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Periodic Trends

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PERIODIC TRENDS OF THE ELEMENTS. Chemical properties. Physical properties. Atomic radii. Ionic radii. Ionization energy. Electron affinity. Electronegativity – PowerPoint PPT presentation

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Title: Periodic Trends


1
Periodic Trends or in other words let the chart
do the work.
2
PERIODIC TRENDS OF THE ELEMENTS
PERIODIC TABLE ? I INVENTED IT !
  • Chemical properties
  • Physical properties
  • Atomic radii
  • Ionic radii
  • Ionization energy
  • Electron affinity
  • Electronegativity
  • Metallic nonmetallic character

Mendeleev
3
DEFINITIONS OF PERIODIC PROPERTIES
  • Chemical properties refers to the tendency of
    atoms to combine with other elements and / or
    molecules to form compounds.
  • Physical properties refers to characteristics
    such as density, physical state (solid, liquid or
    gas), electrical and thermal conductivity,
    malleability (can be hammered into shape),
    ductility (can be stretched), color, luster
    (shininess), brittleness, etc.

4
Definitions of periodic properties (contd)
  • Atomic radii is measured as one half of the
    distance between the nuclei of two adjacent
    similar atoms (it measures atomic size)
  • Ionic radii is measured as one half of the
    distance between the nuclei of two adjacent
    similar ions (it measures ion size)
  • Ionization energy measures the energy needed to
    remove an electron from a free atom in the gas
    state (it measures how tightly electrons are
    bound to an atom)
  • Electron affinity measures the energy released
    when an electron is added to a free atom in the
    gas state ( it measures how well atoms attract
    electrons)

5
Definitions of periodic properties (contd)
  • Electronegativity measures the electron
    attracting ability of an atom when it is bonded
    to another atom
  • Metallic character measures the tendency of an
    element to act as a metal in things such as
    conductivity, tendency to lose electrons,
    shininess, maebility and ductility
  • Nonmetallic character measures the tendency of an
    element to act as a nonmetal in things such
    nonconductivity, tendency to gain electrons, low
    luster and brittleness.

6
General organization of the periodic table
  • Columns (families or groups) contain elements
    with similar valence electron configurations
    ns1, ns2, ns2 np3, etc. and therefore similar
    chemical properties
  • Rows (periods or series) contain elements with
    valence electrons at the same energy level (n1,
    n2, n3, etc.)
  • Blocks of elements contain atoms with the same
    valence electron orbital type (s, p, d or f)

7
Orbital Blocks on the Periodic Table
metals
non metals
I N E R T G A S
s B L O C K
p BLOCK
d BLOCK
f BLOCK
METALLOIDS
8
Common Chemical Families Their Properties
  • Column I (Alkali Metals) Li Na K Rb Cs and
    Fr
  • Form 1 cations
  • Are highly metallic
  • Reaction readily and rapidly with water to form
    hydroxides and hydrogen gas
  • React with the halogens to form ionic salts with
    the formula type MX. For example NaCl or KBr.
  • Valence Electrons are ns1
  • Column II (Alkaline Earth Metals) Be Mg Ca Ba
    Sr and Ra
  • Form 2 cations
  • Are highly metallic
  • Reaction readily and rapidly with water to form
    hydroxides and hydrogen gas
  • React with the halogens to form salts with
    formula type MX2 (for example MgBr2 or BaCl2)
  • Valence Electrons are ns2

9
Common Chemical Familes Their Properties page 2
  • Column VII (Halogens) F2 Cl2 Br2 I2 and At2
  • All are diatomic elements (occur as a
  • molecule consisting of two atoms
  • Are highly nonmetallic
  • React ready and rapidly with metals to form salts
  • Occur in all three phases at room temperature F2
    and Cl2 are gases, Br2 is a liquid and I2 is a
    solid.
  • Valence Electrons are ns2 np5

10
Common Chemical Families Their Properties page 3
  • Column VIII (Noble Gases or inert gases)
  • All are unreactive under ordinary conditions
  • All have completed outer energy levels
  • All are gases at room temperature and pressure
  • Their electron configuration is ( ns2 np6 )
    This is a filled outer shell, which the other
    atoms attempt to achieve through chemical
    reactions and bonding.

11
CHEMICAL FAMILIES
Alkali metals
Halogens
Alkaline Earth Metals
I N E R T G A S E S
C C O O L L M M U U N N I
II
C
O L
U M
N
V
I I
Transitional Metals
LANTHANIDE SERIES ACTINIDE SERIES
12
What factors determine the periodic trends of the
elements ?
  • The number of protons and electrons an atom
    contains. (more protons create a greater nuclear
    charge which attracts electrons more strongly)
  • Distance separating the outer electrons (valence
    electrons) and the nucleus. (when electrons are
    closer to the nucleus they are held more tightly)
  • 3. Pairing of electrons in the outer energy
    level orbitals. (paired electrons are more stable
    than unpaired electrons)

13
Periodic trends Atomic Radii
  • As we move across a row (period) from left to
    right on the periodic table, atoms become
    smaller as the atomic number becomes larger.
    This increase in nuclear charge allows the
    nucleus to pull in the electrons more tightly
    and thereby reduce atomic size (radius).
  • As we move down a column on the periodic table,
    elements contain more electrons and more energy
    levels become populated resulting in an increase
    in atomic size (radius). Remember that
    completing a period on the periodic table
    results in a completed energy level within the
    atom.

14
Periodic Trends Atomic Radii
  • SIZE OF ATOMS DECREASES
  • S
  • I
  • Z
  • E
  • I
  • N
  • C
  • R
  • E
  • A
  • S
  • E
  • S

15
Periodic Trends - Ionic Radii
Metal atoms lose electrons to form ions so they
are smaller than the original atoms. In a sodium
atom, the last electron is in the 3rd energy
level, but when it is lost, the outer shell is
now the second level. Nonmetal atoms gain
electrons to form ions. Since the number of
protons doesnt change, the electrons cant be
held as tightly and they are able to move further
away. Anions in the same period are larger
than cations in that row.
16
Periodic Trends Ionic Radii
  • Na1 has 2
    filled energy levels

  • K1 has 3
    filled energy levels

8e
2e
11
8e
8e
2e
19
17
Periodic Trends Ionic Radii

The fluoride ion has 9 protons attracting the
the 10
electrons so there is more attraction

The nitride ion (N-3) has only 7 protons
attracting 10 electrons so the electrons are not
held as tightly and the ion is larger than the
fluoride ion.





8e
2e
9
8e
2e
7
18
Periodic Trends Ionic Radii
?
  • Why is this important ?
  • Excellent question glad you asked.
  • Explaining a lot of bonding and chemistry deals
    with two ideas.
  • Electronic Factors Chemistry occurs because of
    the charge factors, this is why it is NaCl and
    not NaCl2
  • Steric Factors Chemistry occurs because of the
    size of molecules, just like in our macro world.
    Medicines and body chemistry use this a lot.

19
Periodic Trends Ionic Radii
  • SIZE OF IONS INCREASES
  • S
  • I
  • Z
  • E
  • I
  • N
  • C
  • R
  • E
  • A
  • S
  • E
  • S

20
Periodic Trends Ionization Energy
  • The outer electron for the sodium
    ion is in the 3rd energy level so it is
    closer to the nucleus than the outer
    electron for the potassium ion which
    is in the 4th energy level. The
    further away from the nucleus, the
    less force of attraction and the
    easier it is to remove
    that electron.


1e
8e
2e
11
1e
18e
8e
2e
19
21
Periodic Trends Ionization Energy
Fluorine has two options to become an
ion. It can gain 1 e-1 to have a filled
second shell or it can lose 7e-1 to have a
filled first shell. It will always be the
choice that deals with the lower number
of electrons.
Sodium
can either lose a single electron or gain
7 electrons to achieve a filled outer
shell. Remember that it is always
deal with the lower number of
electrons so Sodium will lose one
electron and in this case form a
positive ion.







7e
2e
9
1e
8e
2e
11
22
Periodic Trends Ionization Energy
  • Successive Ionization Energies
  • For a sodium ion to be formed, only one electron
    has to be removed, and it takes energy to remove
    the electron because it doesnt just happen on
    its own. This energy is called the Ionization
    Energy.
  • For a magnesium ion to be formed, two electrons
    have to be removed. It would be simple if we can
    simply double the Ionization Energy but it
    doesnt work that way. We need to have an
    Ionization Energy for the 2nd electron. After
    the first electron, the others are called
    successive ionization energies.
  • Different factors affect the value of the these
    energies and will be described on the next slide.

23
Periodic Trends Ionization Energy Successive
Ionization Energy
  • Distance from the nucleus. The closer the e-1
    is to the nucleus the more energy is needed to
    remove it.
  • Proton to electron ratio. The more protons
    present in the nucleus, the more attraction and
    the more energy is needed to remove an electron.
  • Filled and half-filled sets of orbitals. Filled
    and half-filled sets of orbitals are more stable
    and therefore take more energy to remove an
    electron from them.
  • Shielding effect. Since magnetic attraction
    cant move through magnetic material, electrons
    in inner shells, can block the attraction to the
    electrons in the outer shell, making it easier to
    remove those electrons.

24
Periodic Trends Ionization Energy
  • Across a row, the Ionization Energy increases.
  • Down a column, the ionization energy decreases.

25
Periodic Trends Ionization EnergySuccessive
Ionization Energies
  • The Blue line occurs when the next electron being
    removed is from an inner shell.

26
Periodic Trends Electron Affinity
  • Elements which are more nonmetallic (to the left
    up on the periodic table) have greater electron
    affinity.
  • Those which are more metallic (to the right
    down on the periodic table) have lower electron
    affinity.

27
Periodic Trends Electron Affinity
  • Across each row, electron affinity increases.
  • Down a column, electron affinity decreases.

28
Periodic Trends - Electronegativity
  • Electronegativity is the ability of an atom to
    grab onto the shared electrons in a bond.
  • Flourine has the highest electronegativity value
    of 4.0
  • This concept was developed by Linus Pauling (2
    unshared Nobel Prizes). It was devised by
    comparing the different bonds to each other, then
    when fluorine was discovered to have the highest
    value he assigned it a value of 4. There are no
    units !
  • As you move away from Fluorine, the values get
    lower.

29
Periodic Trends Electronegativity (En)
  • Across each row, electronegativity increases.
  • Down a column, electronegativity decreases.
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