Download IONS OF THE FIRST 20 ELEMENTS

Chemistry: The Physical Setting with Miss Lavonne
Full Name: ____________________________________________________ 11.24.09 (Tu)
REVIEW SHEET: IONS OF THE FIRST 20 ELEMENTS
CATION (+)
Atomic
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ANION (-)
Symbol
Name
of
Element
Ion
Name
of Ion
(“___________ ion”)
H
Hydrogen
H+
Hydrogen ion
Li
Be2+
Na
Aluminum ion
Calcium
Atomic
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Summary Notes: The Periodic Table and Its Elements
Symbol
Name
of
Element
Ion
Name
of Ion
(_________--ide ion)
N
Nitrogen
N3-
Nitride ion
S
12.2.09 (W)
Group 1 Elements: The Alkali Metals
The alkali metals, found in group 1 of the periodic table (formerly known as group IA), are very reactive metals that do not occur freely
in nature. These metals have only one electron in their outer shell. Therefore, they are ready to lose that one electron in ionic bonding
with other elements. As with all metals, the alkali metals are malleable, ductile, and are good conductors of heat and electricity. The
alkali metals are softer than most other metals. Cesium and francium are the most reactive elements in this group. Alkali metals can
explode if they are exposed to water.
The Alkali Metals are: Lithium, Sodium, Potassium, Rubidium, Cesium, and Francium
Group 2 Elements: The Alkaline Earth Metals
The alkaline earth elements are metallic elements found in the second group of the periodic table. All alkaline earth elements have an
oxidation number of +2, making them very reactive. Because of their reactivity, the alkaline metals are not found free in nature.
The Alkaline Earth Metals are: Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium
Groups 3 – 12 Elements: The Transition Metals
The 38 elements in groups 3 through 12 of the periodic table are called "transition metals". As with all metals, the transition elements
are both ductile and malleable, and conduct electricity and heat. The interesting thing about transition metals is that their valence
electrons, or the electrons they use to combine with other elements, are present in more than one shell. This is the reason why they
often exhibit several common oxidation states. There are three noteworthy elements in the transition metals family. These elements are
iron, cobalt, and nickel, and they are the only elements known to produce a magnetic field.
The Transition Metals are: Scandium, Titanium, Vanadium, Chromium, Manganese, Iron, Cobalt, Nickel, Copper, Zinc, Yttrium, Zirconium, Niobium,
Molybdenum, Technetium, Ruthenium, Rhodium, Palladium, Silver, Cadmium, Hafnium, Tantalum, Tungsten, Rhenium, Osmium, Iridium, Platinum, Gold, Mercury,
Rutherfordium, Dubnium, Seaborgium, Bohrium, Hassium, Meitnerium, Ununnilium, Unununium, and Ununbium
Groups 13 – 15 (3 - 5) Elements: The “Other Metals”
The 7 elements classified as "other metals" are located in groups 13, 14, and 15. While these elements are ductile and malleable, they
are not the same as the transition elements. These elements, unlike the transition elements, do not exhibit variable oxidation states,
and their valence electrons are only present in their outer shell. All of these elements are solid, have a relatively high density, and are
opaque. They have oxidation numbers of +3, ±4, and -3.
The "Other Metals" are: Aluminum, Gallium, Indium, Tin, Thallium, Lead, and Bismuth
The “Stair-step Line” Elements: The Metalloids
Metalloids are the elements found along the stair-step line that distinguishes metals from non-metals. This line is drawn from between
Boron and Aluminum to the border between Polonium and Astatine. The only exception to this is Aluminum, which is classified under
"Other Metals". Metalloids have properties of both metals and non-metals. Some of the metalloids, such as silicon and germanium, are
semi-conductors. This means that they can carry an electrical charge under special conditions. This property makes metalloids useful in
computers and calculators
The Metalloids are: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, and Polonium
Groups 14 – 16 (4 – 6) Elements: The Nonmetals
Non-metals are the elements in groups 14-16 of the periodic table. Non-metals are not able to conduct electricity or heat very well. As
opposed to metals, non-metallic elements are very brittle, and cannot be rolled into wires or pounded into sheets. The non-metals exist
in two of the three states of matter at room temperature: gases (such as oxygen) and solids (such as carbon). The non-metals have no
metallic luster, and do not reflect light. They have oxidation numbers of ±4, -3, and -2.
The Non-Metal elements are: Hydrogen, Carbon, Nitrogen, Oxygen, Phosphorus, Sulfur, and Selenium
Group 17 (7) Elements: The Halogens
The halogens are five non-metallic elements found in group 17 of the periodic table. The term "halogen" means "salt-former" and
compounds containing halogens are called "salts". All halogens have 7 electrons in their outer shells, giving them an oxidation number
of -1. The halogens exist, at room temperature, in all three states of matter:



Solid- Iodine, Astatine
Liquid- Bromine
Gas- Fluorine, Chlorine
The Halogens are: Fluorine, Chlorine, Bromine, Iodine, and Astatine
Group 18 (8) Elements: The Noble Gases
The six noble gases are found in group 18 of the periodic table. These elements were considered to be inert gases until the 1960's,
because their oxidation number of 0 prevents the noble gases from forming compounds readily. All noble gases have the maximum
number of electrons possible in their outer shell (2 for Helium, 8 for all others), making them stable.
The Noble Gases are: Helium, Neon, Argon, Krypton, Xenon, and Radon
The Rare Earth Elements
The thirty rare earth elements are composed of the lanthanide and actinide series. One element of the lanthanide series and most of
the elements in the actinide series are called trans-uranium, which means synthetic or man-made. All of the rare earth metals are found
in group 3 of the periodic table, and the 6th and 7th periods. The Rare Earth Elements are made up of two series of elements, the
Lanthanide and Actinide Series.
The Rare Earth Elements are: Lanthanum, Cerium, Praseodymium, Neodymium, Promethium, Samarium, Europium, Gadolinium, Terbium, Dysprosium,
Holmium, Erbium, Thulium, Ytterbium, Lutetium, Actinium, Thorium, Protactinium, Uranium, Neptunium, Plutonium, Americium, Curium, Berkelium, Californium,
Einsteinium, Fermium, Mendelevium, Nobelium, and Lawrencium
Important Trends in The Periodic Table
The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and
understood by analyzing the electron configurations of the elements. Elements tend to gain or lose valence electrons to achieve stable
octet formation. Stable octets are seen in the inert gases, or noble gases, of Group VII of the periodic table. In addition to this activity,
there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens,
the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus
and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to
the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to
the nucleus) increases downward within each group. These trends explain the periodicity observed in the elemental properties of
atomic radius, ionization energy, electron affinity, and electronegativity.
Atomic Radius (AR)
The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching each
other. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the
largest atomic radii are located in Group I and at the bottom of groups.
Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Electrons within a shell cannot
shield each other from the attraction to protons. Since the number of protons is also increasing, the effective nuclear charge increases
across a period. This causes the atomic radius to decrease.
Moving down a group in the periodic table, the number of electrons and filled electron shells increases, but the number of valence
electrons remains the same. The outermost electrons in a group are exposed to the same effective nuclear charge, but electrons are
found farther from the nucleus as the number of filled energy shells increases. Therefore, the atomic radii increase.
Ionization Energy (IE)
The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The
closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy
will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is
the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive
ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energies
increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group
(increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.
Electron Affinity (EA)
Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a
gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about
the electron affinities of certain groups in the periodic table. The Group IIA elements, the alkaline earths, have low electron affinity
values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high
electron affinities because the addition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases,
have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other
groups have low electron affinities.
Electronegativity (E-)
Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an
atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy. Electrons with low ionization
energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high
ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus. In a group, the
electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus
(greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly
electronegative element is fluorine.
Summary of Periodic Table Trends
Moving Left  Right
Atomic Radius (AR) Decreases
Ionization Energy (IE) Increases
Electronegativity (E-) Increases
Moving Top  Bottom
Atomic Radius (AR) Increases
Ionization Energy (IE) Decreases
Electronegativity (E-) Decreases