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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
  • Atomic Radius
  • 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
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