Chapter 6.1 Notes
Developing the Periodic Table
• Scientists studied elements and found that there are groups of elements that have similar
chemical and physical properties.
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Is this a coincidence? No, it is not. It turns out that elements that have the same number of
electrons in their outer energy levels have similar properties. So knowing the electron
configuration of an element will help us predict its properties.
Early Attempts at a Periodic Table
• Beginning in 1817, Johann Dobereiner (German chemist) grouped elements of similar
properties into groups of three. These were called triads.
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In 1863, John Newlands (English chemist) arranged the elements in order of increasing atomic
mass and noted that there appeared to be a repetition of similar properties every eighth
element. This arrangement was called the law of octaves.
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Six years later, Dmitri Mendeleev (Russian chemist) proposed a similar ordering by increasing
atomic mass but allowed the number of elements in a horizontal row (period) to vary. He
noticed blank spots in his table and accounted them to elements not yet discovered. He was
able to predict the properties and atomic masses of several elements that were unknown at the
time. His predictions have been found to be very nearly correct.
Modern Periodic Table
• There was a problem with Mendeleev’s table of elements. Certain elements seemed to be in
the wrong columns. Switching their positions put them in the columns where they belonged,
according to their properties. He thought perhaps the atomic masses of these elements had
been poorly measured.
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In 1913, a man named Henry Moseley found the reason for these apparent exceptions. He
found that the order was due to increasing atomic number and not because of increasing
atomic mass. This is described by the periodic law which states the properties of the elements
are a periodic function of their atomic numbers.
Let’s take a look at the periodic table now.
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How many rows are there? These are called periods.
How many columns are there? These are called groups.
How is the table organized? Metals, nonmetals and groups/families.
What is the periodic table composed of mostly, metals or nonmetals?
Properties of Metals and Nonmetals
Metals:
• Hard and shiny solids
• Conduct heat and electricity
• Malleable, easy to bend and stretch
• Tend to lose e- when forming chemical bonds
Nonmetals:
• Usually gases
• Dull and brittle
• Do not conduct heat and electricity
• Tend to gain (or share) e- when forming chemical bonds
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General Rule of Thumb: Elements with three or fewer electrons in the outer level are
considered to be metals. Elements with five or more electrons in the outer level are considered
to be nonmetals.
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Elements that have properties of both metals and nonmetals are called metalloids (sometimes
also called semi-metals). There are 8 of them. Label them on your Periodic Table.
Chapter 6.2 Notes – Classification of the Elements
• Valence Electrons:
o Elements within a specific group of the periodic table have the same number of valence
electrons.
o Example: group 14 elements all have 4 valence electrons.
o Remember that the outer most electrons (valence e-) determine an elements chemical
and physical properties. These are the electrons that fill the s and p orbitals (8 max).
o When an atom has 8 e- in its outer level that will render it unreactive. This is called the
octet rule.
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The s-, p-, d- and f- Block Elements: Now elements with similar electron configurations can be
placed in the same column (group) of the periodic table.
o Elements that end in an s-orbital fill groups 1 and 2.
o Elements that end in a p-orbital fill groups 13-18 (with the exception of He)
o Elements that end in a d-orbital fill groups 3-12.
o Lastly, elements that end in a f-orbital fill the lanthanoid and actinoid series at the
bottom of the periodic table.
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Practice Problems: pg 162 #’s 7-13
Chapter 6.3 Notes – Periodic Trends
• Atomic Radius – the size of an atom depends on the electron cloud surrounding the nucleus but does not
have a clearly defined edge. So the atomic size is defined by how closely an atom lies to a neighboring
atom. The atomic radius is defined as ½ the distance between adjacent nuclei that are bonded together.
The general trend is that atomic radius increases from left to right and top to bottom in the periodic
table.
o See Fig 6-11 on pg 163.
• Ionic Radius – this is the radius of an atom after it has gained or lost one or more electrons. Remember
that the #P do not equal #e- for an ion.
o Cations – lose electrons and become smaller then the parent atom.
o Anions – gain electrons and become larger than the parent atom.
o See Fig 6-13 and Fig 6-14 on pg 166.
• Ionization Energy – this is the amount of energy required to remove an electron from an atom. For
metals, this is relatively easy to do since they want to form positive ions thus a small IE. For nonmetals
that want to gain more electrons, not lose electrons, this is harder to do thus a higher IE. The general
trend is that ionization energy increases from bottom to top and from left to right in the periodic table.
o The octet rule states that atoms tend to gain, lose or share electrons in order to acquire a full set
of 8 valence electrons.
o See Fig 6-16 on pg 167.
• Electronegativity – an elements relative ability of it’s atoms to attract electrons in a chemical bond.
Metals normally have a low EN while nonmetals have a high EN. The general trend is that EN increases
from bottom to top and from left to right in the periodic table.
o See Fig 6-18 on pg 169.
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Practice Problems: pg 165 #’s 16-18, pg 169 #’s 19-22
Misc Info - Relative Stability
• Rule of Thumb: An atom having a filled or half filled sublevel is slightly more stable (less
reactive) than an atom without a filled or half filled sublevel.
Examples:
3d6 (↑↓) (↑ ) (↑ ) (↑ ) (↑ )
3d5 (↑ ) (↑ ) (↑ ) (↑ ) (↑ )
4p3 (↑ ) (↑ ) (↑ )
4p6 (↑↓) (↑↓) (↑↓)
-less stable
-more stable
-less stable
-more stable
Also note exceptions to electron configs for elements Cr and Cu