Periodic Table:
Dmitri Ivanovich Mendeleev produced remarkably his version of the periodic table of elements in the 1800’s.
Mendeleev periodic table appeared in his work "On the Relationship of the Properties of the Elements to their
Atomic Weights" in 1869. Mendeleev placed many elements out of order based on their accepted atomic
weights at the time.
Mendeleev predicted the existence and properties of unknown elements which he called eka-aluminum, ekaboron, and eka-silicon. The elements gallium, scandium and germanium were found later to fit his predictions
quite well.
Periodic Law: Many of the physical and chemical properties of elements repeat in a periodic manner.
Periodic Trends:
Atomic Radius
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
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.
Electronegativity
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
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Atomic Radius Decreases
Ionization Energy Increases
Electron Affinity Generally Increases (except Noble Gas Electron Affinity Near Zero)
Electronegativity Increases
Moving Top → Bottom
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Atomic Radius Increases
Ionization Energy Decreases
Electron Affinity Generally Decreases Moving Down a Group
Electronegativity Decreases
Ions:
Cations and anions are both ions. The difference between a cation and an anion is the net electrical charge of
the ion.
Ions are atoms or molecules which have gained or lost one or more valence electrons giving the ion a net
positive or negative charge.
Cations are ions with a net positive charge.
Silver: Ag+, hydronium: H3O+, and ammonium: NH4+
Anions are ions with a net negative charge.
Examples: hydroxide anion: OH-, oxide anion: O2-, and sulfate anion: SO42Anions are Bigger, Cations are smaller.
anions are negatively charged species because they have gained an extra electron, which is negatively charged.
The more electrons an atom has the more the outer electrons are shielded from the pull that the positive
nucleus has on them, therefore when an atom gains an electron the nucleus can not pull in a strongly and the
resulting radius, specifically ionic radius is larger.
cations have lost an electron and therefore do not have as many electrons to shield the outer shell electrons
from the attraction of the nucleus, the electron cloud is pulled closer and they are smaller
The Octet Rule:
The octet rule states that elements gain or lose electrons to attain an electron configuration of the nearest noble
gas. Here is an explanation of how that works and why elements follow the octet rule.
The Octet Rule
Noble gases have complete outer electron shells, which make them very stable. Other elements also seek
stability, which governs their reactivity and bonding behavior. Halogens are one electron away from filled energy
levels, so they are very reactive. Chlorine, for example, has seven electrons in its outer electron shell. Chlorine
readily bonds with other elements so that it can have a filled energy level, like argon. +328.8 kJ per mole of
chlorine atoms are released when chlorine acquires a single electron. In contrast, energy would be required to
add a second electron to a chlorine atom. From a thermodynamic standpoint, chlorine is most likely to
participate in reactions where each atom gains a single electron. The other reactions are possible, but less
favorable. The octet rule is an informal measure of how favorable a chemical bond is between atoms.
Why Do Elements Follow the Octet Rule?
Atoms follow the octet rule because they always seek the most stable electron configuration. Following the octet
rule results in completely filled s- and p- orbitals in an atom's outermost energy level. Low atomic weight
elements (the first twenty elements) are most likely to adhere to the octet rule.
Properties of Metals:
Examples of Metals
Most of the elements on the periodic table are metals, including gold, silver, platinum, mercury, uranium,
aluminum, sodium and calcium. Alloys, such as brass and bronze, also are metals.
Location on the Periodic Table
Metals are located on the left side and the middle of the periodic table. Group IA and Group IIA (the alkali
metals) are the most active metals. The transition elements, groups IB to VIIIB, are also considered metals. The
basic metals are the element to the right of the transition metals. The bottom two rows of elements beneath the
body of the periodic table are the lanthanides and actinides, which are also metals.
Properties
Metals are shiny solids are room temperature (except mercury, which is a shiny liquid element), with
characteristic high melting points and densities. Many of the properties of metals, including large atomic radius,
low ionization energy, and low electronegativity, are due to the fact that the electrons in the valence shell of a
metal atoms can be removed easily. One characteristic of metals is their ability to be deformed without breaking.
Malleability is the ability of a metal to be hammered into shapes. Ductility is the ability of a metal to be drawn
into wire. Because the valence electrons can move freely, metals are good heat conductors and electrical
conductors.
Summary of Common Properties
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Shiny 'metallic' appearance
Solids at room temperature (except mercury)
High melting points
High densities
Large atomic radii
Low ionization energies
Low electronegativities
Usually, high deformation
Malleable
Ductile
Thermal conductors
Electrical conductors