# Periodic Tablerev2018

```Almost Everything you need to know about the Periodic
Table of the Elements (PTE), and periodic law…
Introduction (you may actually want to start with page 2, Periodic Table
History, if you don’t understand this part; but come back to it at the end)
Concentrate on understanding the underlying concepts and you won’t
need to memorize anything.
Focus: in chemistry, it’s almost always about the electrons!
Periodic Law: When the elements are arranged in order of increasing
atomic number, there is a periodic repetition of their properties.
(For example, the atom with the largest radius in any period (row of elements)
always occurs in group 1, and the size of the atoms tends to decrease as the atomic
number increases in any period; the trend is that atomic radius tends to decrease
as you go from left to right (increasing atomic number) in a period.)
The electrons are held in place because of their attraction to the nucleus. The
periodic trends of the properties of elements are a direct result of the variations in
the attraction of the nucleus for the electrons, particularly the valence electrons, of
the atom.
The strength of the attraction depends on three factors
1) charge of the nucleus (aka nuclear charge): more protons results in stronger
attraction for the electrons in a shell)
2) distance between the nucleus and the valence shell ( just like with magnets, as
they get further away from each other, the attraction of protons for electrons decreases:
the force is inversely proportional to the distance between the charges SQUARED!)) and
3) shielding effect: determined by the number of inner electrons between the
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valence shell and the nucleus. Although the protons in nucleus attract the valence
electrons, all of the inner shell electrons REPEL the valence electrons: we use the term
“shielding” to indicate that the inner electrons shield the valence electrons from feeling
the full effect of the positively charged nucleus.
The reasons for virtually any periodic trend can be described by
1) an increase in nuclear charge (number of protons in the nucleus) which
causes the valence electrons to be held more strongly. (This is the
dominant factor as you go across a row,)
OR
2) the increase in distance between the nucleus and the outer shell of
electrons, and increased amount of shielding by the electrons in the inner filled
principle energy levels, which results in a decreased attraction between the protons
(nucleus) and valence electrons. (These are the dominant factors as you go down a
family/ group/column)
When asked to explain the causes of trends across periods or down families, always
think of the factors described above. When trying to figure out the trends in atomic
DON’T GUESS: USE TABLE S
Periodic Table History
1. Mendeleev---The Russian who first organized the elements into what was to become
the periodic table. He arranged the elements in families or groups, based on similar
chemical properties and atomic masses. Chemical reactivity is determined
primarily by the electron configuration ( in particular the number of valence
electrons.) The outermost shell of electrons is the valence shell. The families or
groups of the periodic table are organized in vertical columns, and every member of
a column has the same number of valence electrons. For this reason, elements in
the same column tend to have similar properties and reactivity. They will also
have similar formulas when reacting: All group 2 metals form 1:1 oxides: BeO,
MgO, CaO, SrO, BaO, etc.
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2. Elements are the most stable when they have a full outer shell…and the magic
number is 8 electrons! Magic eight ball, eight is enough, however you want to
remember it, understand that all chemical reactions take place to fulfill the Octet
Rule…attaining a full outer shell through transfer of electrons from one element to
another (ionic bonds, like sodium chloride, Na+Cl-) or by sharing electrons between
more than one element (covalent bonds like CO2). Mendeleev’s PTE arrangement
was revised by Thomas Mosely, and elements are listed in order of atomic number
(number of protons), not atomic mass in the modern periodic table of the elements.
Exception to Octet Rule: H, He, Li, Be , B will adopt stable configurations with only 2
electrons since they will fill only the first shell.
3. Metals vs. Non-metals The most fundamental difference in the elements is the
distinction between metals and non-metals. Metals and non-metals are separated in
the periodic table by the “staircase”, a bold stepped division starting between boron
and aluminum and proceeding at a diagonal through the elements polonium and
astatine. Elements to the left of the ladder are metallic and those to the right of the
ladder are non-metallic. Elements that border the ladder are known as semimetals or metalloids, and have properties in between those of metals and non-metals.
boron, silicon, germanium, arsenic, antimony, tellurium and polonium. Most of the
elements in the periodic table are metallic.
4. Properties of Metals: Metals are characterized by “loose” valence electrons.
Remember the M e in metals stands for “Mobile Electrons”! All of the properties of
metals, ranging from their luster or shininess, their malleability (can be hammered
into shape) and ductility (capable of being drawn out into thin wire) to their good
conductivity of thermal and electrical energy are due to the “loosely held electrons”
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of metals. Metals in the elemental state are often characterized as “positive ions in a
sea of free electrons.” Metals typically have few electrons in their valence shell
(less than 4), and do not have a high enough nuclear charge to be able to fill the shell
by attracting more electrons. Metals have only one trick in chemical reactions: they
lose electrons to form positive ions. Metals obtain a full (8 is enough) outer shell (a
noble gas electron configuration) by losing (transferring out) electrons in the valence
shell, to drop back to the underlying completed shell. For this reason, metals form
positive ions (cations) which are smaller than their atoms.
5. Ionization energy is the minimum amount of energy it takes to remove the most
loosely held electron from an atom in the gaseous state. In essence, the electron is
promoted to such a high energy (distance from nucleus) that it is no longer held by
the atom, and is separate from it.
Na(g) + ionization energy  Na+ + eMetals have low ionization energies.
Since metals lose electrons to become
positive ions, their ionic radii are smaller than their atomic radii. Metallic
properties increase as you go down a group’s column. This is because the outer
electrons become “looser” as they are further from the positive charge of the nucleus.
In addition to the decreased electrical attraction due to the valence electrons being at
greater distance from the nucleus, the electrical attraction between the protons and the
valence electrons is further reduced by the presence of the completed inner shells of
electrons which shield and repel the valence electrons from the nucleus. These two
factors, increased distance, and shielding by inner electrons result in “looser”
valence electrons and a corresponding increase in metallic properties as you
move down a column or family of elements. Thus, ionization energies decrease as
you go from lithium to francium. However, ionization energies tend to increase as
you go across a period. This is because the same valence shell is being attracted by
an increasing number of protons while the amount of shielding (number of inner
electrons) remains constant. Because we associate metallic properties with “loosely
held” electrons, metallic properties decrease as you go across a row, but increase as
you go down a column of the PTE.
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6. Remember that many of the periodic trends: atomic radii (size), ionization energies,
electronegativity, can be looked up in Table S of the Chemistry Reference Tables.
DON’T GUESS: USE TABLE S
7. Two important family names in the metals are the alkali metals, located in group 1,
but not including hydrogen, and the alkaline earth metals, located in group two.
Both of these groups of metallic elements are so reactive that the elements are never
found unreacted in nature. The elements are found only as compounds, typically
salts. The free elements can be produced from the salts by electrolysis which uses
electrical energy to force electrons back into the metals to decompose the compound
and obtain the metals in their elemental state. Metallic properties tend to decrease
as you move across a row or period. This is due to increased numbers of protons
in the nucleus increasing the attractive forces on the valence electrons. Thus, the
electrons are less “loose”, ionization energies go up as you go across a row, and all
other metallic properties decrease. Atomic radii tend to decrease as you go across a
row, again indicating greater attractive forces pulling the electrons in “tighter” as the
nuclear charge increases.
8. The transition metals are located in a wide band in the center of the periodic table.
Transition elements are characterized by 1) the presence of d-sublevel electrons, 2)
the tendency to have multiple oxidation states(more than one possible charge), 3) the
tendency to form colored compounds, and 4) the tendency to have more than one
shell(principle energy level) of electrons participating in the reaction. Since there are
5 orbitals in the d-sublevel, each of which can hold two electrons, there are ten
elements across the transition elements.
9. The two rows of elements separated out at the bottom of the periodic table are the
lanthanides and the actinides. They are also known as the inner transition
elements. The f-sublevels are filling as you move across the inner transition
elements. Since there are 7 orbitals in the f-sublevel, these rows each contain
fourteen elements. The lanthanoid elements have also been called the “rare earth
elements” All of the elements with more than 82 protons have no stable isotopes; they
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are all radioactive. All of the elements with more than 92 protons (transuranic (past
uranium) elements are manmade elements or the products of nuclear reactions)
10. Properties of Nonmetals Nonmetals are characterized by tightly held valence
electrons, (high ionization energies) The tightly held electrons of nonmetals account
for their poor conductivity, and their tendency to be dull (not shiny) and brittle(not
formable).Nonmetals have four or more valence electrons, and typically will form
covalent or ionic bonds to complete the valence shell. Nonmetals’ tendencies to
gain electrons are reflected in their high ionization energies and high
electronegativity. If nonmetal atoms gain electrons to form negative ions, their ionic
radii are larger than their atomic radii. Metal atoms lose electrons to form positive
ions which are smaller than the starting atoms: Nonmetal atoms may gain electrons
to form negative ions which are bigger than the starting atom.
11. Electronegativity is a measure of an atom’s tendency to attract electrons that are
shared in a covalent bond with another element. Unlike ionization energies,
which can be measured directly, and expressed in kilojoules per mole,
electronegativity is a relative, unit-less scale, which gives only the atom’s
tendency to attract a shared pair of electrons relative to another atom. The most
electronegative element is fluorine, which was assigned a value of 4.0. Metals,
which do not gain electrons have low electronegativity values. Since the group 18
elements tend not to form bonds, there are no electronegativity values for He, Ne, or
Ar. Since an atom’s ability to attract an electron is a measure of how well the nucleus
of the atom can draw in (attract) another electron, electronegativity values decrease as
you go down a family, as the nucleus is shielded by the inner electrons, and the new
electron must be added to an orbital at a greater distance from the nucleus.
12. Groups 14( four valence electrons), 15 (five valence electrons) and 16 (six valence
electrons) are named for the first member of the groups, and are known as the carbon
family, nitrogen family and the oxygen family, respectively. (There actually are
some other names, but they are not well known or used). These families show a
progression from nonmetals to metalloids to metals as you go down a group.
13. Group 17 is an important family in the nonmetals and is known as the halogen group.
The halogen group is interesting in that it is the only group in the periodic table that
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contains elements in all three common phases of matter at STP. (Standard
temperature and pressure are 0C and 1 atmosphere of pressure). Fluorine and
chlorine are gases at STP, bromine is a liquid, and iodine and astatine are solids. The
halogens form diatomic molecules: F2, Cl2, Br2, and I2. These molecules have
nonpolar covalent bonds and have only weak van derWaals forces (induced dipoleinduced dipole) between the molecules. However, the strength of the van der Waals
intermolecular forces increases as molecule size increases, which is the reason the
smallest (F2, Cl2) are gases, Br2 is a liquid and I2 is a solid at room temperature. The
halogens have seven valence electrons and need to gain only one valence electron to
fill their outer shell and form negative one ions. They halogens are so reactive that
they are never found unreacted in nature. The elements are found only as compounds,
typically salts. The free elements can be produced from the salts by electrolysis
which uses electrical energy to remove extra electrons from the nonmetals to
decompose the compound and obtain the halogens in their elemental state.
14. Group 18 The Noble gases
Group 18 elements are characterized by eight valence
electrons. Since these elements have a complete valence shell in their elemental state,
they have little tendency to participate in reactions. Thus they have high ionization
energies and low electronegativities. (He, Ne, Ar have no electronegativity values
listed because they do not form bonds!) Although they have also been known as
“inert” gases, the higher elements of the family can be compelled to form compounds
with other elements under forcing conditions. Therefore, the term inert gas has fallen
out of favor with chemists.
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