of each of the following elements.

2. What ions will they form?

3. When ions, how many electrons does each have? How many protons?

4. Predict the relative diameters of the members of this isoelectronic series.

Element Electron config Ion Ion # e- Ion # p+

O

F

Ne

Na

Mg

Prediction (smallest to largest):

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2. Atomic Radius

½ distance between nuclei

a) Trend down a GROUP: ↑

i. larger atoms – valence e-’s are farther away from nucleus

ii. shielding effect – the number of e-’s between the nucleus and valence e-

’s helps keep the valence e-’s farther away from the nucleus, thus ↓ the

pull of the nucleus on the valence e-’s.

b) Trend across a PERIOD: ↓ (same principal energy level)

i. for every added e-, one more p+

↑ pull on outer e-’s by nucleus

ii. not as noticeable in periods with heavier elements

(inner e-‘s shield the valence e-’s greater distance

between nucleus and valence e-’s)

iii. shielding effect is constant across a period, as e-’s are added only to the

valence, or outermost energy level

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Atomic Radii

1. Which groups and periods of elements are shown in the table of atomic radii?

______________________________________________________________________

2. In what unit is atomic radius measured? __________Express this unit in m __________

3. What are the values of the smallest and largest atomic radii shown? What elements have

these atomic radii? _______________________________________________________

4. What happens to atomic radii within a period as the atomic number increases?

______________________________________________________________________

5. What accounts for the trend in atomic radii within a period?

______________________________________________________________________

______________________________________________________________________

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6. What happens to atomic radii within a group? ____________________________

7. What accounts for the trend in atomic radii within a group?

_______________________________________________________________________

_______________________________________________________________________

8. a) Divide the atomic radius of Cs by the atomic radius of Li and round to 2 significant

figures. Cs:Li _______________

b) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant

figures. Cs:Rn ______________

c) Summarize your findings about a) and b) here: _____________________________

___________________________________________________________________

3. Ionic Radius

a) Cations (+) are always SMALLER than the original atom

(↓ # e-’s ↓ repulsion and greater pull on the valence e-’s by the nucleus

remaining e-’s shrink to nucleus)

b) Anions (-) are always LARGER than the original atom

(↑ # e-’s ↑ repulsion shell swells)

Periodic Table

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Ionic Radius

1. In this table of ionic radii, how is the charge of the ions of elements in groups 1A-4A

related to the group number? _____________________________________________

2. a) Divide the radius of Cs with the radius of its ion: ____________________

b) Divide the radius of Li with the radius of its ion: ____________________

c) Divide the radius of Be with the radius of its ion: ____________________

d) Divide the radius of B with the radius of its ion: _____________________

e) Summarize your findings about a)-d) here: _________________________

___________________________________________________________________

3. a) Divide the radius of the F ion with the radius of the neutral F atom:

b) Divide the radius of the O ion with the radius of the neutral O atom:

c) Divide the radius of the N ion with the radius of the neutral N atom:

d) Summarize your findings about a)-c) here: ___________________________________

_____________________________________________________________________

e) Compare and contrast 2 e) and 3 d) ________________________________________

____________________________________________________________________

Periodic Table

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4. Ionization Energy Definition: the energy required to remove an electron from an atom in the gas phase

(in J or kJ)

a) Successive ionization energies for each atom (since > 1 electron can be removed) Removing each subsequent electron requires more energy

Diagram - removing successive electrons from Be:

Ionization Energies of Na, Mg, and Al (in kJ/mol)

Successive ionization energies (kJ/mol)

Element First Second Third Fourth

Na

496 4,562 6,912 9,543

Mg

738 1,451 7,733 10,540

Al

578 1,817 2,745 11,577

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1. What happens to the values of the successive ionization energies of an element?

___________________________________________________________________

2. How is a jump in ionization energy related to the valence electrons of the element?

___________________________________________________________________

___________________________________________________________________

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1. What is meant by first ionization energy? _______________________________________

_______________________________________________________________________

2. Which element has the smallest first ionization energy? The largest? What are their

values? ________________________________________________________________

3. What generally happens to the first ionization energy of the elements within a period as

the atomic number of the elements increases? ________________________________

4. What accounts for the general trend in the first ionization energy of the elements within a

period? ________________________________________________________________

______________________________________________________________________

5. Based on the graph, rank the group 2A elements in periods 1-5 in decreasing order of first

ionization energy. _________________________________________________________

8. What generally happens to the first ionization energy of the elements within a group as the

atomic number of the elements increases? _____________________________________

9. What accounts for the general trend in the first ionization energy of the elements within a

group? _________________________________________________________________

_______________________________________________________________________

b) Summary of trends in first ionization energies: trend down a GROUP: ↓ trend across a PERIOD: ↑

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5. Electronegativity = how much one atom pulls on another atom’s electrons in a bond

∴ only refers to atoms in a bond (molecule or compound)