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## Periodic Table Trends

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### Title: Periodic Table Trends Author: Donald Sandberg Last modified by: Donald Sandberg Created Date: 10/24/2011 4:47:17 PM Document presentation format – PowerPoint PPT presentation

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Title: Periodic Table Trends

1
Periodic Table Trends
2
Periodicity
• A regular pattern in the properties of elements
• Also called a TREND
• Some properties are similar due to the valence
e-s ? specifically in groups of element
• Remember ?? the main group elements, Roman
numerals on the top tells us the number of
valence e-

3
• some properties of elements are related to the
positive attraction of the nucleus for electrons
called nuclear effective charge
• This attraction is due to electrostatic
attraction and is dependent on quantity of the
charge as well as the distance separating the
charge
• Example

distance
quantity
4
• Cant know the exact size of an atom? can only
look at relative size
• Atomic radius is defined as the closest distance
to which one atom will approach another atom of
any other size

5
• In a period, atomic radius decreases as atomic
number increases (from left to right in a row)
• WHY?
• As atomic number increases, the increasing number
of protons attract the valence electrons closer
to the nucleus ?decreasing the size of electron
cloud
• Elements in the same period have valence e- in
the same valence shell ? the effective nuclear
charge is felt more by the valence e- across a
period

6
• In a group, atomic radius increases as atomic
number increases (from top to bottom)
• WHY?
• As atomic number increases in a group, each
successive element has its valence electron
farther from the nucleus (since they are placed
in a higher energy level with each successive
element)
• ??the effective nuclear charge has less of an
effect on the valence e- and the size of the atom
increases

7
Summary
increases
• Atomic radius from top to bottom in any
group
• Atomic radius from left to
right in any period

decreases
8
Ionization Energy and Trends
• Ionization energy (I.E.)
• Amount of energy required to remove an electron
from a gaseous atom
• I.e. X(g) I ? X(g) e-
• it is an endothermic process
• The amount of energy required to remove ONE e-
from a gaseous atom is called the first
ionization energy (I1)
• The amount of energy required to remove a SECOND
e- from a positive ion is called the second
ionization energy (I2)

9
• Each successive ionization requires more energy
because as electrons are being removed, the
nuclear effective charge has a greater pull on
the electrons
• In a period, I.E. increases as atomic number
increases
• WHY?
• As the nuclear charges increases across a period,
it becomes more difficult to remove
electrons?more energy is required to remove them

10
• In a group, I.E. decreases as atomic number
increases
• WHY?
• With every successive element in a group there is
an additional energy level of electrons so the
electrons are at a greater distance from the
nucleus
• This means it is easier to remove electrons since
not much energy is required to remove them
• Also, the shielding effect of the kernel electron
makes the valence e- feel less attracted to the
nucleus

11
Summary
increases
• I.E. across a
period
• I.E. down a group

decreases
12
Electron Affinity and Periodicity
• Electron affinity (E.A)
• The amount of energy released when a gaseous atom
accepts an electron to form an anion. (a
negatively charged atom)
• i.e. X(g) e- ? X-(g) Energy
• It is an exothermic process
• Adding the first e- is an EXOTHERMIC process
• Adding a second e- is ALWAYS an ENDOTHERMIC
process

13
• Adding 2e- at the same time would still be an
ENDOTHERMIC process
• I.e. O(g) 2e- energy ? O2-(g)
• In general, Electron Affinity
as atomic number increases in a
PERIOD
• The Electron Affinity
as atomic number increases in a GROUP
• Electron affinity is the OPPOSITE of an atoms
ionization energy

increases
decreases
14
• WHY?
• The more attracted the element is for electrons,
the more energy it will release when it accepts
an e-
• MOST reactive ?Fluorine
• (in terms of E.A)

15
Electronegativity
• The measure of the ability of an atom to attract
SHARED electrons to itself
• Electronegativity is much higher for non-metals
than metals
• In a period, the electronegativity will increase
as the atomic number increases
• In a group, the electronegativity will decrease
as the atomic number increases

16
Why?
• The stronger the pull of the nucleus for
electrons, the higher the value of
electronegativity

17
Reactivity
• The alkali metals are the most reactive metals
and the halogens are the most reactive non-metals
• Reactivity increases as one moves away from the
group containing Zn on the periodic table
• (transition metals are fairly unreactive when
compared to the main group elements)
• For metals, reactivity increases down a group but
for non-metals, increase going up a group

18
Metallic Character
• Metals are located on the left side of periodic
table and non metals are on the upper right side
of the table
• Elements tend to posses more metallic character
as one moves to the lower left corner of the
periodic table
• Fr is the most metallic element and He is the
most non-metallic element

19
• According to this trend, which transition metal
is the least metallic?
• Zn