Chapter 5 – The Periodic Law
I. History of the Periodic Table
• A. 60 elements were known by 1860
• 1. Chemists had to memorize the properties
of all these
• 2. No method existed for determining mass or
number of atoms of an element in a
compound at this time
• 3. Different chemists used
• different masses and got
• different compositions
• for the same compound
• B. In September of 1860, the First
International Congress of Chemists was
held in Germany
• 1. It settled the issue of atomic mass
• 2. Cannizzaro presented a method for
accurately measuring the relative mass of
atoms
• 3. A search was begun to
• find relationships between
• atomic mass and other
• properties of elements
Cannizzaro
reaction
• C. Mendeleev and Chemical Periodicity
• 1. Demitri Mendeleev put the name or each
element and all the information he knew
about it on a card
• a. Atomic mass
• b. Physical properties
• c. Chemical properties
• 2. He arranged the cards according to the
properties and looked for trends or patterns
• a. When arranged by atomic mass,
properties repeated at regular intervals
• b. This kind of pattern is called periodic
• 1) days of the week 2) months of the year
• 3. Mendeleev crated a table to show these
patterns, and it became the 1st Periodic Table
of Elements
• a. When arranged by repeating properties,
occasionally the atomic mass did not follow
the correct increasing sequence
• b. Also, he left blank spaces, where he
predicted new elements would
• be discovered, to fill in missing
• properties or gaps in the table
• 1) scandium
• 2) gallium
• 3) germanium
• D. Moseley and the Periodic Law
• 1. Henry Moseley examined the spectra of 38
different metals
• a. He discovered a new pattern
• b. The elements fit into more exact patterns
when arranged by their atomic numbers
(number of protons)
• 2. Developed the definition for atomic #
• a.Followed Mendeleev’s principle for repeating
properties in stating the periodic law
• b.Periodic Law – the physical and chemical
properties of the elements are periodic
• E. The Modern Periodic Table
• 1. The Periodic Table is an arrangement of
the elements in order of their atomic
numbers, so that elements with similar
properties fall in the same column or group
• 2. Noble Gases
• a. discovered later than most other elements
because of their inactivity
• b. placed in a new group after Group VII (17)
Helium
Neon
Argon
Krypton
Xenon
Radon
• 3. The Lanthanides (lanthanoids)
• a. 14 elements from atomic number 58 to
atomic number 71
• b. Chemical and physical properties are so
similar that it is difficult to separate them into
individual elements
• 4. The Actinides (actinoids)
• a. 14 elements from atomic number 90 to
atomic number 103
• b. Are pulled to the bottom of the periodic
table, along with the lanthanides, in order to
save space
• c. Lanthanides and actinides are also known
as the rare earth elements
• 5. Periodicity
• a. All elements in any particular column of
group have similar chemical and physical
properties
• b. The similarities are due to the number of
electrons in the outer level only of the atom
• c. Outer level electrons determine how atoms
behave chemically
II. Electron Configuration & the
Periodic Table
• A. The Noble Gases rarely undergo a
chemical reaction
• 1. This is the result of their electron
configuration
• 2. Their highest energy levels contain stable
octets (except for Helium which has only one
energy level that is completely filled with two
electrons
• 3. An atom’s highest energy level
determines its chemical properties
• B. Periods and Blocks of the Periodic Table
• 1. Elements are arranged vertically on the
periodic table in groups
• 2. Elements are arranged horizontally on the
periodic table in periods
• 3. There are 8
• groups and 7
• periods of
• elements on
• the periodic
• table
• 4. The group number tells how many
electrons are in the highest (or outer) energy
level
• 5. The period number tells how many energy
levels the atom has
• a. The 1st period has only one sublevel (s),
which can hold only 2 electrons; therefore, this
period can have only 2 elements
• b. The 2nd period has two sublevels (s & p,
and p has 3 orientations); therefore, this
period can hold 8 total electrons and 8
elements
• c. Period 3 is the same as Period 2, with 8
elements and 8 electrons
• d. Periods 4 & 5 have 2 electrons in the s
orbital, 10 electrons in the d orbitals, and 6
electrons in the p orbitals, for a total of 18
elements in these periods
• e. Periods 6 & 7 have 2 electrons in the s
orbital, 14 electrons in the f orbitals, 10
electrons in the d orbitals, and 6 electrons in
the p orbitals for a total of 32 elements in
these periods
• C. The s block elements (Groups I & II)
• 1. The elements in Group I are called alkali
metals
• 2. They all have one electron in their outer
level, as well as other similar properties
• a. silvery
• b. soft
francium
• c. extremely reactive (never free in nature)
• d. must be stored in kerosene or oil
lithium
Sodium or potassium
rubidium
cesium
• 3. The elements in Group II are called the
alkaline earth metals
• 4. They all have 2 electrons in their outer level
and have similar properties
• a. harder, stronger, & denser than alkali
radium
metals
• b. have higher melting points than alkali
metals
• c. less reactive than Group I, but still never
found free in nature
beryllium
barium
• D. Hydrogen & Helium
• 1. Hydrogen has 1 electron and Helium has 2
electrons; therefore, each of them has only
one orbital, so they are placed in the s block
• 2. Neither of them, however, are metals like
the other elements in Groups I and II
• E. d block elements (between Groups II – III)
• 1. These are also known as the transition
elements
• 2. They are less reactive than other metals
• F. p block elements (Groups III-VIII)
• 1. These are called the main group elements
along with the s block
• 2. These are mostly nonmetals & metalloids
• 3. The metalloids are boron, silicon,
germanium, arsenic, antimony, tellurium,
polonium, and astatine
• 4. The elements in Group VII are called the
halogens
iodine
fluorine
bromine
astatine
• G. f block elements (lanthanides & actinides)
• 1. Also known as the rare earth elements
• 2. Lanthanides should fit in period 6 between
56 & 72
• 3. Actinides should fit in period 7 between 88
nobelium
& 104
lanthanum
promethium
uranium
actinium
erbium
plutonium
einsteinium
III. Electron Configuration &
Periodic Properties
• A. Atomic Radii
• 1. Ideally, the size of an atom should be from
its nucleus to the outer edge of its highest
orbital, but the outer edge of the orbital is not
exact due to the electron’s fluctuating levels of
energy
• 2. Atomic radius – one half
• the distance between the
• nuclei of identical atoms that
• are bonded together
• B. Period Trends
• 1. There is a gradual decrease in atomic radii
across the 2nd period from lithium to neon
• 2. The trend toward smaller atoms across a
period (from left to right) is caused by the
increasing positive charge of the nucleus
• 3. Electrons (-) are pulled closer to the
nucleus(+) as the nucleus becomes more
positive
• C. Group Trends
• 1. The radii of the atoms of main group
elements increase as you go from the top to
the bottom of a group
• 2. This is due to the addition of one energy
level per atom as you progress down the
group
• 3. This does not always
• hold true, however, for
• those atoms preceded
• by the d block elements
• (Ex.- Al & Ga)
• D. Ionization Energy
• 1. An electron can be removed from an atom
if enough energy is supplied
• atom A + energy → A+ (positive ion A) + e• 2. An ion is an atom or group of bonded
atoms that have a positive or negative charge
• 3. The process that results in the formation of
an ion or ions is called ionization
• 4. The energy required
• to remove one electron
• from a neutral atom is
+
• called ionization energy
• E. Period Trends
• 1. Ionization energy for the main group
elements increases as you move from left to
right across the periodic table
• 2. This is due to a greater number of protons
having a greater attraction for the electrons
• F. Group Trends
• 1. For main group elements, ionization energy
usually decreases as you move down the
group
• 2. This is due to the outer electrons being
further and further from the attraction of the
nucleus as each new energy level is added
• G. Removing Electrons from Positive Ions
• 1. The amount of energy required to remove
the first electron from an atom is referred to
as the first ionization energy (IE1)
• 2. IE1 will be less than IE2 or IE3, etc.
because as each electron is removed, the
attractive force of the nucleus becomes
greater and greater in relationship to the
electrons
• 3. Therefore, IE2 is
• greater than IE1 and
• IE3 is greater than
• IE2 and so on
• H. Electron Affinity
• 1. The energy change that occurs when an
electron is acquired by a neutral atom is called
electron affinity
• 2. Most atoms release energy when they
acquire an electron (A + e- → A- + energy)
• 3. Some atoms must be “forced” to take on
another electron, and energy
• must be added in order to
• make this happen
• (A + e- + energy → A-)
• This type of ion will be
• unstable
• I. Trends
• 1. In general electrons are gained more
easily as you go from left to right on the table
(across a period)
• 2. An exception is the noble gases, which
require huge amounts of energy to add
electrons, since they already have a full octet
• 3. Usually it is more difficult to add electrons
as you move down a group, because as you
move down the group the electrons are getting
further from the attractive force of the nucleus
• 4. It is also more difficult to add a second
electron to ion that is already negative
• J. Ionic Radii
• 1. A positive ion is a cation, and its radius
will be smaller than its atom’s radius
• 2. A negative ion is an anion, and its radius
will be larger than its atom’s radius
• 3. Atomic radii tend
• to decrease across
• a period, and
• increase down a
• group
• K. The electrons available to be gained, lost,
or shared are called valence electrons
• L. A measure of the ability of an atom to
attract electrons is called electronegativity
• 1. Electronegativity tends to increase across
a period
• 2. Electronegativity tends to decrease down a
group
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LAB – Flame Test
1. How can a flame test be used as a confirmatory test for an
unknown metal?
2. What is one purpose for which these metal ions are used?
3. What kind of ions color the flame when glass is melted?
4. Why do ions produce colors when placed in a flame?
5. What condition is the atom in before it is placed in the
flame?
6. After placing the atom in the flame, what condition is it then
in?
7. What condition is the atom in when a color is produced in
the flame?
8. When atoms produce colors in a flame, what series in the
electromagnetic spectrum is this called?
9. How was the color of the flame affected when viewed
through the cobalt glass?
10. Imagine and describe a scenario in forensics in which a
flame test could be used for identification of something.