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Chapter 6
The Periodic Table and
Periodic Trends
The Periodic Table
1. Mendeleev - arranged elements in order of
increasing atomic mass.
- arranged elements with most
similar properties in columns
2. Moseley - used x-rays to determine the
atomic number of the elements.
-elements arranged by atomic number
show a clear “periodicity” of properties.
Periodic Law
• There is a periodic repetition of the
element’s chemical and physical properties
when they are arranged in order of
increasing atomic number. This repetition
is known as the Periodic Law.
Arrangement of the Periodic Table
1. Period - horizontal row of elements
- period # = # of energy levels an
element has
2. Group (or family) - vertical column of elements
- Number (AorB) above a column is its
family designation.
- elements in a group have much more similar
properties than elements in a period.
Group A Elements
(Representative Elements)
contain metals, non-metals and metalloids
• s and p blocks.
• Group # = # of electrons in the outermost
energy level.
• All shells beneath the outermost have stable octet.
– Except: Helium (He), Lithium (Li) &
Beryllium (Be). First energy level is
full with 2 electrons.
Group B Elements
(Transition Metals)
• d block.
• May have more than 1 ionic charge because
electrons in the shell beneath the outermost
get involved in bonding.
– Shells beneath the outermost may have more
than 8 electrons.
Classification of the Elements
• There are 3 classes of elements; metals,
nonmetals and metalloids.
• Metals - make up about 80% of the table
– Good conductors of heat and electricity
– Have a high luster or sheen. (shiny)
– Solids at room temp. except for Mercury (Hg)
– Ductile; drawn into wires
– Malleable; hammered into thin sheets.
Nonmetals
• Found in the upper right corner of the table
• Properties of nonmetals are not as similar as
they are for metals.
• Most are gases at room temperature.
• Some are solids like sulfur and phosphorus
• Bromine is a dark red liquid.
• Generally have properties opposite of the
metals.
• Not good conductors (carbon exception),
dull and brittle if they are solids.
Metalloids
• Along the staircase that separates the metals
and nonmetals.
• Have properties that are similar to metals
and nonmetals under certain conditions.
• Aluminum (Al) is a metal.
Classification of Elements
Continued
1. Metals – lose electrons to form positive
ions (cations) when they bond.
S block:
a) Group 1A (ns1) - Alkali or active metals
- form +1 ions
- very reactive, stored in kerosene
S block:
a) Group 2A (ns2) -
Alkaline earth metals
- form +2 ions
Note – Elements with the same outer shell electrons
configuration have the same physical and chemical
properties. (same number of valence electrons.)
Metals (continued)
• D block – these metals are all similar
because they all have the same s2
electron configuration in their
outermost shell.
– Transition Metals – Group B metals
– May have more than 1 ionic charge
(designated by Roman #)
Metals (continued)
• Hund’s Rule – 1/2 filled and
completely filled sublevels have
the lowest energy are the most
stable, and therefore the preferred
electron arrangement.
Write the electron configuration for
Fe
• # 26
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Forms Fe+2 when it loses 2 electrons
Forms Fe+3 when this additional electron is lost
the ½ filled is d is more stable.
Write the electron configuration for
Cu
• # 29
1s2 2s2 2p6 3s2 3p6 4s2 3d9
Forms Cu+2 when it loses 2 electrons
Hund’s Rule Shift: 1 e- shifts from 4s to 3d.
The 3d is now completely filled and Cu will
only lose 1 electron to form Cu+1.
Metals (continued)
• F block – Inner Transition Metals
– 2 rows:
• 1st row is Lanthanide series – “rare
earth” elements (natural).
• 2nd row is Actinide series – mostly
synthetic (manmade).
–Parentheses around mass # means
radioactive.
Write configurations for #58 and #95.
Then write the condensed
configurations.
• Ce # 58
•
1s2 2s2 2p6 3s2 3p6
2 4s2 3d10 4p6 5s2 4d10 2
6 6s2 5d1 4f 1
8
5p
8
18
19
9
2
same
Am # 95
1s2 2s2 2p6 3s2 3p6
4s2 3d10 4p6 5s2
4d10 5p6 6s2 4f 14
18 5d10 6p6 7s2 6d1 5f6
32 Inner transition metals
24 are most similar b/c
9 their last 2 energy level
2 config. are identical.
2. Semi-metals
• Located on either side of the “stairs” in the
p block separating the metals from the nonmetals.
• Have properties of both metals and
nonmetals.
– Aluminum (Al) is a metal
3. Non-metals
•
in upper right corner of p block
(except last column).
• gain electrons to achieve stable octet
• form negative ions - anions
• group 5, 6, and 7 gain 3, 2, or 1 electrons
Halogens - Group 7A (ns2p5)
Gains 1 e- and has a –1 charge
Question: Which are more similar?
F, Cl, and Br or N, O, and F
Answer: F, Cl, and Br
Because in the same group or column,
they have the same outer shell electron
configuration, therefore react similarly.
4. Noble Gases
2
6
(ns p )
• Last column on the right
in p block.
-2
i.e. Sulfide ion (S ) ion
• All but He already have their stable 8 in the
is isoelectric
to Argon
(Ar)
outer
shell. Therefore
they
are atom.
called:
– Inert – means non-reactive
• Isoelectric- same electron configuration
as a noble gas.
– Atoms react to become isoelectric with a
noble gas.
Periodic Trends
• Trends or patterns on the Periodic
Table are all basically the result of
atomic size and whether an atom of an
element is a metal or a non-metal.
1. Atomic size (atomic radii) –
size of the atom
• Nuclear charge – pull of protons in
nucleus on electrons
– increase number of protons = increase
the nuclear charge.
Atomic Size
}
Radius
• Atomic Radius = half the distance between two nuclei
of a diatomic molecule.
1. Atomic size (atomic radii) –
size of the atom
• Across a period – decreases due to
increasing nuclear charge which pulls
the electrons on the same energy level in
closer to the nucleus.
• Down a group – increases due to the
addition of energy levels which are
farther from the nucleus.
1. Atomic size
Across a period:
Greater nuclear charge – same energy level
Down a group:
More energy levels
2. Metallic vs. Non-metallic Character
a) Metallic Character – tendency to lose electrons
- larger atom – easier to lose electrons.
Across a period –decreases because the
smaller the atom, the stronger the pull
from the nucleus so the harder it is to lose
electrons.
– Also the # of e- to lose is increasing as you
approach the non-metals.
2. Metallic vs. Non-metallic Character
a) Metallic Character – tendency to lose electrons
- larger atom – easier to lose electrons.
• Down a group – increases because larger
atoms have less pull from nucleus on
outer electrons.
– More energy levels easier to lose electrons.
2. Metallic vs. Non-metallic Character
b) Non-metallic Character – tendency to gain e- smaller atom – easier to gain electrons
Across a period –increases because as
atomic size decreases its easier to gain
electrons due to increased pull from
nucleus.
- also needs to gain fewer electrons.
2. Metallic vs. Non-metallic Character
b) Non-metallic Character – tendency to gain e- smaller atom – easier to gain electrons
Down a group – decreases because
atoms are larger (more energy
levels).
- less pull from nucleus- harder to
gain electrons.
3. Shielding Effect
Each additional energy level acts as
another layer of electrons that shields
the outermost electrons from the pull
of the nucleus.
3. Shielding Effect
• Across the period – unchanged; on the
same energy level
– With each addition electron, there is an
additional proton.
• Down a group – increases due to
additional energy levels between the
nucleus and outer electrons.
4. Ionization Energy
• Amount of energy required to remove
an electron from an atom.
– Metals have low ionization energies
because they lose electrons easily.
– Non-metals have high ionization
energies because they do not lose
electrons easily.
4. Ionization Energy
• For group A metals: as electrons are
removed 1 by 1 from an atom, there
is a steady rise in the ionization
energy until the structure of a noble
gas is attained.
• Then there is a dramatic rise when a
stable octet is disturbed.
4. Ionization Energy
• 1st ionization energy – amount of energy
needed to remove the 1st electron from
an atom.
• 2nd ionization energy – amount of energy
needed to remove the 2nd electron from
an atom.
• 3rd ionization energy – amount of energy
needed to remove the 3rd electron from
an atom.
4. Ionization Energy
2
1st ion.
8
1 Na
119
2 1s2 2s2 2p6 3s1
8
176
2 Mg
1s2 2s2 2p6 3s2
2
8 Al
138
3 1s2 2s2 2p6 3s2 3p1
2nd ion.
3rd ion.
4th ion.
1090
1652
2281
347
1848
2519
434
656
2767
4. Ionization Energy
• Across a period – 1st ionization energy
increases (across group A) because nuclear
charge is increasing, holding electrons in
tighter and closer to nucleus.
– For transition metals (pds 4, 5, & 6) in D
block there is a sharp drop in ionization
energy between the last transition element
and where the next energy level’s p
sublevel begins because the p sublevel is
farther from the nucleus – becomes easier
to lose electrons.
4. Ionization Energy
• Down a group – 1st ionization energy
decreases because larger atoms have less
pull on outer electrons due to increased
shielding effect (see # 3).
5. Electronegativity or
Electron Affinity (attraction)
• Energy change due to the addition
of an electron to an atom.
Halogens
group
7A have
the
• Metals
haveinlow
electron
affinities.
highest electron
because
• Non-metals
have affinity
high electron
they onlybecause
need toby
gain
1 electron.
affinities
gaining
electrons, they achieve the stable
configuration of a noble gas.
5. Electronegativity or
Electron Affinity (attraction)
• Across a period – increases because
smaller atoms have a greater attraction
for electrons and also approaching the
non-metals that want to gain.
• Down a group – decrease because larger
atoms have less attraction to gain
electrons (outer shell farther from
nucleus).
6. Ionic Size
• Ion – atom with unequal # of protons and
electrons and has charge.
• Metals lose electrons to form positive
ions that are smaller than their respective
atoms because when outer electrons are
lost, the protons in the nucleus have a
great pull on the remaining electrons.
6. Ionic Size
• Ion – atom with unequal # of protons and
electrons and has charge.
• Non-metals gain electrons to form
negative ions that are larger than their
respective atoms because when electrons
are gained, there are less protons than
electrons, so the electrons are not held as
tightly.
6. Ionic Size
• Across a period – gradual decrease in the
size of positive ions as they lose more
electrons and then at group 5 begin the
large negative ions that gradually
decrease in size as you continue across.
LOSE electrons
1
2
3
4
GAIN electrons
3
2
1
7. Ionic Size
• Down a group – increases due to
additional energy levels
– Holds true for both positive and
negative ions.
Summation of Periodic Trends
Atomic size
Metallic
character
Non-metallic
character
Going across
a period.
Going down a
group.
decreases
increases
decreases
increases
increases
decreases
Summation of Periodic Trends 2
Going across
a period.
Nuclear
charge
Shielding
effect
Speed of
reaction –
metals
increases
Going down a
group.
increases
no change
increases
decreases
increases
Summation of Periodic Trends 3
Going across
a period.
Speed of
reaction –
non-metals
Ionization
energy
Electron
affinity
Going down a
group.
increases
decreases
increases
decreases
increases
decreases
Summation of Periodic Trends 4
Going across
a period.
Going down a
group.
Ionic size of Decreases thru increases
+ ions(cation) group 4A
Ionic size of At group 5,
increases
- ions(anions) large increases
then a gradual
decrease