Transcript Periodic Trend Practice I

The Periodic Law
Chapter-5
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Main Concepts
Main Concepts:
• Development of the Periodic Table
• Electron Configurations and the Periodic Table
• Periodic Trends
• Ionization Energy
• Electron Affinity
• Electronegativity
• Atomic Radius
• Ionic Radius
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The Periodic Table
By 1860, 63 elements had been discovered and
scientist started to notice patterns with the
elements.
Stanislao Cannizzaro developed a method for
measuring atomic masses based on the work of
Avogadro. This lead to the standardization of
atomic masses.
In 1869, Dmitri Mendeleev organized the
elements into a periodic table based on atomic
masses and grouped the elements by
similarities in their chemical and physical
properties.
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The Periodic Table
Mendeleev
demonstrated that
elements have
predictable
properties that
followed a repeating
pattern.
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The Periodic Table
Mendeleev predicted the existence and
properties of new elements and pointed out
accepted atomic weights that were in error.
He 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.
Mendeleev was so confident in his table that
he used it to predict the physical properties of
three unknown elements.
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The Periodic Table
These three unknown elements (Sc, Ga, and Ge)
were discovered between 1874-1885 and were
very close to the actual values he predicted
leading to the general acceptance of his periodic
table.
When discrepancies arose between atomic
masses and properties of other elements (Ar/K,
Co/Ni, Th/Pa, Te/I), these elements in question
were grouped by similar properties.
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The Periodic Table
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The Periodic Table
In 1913, Henry Moseley discovered that the
positive charge of the nucleus increased by
one unit from one element to the next.
He proposes that the elements in the periodic
table be arranged in order of increasing atomic
number instead of atomic mass.
Mosely’s work provided experimental proof for
Mendeleev’s ordering of the periodic table by
properties.
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The Periodic Table
Moseley’s work lead to
the modern definition
of atomic number and
that atomic number
and not atomic mass is
the basis for the
organization of the
periodic table.
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The Periodic Table
1894-1910 William Ramsay
discovers argon, neon,
and isolates helium for
the first time.
To fit these elements into
the Periodic Table,
Ramsey proposed a new
group between 17
(fluorine family) and
Group I (lithium family).
By 1910 all the noble
gases are discovered.
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Discovery of the Noble Gases
Sir William Ramsay
Periodic Trends
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Energy Contributions Within the Atom
1) Kinetic energy of the electrons as they
move around the nucleus
2) Attraction between protons in the nucleus
and the electrons (potential energy).
3) Repulsion between electrons in the electron
cloud (potential energy).
• Many of the properties of elements depend
not only on their electron configurations but
also on how strongly their outer electrons
are attracted to the nucleus.
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Nuclear Charge
• In atoms having more than
one electron, electrons are
both attracted to the
nucleus and repelled by
other electrons.
The nuclear charge that an
electron experiences
depends on both factors.
Effective Nuclear Charge
• The effective nuclear charge, is the net positive
charge experienced by an electron. The term
"effective" is used because shielding of inner
electrons prevents valence electrons from
experiencing the full nuclear charge.
• The “shielding effect” describes the decrease
in attraction between an electron and the
nucleus in any atom with more than one
electron shell.
• The shielding effect describes the balance
between the pull of the protons on valence
electrons and the repulsion forces from inner
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electrons.
Shielding
Trend: Shielding increases across a period by
sublevel and down a group by principal
quantum number.
• No shielding in a 1s-sublevel, s-sublevel shields
p-sublevel shields d-sublevel…
shielding remains constant within a sublevel.
• Although shielding increases across a period,
the effective nuclear charge increases more
rapidly due to added protons in the nucleus
from one element to the next. This causes the
electron cloud to be drawn in reducing the
atomic radius.
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Effective Nuclear Charge
• Many of the properties of atoms depend on
their electron configurations and how
strongly their outer electrons are attracted to
the nucleus.
• The higher the effective nuclear charge the
more tightly electrons are held within the
atom.
• Electrons are held less tightly in larger
orbitals due to the shielding effect.
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The Periodic Law and Periodicity
• Periodic Law: The physical and chemical
properties of the elements are periodic
functions of their atomic number.
• [Periodicity]: When elements are arranged in
order of increasing atomic number, elements
with similar properties recur at regular
intervals.
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Periodic Trends
Ionization Energy:
the energy required to remove one or more
electrons from a gaseous atom or ion to form a
cation.
Electron Affinity:
the energy change that occurs when one or
more electrons are added to a gaseous atom to
form an anion.
Electronegativity:
measures the ability of an atom to attract
electrons.
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Periodic Trends
Atomic Radius:
one half the distance between the nuclei of
adjacent atoms.
Ionic Radius:
one half the diameter of an ion in an ionic
compound.
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Periodic Trends
• Energy (ionization energy, electron affinity,
and electronegativity) increases diagonally
to fluorine (F): left to right across a period
and bottom to top up a group.
• Size (atomic and ionic radius) increases
diagonally to francium: top to bottom down
a group and right to left across a period.
• Metallic properties increase to francium and
properties of nonmetals increase to fluorine.
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Periodic Trends
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Periodic Trends: Ionization Energy
Ionization Energy:
• the energy required to remove one or more
electrons from a gaseous atom or ion to form a
cation: measures how strongly an atom holds on
to its electrons. Ionization energy is always
endothermic and forms cations.
First ionization energy:
• A + energy  A+ + e• Na + energy  Na+ + e23
Ionization Energy
Second ionization energy:
• A+ + energy  A+2 + e• Mg+ + energy  Mg+2 + e- or
• Mg + energy  Mg+2 + 2eOxidation: reactions that remove one or more
electrons from a substance. In oxidation, the
charge of the substance increases
mathematically (becomes more positive).
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Periodic Trends: Ionization Energy
• Removal of 1st electron = 1st ionization energy
• Removal of 2nd electron = 2nd ionization energy
• Removal of 3rd electron = 33d ionization energy
• Ionization energy increases as successive
electrons are removed from ions resulting in
higher positive charges.
• After all valence electrons have been removed,
the ionization energy takes a quantum leap in
magnitude.
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Ionization Energy of Carbon
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Ionization Energy of Sodium (kJ/mol)
1st
2nd
3rd
4th
5th
6th
7th
8th
9th
10th
11th
496
4,562
6,910
9,543
13,354
16,613
20,117
25,496
28,932
141,362
159,075
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Ionization Energy
Ionization Energy (kJ/mol)
1st
2nd
• Li
520
7300
• Na
496
4562
• K
419
3051
3rd
11815
6912
4411
• Mg
• Ca
738
590
1451
1145
7733
4912
10540
• Al
578
1817
2745
11578
4th
9544
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Ionization Energy
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Ionization Energy
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d- block Ionization Energy
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Ionization Energy
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Practice: Ionization Energy
Choose the elements for the following energies:
Ne, Ga, C, Ba, P, O, W, Pb, Se, I, Rb
1st
899
2nd
1757
1086 2350
1314 3390
3rd
14850
4th
21005
5th
34000
4620
6220
38000 47261 56,800
7470
11000
5300
6th
42000
13000
7th
55000
71320
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Practice: Ionization Energy
Choose the elements for the following energies:
Ne, Ga, C, Ba, P, O, W, Pb, Se, I, Rb
1st
2nd
578 1820
3rd
2750
4th
11600
5th
14800
6th
18400
7th
22600
1012 1904
2910
4960
6240
21000 24400
2080 3950
6120
9370
12220
15000
18000
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Practice: Ionization Energy
Choose the elements for the following energies:
Ne, Ga, C, Ba, P, O, W, Pb, Se, I, Rb
1st
2nd 3rd 4th 5th
6th
7th
8th
1680 3375 6045 8408 11020 15160 17860 92010
403 2632 3860 5200 8240 12100 18760 24360
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Electron Affinity
• Electron Affinity: the energy change that
occurs when one or more electrons are added
to a gaseous atom to form an anion: measures
how strongly atoms attract additional
electrons.
Periodic trends for electron affinity are less
predictable than those for ionization energy.
The energy released is expressed as a
negative value (exothermic). The greater the
negative value the easier it is for an atom to
gain an electron.
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Electron Affinity
The first electron affinity may be either
endothermic or exothermic.
• A + e-  A- + energy
• F + e-  F• A + e- + energy  A• Be + e- + energy  Be-
exothermic
ΔH = - 328 kJ/mol
endothermic
ΔH = + 53 kJ/mol
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Electron Affinity
The second electron affinity is always positive and
endothermic because it is always more difficult
to add an electron to an already negatively
charged ion.
• A- + e- + energy  A-2 second electron affinity
• O + e-  O- + energy
ΔH = -141 kJ/mol
O- + e- + energy  O-2
ΔH = +844 kJ/mol
Reduction Reactions: reactions that add one or
more electrons to a substance; the charge of the
substance decreases mathematically.
O + 2e-  O-2
more negative
Cu+2 + e-  Cu+
less positive
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Trend: Electron Affinity
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Trend: Electron Affinity
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Electronegativity
• Electronegativity: the measure of the ability of
an atom in a chemical compound to attract
electrons.
• Developed by Linus Pauling, electronegativity
is related to the ionization energy and electron
affinity of an atom in a molecule.
• Elements with a high electronegativity have a
greater tendency to attract electrons than
elements with a low electronegativity.
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Electronegativity
• Fluorine (F): the most
electronegative element is
arbitrarily assigned an
electronegativity value of 4.
In general, nonmetals are more
electronegative than metals.
• Electronegativity is used to determine bond
polarity which describes the degree of sharing of
electrons between atoms.
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Electronegativity
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Atomic Radius
Atomic Radius: one half the distance between
the nuclei of identical atoms joined in a
molecule. The electron cloud determines the
size of the atom but the boundary varies.
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Atomic Radius
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8.3
Atomic Radius
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Atomic Radii
8.3
Ionic Radius
Ionic Radius: is one half the diameter of
an ion in an ionic compound.
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Ionic Radius
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Ionic Radius
Cation: positive ion
formed by metals
giving up electrons.
Formation of a cation
by the loss of an
electron/s leads to a
decrease in radius.
The remaining
electron cloud is
smaller and closer to
the nucleus.
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Ionic Radius
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Ionic Radius
• Anion: a negative ion formed by a nonmetal
gaining electrons. Formation of an anion by a
gain in electrons leads to an increase in ionic
radius. The electron cloud spreads out
because there is a smaller attractive force on
the electrons from the nucleus and greater
repulsion forces between them.
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Ionic Radius
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Atomic/Ionic Radius
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Atomic/Ionic Radius
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Ionic Radius
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Trend Practice
1) List in order of increasing charge:
C (Z = 6), Al (Z = 13), S (Z = 16), Zn (Z = 30)
2) Largest size: Mn (Z = 25) or Mo (Z = 42)
Zr (Z = 40) or Sc (Z = 21)
O (Z = 8) or S (Z = 16)
Al (Z = 13) or Mg (Z = 12)
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Trend Practice
2) Largest size:
sulfur ion or sodium ion
magnesium atom or magnesium ion
iron ion or phosphorus ion
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Periodic Trend Practice
3) Cr (z = 24)
Al (z = 13)
Smallest atomic radius?
Ti (z = 22)
Smallest electronegativity?
Largest ionic radius?
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Periodic Trend Practice
4) I (z = 53)
N (z = 7);
S (z = 16)
Smallest atomic radius?
Largest electronegativity?
Largest ionic radius?
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Periodic Trend Practice
5)
Mn
Mg
Ge
Ba.
Which element/s has electronegativity greater
than Zn?
Which element would most easily form an anion?
Which element has the smallest electron affinity?
Which element has the largest atomic radius?
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Periodic Trend Practice
6)
Ni
Cl
Li
Fe
Largest 1st ionization energy?
Largest 2nd ionization energy?
Largest ionic radius?
Most easily forms a cation?
Arrange in order of increasing size.
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Periodic Trend Practice
6)
Ni
Cl
Li
Fe
Which is larger Ni atom or Ni ion?
Which is smaller Cl atom or Cl ion?
Which is larger Ni ion or Cl ion?
Which is monatomic
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Periodic Trend Practice
7) Equation for the second ionization energy
of Al.
8) Equation for the first electron affinity of Al.
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Trend Practice
9) List in order of
10) List in order of
decreasing atomic
decreasing 1st ionization
size:
energy:
a) Ca, Mg, Sr
a) Kr, He, Ar
b) Br, Rb, Kr
b) K, Ca, Rb
c) K, Ga, Ca
c) Sb, Te, Sn
d) Sr, Ca, Rb
d) I, Xe, Cs
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Electron Configuration Practice
Identify the element or ending electron configuration.
1) 3p4
2) 5p1
3) xenon (Z = 54 )
4) 5s1
Electron Configuration Practice
Identify the element or ending electron configuration.
5) Vanadium (Z = 23)
6) 5f3
7) 3d8
8) lawrencium (Lr) (Z = 103)
Electron Configuration Practice
Identify the element or ending electron configuration.
1) iridium (Z = 77)
2) 5s2
3) 3p5
4) 4p6
Electron Configuration Practice
Identify the element or ending electron configuration.
5) 3f6
6) 5f2
7) 6d5
8) 4s23d1
Electron Configuration Practice
a) 2s22p5,
b) 5s24d105p5,
c) 2s22p2,
d) 6s25d106p5, e) 2s22p4
1) Which element/s are in the same block?
2) Which element/s are in the same period?
3) Which element/s are in the same group?
4) Which element has the highest electron affinity?
Electron Configuration Practice
a) 2s22p5,
b) 5s24d105p5,
c) 2s22p2,
d) 6s25d106p5, e) 2s22p4
5) Which element/s would form a -1 ion?
6) Which element/s would form a -4 ion?
7) Which element/s would form a +3 ion?
8) Which element/s would form a -2 ion?
Electron Configuration Practice
a) 2s22p5,
b) 5s24d105p5,
c) 2s22p2,
d) 6s25d106p5, e) 2s22p4
9) Which element has the highest electronegativity?
10) Is the atomic or ionic radius larger for “e”?
11) Which element/s have seven valence electrons?
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Electron Configuration Practice
Given: 1s22s22p63s23p2
1) Group number?
2) Period number?
3) Block?
4) Type of element?
5) Charge of the most common ion?
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Electron Configuration Practice
Given: 1s22s22p63s23p64s23d104p65s24d105p66s24f5
1) Name of element?
2) Period element is in?
3) Block
4) Group number?
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Electron Configuration Practice
Given: 1s22s22p63s23p64s23d104p65s24d105p66s24f5
5) Energy level?
6) Type of element?
7) Charge of the most common ion?
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