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Elements with the same kind of outer shell make bonds in similar
ways.
Alchemists started to look for patterns among the various elements,
especially the ratios of elements in molecules.
Newlands noticed that when (some) of the elements are listed in
ever increasing mass, they tended to repeat their properties every
eighth time. He called this observation his "Law of Octaves".
This is the evidence behind Lewis's "octet rule".
A Russian Alchemist named Dimitri Mendeleev developed a slightly
better theory about the chemical periodicity, but he still arranged
the elements according to their atomic mass, instead of their
atomic number! However, this is pretty close for the smaller elements
because it is a good guess that the most common isotope has the
same number of protons as neutrons.
Moseley's X-ray work allowed him to count the protons and arrange the
elements into increasing atomic number. When Alchemists arranged
the elements according to Moseley's atomic NUMBER (instead of
atomic mass), the elements with similar electronic configuration
lined up into Groups.
The Periodic Table helps us to keep track of the many different elements and their properties.
Each column, running from top to bottom, is called a Group
and contains elements with the same type of outer shell. Therefore
elements within a Group often have similar chemical properties
because they have similar electronic structures. By knowing the
behaviour of one element in a Group, you can make a good guess
that the others will behave the same.
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The first 20 elements can be remembered by learning this (stupid)
poem:
"Hi. Here Little Beggar Boys Catch Newts Or Fishes.
Never Nab Maggots Alone Since Poisonous Substances Clog Arteries
Keeping Calm."
Forms of an element having different properties (like atomic size, arrangement or molecular configurations) are called allotropes (pronounced "al-oh-tropes"), or allotropic forms of that element.
The atomic radius is the distance from the atom's nucleus
to its outermost shell. The outermost shell is where that atom
will run up against another atom's shell. This is what we normally think of as the size of the normal atom.
The Periodic Table acts as a guide to understanding the size of
each atom.
An angstrom (pronounced "ang-strum") is 10 billionth
of a meter, or 0.0000000001 of a meter.
Alchemists use X-rays to figure out the distance between two atoms
in a bond. The exact distance will vary from one kind of molecule
to another and even from one kind of arrangement to another, but
an average size can be assigned to each atom.
As you go DOWN the GROUP the atoms get BIGGER because each new
Period adds a new shell.
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Ionization energy is the energy needed to REMOVE electrons.
The first ionization energy is the energy needed to remove the
least tightly bound electron from a neutral atom, creating a cation (X+).
One electron volt (eV) is the energy gained by an electron when it "falls" through a potential (difference) of 1 volt (V). It takes 13.6 eV to remove the electron from a hydrogen atom, so hydrogen's (only) electron has an ionization energy of 13.6 eV. Ionization energies are smallest at the lower left of the Table (near Cs) and largest at the upper right (around He). Note: The Group III elements break away from this pattern. |
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Electron affinity is a measure of how well an atom ACCEPTS
electrons to become an anion (X-). This is the opposite of ionization energy, but uses a whole different Table because the production of the two ions is caused by different affects.
Elements with high electron affinity (close to fluorine in the Table) readily accept an electron. Those with low electron affinity are reluctant to take an electron. The strongest diatomic compounds created by ionic (electrovalent) bonds are those made from one atom with a very low ionization energy (like Na) and the other atom with a very high electron affinity (like Cl). The atom with low ionization energy will readily lose its electron, becoming the cation. The atom with the high electron affinity will readily accept an electron, becoming the anion. Electrostatic forces make the bond between them. |
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Linus Pauling introduced the idea of "electronegativity"
to simplify how we relate an atom's power to ATTRACT electrons
in molecules. Electronegativity affects what kind of bond(s)
an element can make, so it is very useful.
Electronegativity increases as we move left to right across a Period and decreases as we move down a Group. This is the opposite of atomic size. |
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Summary of Periodic Properties
As you move across a Period the increase in nuclear charge makes
the nucleus more electron attracting, causing higher electronegativity
and ionization energy but smaller size.
As you go down a Group the additional electron shells move (and
shield) the outer electrons from the nucleus causing lower electronegativity
and ionization energy and a bigger size.
The noble gases are stable, inert and monatomic.
HYDROGEN is an unusual element.
It has some properties like Group I elements because it has a single electron in its outer (only) shell, but it can also behave like a Group VII element as it tries to complete its outer (only) shell. Hydrogen can form covalent bonds or ionic bonds of either ion, depending on the conditions it is in. It can form ionic bonds as an anion (H-) called a hydride (pronounced "hi-dried") by accepting an extra electron to complete its K-shell. The exact position of hydrogen in the Table is not as important as knowing that it can behave in a wide variety of ways depending upon the OTHER atom it is bonding to. |
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Group I elements are the "ALKALI METALS".
These elements have a single electron in their outer shell and form cations readily. They rapidly form hydroxide (something-OH), when placed in water or moist air. This reactivity increases as you go down the Group going from a slight tarnish for lithium (Li) to an explosion with cesium (Cs). These hydroxides are strongly "alkaline" (having an excess amount of hydroxide ions, OH-). They have excellent metallic behavior, including high electrical conductivity (ability to pass along electrons), high heat conductivity and they are shiny ("have a high luster"). They are also malleable (can be hammered or pressed into thin sheets) and ductile (can be drawn into a wire). All these metallic properties increase as you move down the Group. All these elements are very soft, silvery metals. The softness increases down the Group. Their chemical property (alkali) and physical properties (that of a metal) cause them to be called the "alkali metals". |
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Group II elements are the "ALKALINE METALS".
These have two electrons in their outer shell and form cations rather well. They have very similar properties to their neighbors on their left (the Group I elements) but not so strong. Their reactivity, softness and conductivity (along with other metallic properties) increase down the Group (like Group I) They cause water to have excess hydroxide ions (OH-) but not as alkaline as the Group I elements. They have good metallic behaviour, but not as good as the Group I elements. Therefore, they are called the alkaline (notice the "ne" at the end) metals. |
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Metals in the Groups III and onwards are the "POST-TRANSITIONAL METALS".
They are arranged in a staircase with aluminum (Al) near
the top (Group III, Period 3) and polonium (Po) at the bottom
(Group VI, Period 6). As you go down the Groups the metallic properties
become stronger. As you move across a Period the metallic properties
become weaker. These two effects produce a staircase.
Aluminum is the most abundant metal on earth (or the moon) and a very important element to Man in the 20th century. In nature it is found as aluminum oxide because aluminum is VERY reactive with oxygen. Most of it is the ore "bauxite". A man named Hall found a way to make pure aluminum using electricity to push electrons into the bauxite, releasing the oxygen. Pure aluminum quickly "rusts" again, but the "rust" produces a protective film (like tarnish on alkali metals) so the aluminum underneath is protected from the air. |
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Slightly above the Metals in the Group III and onwards are "SEMIMETALS
AND METALLOIDS".
The elements that carpet the metal staircase, are "semimetals"
because they have metallic behaviors under some conditions. For
example, graphite is a semimetal because it carries electrons
in directions parallel to the carbon sheets, but not in other
directions. The electrical properties of the semimetals can often
be increased by adding tiny amounts of metals, producing semiconductors.
Most semimetals also display (slightly and under certain conditions)
the chemical properties of a metal, such as the ability
to produce alkaline solutions, and are thus also called "metalloids".
Elements are not really defined as metalloids or semimetals. Instead
an element's properties are described as metalloid (if
it has the chemical properties of a metal) or semimetal (if it
conducts electricity and has other physical properties of a metal).
It's a "fuzzy" part of the Table.
Nonmetals in Group IV and onwards are the "POST-TRANSITIONAL
NONMETALS".
These elements are bad conductors of electricity and heat. (Except
carbon as graphite, but that is because of graphite's molecular
structure, not carbon's atomic properties)
Because these elements have no metal properties, they are called
"nonmetals".
The Group IV elements have four electrons in their outer shell.
The top two Group IV elements (carbon and silicon) share their
outer electrons in covalent bonds (they have a covalency of 4),
but the others are metals.
Carbon is so good at forming covalent bonds that it has a chemistry of its own, Organic Chemistry. Organic molecules use carbon as the base of their structures and these can be very large and complex. Pure carbon has three allotropes. 1) Graphite is layers (sheets) of carbon. These layers slide over each other very easily and some are left behind when dragged across a sheet of paper. Graphite is pencil lead. 2) Diamonds are three dimensional crystals of carbon. They are the hardest substance known. 3) Fullerenes are carbons arranged in cages, shells and tubes. The most common fullerene is a sphere of 60 carbons (a buckyball). |
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The nonmetals in Group V include nitrogen (N) and phosphorous
(P).
Nitrogen (N2) makes up 80% of the air and is a VERY stable molecule - relatively inert! Pure phosphorous has two principal allotropes. When phosphorous atoms are arranged in a crystal, it is a dangerous (poisonous and spontaneously flammable) substance called "white phosphorous". But "red phosphorous" is a safe amorphous allotrope (like glass, with no repeating arrangement of atoms).
The nonmetals in Group VI include oxygen (O) and sulfur (S).
Note: the four elements at the top of Groups V and VI: nitrogen (N), phosphorous (P), oxygen (O) and sulfur (S): can (under certain conditions) make unusual numbers of bonds in unusual ways. |
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The Group VII elements are the "halogens" (from the
Greek meaning "salt-producer").
All except astatine (At) can form diatomic, elemental molecules by sharing a pair of electrons between them (a covalent bond). These elemental forms are VERY reactive and that reactivity decreases as you go down the Group. Because the elemental forms are so active, they are rarely found free in nature. Instead they are found combined with other elements forming compounds. Fluorine is the most reactive member and forms the most stable compounds. |
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The Group VIII elements are (monatomic) gases called the "noble gases".
All the other elements are a wee bit confusing. They are called the "TRANSITION ELEMENTS", but the ones with f orbitals are sometimes called the "INNER TRANSITION ELEMENTS". It depends on whom you talk to and how specific you want to be in your discussion.
Their strange shells allow the transition elements to conduct
electrons very easily so they are metals and sometimes called the transition
metals. Unlike the other (pre-transition) metals that are soft, the transition
metals are mostly hard, brittle and have high melting temperatures.
Mercury (Hg) is an exception, because it is a liquid at room temperatures
(but it would have all these metal properties if cooled to a solid).
The weird outer shells of the transition elements cause them
to be smaller than expected, produce covalent bonds between ions and have a variety
of stable ionic states.
There are a lot of important elements among them. Iron (Fe), tin
(Sn), copper (Cu), silver (Ag) and gold (Au) have been important
elements in creating our civilization.
All the transition metals have d orbitals as their second outermost orbitals. They start right after Group II and end right before Group III. |
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The fourth Period is where the transition metals start with scandium (Sc). Scandium is the first atom to have a d orbital. It's the first transition metal and zinc (Zn) is the last transition metal in Period four. Yttrium (Y) starts the transition elements in the fifth Period and cadmium (Cd) finishes that Period of transition elements. Lanthanum (La) starts the transition elements in Period six and the last transition element of Period six is mercury (Hg). Actinium (Ac) is the only transition element in the seventh Period.
Two inner transition element series are hidden within the transition elements starting in the sixth Period.
The lanthanoid series starts with cerium (Ce), the first element to use f orbitals. It ends with lutetium (Lu). The actinoid series is in the seventh Period and starts with thorium (Th). |
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An atom's electron structure can often be read from the Table. The outer orbitals (found in the outer shell) affect the shape and behavior of the atoms. |
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