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Section 6-1 Notes
Organizing the Elements
Organizing the Elements
 As new elements were discovered chemists
needed to find a logical way to organize them
 Properties of elements were used to sort them in to
groups
 In 1829, J.W. Dobereiner published a
classification system in which the elements
were grouped into triads
 A group of 3 elements with similar properties
 Ex. Chlorine, Bromine, and Iodine
 One element in each triad tended to have
properties with values that fell midway
between those of the other two elements
 Problem - some of the elements could not be
grouped into triads
Mendeleev’s Periodic Table
 1869, Dmitri Mendeleev published a table
of elements
 The organization he chose was a periodic
table, based on a set of repeating
properties of the 60 known elements.
 Mendeleev arranged the elements in order
of increasing atomic mass – see fig. 6.3
 He added ? because:
 1st because he knew that Bromine needed to be
with Chlorine & Iodine.
 2nd because he predicted that elements would be
discovered that fit in those spaces
Periodic Law
 Mendeleev arranged his element based on
increasing atomic mass, but he placed
tellurium(127.6amu) before iodine (126.9amu)
in his table.
 He did this because it placed iodine in the same
group as chlorine and bromine.
 He knew that these elements had similar properties
 1913, Henry Moseley, British physicist,
determined an atomic number for each
known element
 By using the atomic number, it makes sense
that Tellurium should come before Iodine
 Therefore, in the modern periodic table,
elements are arranged in order of increasing
atomic number not mass
Periodic Law
 There are seven rows (periods) in the table
 Each period of the table corresponds to a
principal energy level
 The elements within a column (group) have
similar properties
 These properties repeat in each period, a
pattern known as the periodic law
 When elements are arranged in order of
increasing atomic number, there is a
periodic repetition of their physical and
chemical properties
Metals, Nonmetals and Metalloids
 The elements have been grouped into 3
broad classes based on their
characteristics
 Metals
 Nonmetals
 Metalloids
Properties of Metals
 Good conductors of heat and electricity
 High luster or sheen
 Solid at room temperature
 Exception - Mercury
 Ductile - can be drawn into wires
 Malleable - can be hammered into thin sheets
Properties of Non Metals
Most are gases at room temperature
Sulfur and Phosphorus are solids
Bromine is a dark-red liquid
Poor conductors of heat and
electricity
Carbon is an exception
Solid nonmetals tend to be brittle
Metalloids
 Metalloids generally have properties
similar to those of metals and nonmetals
 Depending on the conditions a metalloid
can act as a metal or a nonmetal
 Example
 Pure silicon is a poor conductor of electric
current, similar to nonmetals
 But, when mixed with boron it is a good
conductor of electric current, similar to metals
Section 6-2 Notes
Classifying the Elements
Squares in the Periodic Table
 The periodic table displays the symbols
and names of the elements, along with
information about the structure of their
atoms.
 Each square include the symbol, atomic
mass, and atomic number of the element
 There is a vertical column that which lists
the number of electrons in each energy
level
 Also some tables provide a color
coordinated chart to distinguish some
specific groups of elements
Groups of Elements
 Alkali metals- group 1A elements
 Arabic for “the ashes”, Na & K are common in
wood ashes
 Alkaline earth metals- group 2A elements
 Halogens- group 7A elements
 Hals- Greek for salt genesis- Latin for to be
born
 These elements can be produced from their
salts
 Chalcogens- group 6A elements
 Noble Gases – group 8A elements
Electron Configurations in Groups
 Elements can be sorted into:
 Noble gases
 Representative elements
 Transition metals
 Inner transition metals
 Based on their electron configurations.
 The noble gases all have a full outer shell
so they rarely take part in a reaction
The Representative Elements
 Groups 1A-7A are the representative
elements
 They are referred to as representative
elements because they display a wide
range of physical and chemical properties
 Some are metals, metalloids, and
nonmetals
 Most are solids, some are gases and one
is a liquid at room temperature
 The s and p sublevels are the highest
occupied energy levels that are being filled
Transition Elements
The B groups on the periodic table
are known as transition elements
These elements are all metals
These elements are characterized by
the presence of electrons in the d
sublevel
Inner transition metals are
characterized by the presence of
electrons in the f sublevel
Section 6-3 Notes
Periodic Trends
Trends in Atomic Size
 To determine atomic size we use two
atoms of the same element joined
together
 Since the atoms in each molecule are
identical, the distance between the
nuclei of the atoms can be used to
estimate the size of the atoms
 The atomic radius is ½ the distance
between the two nuclei
 In general, atomic size increases from
top to bottom within a group
 Atomic size decreases from left to right
across a period
Group Trends in Atomic Size
What happens to the # of protons we
move through a group (column)?
The increase in positive charge draws
the electrons closer to the nucleus
The increase in # of occupied orbitals
shields electrons in the highest energy
level from the attraction of protons
 Called the shielding effect
The shielding effect is greater then the
effect of the increase in nuclear charge
As a result, the atomic size increases
Ions
 An ion is an atom or group of atoms that
has a positive or negative charge
 Positive and negative ions form when
electrons are transferred between atoms
 An ion with a positive charge is called a
cation
 Na+, Ca2+
 An ion with a negative charge is called an
anion
 Cl-, O2-
Trends in Ionization Energy
 The energy required to remove an electron
from an atom is called ionization energy
 Always measured when the element is in its
gaseous state
 First ionization energy tends to decrease
from top to bottom
 Recall the shielding effect
 Tends to increase from left to right across a
period
Group Trends in Ionization Energy
Recall that the atomic size increases
as the atomic number increases
within a group
As the size increases, nuclear charge
has a smaller effect on the electrons
in the highest energy level.
Therefore, it takes less energy to
remove an electron, causing the first
ionization energy to be lower
Periodic Trends in Ionization Energy
Energy generally increases as you
move left to right due to an increase
in nuclear charge and the shielding
effect remaining the same for the
period
Since there is an increase in the
attraction of the nucleus for an
electron, it takes more energy to
remove an electron from an atom
Trends in Ionic Size
 Cations are always smaller than the
atoms from which they form.
 Anions are always larger then the atoms
from which they form
 When a sodium atom loses an electron,
the attraction between the remaining
electrons and the nucleus is increased
 As a result, the electrons are drawn closer,
making the size smaller
 This is opposite for nonmetals that gain
electrons because the attraction is
decreased
 More negative charge then positive charge
Trends in Electronegativity
 Electronegativity is the ability of an atom of
an element to attract electrons when the
atom is in a compound
 Scientists use factors such as ionization
energy to calculate values for
electronegativity
 In general, electronegativity values decrease
from top to bottom within a group
 The values tend to increase from left to right
across a period
 Cesium is the least electronegative element
 Fluorine is the most electronegative element
 The electronegativity of transition metals is
irregular