Metals and Non-Metals
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All elements can be classified as either metals, non-metals, or semimetals
(metalloids)
Metals: elements that have small numbers of electrons in the electronic
shell of highest principal quantum number. Removal of an electron(s) from a
metal atom occurs without great difficulty, producing a positive ion (cation).
Metals generally are malleable, and ductile solids with a lustrous appearance
and an ability to conduct heat and electricity. Found to the left of the
“diagonal line” on the Periodic Table
Nonmetals: elements whose atoms tend to gain small numbers of electrons
to form negative ions (anions) with the electron configuration of a noble gas.
Nonmetal atoms may also alter their electronic configuration by sharing
electrons. Nonmetals are mostly gases, liquid (bromine), or low melting solids
and are very poor conductors of heat and electricity. Found to the right of
the “diagonal line” on the Periodic Table
Semimetals (metalloids): elements that display both metallic and nonmetallic
properties under appropriate conditions (e.g., Si, Ge, As). Elements that
touch the “diagonal line” on the Periodic Table.
Semiconductor: Any of various solid crystalline substances having electrical
conductivity between that of a conductor and that of an insulator in being
nearly as great as that of a metal at high temperatures and nearly absent at
low temperatures (e.g. silicon, germanium).
Families of Elements
Families of elements display similar chemical and physical properties
Group IA- Alkali Metals
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Low densities; large atomic sizes
Low ionization energies
Readily lose a single valence electron; good reducing agents
Most reactive known metals; larger atoms most reactive
Conduct heat and electricity
Some Physical Properties of the Alkali Metals
Element
Melting Point
(oC)
Boiling Point
(oC)
Density
(g/cm3)
Atomic Radius
(nm)
Lithium
179
1336
0.53
0.123
Sodium
98
883
0.97
0.157
Potassium
64
758
0.86
0.203
Rubidium
39
700
1.53
0.216
Cesium
28
670
1.90
0.235
Group IIA - Alkaline Earth Metals
 No clear trends in melting points, boiling points or density
 Higher melting points, and are harder, than alkali metals
 When oxidized, lose two valence electrons
Some Physical Properties of the Alkaline Earth Metals
Element
Melting Point
(oC)
Boiling Point
(oC)
Density
(g/cm3)
Atomic Radius
(nm)
Beryllium
1280
1500
1.86
0.089
Magnesium
651
1107
1.75
0.136
Calcium
851
1487
1.55
0.174
Strontium
800
1366
2.60
0.191
Barium
850
1537
3.59
0.198
Group VIIA - Halogens
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Fluorine and chlorine are gases at room temperature; bromine is a liquid;
iodine is a solid
Bromine vaporizes easily
Iodine sublimes easily upon heating
Very reactive nonmetals
Fluorine is the most electronegative and the most reactive element known
Strongest elemental oxidizing agents; gain one electron to become anions
Size, mass and reactivity decrease as atomic number increases within
this family
Some Physical Properties of the Halogens
Element
Melting Point
(oC)
Boiling Point
(oC)
Atomic Radius
(nm)
Fluorine
-220
-188
0.071
Chlorine
-101
-35
0.099
Bromine
-7.2
58.8
0.114
Iodine
114
184
0.133
Noble Gases
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Initially thought to be inert; end in ns2p6 electronic configuration
Xenon compounds have now been synthesized; examples: XeF2, XeOF2,
XeO3, H4XeO6
Ions
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Last day we saw that the valence electrons of elements in the same
family have the same electronic configuration (with different principal
quantum numbers)
Example: All Alkaline Earth Metals have an electronic configuration that
ends in s2; all halogens end in s2p5; all Noble Gases end in s2p6
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Group IA and IIA metals lose electrons to form cations. They lose
electrons to form ions with the same electronic configuration as the
nearest Noble Gas.
Metals lose their valence electrons to form ions.
Example: Na  1s22s22p63s1 or [Ne] 3s1 ; Na+  1s22s22p6 = [Ne]
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Non-metals gain electrons to form anions. They gain electrons to form
ions with the same electronic configuration as the nearest Noble Gas.
Non-metals gain valence electrons to form ions.
Example: O  1s22s22p4 or [He] 2s22p4 ; O2-  1s22s22p6 = [Ne]
Periodic Variation
The Size of Atoms
Atomic Radius: difficult to define since the probability of finding an
electron decreases with increasing distance from the nucleus, yet never
drops to zero! The distance between two nuclei of bonded atoms can be
determined. Atomic radius is normally based on this measurement.
Variation of Atomic Radii Within a Group of the Periodic Table
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The higher the principal quantum number of an electronic shell, the
further from the nucleus will significant electron density still exist.
This leads to the general trend: The more electronic shells in an atom, the
larger an atom. Atomic radius increases from top to bottom through a
group of elements.
Variation of Atomic Radii Within a Period of the Periodic Table
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Across a period, electrons are added to the valence shell (this excludes
the transition metals). Protons are also added to the nucleus. Electrons
in subshells partially shield the valence electrons from the attractive
force of protons in the nucleus. Across a period, however, this shielding
does not increase, but the number of protons in the nucleus does!
This leads to the general trend: Atomic radii decrease from left to right
through a period of elements.
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The transition metals also follow this general trend, however, across a
period, electrons are added to an inner shell, d-orbital. This means that
the atomic radii do not decrease as significantly from left to right
through a period for the transition metals as they do for the main group
elements.
The Size of Ions
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Cations are smaller than the atoms from which they are formed. The
nucleus is able to draw the remaining electrons closer.
For cations with the same number of electrons (isoelectronic), the more
positive the charge, the smaller the ionic radius.
Example: Na+ is smaller than Mg2+. Both have 10 electrons.
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Anions are larger than the atoms from which they are formed. Electrons
are not held as tightly by the nucleus and repulsions among electrons
increase.
For isoelectronic anions, the more negative the charge, the larger the
ionic radius.
Example: S2- is larger than Cl-. Both have 18 electrons.
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Ionization Energy
Ionization Energy (I): the quantity of energy that a gaseous atom must
absorb so that an electron is stripped from the atom. The electron that is
lost is the one that was the most loosely held.
Example:
Mg (g)  Mg+ + eMg+ (g)  Mg2+ + e-
I1 = 738 kJ
I2 = 1451 kJ
The general trend for ionization energy is: Ionization energies decrease as
atomic radii increase.
Electron Affinity
Electron Affinity (EA): measure of the energy change that occurs when a
gaseous atom gains an electron.
Example:
O (g) + e-  O- (g) EA = -141.0 kJ (Energy is given off!)
O- (g) + e-  O2- (g) EA = +744 kJ (Energy must be added since
there is repulsion between Oand the negative electron)
The general trend for electron affinity is: Smaller atoms to the right of
the periodic table tend to have large, negative electron affinities (most
easily gain an electron.)
To Summarize: