•Originally constructed to represent the patterns observed
in the chemical properties of the elements.
•First chemist to recognize patterns was Johann Dobereiner
(1780-1849).
•Noticed several groups of three elements had similar
properties, for example, chlorine, bromine, and iodine.
•Tried to expand his model of triads but it was severely
limited.
•Next notable attempt was made by John Newlands in 1864.
•Suggested that elements should be arranged in octaves.
•This was based on the idea that certain properties seemed
to repeat for every eighth element.
•Model did attempt to group based on properties but not
generally successful.
•Present form of periodic table conceived by Julius Lothar
Meyer (1830-1985) and Dmitri Mendeleev (1834-1907).
Meyer
Mendeleev
•Mendeleev is given most of the credit because he
emphasized the table could be used to predict the existence
and properties of unknown elements.
•He published his table in 1872.
•Mendeleev predicted the existence and properties of the
elements gallium, scandium, and germanium from gaps in
his periodic table.
•Germanium was discovered in 1886 and his predicted
values and those observed are in excellent agreement.
•Mendeleev was also able to predict atomic masses of
several elements, including indium, beryllium and uranium.
•Mendeleev’s table was almost universally adopted and
remains one of the most valuable of a chemist’s tools.
•The fundamental difference between Mendeleev’s table
and the modern periodic table is the modern table uses
atomic number to order the elements rather than atomic
mass.
•Valence electrons are the electrons in the outermost
principal quantum level (outermost energy level) of an
atom.
•Electron configuration for nitrogen: 1s22s22p3
•The valence electrons for nitrogen are the 2s and 2p
electrons; therefore, nitrogen has five valence electrons.
•Valence electrons are important because they are involved
in bonding.
•Core electrons are the inner electrons.
•Elements with the same valence configuration show similar
chemical behavior.
•Groups 1, 2, 13-18 are often called the main-group or
representative elements.
•Every member of these groups has the same valence
electron configuration.
•Predicting the valence electron configurations of the
transition metals, the lanthanides, and the actinides is
somewhat more difficult because of the many exceptions.
•There are observed trends in
several important atomic
properties: ionization energy,
electron affinity, and atomic
size.
•The effective nuclear charge is the pull that an
electron “feels” from the nucleus.
•Effective Nuclear Charge (Zeff) = # protons - # core
electrons
•The closer an electron is to the nucleus, the more
pull it feels.
•As effective nuclear charge increases, the electron
cloud is pulled in tighter.
•Ionization energy is the energy required to remove an
electron from a gaseous atom or ion:
X (g) → X+ (g) + e-
•Consider the energy required to remove several electrons
from aluminum in the gaseous state.
Al (g) → Al+ (g) + eI1 = 580 kJ/mol
Al+ (g) → Al2+ (g) + eI2 = 1815 kJ/mol
Al2+ (g) → Al3+ (g) + eI3 = 2740 kJ/mol
Al3+ (g) → Al4+ (g) + eI4 = 11,600 kJ/mol
Al (g) → Al+ (g) + eAl+ (g) → Al2+ (g) + e-
I1 = 580 kJ/mol
I2 = 1815 kJ/mol
•The highest energy electron (the one bound least tightly is
removed first.
•I1 is the first ionization energy and for aluminum, this
electron comes from the 3p orbital ([Ne]3s23p1).
•I2 is the second ionization energy and this electron comes
from the 3s orbital.
•Why is I1 smaller than I2?
•The first electron is removed from a neutral atom and the
second is removed from a 1+ ion.
•The increase in positive charge binds the electrons more
firmly and it takes more energy to remove an electron.
Al (g) → Al+ (g) + eAl+ (g) → Al2+ (g) + eAl2+ (g) → Al3+ (g) + eAl3+ (g) → Al4+ (g) + e-
I1 = 580 kJ/mol
I2 = 1815 kJ/mol
I3 = 2740 kJ/mol
I4 = 11,600 kJ/mol
•Why is I4 so high?
•The fourth electron is “core” electron (Al3+ = 1s22s22p6) and
core electrons are bound more tightly than valence electrons.
•In general as we go across a period from left to right, the
first ionization energy increases.
•Reason: increase in effective nuclear charge (more
protons in nucleus) felt by the valence electrons across a
period.
•Causes the valence electrons to be held more tightly,
which makes it more difficult to remove them.
•Note: there are exceptions in ionization energy trends in
going across a period. Due to shielding and electron
repulsions.
•First ionization energy decreases in going down a group.
•Reason: going down a group the electrons being removed
are, on average, farther from the nucleus.
•As n increases, the size of the orbital increases, and the
electrons are farther from the nucleus, and thus are easier
to remove.
•Electron affinity is the energy change associated with the
addition of an electron to a gaseous atom:
X (g) + e- → X- (g)
•If the addition of the electron is exothermic, the
corresponding value for electron affinity will carry a
negative sign.
•The incoming electron experiences an attraction to the
nucleus, which causes the potential energy to be lowered
as the electron approaches the atom.
•The trends in electron affinity are similar to those for
ionization energy.
•Electron affinity becomes more exothermic from left to
right across a period. A valence shell that holds its electrons
tightly will also tend to bind an additional electron tightly.
•Electron affinity becomes less negative down a group. A
valence shell that loses electrons easily (low IE) will have
little attraction for additional electrons (small EA).
•Note: there are exceptions.
•The radius of an atom (r) is defined as half the distance
between the nuclei in a molecule consisting of identical
atoms.
•For nonmetallic atoms that do not form diatomic
molecules, the atomic radii are estimated from their various
covalent compounds.
•The radii for metal atoms (metallic radii) are obtained from
half the distance between metal atoms in solid metal
crystals.
•Atomic radii decrease in going from left to right across a
period.
•Due to increasing effective nuclear charge in going from
left to right. Valence electrons are drawn closer to the
nucleus, decreasing the size of the atom.
•Atomic radius increases down a group, because of the
increases in the orbital sizes in successive principal quantum
levels.
•Negative ions are always larger than the atoms from which
they are formed.
•When electrons are added to an atom, the mutual repulsions
between them increase.
• The causes the electrons to push apart and occupy a larger
volume.
•Positive ions are always smaller than the atoms from which
they are formed.
•When electrons are removed from the valence shell, the
electron-electron repulsions decrease, which allows the
remaining electrons to be pulled closed together around the
nucleus.
Electronegativity
•Valence electrons hold atoms together in
chemical compounds.
•In many compounds, the negative charge of
the valence electrons is concentrated closer to
one atom than to another.
•This uneven concentration of charge has a
significant effect on the chemical properties of a
compound.
Electronegativity
•Electronegativity is a measure of the ability of an atom
in a chemical compound to attract electrons (the most
electronegative element is fluorine).
•Electronegativity increases across each period.
•Electronegativity decreases or stays the same down a
group.