The History of the Modern

Periodic TableSee separate slide show for Periodic Table History

Periodic Law• When elements are arranged in order of

increasing atomic #, elements with similar properties appear at regular intervals.

0

50

100

150

200

250

0 5 10 15 20

Ato

mic

Ra

diu

s (

pm

)

Atomic Number

Chemical ReactivityFamilies Similar valence e- within a group

result in similar chemical properties

•Alkali Metals•Alkaline Earth Metals•Transition Metals•Halogens•Noble Gases

Periodic Table Reveals Periodic Trends

• Effective Nuclear charge

• atomic size or radius

• ionization energy

• electron affinity

• electronegativity

• metallic character

• Reactivity

• bonding characteristics

• crystal configurations

• acidic properties

• densities

• Melting/Boiling points

Electron screening or shielding

• Electrons are attracted to the nucleus• Electrons are repulsed by other electrons• Electrons would be bound more tightly if

other electrons weren’t present.• The net nuclear charge felt by an electron is

called the effective nuclear charge ( Zeff ).

Quantum Mechanical Model

Zeff is lower than actual nuclear charge.

Zeff increases toward nucleus ns > np > nd > nf

This explains certain periodic changes observed.

Effective Nuclear Charge ( Zeff)

• The effective nuclear charge acting on an electron equals the number of protons in the nucleus, Z, minus the average number of electrons, S that are between the nucleus and the electron in question.

Zeff = # protons # shielding electrons

Zeff = attractive forces repulsive forces Zeff = Z S

For Example, Lithium vs. Carbon

Li Zeff = 3 2 = 1

C Zeff = 6 2 = 4

So, carbon has a much smaller atomic radius compared to lithium: Rcarbon =77

pm Rlithium = 152 pm

When moving across a row:The greater the Zeff value, the smaller the atom’s radius.

Trend #1 Atomic Radii

1

2

3

4 5

6

7

Increases to Left and Down

•Why larger going down?

•Why smaller to the right?

•Higher energy levels have larger orbitals

•Shielding - core e- block the attraction between the nucleus and the valence e-

• Increased nuclear charge without additional shielding pulls e- in tighter

Practice…

• Referring to a periodic table, arrange the following atoms in order of increasing size:– Phosphorus– Sulfur– Arsenic– Selenium

• S < P < Se < As

Atomic radii

The Periodic Table & Radii

Periodic Trend is Due to Effective Nuclear Charge

Atomic Radii vs. Zeff:

Trends in Ionic Radii

• Using your knowledge of Zeff, how would the size of a cation compare to neutral atom? Anion?

Trends in Ionic Radii

• The cation of an atom decreases in size.

• The more positive an ion is, the smaller it is because Zeff increases

• The anion of an atom increases in size.

• The more negative an ion, the larger it is because Zeff decreases.

Cations lose electrons, become smaller

Anions gain electrons, become bigger

Ion Radii

1

2

3

4 5

6

7

+3 +4 -3 -2 -1

Increases downIncreases moving across, but depends if cation OR anion

Ions and Ionic Radii

Practice…• Arrange the following atoms and ions in order

of decreasing size: – Mg2+

– Ca2+

– Ca• Which of the following ions is the largest:

– S2-

– S– O2-

Practice…• Arrange the following ions in order of decreasing

size:– S2-

– Cl-

– K+

– Ca2+

• Which of the following ions is the largest?– Rb+

– Sr2+

– Y3+

Trend in Ionization Energy

• Ionization NRG is the NRG required to remove an electron from an atom

Successive Ionization NRG

• Ionization energy increases for successive electrons from the same atom.

*Notice the large jump in ionization energy when a core e is removed.

Why do you think there is such a big jump for Mg3+?

• The smaller the atom, the higher the ionization energy due to Zeff

• Bigger atoms have lower ionization NRG due to the fact that the electrons are further away from the nucleus and therefore easier to remove.

Increases

Dec

reas

es

Practice…• Which of the following elements would

have the highest second ionization energy? Justify your answer.–Sodium, Sulfur, or Calcium

• Which will have the greater third ionization energy, Ca or S? Justify your answer.

Practice…• Referring to a periodic table, arrange the

following atoms in order of increasing first ionization energy (Ne, Na, P, Ar, K) Justify your answer.

• Based on the trends discussed in this section, predict which of the following atoms (B, Al, C or Si) has the lowest first ionization energy and which has the highest first ionization energy.

Electron Affinity

• The energy change associated with the addition of an electron

• Tends to increase across a period• Tends to decrease as you go down a group• Abbreviation is Eea, it has units of kJ/mol. Values are

generally negative because energy is released.• Value of Eea results from interplay of nucleus

electron attraction, and electron–electron repulsion.

Ionization NRG vs. Electron Affinity• Ionization energy measures the ease with

which an atom loses an electron • Electron affinity measures the ease with

which an atom gains an electron

Electron Affinity

Trends in Electronegativity

• tendency for an atom to attract electrons when it is chemically combined with another atom.

• decreases as you move down a group• increases as you go across a period from

left to right.

Trend #5 Metallic Character• The metallic character of atoms can be related

to the desire to lose electrons.

• The lower an atom’s ionizatoin energy, the

greater its metallic character will be.

• On the periodic table, the metallic character of

the atoms increase down a family and

decreases from left to right across a period.

Metals Nonmetals

• Shiny Luster• Various colors (most

silvery)• Solids are malleable and

ductile• Good conductors of heat

and electricity• Most metal oxides are

ionic solids that are basic• Tend to form cations in

aqueous solution

• No luster• Various colors• Brittle solids• Poor conductors of heat

and electricity• Most nonmetal oxides

are molecular substances that form acidic solutions

• Tend to form anions or oxyanions in aqueous solution

Metallic Character

1

2

3

4

5

6

7

Increases moving down and across to the left

Fr

Cs Ba

Ra

Lower left corner -- elements mostlikely to lose their valence electrons

Rb

Metals and Nonmetals

• Low ionization energies of metals means they tend to form cations (positive ions) relatively easily

• Due to their electron affinities, nonmetals tend to gain electrons when they react with metals.

# 6 Melting/Boiling Points

• Highest in the middle of a period (generally).

1

2

3

4 5

6

7

Some Important Properties of Alkali Metals

• Soft metallic solids• Easily lose valence electrons (Reducing

Agents)– React with halogens to form salts– React violently with water

• Large Hydration NRG– Positive ionic charge makes ions attractive to

polar water molecules

Alkaline Earth Metals…• Harder and more dense than Alkali Metals• Less reactive than alkali metals (lower first

ionization energies)• Reactivity increases as you move down the

periodic table.

The Halogens…

• “Salt Formers”• Melting and Boiling Points increase with

atomic number.• Highly negative electron affinities• Tendency to gain electrons and form halide

ions

Noble Gases …

• Monoatomic ions• Gases at room temperature• Large 1st ionization energies• “Exceptionally” unreactive

Practice…

• Look at Sample Integrative Exercise 7 on page 264