Chapter 7
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1st developed by Dmitri Mendeleev (Russia) &
Lothar Meyer (Germany) on the basis of the
similarity in chemical and physical properties
Mendeleev …
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started by organizing elements by increasing mass.
Recognized a repetition of pattern.
Placed elements by same column  same properties
Predicted correctly about the existence of new elements
Henry Moseley
◦ established that each element has a unique atomic
number, which added more order to the periodic table
◦ Identified the atomic number with the # of protons in the
nucleus of the atom & the # of electrons in the atom.
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Atoms aren’t hard spheres with well-defined
shells of electrons
The edges of atoms are a bit “fuzzy”
The quantum mechanical model of the atom
supports the notion of electron shells: certain
distances from the nucleus at which there is a
higher likelihood of finding an electron
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The size of an atom can be gauged by its bonding atomic
radius, based on measurements of the distances separating
atoms in their chemical combinations with other atoms
Measure the atomic radius from the center of the nucleus to
the outermost electron.
Atom size increases going down a group.
Atomic size decreases going left to right across the period.
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Ionization energy – the minimum energy
required to remove an electron from the
ground state of the isolated gaseous atom or
ion
1st ionization energy (I1) – The energy needed
to remove the first electron from a neutral
atom, forming a cation
2nd ionization energy (I2) – the energy needed
to remove the second electron
The greater the ionization energy, the harder
it is to remove an electron
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HIGH ionization
energy means the
atom hold onto the
electron tightly and a
lot of energy is need
to pull it off
LOW ionization energy
means the atom holds
onto the electron
loosely so breaking it
apart doesn’t require
much energy
Periodic Trends in Ionization Energies
 Ionization energy decreases as you move down a
group.
 Ionization energy increases as you move from left
to right on the periodic table.
 Representative elements show a larger range of
values of I1 than do the transition metal elements
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Electron affinity – the energy change that
occurs when an electron is added to as
gaseous atom
A negative electron affinity = the anion is
stable
A positive electron affinity = the anion is
higher in energy than are the separated atom
and electron. The anion is not stable and will
not form
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If the electron affinity is negative, the atom
releases energy.
Normally, non-metals have a more negative
electron affinity than metals. The exception
is the noble gases.
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Election affinities become more negative as
we proceed from left to right
Halogens have the most negative electron
affinities
The electron affinities of the noble gases are
all positive since the added electron would
have to occupy a new, higher-energy subshell
Electron affinity doesn’t change greatly as we
move down a group. Electron affinity should
become more positive (less energy released).
Metals
Non-Metals
Have a shiny luster; various colors, although Do not have a luster; various colors
most are silvery
Solids are malleable and ductile
Solids are usually brittle; some are hard, and
some are soft
Good conductors of heat and electricity
Poor conductors of heat and electricity
Most metal oxides are ionic solids that are
basic
Most non-metallic oxides are molecular
substances that form acidic solutions
Tend for form cations in aqueous solutions
Tend to form anions or oxyanions in
aqueous solution
Pg. 239 --Table 7.3
Characteristic Properties of Metals and Nonmetals
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Metallic Character - The tendency of an
element to exhibit properties of metals
Metallic character generally increases going
down a column and decreases going from
left to right across a period
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Metals conduct heat & electricity
They are malleable & ductile
Solids at room temp. except mercury(Hg) (it’s
liquid)
Melt at very high temps
Have low ionization energies & are
consequently oxidized (lose electrons) when
they undergo chemical reaction.
Many transition metals have the ability to
form more than one positive ion.
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metal oxide + water  metal hydroxide
◦ Most metal oxides are known as basic oxides
◦ Ex: Na2O (s) + H2O(l)  2NaOH (aq)
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metal oxide + acid  salt + water
◦ Ex: MgO (s) + 2HCl (aq)  MgCl2 (aq) + H20 (l)
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Not lustrous & generally are poor conductors
of heat and electricity
Non-metals commonly gain enough electrons
to fill their outer p sub-shell completely,
giving a noble gas electron configuration.
Molecular substances - Compounds
composed entirely of nonmetals
◦ Ex: oxides, halides, and hydrides
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Melting points are gen. lower than those of
metals
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Nonmetal oxide + water → acid
◦ Most nonmetal oxides are acidic oxides
◦ CO2 (g) + H2O (l)  H2CO3 (aq)
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Nonmetal oxide + base  salt + water
◦ CO2 (g) + 2NaOH (aq)  Na2CO3 (aq) + H2O (l)
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Have properties that are intermediate
between those of metals and nonmetals
Group 1A: The Alkali Metals
Characteristics
◦ Soft metallic solids
◦ Silvery
◦ metallic luster
◦ high thermal and electrical conductivities
◦ Low densities and melting points
◦ Most active metals
◦ Exist in nature only as compounds
Group 2A: Alkaline Earth Metals
 Solids with typical metallic properties
 Harder, more dense, and melt at higher
temperatures when compared to alkali metals
 Very reactive towards nonmetals, but not as
reactive as alkali metals
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Both alkali and alkaline earth metals react
with hydrogen to form ionic substances that
contain the hydride ion, H-
Hydrogen
 Hydrogen is a nonmetal with properties that
are distinct from any of the groups of the
periodic table
 It forms molecular compounds with other
nonmetals, such as oxygen and the halogens
Group 6A: The Oxygen Group
 Most important element in group 6A
 Exists in several allotropic forms (different forms
of the same element in the same state)
 Oxygen is encountered in two molecular forms,
O2 (common form) and O3 (aka ozone)
 Oxygen has a strong tendency to gain electrons
from other elements, thus oxidizing them
 In combination with metals, oxygen is usually
found as the oxide ion, O2-, although salts of the
peroxide ion, O22-, and superoxide ion, O2-, are
sometimes formed
Sulfur!!
 2nd more important element in group 6A
 Also exists in several allotropic forms
 Elemental sulfur is more commonly found as
S8 molecules
 In combination with metals, it is more often
found as the sulfide ion, S2-
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Nonmetals that exist as diatomic
molecules
There melting and boiling points
increase as you go down the
column
Have the most negative electron
affinities of the elements
Their chemistry is dominated by a
tendency to form 1- ions,
especially in reactions with metals
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Group 8A: The Noble Gases aka inert gases
Nonmetals that exist as monoatomic gases
Very unreactive since they have completely filled s
and p subshells. Have the complete octet
Have large 1st ionization energies
Only the heaviest noble gases are known to form
compounds, and they do so only with very active
nonmetals, like fluorine