1. Metals, Nonmetals, & Metalloids, Families, Periodic Trends
The Periodic Table
Metals, Nonmetals, & Metalloids, Families, Periodic Trends

2. Periodic Properties of the Elements
1869: Dmitri Mendeleev and Lothar Meyer
publish identical tables. Mendeleev gets the
credit and becomes “Father of Periodic Table”.
1913: Henry Moseley develops
concept of atomic numbers.
Modern periodic table arranged in order of increasing atomic number.

4. Characteristics of Metals
Most elements are metals
Have a shiny luster; most are silvery
All solids at room temperature except Hg
Are malleable (thin sheets) and ductile (wires)
Many have a very high melting point
(1900°C for chromium)
Good conductors of heat and electricity
Form cations
Metal + nonmetal = ionic cmpd

5. Characteristics of Nonmetals
Do not have luster; various colors
Poor conductors of heat and electricity
Melting points are generally lower than those of metals
May be solids, liquids, or gases at room temp
Have high electron affinity; form anions
Compounds of nonmetals with metals are ionic
Compounds with only nonmetals are molecular

6. Characteristics of Metalloids
Include the elements B, Si, Ge, As, Sb, Te, At
Form the division line between metals and nonmetals on the periodic table
Have intermediate properties.
Si and Ge are used in integrated circuits and computer chips.

7. Parts of Periodic Table
Representative elements: Group A elements
Transition elements: Group B elements (d-block)
Inner Transition (Rare Earth) elements (f-block)
Group 1A: Alkali metals; Group 2A: Alkaline Earth metals; Group 6A: Chalcogens; Group 7A: Halogens; Group 8A: Noble Gases

8. Hydrogen Does not belong to any group Diatomic gas at RT
Metallic at extreme pressures
Usually forms covalent bonds
Hydrogen ion (1+ charge) or hydride ion (1- charge)

9. What is Electron Shielding?
Electrons in any orbital will partially shield an e- in any other orbital.
Core e-’s shield or screen outer e-’s from the full charge of the nucleus.
Electrons in the same shell repulse each other.

10. What Determines Effective Nuclear Charge (Zeff)?
The Zeff felt by outer electrons is determined by: Zeff = proton # - core electron #
Effective nuclear charge increases as nuclear charge increases.
Effective nuclear charge decreases as the e- moves farther from the nucleus.

11. Periodic Trend 1: Atomic Radius
Atomic radius decreases
up a family (bottom to top
of periodic table).
WHY?
Lower-numbered energy levels are closer to the nucleus.
Atomic radius decreases
across period from
left to right.
WHY?
Shielding stays constant across a period so Zeff increases. Electrons move closer to the nucleus

13. Periodic Trend 2: Ionic Radius
CATIONS
Radius of a cation is smaller than its parent atom.
WHY?
Outermost electrons leave to form a cation.
b. New outermost electrons are in a lower energy level, closer to the nucleus.
ANIONS
Radius of an anion is larger than its parent atom.
WHY?
a. Electrons are added to outer shell to form an anion.
b. More electron-electron repulsion.

15. Ionic Radius Trend Up a Family
Cation radii decrease
up a family (bottom to top).
WHY?
Atomic radii decrease up a family, and cations are smaller than their parent atoms.
Follows same pattern as atomic radius!
Anion radii decrease
up a family (bottom to top).
WHY?
Atomic radii decrease up a family, and anions are larger than their parent atoms.
Follows same pattern as atomic radius!

17. Ionic Radius Trend Across a Period
The ionic radius trend changes across a period as we move from metals to nonmetals.
The radii of metallic ions decrease while the radii of nonmetallic ions increase across a period.
This causes a wave-like pattern on the table.

18. Ionization Energy *** The greater the ionization energy,
Ionization Energy is the minimum amount of energy needed to remove an electron from the ground state of an isolated gaseous atom or ion. Formation of cations!
First Ionization Energy (I1) removes the first electron from a neutral atom:
Na(g)  Na+(g) + 1e-
Second Ionization Energy (I2) removes the second electron from an ion:
Na+(g)  Na2+(g) + 1e-
*** The greater the ionization energy,
the more difficult it is to remove an electron.

19. Variations in Successive Ionization Energies
Ranking of ionization energy:
I1 < I2 < I3
Every element shows a large
increase in ionization energy
when electrons are removed
from its noble gas core.

21. Periodic Trend 3: Ionization Energy
Ionization energy increases as we go up a family (bottom to top).
Ionization energy increases as we go from left to right across a period.
WHY?
As atom gets smaller, it is harder to remove electrons because they are more attracted to the nucleus.
WHY?
As Zeff increases and the atom gets smaller, it is harder to remove electrons.

24. Periodic Trend 4: Electronegativity
Electronegativity: the ability of atoms in a molecule to attract electrons to itself.
Greater electronegativity = greater ability to attract electrons.
Values from 0.7 for Cs to 4.0 for F.
Metals are least electronegative, halogens most electronegative.