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11/30/15
Chapter 5:
The Periodic Law
Section 5.1
Objectives:
•  Explain the roles of Mendeleev and Moseley in
the development of the periodic table.
•  Describe the modern periodic table.
•  Explain how the periodic law can be used to
predict the physical and chemical properties of
elements.
•  Describe how the elements belonging to a group
of the periodic table are interrelated in terms of
atomic number.
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History of the Periodic Table - Mendeleev
•  What is the meaning of the word periodic?
•  Russian chemist Dmitri Mendeleev noticed that when
elements were arranged in order of increasing atomic
mass, certain similarities in their chemical properties
appeared at regular intervals.
•  In 1869, he published a table in which elements of similar
properties were grouped together – a periodic table of the
elements.
•  Based on his table, he was able to predict the existence of
elements not yet discovered.
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Properties of Some Elements Predicted By Mendeleev
Elements discovered later:
Since Mendeleev’s time, more than 40 new elements have been
discovered/synthesized. Each can be placed in a group of other
elements with similar properties.
NOBLE GASES
•  1894 argon (Ar) was discovered.
•  It was previously unknown because it is completely unreactive.
•  1895 helium (He) was discovered.
•  These discoveries led to a new column (group) being added to the
periodic table.
LANTHANIDES
•  14 elements
•  Atomic numbers 58-71
To save space, the
•  Belong in Period 6
lanthanides and actinides are
ACTINIDES
usually set off below the main
•  14 elements
portion of the periodic table.
•  Atomic numbers 90-103
•  Belong in Period 7
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History of the Periodic Table - Mosley
•  In 1911, English scientists Mosley and Rutherford recognized
that elements fit into patterns better when they were arranged
according to increasing number of protons, rather than
increasing atomic mass.
•  Mosley’s work led to the modern definition of atomic number.
•  When arranged by atomic number, elements with similar
properties appear at regular intervals.
•  Periodic law: the physical and chemical properties of the
elements are periodic functions of their atomic numbers
•  The periodic table is an arrangement of the elements by
increasing atomic number so that elements with similar
properties fall in the same column, or group.
Section 5.2
Electron Configuration and the
Periodic Table
Objectives:
•  Describe the relationship between electrons in sub-levels
and the length of each period of the periodic table.
•  Locate and name the 4 blocks of the periodic table.
Explain the reasons for these names.
•  Discuss the relationship between group configurations
and group numbers.
•  Describe the locations in the periodic table and the
general properties of the alkali metals, the alkaline-earth
metals, the halogens, and the noble gases.
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•  A row across on the Periodic Table is called a period.
•  There are 7 periods on the Periodic Table.
•  The length of each period is determined by the sublevels
being filled by elements in that period.
Relationship Between Periodicity and
Electron Configurations
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You can tell which period an element is in by looking at its electron configura8on. The number of the highest main energy level is the period number for that element. In which period will you find the following elements? Use the periodic table to check your answer. 1.  C -­‐ 1s22s22p2 2.  K – 1s22s22p63s23p64s1 3.  Ag – 1s22s22p63s23p64s23d104p65s24d9 Based on the electron configura8on of the elements, the periodic table can be divided into four blocks, the s, p, d, and f blocks. The name of each block is determined by whether an s, p, d, or f sublevel is being filled. 6
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The s-Block Elements:
Groups 1 and 2
•  Group 1
– 
– 
– 
– 
– 
Group configuration: ns1 (n= highest main energy level)
Group name: Alkali metals
Soft, silver in color
Extremely reactive metals; most reactive of all the metals
Not found in nature as free elements due to this. They will be in
compounds.
–  They react vigorously with water and most nonmetals.
•  Group 2
– 
– 
– 
– 
Group configuration: ns2
Group name: Alkaline-earth metals
Harder, denser, stronger, than group 1 metals
Very reactive metals; too reactive to be found as free elements
Special Cases: H and He
–  H has group configuration of ns1 but does not share
the properties of Group 1.
–  He has group configuration of ns2 but is part of
Group 18. (This is because helium has a full outer
energy level which gives it the unreactive nature of
a Group 18 element.)
Hydrogen and helium are the simplest and most abundant
elements in the universe. Hydrogen makes up about 76% of the
mass of the universe. Helium makes up about 23% of the mass
of the universe.
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The d-Block Elements:
Groups 3 – 12
•  Groups 3 – 12
–  Block name: transition elements
–  All are metals
–  Transition elements become less
reactive as you move right across
a period
Palladium, platinum, and gold
are among the least reactive
of all elements
•  Group 13
–  Group configuration: ns2np1
–  Group name: Boron family
•  Group 14
–  Group configuration: ns2np2
–  Group name: Carbon family
•  Group 15
–  Group configuration: ns2np3
–  Group name: Nitrogen family
•  Group 16
–  Group configuration: ns2np4
–  Group name: Oxygen family
•  Group 17
–  Group configuration: ns2np5
–  Group name: Halogens
–  Most reactive non-metals
•  Group 18
–  Group configuration: ns2np6
–  Group name: Noble Gases
–  Unreactive gases
The p-Block Elements:
Groups 13 – 18
The p-block elements together with
the s-block elements are often called
main-group elements.
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The f-Block Elements:
Lanthanides and Actinides
•  Lanthanides
–  14 elements filling the 4f sublevel
–  all metals
•  Actinides
–  14 elements filling the 5f sublevel
–  First 4 actinides are found in nature,
the rest are man-made
–  All 14 actinides are radioactive
To save space, the lanthanides
and actinides are usually set
off below the main portion of
the periodic table.
Section 5.3
Electron Configuration and Periodic
Properties
Objectives:
•  Define atomic radii, ionization energy, electron affinity,
and electronegativity.
•  Compare the periodic trends of atomic radii, ionization
energy, and electronegativity, and state the reasons fro
these variations.
•  Define valence electrons, and state how many are
present in atoms of each main-group element.
•  Compare the atomic radii, ionization energies, and
electronegativities of the d-block elements with those of
the main-group elements.
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Copy the following table into your notes:
Period
Trend
Group
Trend
Atomic
Radius
Ionization
Energy
Electron
Affinity
Electronegativity
Atomic Radius
•  The boundaries of an atom are fuzzy, and an atom’s
radius can vary under different conditions.
•  Atomic radius is defined as one-half the distance
between the nuclei of identical atoms that are bonded
together.
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Atomic Radius
•  Atomic radius tends to decrease as you move left to
right across a period.
•  The increasing positive charge of the nucleus holds
the electrons more tightly.
•  Atomic radius tends to increase as you move down
a group.
•  The increasing size of the electron cloud (more
energy levels) causes the atom to be larger.
H
An atom of uranium (U) is only about 3x
larger than an atom of hydrogen even
though U is 238 more massive!!
U
Periodic Trends of
Atomic Radius
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Sample Question:
Of the elements magnesium (Mg), chlorine,
(Cl), sodium (Na), and phosphorus (P), which
has the largest atomic radius? Explain your
answer in terms of trends of the periodic table.
Sodium has the largest atomic radius. All of the
elements are in the third period. Of the four, sodium
has the lowest atomic number and is the first element
in the period. Atomic radius decreases across a
period.
Ionization Energy
•  An ion is an atom (or group of bonded atoms) with a
positive or negative charge.
•  Sodium (Na), for example, easily loses its 3s1
electron to form Na+.
•  Any process that forms an ion is referred to as
ionization.
•  The energy required to remove one electron from a
neutral atom of an element is the ionization energy,
IE (or first ionization energy, IE1). Ionization energy is
measured in kJ/mol.
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Ionization Energy (Main Group Elements)
• Ionization energy generally increases across the
periods.
•  This increase is caused by increasing nuclear charge (more
protons in the nucleus).
•  A higher charge in the nucleus attracts the electrons in the same
energy level more strongly and makes it more difficult to remove
them.
• Ionization energy generally decreases down the
groups.
•  As you move down a group, the electrons are in higher energy
levels, farther from the pull of the protons in the nucleus.
•  The electrons are removed more easily.
Ionization Energy
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Sample Question
Consider two main-group elements, A and B. Element A
has a first ionization energy of 419 kJ/mol. Element B
has a first ionization energy of 1000 kJ/mol.
a)  Which element is more likely to be in the s block?
The p block?
b)  Which element is more likely to form a positive ion?
•  Element A has a very low ionization energy, which means that
atoms of A lose electrons easily. It is most likely to be an s-block
metal because ionization energies increase across the periods.
•  Element B has a very high ionization energy which means that
atoms of B have difficulty losing electrons. Therefore element B
would most likely lie at the end of a period in the
p block.
•  Element A is more likely to form a positive ion because it has
a much lower ionization energy than element B does.
Electron Affinity
•  The energy change that occurs when an electron is
gained by a neutral atom is called the atom’s electron
affinity. (How strongly can an atom attract another
electron?)
•  Electron affinity generally increases across periods.
•  Increasing nuclear charge (more protons) attracts
the electrons in the same sublevel more strongly.
•  Electron affinity generally decreases down groups.
•  The larger an atom’s electron cloud is, the farther
away its outer electrons are from its nucleus.
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Valence Electrons
•  Atoms lose, gain, or share the electrons in order to
become more stable. Chemical compounds are formed
when atoms lose, gain, or share electrons.
•  The electrons that are available to be lost, gained, or
shared are those in the highest energy levels. They are
referred to as valence electrons.
•  Valence electrons are often located in incompletely filled mainenergy levels.
•  example: the electron lost from the 3s sublevel of Na to form Na+
is a valence electron.
•  For main-group elements, the valence electrons are the
electrons in the outermost s and p sublevels. The inner
electrons are in filled energy levels and are held too tightly
to be available for compound formation.
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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.
•  Electronegativity is a measure of the ability of an atom
in a chemical compound to attract electrons from another
atom in the compound.
•  Electronegativity tends to increase across periods,
and decrease or remain about the same down a
group.
Fluorine is the most electronegative element on the
periodic table.
Sample Question:
Of the elements gallium (Ga), bromine (Br), and
calcium (Ca), which has the highest
electronegativity? Explain your answer in terms of
periodic trends.
All of these elements are in the fourth period. Bromine is
farthest to the right in the period. Therefore, bromine
should have the highest electronegativity because
electronegativity increases across the periods.
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