Transcript Chapter 7.1: Periodic Properties of the Elements

The Periodic Table
Metals, Nonmetals, &
Metalloids, Families, Periodic
Trends
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.
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)have a very high melting point
• Many
(1900°C for chromium)
• Good conductors of heat and electricity
• Form cations
• Metal + nonmetal = ionic cmpd
Characteristics of Nonmetals
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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
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.
Parts of Periodic Table
• Representative elements: Group A elements
• Transition elements: Group B elements (dblock)
• 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
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)
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.
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 emoves farther from the nucleus.
Periodic Trend 1:
Atomic Radius
Atomic radius decreases
up a family (bottom to top
of periodic table).
Atomic radius decreases
across period from
left to right.
WHY?
Lower-numbered energy
levels are closer to the
nucleus.
WHY?
Shielding stays constant
across a period so Zeff
increases. Electrons move
closer to the nucleus
Periodic Trend 2:
Ionic Radius
CATIONS
Radius of a cation is smaller
than its parent atom.
ANIONS
Radius of an anion is larger
than its parent atom.
WHY?
a. Outermost electrons leave
to form a cation.
b. New outermost electrons
are in a lower energy level,
closer to the nucleus.
WHY?
a. Electrons are added to
outer shell to form an anion.
b. More electron-electron
repulsion.
Ionic Radius Trend Up a Family
Cation radii decrease
up a family (bottom to top).
Anion 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!
WHY?
Atomic radii decrease up a
family, and anions are larger
than their parent atoms.
Follows same pattern as
atomic radius!
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.
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.
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.
Periodic Trend 3:
Ionization Energy
Ionization energy increases
as we go up a family
(bottom to top).
WHY?
As atom gets smaller,
it is harder to
remove electrons
because they are
more attracted to
the nucleus.
Ionization energy increases
as we go from left to right
across a period.
WHY?
As Zeff increases and
the atom gets
smaller, it is harder
to remove electrons.
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.