This document discusses the periodic table and periodic trends of elements. It explains how early scientists like Deboreiner, Newlands, and Mendeleev contributed to the development of the periodic table by arranging elements by atomic mass. The periodic table arranges elements by atomic number and shows trends in properties across periods and down groups, including decreasing atomic radius and ionization energy but increasing atomic radius. Metallic character decreases across periods but increases down groups.

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4. Metal an atom with 1-3 extra
valence electrons.
Shiny
Dense
Malleable
Ductile
Electrical conductors
Thermal conductors
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5.  In 1817,German scientist named Deboreiner realized
that there's is a connection between relative atomic
mass and chemical properties
 In 1829,he arranged many elements into group of
three, called Deboreiner’s triads in which the relative
mass of the middle element about the average of
masses of the other two. However only few elements
form the group of three.
 In 1864, the British scientist called Newlands,
Arranged elements in order of relative atomic mass
and gave each one number, Hydrogen the lightest
element was no. 1, lithium 2, helium was not known at
that time and so on.
 He was the first scientist to list elements in numerical
order. BUT his law didn’t apply after 17 elements
(Ca).
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6.  In 1869 Mendeleef, a Russian scientist, stated that the
properties of element are a periodic in function of their
atomic masses.
 This periodic law means that if the elements are arranged
in order of increasing their relative mass, element of similar
properties appear in regular intervals . This kind of
repetition is periodic
 Mendeleef ‘s periodic table was in complete because many
elements (e.g. noble gases were not known that time.
 In 1913 an English research student called Moseley. He
found that atomic number is more important than the
relative atomic mass and that there is no exceptions in
periodic table
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7. Periodic Table of an element
 Is the arrangements of elements in order of their proton numbers,
Horizontal raw are period and the vertical column are groups
Characteristics of Periodic Table
 There are 8 groups and 7 periods
 Period 1 contain only Hydrogen and Helium. Hydrogen is placed both in
group one and group 7 because its properties resembles of those of
alkali metals and halogens
 Period 2 and 3 (Short periods) contains 8 element each
 Period 4 and 5 ( Long periods) each contains 18 elements. 8 elements
of these correspond to the 8 element in the short period. The other 10
are called transition elements . Which show similar properties before
them and also some properties similar to those following them.
 Period 6 and 7 contain 32 and 17 elements
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8. Periodicity of properties of the element.
 Physical properties.
 Physical properties of elements in the same group vary gradually and regularly.
 All group one alkali metals have low densities, low melting point and high
conductivities.
 In group 7 halogens, The melting points, boiling points and the densities rise
regularly. The first member of the group e.g lithium and fluorine is usually
slightly different form the rest of the group)
Chemical properties.
 Vary gradually for the same group of elements.
 Of the alkali metals potassium is more electropositive than sodium which is more
electropositive than lithium.
 Of halogens fluorine is most electronegative metals and iodide is least.
 Metals are on the left hand side of the table and non metals are on the right.
 In group 1 metals are reactive as their relative atomic mass increases.
 In group 7 the halogen are less reactive when their masses increase.
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9. Families of elements.
 The periodic table is accounted for the regular
arrangement of electrons in atoms of elements.
The periodicity of electron configuration leads to
periodicity of chemical properties which depends
on configuration of the outermost orbit.
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10. Noble gasses configuration.
 The outer most shell of electron with two electrons
in case of helium and eight for other elements
 Are very stable and the electrons are not used in
chemical reactions.
 The noble gasses are very un reactive. Noble gas
elements are Ne (2:8) Ar (2:8:8)
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12. GROUP 1 & 2 METALS.
 An alkali metal atoms
◦ Has one electron more than noble gas structure. The electron is readity detached and the element
are therefore very electropositive . These include: Li (2:1) Na,(2:8:1) K, (2:8:8:1)
 Alkali earth metals
◦ Have two electrons more than noble gas these can readity be detached and the metals are
electropositive. Examples of alkal earth metals are Be (2:2), Mg (2:8:2), Ca, (2:8:8:2)
 They show relative weak metalic bonding because they have only one metalic bonding.
 They are soft, can be cut with knife.
 BP and MP are low.
 Low standard heat enthalples
 Low densities.
 Group 2 with two valence electron they show stronger metalic bonding.
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13. periodic trends
 are the tendencies of certain elemental
characteristics to increase or decrease as one
progresses along a row or column of the periodic
table of elements.
 Elemental characteristics in periodic table are
 Atomic radius
 Electronegativity
 Ionization energy
 Metallic characters
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14. The atomic radius
 is the distance from the atomic nucleus to the outermost
stable electron orbital in an atom that is at equilibrium.
 The atomic radius tends to decrease as one progresses
across a period because the effective nuclear charge
increases, thereby attracting the orbiting electrons and
lessening the radius.
 The atomic radius usually increases while going down a
group due to the addition of a new energy level (shell).
However, diagonally, the number of protons has a larger
effect than the sizeable radius. For example, lithium (145
pm) has a smaller atomic radius than magnesium (150
pm).
 Atomic radii decrease left to right across a
period, and increase top to bottom down a group.
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15. • Atomic Radius
• size of atom
• Ionization Energy © 1998 LOGAL
• Energy required to
remove one e- from a
neutral atom.
(gaseous state)
© 1998 LOGAL
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16. 250
200
150
100
50
0
0 5 10 15 20
Atomic Number
Atomic Radius (pm)
Li
Ar
Ne
K
Na
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17. 1
2
3
4
5
6
7
• Increases going down and Decreases going
right.
Decreases!
I
ncr
eas
es
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18. • Why larger going down?
• Higher energy levels have larger orbitals
• Shielding - core e- block the attraction between the
nucleus and the valence e-
• Why smaller to the right?
• Increased nuclear charge without additional
shielding pulls e- in tighter
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19. • Ionic Radius
• Cations (+)
• lose e-
• smaller
© 2002 Prentice-Hall, Inc.
• Anions (–)
• gain e-
• larger
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20. Effective nuclear charge concept
 The shielding effect that electrons close to the nucleus
have an outer shell of an electrons in many electrons
atoms.
 The presence of shielding electrons reduces the
electrostatics attraction between the positively charged
protons in the nucleus and the outer electrons.
 Moreover, the repulsive forces between the electrons in
many electrons atom further offset the attractive force
exerted by the nucleus the concept of effective nuclear
charge
 Effective nuclear charge increases from left to right across
the period and from the bottom to top in a group of
representative elements
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21. Ionization energy
 is the minimum energy required to remove one electron from each
atom in a mole of atoms in the gaseous state.
 Trend-wise, ionization potentials tend to increase while one
progresses across a period because the greater number of protons
(higher nuclear charge) attract the orbiting electrons more strongly,
thereby increasing the energy required to remove one of the
electrons.
 As one progresses down a group on the periodic table, the ionization
energy will likely decrease, since the valence electrons lie are farther
away from the nucleus and experience a weaker attraction to the
nucleus' positive charge.
 There will be an increase of ionization energy from left to
right of a given period and a decrease from top to bottom.
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22. 2500
2000
1500
1000
500
0
Ar
Ne
He
0 5 10 15 20
Atomic Number
1st Ionization Energy (kJ)
Li Na K
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23. • Decreases going down and Increases going
right.
1
2
3
4
5
6
7
Increases!
Decr
eas
es
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24. • Why opposite of atomic radius?
• In small atoms, e- are close to the nucleus where the attraction is stronger, therefore they are more difficult to remove
• Which family has the highest IE and why?
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25. Electron affinity
 Electron affinity of an atom can be described either as the energy gained by an
atom when an electron is added to it, or conversely as the energy required to
detach an electron from a singly-charged anion.
 The sign of the electron affinity can be quite confusing, as atoms that become
more stable with the addition of an electron (and so are considered to have a
higher electron affinity) show a decrease in potential energy; i.e. the energy
gained by the atom appears to be negative.
 As one progresses from left to right across a period, the electron
affinity value will decrease, i.e. the actual electron affinity of the
atom will increase , due to the larger attraction from the nucleus, and the
atom "wanting" the electron more as it reaches maximum stability.
 Down a group, the electron affinity decreases because of a large
increase in the atomic radius, electron-electron repulsion and the shielding
effect of inner electrons against the valence electrons of the atom.
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26. Electronegativity
 is a measure of the ability of an atom or molecule to attract
electrons in the context of a chemical bond.
 The type of bond formed is largely determined by the
difference in electronegativity between the atoms involved.
 Trend-wise, as one moves from left to right across
a period in the periodic table, the
electronegativity increases due to the stronger
attraction that the atoms obtain as the nuclear charge
increases.
 Moving down a group, the electronegativity
decreases due to the larger distance between the
nucleus and the valence electron shell, thereby decreasing
the attraction, making the atom have less of an attraction
for electrons or protons.
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27. • Electronegativity
• Decreases going down and Increases going
1
2
3
4
5
6
7
right.
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28. • Noble gases omitted because they do not
form many compounds.
• When F combines with another element, it
attracts e- strongly. Tug-of-War!!
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29. Metallic character
 refers to the chemical properties associated with elements
classified as metals.
 These properties, which arise from the element's ability to lose
electrons, are: the displacement of hydrogen from dilute acids;
the formation of basic oxides; the formation of ionic chlorides;
and their reduction reaction, as in the thermite process.
 As one moves across a period from left to right in the
periodic table, the metallic character decreases , as the
atoms are more likely to gain electrons to fill their valence shell
rather than to lose them to remove the shell.
 Down a group, the metallic character increases, due to
the lesser attraction from the nucleus to the valence electrons (in
turn due to the atomic radius), thereby allowing easier loss of the
electrons or protons.
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30. • Melting/Boiling Point
• Highest in the middle of a period.
1
2
3
4
5
6
7
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31. • Which atom has the larger radius?
• Be or Ba
• Ca or Br
Ba
Ca
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32. • Which atom has the higher 1st Ionization Energy?
• N or Bi
• Ba or Ne
N
Ne
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33. • Which atom has the higher melting/boiling point?
• Li or C
• Cr or Kr
C
Cr
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34. • Which particle has the higher electronegativity?
• K or Ca
• Al or Ga
Ca
Al
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35. THANK YOU FOR YOUR
ATTENTION!!!
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