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CH – 5 THE PERIODIC LAW
I. HISTORY OF THE PERIODIC TABLE
A. Antoine Lavoisier (1790s)
1. Compiled list of 23 known elements
B. Johann Dobereiner (1817)
1. Found that several groups of three elements have
similar properties and called it triads.
2. Second element has an atomic mass halfway
between 1st and 3rd elements
EXAMPLES OF DOBEREINER'S TRIADS
TRIAD 1
Name
TRIAD 2
TRIAD 3
Atomic
Atomic
Atomic
Name
Name
Mass
Mass
Mass
First Element
Calcium
40.1 Chlorine
35.5 Sulfur
Third Element
Barium
137.3 Iodine
126.9 Tellurium 127.6
Average
88.7 Average 81.2 Average 79.9
Second Element
Strontium 87.6 Bromine
79.9 Selenium
32.1
79.0
C. Stanislao Cannizzaro (1860) – presented an
accurate method for measuring relative masses of
atoms at the First International Congress of Chemists in
Karlsruhe, Germany
1. This initiated a search for relationships between
atomic mass and properties of elements
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D. Emile Beguyer de Chancourtois – Telluric Screw
1. Positioned the elements according to increasing
atomic weight along a spiral inscribed on the surface
of a cylinder and inclined at 45° from the base.
E. John Newlands (1863-1864)
1. Arranged the elements in order of their increasing
atomic masses
2. Noted that there is repetition of similar properties
every eighth element
3. Placed seven elements in each group, and called it
octaves
4. Law of Octaves – the same properties are repeated
every eighth element
5. Did not work for all elements
NEWLAND'S LAW OF OCTAVES
1
Li
Na
K
2
Be
Mg
3
B
Al
4
C
Si
5
N
P
6
O
S
7
F
Cl
F.Dmitri Mendeleev (1869)
1. Listed 60+ elements in several vertical columns in
order of their increasing atomic mass
2. Noticed a regular recurrence of their physical and
chemical properties
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3. Left blank spaces in the table because there were no
known elements with the appropriate properties at
that time
4.1871 – Predicted the physical and chemical properties
of 3 missing elements
5. Credited with the Periodic Law
G. Lothar Meyer (1870)
1. German chemist, working on element organization at
the same time as Mendeleev; published his work a
year later
H. William Ramsey (1868-1900)
1. Discovery of noble gases
2. Proposed a new group in the periodic table to
accommodate the noble gases discovered
3. Noble gases were hard to detect due to lack of
chemical reactivity
I. Henry Moseley and the Periodic Law
1. Mendeleev had a few elements that did not fit
2. Determined the nuclear charge and atomic number
of elements by analyzing the spectra of 38 metals
3. Rearranged the elements in the table in the order of
their increasing atomic numbers
J. Lanthanides & Actinides
1. Lanthanides were separated and identified in the
early 1900s
a. Hard to isolate because they have very similar
physical and chemical properties
2. Actinides series elements were discovered
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K. Harry D. Hubbard (1924)
1. Modernized Mendeleev’s periodic table
L. Glenn Seaborg (1940)
1. Discovered 10 Transuranium elements
2. Reorganized the table in 1944 by placing the
lanthanide/actinide series at the bottom of the table
M. Modern Periodic Table
1. Periodic Law: The physical and chemical
properties of the elements are periodic functions of
their atomic numbers
2. Periodic table – arrangement of the elements in
order of their atomic numbers so that elements with
similar properties fall in the same column or group
3. Period or Series – horizontal rows of the periodic
table (seven periods)
a. Denote number of principal energy levels
4. Groups or Families – vertical groups
a. Denote # of valence electrons
5. Noble Gases, Lanthanides, Actinides
II. ELECTRON CONFIGURATION AND THE PERIODIC TABLE
A. Sublevel Blocks of the Periodic Table
1. a. Main-group elements – s- and p-block elements
b. Representative elements – Group A elements
1A – 7A groups; s and p sublevels are partially filled
2. s-block – outermost electrons are in the s sublevel
a. Group 1 – Alkali Metals
b. Group 2 – Alkaline Earth Metals
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c. Hydrogen and Helium
2. p-block – Groups 13-18
a. Electrons add to a p level only after the s sublevel
in the same energy level is filled
b. Boron, carbon, nitrogen, oxygen & noble gas
families
c. Halogens (group 17) – salt formers; very reactive
nonmetals
d. Metalloids – brittle solids with some properties of
both metals and nonmetals
e. Noble Gases – inert due to filled s and p orbitals
3. d-block – Groups 3-12 (Group B elements)
a. Transition elements – d-block elements with
typical metallic properties
b. Form colored ions
4. f-block – lanthanides and actinides
a. Lanthanides – 4f sublevel is being filled
b. Actinides – 5f sublevel is being filled
III. ELECTRON CONFIGURATION AND
PERIODIC PROPERTIES
A. Atomic Forces
1. Nuclear Pull – theoretical force equal to Z (atomic #)
2. Electron Repulsion – like charges repel
3. Shielding
a. Inner shells of electrons “shield” the outer shell
electrons from the nucleus” pull
4. Z-effective
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a. Outer electrons feel less than full strength of
nucleus because of electrons between them
B. Shielding Effect – decrease in the attraction between
outer electrons and the nucleus due to the presence of
other electrons between them
1. Period Trend – Constant within a period because
the main energy level remains the same
2. Group Trend – Increases from top to bottom
because of the increase in the main energy level
C. Atomic Radius – defined as one-half the distance
between the nuclei of identical atoms that are bonded
together
1. Period Trends – Decreasing trend across a period.
a. Size become smaller across a period due to
increase in the number of protons in the nucleus
but shielding effect is constant for the electrons
being added to same energy level
2. Group trends – Atomic radii of the main-group
elements increase down a group
a. This is due to increase in the number of main
energy levels, thus decreasing the attraction
between the nucleus and outermost electrons
D. DIATOMIC ELEMENTS – always exist as pairs
Br2, I2, N2, Cl2, H2, O2, F2
E. Ionization Energy – energy required to remove one
electron from a neutral atom of an element; describes
how tightly an electron is held by atom
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1. Ion – atom or group of bonded atoms that has a
positive or negative charge
2. Ionization – process that result in the formation of
an ion
3. First Ionization Energy (IE1) – energy needed to
remove the most loosely held electrons from an atom
Ca + 590 kJ -- Ca+1 + e4. Second Ionization Enrgy (IE2) – amount of energy
needed to remove the second electron
Ca+1 + 1145 kJ -- Ca+2 + ea. IE2 is always greater than IE1 ; More difficult to
remove e- from +ion than a neutral atom
5. Factors Affecting Ionization Energy
a. Nuclear charge – the larger the nuclear charge,
the greater the ionization energy
b.Shielding effect – the greater the shielding effect,
the less the ionization energy
c. Radius – the greater the distance between the
nucleus and outer electrons, the less the ionization
energy
d. Sublevel – an electron from a full or half-full
sublevel requires additional energy to be removed
6. a. IE helps predict whether an element is likely to
form an ionic or covalent compound
b. Metals tend to have low IE and form cations, while
nonmetals tend to have high IE and form anions –
thus causing ionic bonds
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c. Elements with intermediate IE form molecular or
covalent compounds by sharing electrons
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7. Period trends – ionization energies of main group
elements increase across each period
a. Slight increase between IIA and IIIA due to p-sublevels
having a higher energy so electrons are easier to
remove
b. Slight decrease between VA and VIA due to paired
electrons in the p-sublevel. Paired e- have greater
repulsive forces so are easier to remove.
c. Nuclear charge increases within a period, which more
strongly attracts additional electrons in the same energy
level
8. Group trends – among the main group elements,
ionization energies generally decrease down the groups.
a. The size of the atom increases as we go down; thus the
outermost electron is farther from the nucleus with an
increased shielding effect
b. It should be more easily removed and therefore have
lower ionization energy
F. Electron Affinity – energy change that occurs when an
electron is acquired by a neutral atom; atom’s attraction for
additional electron
1. Periodic trends – electron affinity generally increases as
we move from left to right across a period because atoms
become smaller and the nuclear charge increases
a. Increasingly negative left to right across a period
2. Group trends – electron affinity decreases as we move
down a group because of the increasing atomic size, which
decreases attraction between the nucleus and the
outermost electrons
a. Increasingly positive down a group so electrons add
with greater difficulty
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3. Metals have low EA; Nonmetals have high EA
G. Ionic Radii
1. a. All atoms want to have a noble gas configuration
b. Achieve maximum stability with lowest energy
configuration
c. Gain/lose/share electrons to make it happen
2. Cation – positive ion; smaller than neutral atom because
of the loss of outer shell electrons
3. Anion – negative ion
a. Always larger than neutral atoms because the electrons
are not drawn as strongly as they were before the
addition of extra electrons
b. More repulsion among electrons
4. Period Trends – left to right across a period, there is a
decrease in the size of the positive ions
a. From Group 15, negative ions (which are much larger in
size) gradually decrease in size as you move right
b. For d-block and f-block elements, first electrons to be
removed are s-electrons before d- or f- electrons,
resulting in formation of different kinds of ions for the
same element
5. Group Trends – there is an increase in ionic radii with
both anions and cations as you go down each group
a. Outer electrons in both cations and anions are in higher
energy levels from top to bottom of the periodic table,
so there is also a gradual increase of ionic radii
H. Valence Electrons – electrons in the highest energy level of
the atom, which are available to be lost, gained, or shared in
the formation of chemical compounds
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I. Oxidation Numbers – indicates the number of
electrons that would be lost or gained by the atom
1. Metals are found on the left and center of periodic
table
a. Tend to lose electrons, thus they have positive
oxidation numbers
2. Nonmetals are on the right side of the periodic table
a. Tend to gain electrons, thus they have negative
oxidation numbers
J. Electronegativity – a measure of the ability of an
atom in a chemical compound to attract electrons
1. Electronegativity values help predict the type of
bonding that can exist between atoms in compounds.
2. Period Trends – Increases from left to right due to
increase in nuclear pull
a. Metallic elements at the left have very low
electronegativities (Fr & Cs most reactive metals)
b. Non-metallic elements at the right have high
electronegativities (F – most reactive nonmetal)
3. Group Trends – From top to bottom within a group,
electronegativity generally decreases.
K. Oxides
1. Groups 1 & 2 oxides are generally basic
2. Nonmetallic oxides are generally acidic
L. Periodic Properties of d- and f- Block Elements
1. Atomic Radii
2. Ionization Energy
3. Ion Formation and Ionic Radii
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4. Electronegativity
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