1. Periodic Table and Periodicity
• Modern periodic table, periodic law
• d-block elements
• Periodic properties - atomic and ionic radii,
ionization energy, electronegativity and
electron affinity
2. The periodic table is tool that chemists
have developed to organize and
categorize the elements found on Earth.
3. The History of the
Modern Periodic Table
1700 - about 30 elements known to man
4. During the 19th century, chemists
began to categorize the elements
according to similarities in their
physical and chemical properties.
The end result of these studies
was our modern periodic table.
5. Johann Dobereiner
1780 - 1849
Model of triads
In 1829, he classified some elements into
groups of three, which he called triads.
The elements in a triad had similar chemical
properties and orderly physical properties.
(ex. Cl, Br, I and
Ca, Sr, Ba)
6. John Newlands
1838 - 1898
Law of Octaves
In 1863, he suggested that elements be
arranged in “octaves” because he noticed
(after arranging the elements in order of
increasing atomic mass) that certain
properties repeated every 8th element.
7. John Newlands
1838 - 1898 Law of Octaves
Newlands' claim to see a repeating pattern was met with
savage ridicule on its announcement. His classification
of the elements, he was told, was as arbitrary as putting
them in alphabetical order and his paper was rejected
for publication by the Chemical Society.
8. John Newlands
1838 - 1898 Law of Octaves
His law of octaves failed beyond the
element calcium. WHY?
Would his law of octaves work today with
the first 20 elements?
9. Dmitri Mendeleev
1834 - 1907
In 1869 he published a table of
the elements organized by
increasing atomic mass.
10. Lothar Meyer
1830 - 1895
At the same time, he published his own
table of the elements organized by
increasing atomic mass.
12. • Both Mendeleev and Meyer arranged the
elements in order of increasing atomic
mass.
• Both left vacant spaces where unknown
elements should fit.
So why is Mendeleev called the “father of
the modern periodic table” and not Meyer,
or both?
13. • stated that if the atomic weight of an
element caused it to be placed in the
wrong group, then the weight must be
wrong. (He corrected the atomic masses
of Be, In, and U)
• was so confident in his table that he used
it to predict the physical properties of three
elements that were yet unknown.
Mendeleev...
14. After the discovery of these unknown
elements between 1874 and 1885, and the
fact that Mendeleev’s predictions for Sc, Ga,
and Ge were amazingly close to the actual
values, his table was generally accepted.
15. However, in spite of Mendeleev’s great
achievement, problems arose when new
elements were discovered and more
accurate atomic weights determined. By
looking at our modern periodic table, can
you identify what problems might have
caused chemists a headache?
Ar and K
Co and Ni
Te (Tellurium) and I
Th (Thorium) and Pa (Protactinium)
16. Henry Moseley
1887 - 1915
In 1913, through his work with X-rays, he
determined the actual nuclear charge
(atomic number) of the elements*. He
rearranged the elements in order of
increasing atomic number.
*“There is in the atom a fundamental
quantity which increases by regular steps
as we pass from each element to the
next. This quantity can only be the
charge on the central positive nucleus.”
17. Henry Moseley
His research was halted when the British
government sent him to serve as a foot soldier
in WWI. He was killed in the fighting in
Gallipoli by a sniper’s bullet, at the age of 28.
Because of this loss, the British government
later restricted its scientists to noncombatant
duties during WWII.
18. Glenn T. Seaborg
After co-discovering 10 new elements, in
1944 he moved 14 elements out of the main
body of the periodic table to their current
location below the Lanthanide series.
These became known
as the Actinide series.
1912 - 1999
19. Glenn T. Seaborg
He is the only person to have an element
named after him while still alive.
1912 - 1999
"This is the greatest honor ever bestowed
upon me - even better, I think, than
winning the Nobel Prize."
Seaborgium
22. The vertical columns of the periodic table
are called GROUPS, or FAMILIES.
The elements in any group
of the periodic table have
similar physical and
chemical properties!
23. Periodic Law
When elements are arranged in order of
increasing atomic number, there is a
periodic pattern in their physical and
chemical properties.
37. Group 1 - The Alkali Metals
A. Group one is the first column of the periodic
table on the left. The alkali metals group consists
of the elements lithium, sodium, potassium,
rubidium, cesium, and francium.
NaCl
38. Group 1 - The Alkali Metals
B. The alkali metals are very reactive
metals that do not occur freely in nature.
C. These metals have only one electron
in their outer shell. Therefore, they are
ready to lose that one electron in ionic
bonding with other elements.
39. Group 1 - The Alkali Metals
D. All of the group 1 elements are silvery-
colored metals. They are soft, and can
be cut easily with a knife.
E. The alkali metals are malleable and
good conductors of heat and electricity.
They have low densities - Li, Na, and K
are less dense than water.
40. Group 1 - The Alkali Metals
F. Alkali metals can explode if they
are exposed to water.
42. Group 2 -The Alkaline Earth
Metals
A. The alkaline earth metals
are elements found in
the second group of the
periodic table. The
alkaline earth metals
consist of the elements
beryllium, magnesium,
calcium, strontium,
barium, and radium.
Yum! Calcium!
43. Group 2 -The Alkaline Earth
Metals
B. All alkaline earth metals have 2
electrons in their outer energy shell, so
they are very reactive - although not as
reactive as the alkali metals.
C. Because of their high reactivity,
alkaline earth metals are not found
freely in nature.
44. Group 2 - The Alkaline Earth Metals
D. All alkaline
earth metals are
shiny, silvery
white metals. If
heated in a
flame, each of
the metals will
give a different
colored flame.
46. Groups 3 –12 : Transition Metals
A. The 38 elements in groups 3 through
12 of the periodic table are called
“transition metals” or d-Block
elements.
B. The transition metals have different
numbers of valance electrons in their
outer energy level.
47. Groups 3-12: Transition Metals
C. The transition metals have an additional
energy level, or energy shell, that fills with
electrons. Transition metals have valance
electrons in their outer shell and in their d-
shell.
D. All transition elements are malleable, and
conduct electricity and heat. All transition
metals except copper and gold have a silvery
appearance.
48. Groups 3-12: Transition Metals
E. The transition metals iron, cobalt, and
nickel are the only elements known to
produce a magnetic field.
49. Groups 13 –15 : Other Metals
A. The seven elements
classified as other metals
are located in groups 13,
14, and 15. The other
metals include aluminum,
gallium, indium, tin,
thallium, lead, and
bismuth.
Al Recycling
50. Groups 13-15: Other Metals
B. While the other metals have some
similarities to transition metals, they are
not the same. The other metals have
valance electrons only in their outer
shell.
C. All other metals are solid, have a
relatively high density, and are opaque.
51. Stair-Step Elements: Metalloids
A. Metalloids are the elements found
along the stair-step line that
distinguishes the metals from non-
metals. The metalloid elements are
boron, silicon, germanium, arsenic,
antimony, tellurium, polonium.
52. Stair- Step Elements: Metalloids
B. Metalloids have properties of both
metals and non-metals.
C. Some metalloids, such as silicon and
germanium, are semi-conductors. This
means that they can carry an electrical
charge under special conditions. This
property makes metalloids useful in
computers and calculators.
53. Groups 14-16: Non-Metals
A. Non-metals are the
elements in groups 14-
16 of the periodic table.
The non-metals include
hydrogen, carbon,
nitrogen, oxygen,
phosphorus, sulfur, and
selenium.
Carbon Burning
54. Groups 14-16: Non-Metals
B. non-metals are not able to conduct
electricity or heat very well. As opposed
to metals, non-metals are very brittle,
and cannot be rolled into wires or
pounded into sheets.
C. The non-metals exist as gasses or
solids at room temperature.
55. Groups 14-16: Non-Metals
D. The non-metals have no metallic luster
(they’re not shiny), and do not reflect light.
Oxygen tank pressure gauge
57. Group 17 - The Halogens
A. The halogens all have 7 electrons in
their outer-most energy level. The
halogens are very reactive, and are
found in nature as diatomic molecules,
which means that each atom is bonded
to another atom. So halogens are
found in groups of two atoms bonded
together.
58. Group 17 - The Halogens
B. The halogen group includes fluorine,
chlorine, bromine, iodine, and astatine.
Chlorinated Pool
59. Group 17 - The Halogens
C. The physical appearances of the halogens
are very different between elements.
D. Fluorine is a poisonous pale yellow gas,
chlorine is a poisonous pale green gas,
bromine is a toxic and caustic brown liquid,
and iodine is a shiny black solid which forms
a purple vapor upon heating.
61. Group 18 - The Noble Gasses
A. The noble gasses all have full outer
electron shells. That means that noble
gasses will not react with any other
elements.
B. The noble gasses include helium, neon,
argon, krypton, xenon, and radon.
62. Group 18 - The Noble Gasses
C. All elements in group 18 are found at
room temperature as gasses, and they
have very low boiling points.
Reno, NV Neon
64. The Rare Earth Elements
A. All of the rare earth elements are found in
group 3 and periods 6 and 7 of the periodic
table. The rare earth elements, or
lanthanide and actinide series, are mostly
man-made. One of the lanthanides (period
6) and ALL of the actinides (period 7) are
man-made.
65. The Group “A” elements
are called
REPRESENTATIVE ELEMENTS.
66. Atomic Radius: The distance from the centre of the nucleus to the
outermost shell of an atom is called the atomic radius of that atom.
Periodic Properties
67. Two factors must be taken into consideration in explaining this
periodic trend:
(1) Increasing nuclear charge
(2) Increasing shell
• Left to right along a period
69. Ionic radii of common ions
• Increase down a group
• Decrease along a period
70. • Ionization potential (or ionization energy) is the amount of energy
required to remove one or more electrons from the outermost shell
of an isolated atom in the gaseous state.
Atom(g) + IE Positive ion(g) + electron(g)
• Ionization energy is also called as ionization potential because it is
measured as the minimum potential required to remove the most
loosely held electron from the rest of the atom. It is measured in the
units of electron volts (eV) per atom or kilo joules per mole of atoms
(kJ mol-1)
• Thus, the ionization energy gives the ease with which the electron
can be removed from an atom. The smaller the value of the
ionization energy, the easier it is to remove the electron from the
atom.
71. • The minimum energy required to remove the most loosely bound
electron from an isolated gaseous atom is called ionisation energy. It is
also called the first ionisation energy (IE1).
M (g) + IE1 M+
(g) + e-
• The minimum energy required to remove another electron affinity
from the uni positive ion is called second ionisation energy (IE2).
M (g) + IE2 M2+
(g) + e-
• The second ionisation energy (IE2) is greater than the first ionisation
energy
72. Electron affinity is the amount of energy released when an electron is
added to an isolated gaseous atom.
• Electron affinity is the ability of an atom to hold an additional
electron. If the atom has more tendency to accept an electron then the
energy released will be large and consequently the electron affinity will
be high. Electron affinities can be positive or negative. It is taken as
positive when an electron is added to an atom. It is expressed as
electron volts per atom (eV per atom) or kilo joules per mole.
• Electron affinity depends on:
i) Extent of nuclear charge
ii) Size of the atom
iii) Electronic configuration.
73. Electron Affinity of the Halogens
From chlorine to iodine, which ionize by accepting one electron there
is a decrease in the electron affinity or the energy released. The lower
electron affinity of fluorine when compared to chlorine is not fully
understood.
74. Electronegativity is the tendency of an atom to attract
electrons towards itself in a molecule of a compound. The
value of electronegativity of an element describes the
ability of its atom to compete for electrons with the other
atom to which it is bonded. Electronegativity is however
not the property of an isolated atom.
75. Types of Electronegativity
• When the molecule is formed by transfer of electrons (ionic bonding)
the transfer takes place from electropositive atom to electronegative
atom. In the example below, Na is electropositive and Cl is
electronegative.
76. • If the molecule is formed by sharing of electrons (covalent bond)
the bonded pair of electrons shift towards more electronegative atom
resulting in the formation of polar molecule. In the example below,
chlorine atom is more electronegative as compared to hydrogen atom,
resulting in a covalent bond where the shared pair of electron shifts
towards the more electronegative atom. This results in polar
molecules.
The electron pair is more closer to the chlorine atom and so the
molecule gets polarized i.e., the chlorine atom gets a negative charge
while the hydrogen atom gets a positive charge.
