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Chemistry Notes: Chapter 5
1869 Mendeleev created the first periodic table based on relative atomic masses. He found certain chemical
properties appeared at regular intervals. He left several empty spaces in his periodic table but predicted the
properties of these elements by looking at the properties of the elements surrounding the blank space. Later
these three elements were discovered—scandium, gallium and germanium.
1911 Moseley published a periodic table based on atomic numbers. This was the basis for our modern periodic
table.
Periodic Law—the physical and chemical properties of the elements are periodic functions of their atomic
numbers.
Periodic Table—an arrangement of the elements in order of their atomic numbers so that elements with similar
properties fall in the same column or group.
The noble gases were not a part of the early periodic table. They were not yet discovered when Mendeleev
published the first periodic table.
Lanthanide series—14 elements with atomic numbers from 58-71. They belong with period six.
Actinide Series—14 elements with atomic numbers from 90 to 103. They belong with period seven.
The Lanthanide and Actinide series are set apart from the rest of the table to save space.
Groups of elements have similar properties because of the arrangement of their outer shell electrons.
5-2 Electron configuration and the periodic table.
Vertical columns in the periodic table are called groups. There are 18 groups.
Horizontal rows are called periods. There are 7 periods.
The highest energy level indicates the period.
The periodic table can be divided into 4 blocs according to sublevels. (the s, p, d, and f blocks)
“s” block: Groups 1 & 2 Group 1 is the alkali metals. They are silvery, soft and highly reactive.
Consequently, they are not found alone in nature. They react strongly with water to produce H2 gas.
They combine vigorously with nonmetals and most combine with air so are store in kerosene.
Group 2 are the alkaline-earth metals. They are harder, denser, and stronger than the alkali metals. They are
less reactive than Group 1 but still not found free in nature.
Hydrogen is a nonmetal but because it has only 1 electron in the s orbital, it is positioned above the group 1
metals. It doesn’t, however, have the same properties as the other alkali metals.
“d” block elements—often called the transition elements.
They have high luster and are good conductors of electricity.
Are much less reactive than groups 1 and 2.
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“p” block elements—Groups 13-18. (These along with the “s” block elements comprise the main group
elements. Elements in this group fill the p sublevel. The properties vary greatly because the p block contains
some metals, most of the nonmetals and all of the metalloids.
Groups 17 is know as the “halogens”. They are the most reactive nonmetals. Halogen means salt forming.
“f” block elements—Lanthanide and Actinide series. The “f” sublevel is being filled.
Section 5-3:
Atomic radius—since the boundary of an atom is fuzzy, the radius is defined as ½ the distance between the
nuclei of identical atoms that are bonded together.
Atomic radii get larger as we go from top to bottom in a group due to the addition of another energy level.
Atomic radii get smaller as we go from left to right in a series because of the increasing positive charge in the
nucleus pulling the electrons closer.
Ion—atom or group of atoms that has/have a positive or negative charge because of unequal numbers of
protons and electrons.
Ionization—the process of forming an ion by either removing or adding an electron to the neutral atom.
Ionization energy: the energy required to remove an electron from a neutral atom of an element. Usually
listed in kJ/mol. It results in the formation of a “+” ion.
Generally decreases from top to bottom in a group due to increased energy levels being further from the
nucleus.
Generally increase from left to right because of the increased nuclear charge.
Elements can have 1st, 2nd, 3rd, etc. ionization energies as successive electrons are removed from an
atom. Ionization energy becomes higher with each removal of an electron due to an increased net “+” charge.
Ionization energies become very high when removing electrons from a noble gas configuration (octet).
Electron affinity—The energy change that occurs when an electron is acquired by a neutral atom. Most atoms
release energy.
A + e-  A- + energy
If energy is released, it is a given a negative sign and the ion is MORE STABLE than the atom it was formed
from.
Some atoms gain energy when they gain an electron
A + e- + energy  A- If energy is gained or absorbed, the amount of energy is given a + sign and
the ion is “LESS STABLE” (more unstable) than the atom it was made from.
The halogens gain electrons very readily so have large negative affinities. Electron affinities become more
negative across each period in the p block. There are exceptions due to the stability of half filled p sublevels.
In general, electrons are added with greater difficulty as we go down in a group but there are exceptions.
It is always more difficult to add a second electron to an already negatively charged ion, therefore second
electron affinities always require the addition of energy.
Ionic Radii—
A + ion is called a cation.
Because it results from a decrease in outer level electrons, the radius is always smaller than the atom
because the electron cloud gets smaller from protons pulling the electrons closer to the nucleus.
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A – ion is called an anion. This always results in a larger atomic radius because the added “-“ charge
results in electrons not being pulled so close to the nucleus.
Metals tend to form cations and nonmetals tend to form anions. Cation radii decrease across a period because
of the increasing net + charge in the nucleus.
Anion radii also decrease across a period because of increased “+” charge in the nucleus.
Ionic radii generally increase going from top to bottom in a group because of added energy levels.
Valence Electrons—the electrons available to be lost, gained or shared when chemical compounds are formed.
They are in the outer energy level.
Electronegativity—measure of the attraction of an atom for the shared electron in a chemical compound.
Flourine is the most electronegative of all the elements.
Electronegativities increase across each period with some exceptions. Electronegativites increase across each
period with some exceptions. Electronegativities tend to either decrease down a group or stay the same.
D block elements
Atomic radii generally decrease across the periods.
Ionization energies generally increase across the periods.
Generally have low electronegativities and values generally increase as radii decrease.
F Block elements—
Atomic radii generally decrease but there are exceptions./
Ionization energies generally increase across periods.
All have similar electronegativities.
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