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PERIODICITY
Chapter 6
Mrs. Medina
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Searching for an organizing principle
• Year 1700, only 13 elements were known
– Copper, silver, gold, iron…
• Chemists began using scientific methods and the
rate of discovery increased.
• How would chemists know if all the elements had
been discovered?
– They had to be organized into some sort of system.
• Chemists used the properties of elements to sort
them into groups.
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Mendeleev’s Periodic Table
• In 1869, a Russian chemist and teacher, Dmitri
Mendeleev published a periodic table.
• He arranged the elements in order of increasing
atomic mass.
– By doing this, the elements fell into groups because the
properties began to repeat themselves.
• This repetition of properties allowed Mendeleev to
predict the elements gallium and germanium
earning him international fame for his table.
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The Periodic Law
• Mendeleev’s table placed the elements in order of
increasing atomic mass but this arrangement had to be
broken sometimes because otherwise the properties would
not match up.
• Moseley discovered the proton in 1913.
• When the elements were placed in increasing atomic
number, all the repeating properties fell into place without
any problems.
• In the modern periodic table, elements are arranged in
order of increasing atomic number.
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The Periodic Law
• Periodic Law: When elements are arranged in
order of increasing atomic number, there is a
repetition of their physical and chemical
properties.
– Elements within a group (column) have similar
properties
– Properties in a period (row) change as you move
from left to right.
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Metals, NonMetals and Metalloids
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Metals
• Most elements are metals
• Good conductors of heat and electricity
• High luster, or sheen b/c metals reflect light
• All metals are solids at room temperature except
one: mercury (Hg) which is a liquid.
• Ductile – draw into wires
• Malleable – hammer into thin sheets
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Metals
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Nonmetals
• Include the noble gases and hydrogen
• Among nonmetals there is a greater variation in
physical properties
– Most are gases, a few are solids, and one is a liquid.
• They have properties that are opposite of the
metals:
– Poor conductors of heat and electricity
• carbon being an exception
– Solid nonmetals will be brittle (shatter easily)
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Nonmetals
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Metalloids
• Found on the stair-step line between metals and
nonmetals.
• Metalloids have properties that can be similar to
both metals and nonmetals
– Under some conditions, it can behave as a metal and
conduct electricity
– Under other conditions, it may not conduct electricity
like the nonmetals.
• Silicon, a metalloid, is widely used in the
computer chip industry.
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Metalloids
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Squares in the periodic table
• The periodic table displays the symbols and names
of the elements, along with the information about
the structure of their atoms.
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Groups or Families
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•
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Group 1A: Alkali Metals
Group 2A: Alkaline Earth Metals
Group B: Transition Metals
Group 4A – 6A: …
Group 7A: Halogens
Group 8A: Noble Gases
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Families (groups) in the periodic Table
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Electron Configurations in Groups
• Electrons play a key role in determining the
properties of elements.
• There is a connection between the placement on the
periodic table and the electron configuration.
• We can sort elements based on their electron
configuration (See periodic table with electron config)
–
–
–
–
Noble gases
Representative elements
Transition
Inner transition metals
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Noble Gases
• Found in group 8A; also known as inert gases b/c
they rarely take part in a chemical reaction.
• The s and p sublevels of the highest occupied
energy level are completely filled.
– This explains the relative inactivity of the noble gases.
Helium
Neon
Argon
Krypton
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p63d104s24p6
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The representative elements
• Groups 1A, 2A……3A,4A,5A,6A,7A
– Does not include B middle and noble gases
• Why are they the representative elements?
– They display a wide range of physical and chemical
properties
– Some are metals, some are metalloids, and some are
nonmetals
– Most are solids, few are gases, and 1 is a liquid.
• The s and p sublevels of the highest occupied energy
level are not filled.
– They participate in reactions when they are able to lose or
gain electrons to have full outer energy levels.
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The representative elements
Element
Electron Configuration
Lithium
1s22s1
Sodium
1s22s22p63s1
Potassium
1s22s22p63s23p64s1
Carbon
1s22s22p2
Silicon
1s22s22p63s23p2
Germanium 1s22s22p63s23p63d104s24p2
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Transition and Inner Transition Elements
• Transition Metals
– Group B displayed in the middle of the main body of the
periodic table
– S sublevels of the higher energy level and a nearby d sublevel
contain electrons
• Inner Transition Metals
– Appear below the main body of the periodic table
– S sublevels of the higher energy level and a nearby f sublevel
contain electrons
Element
Electron Configuration
Titanium
1s22s22p63s23p64s23d2
(transition)
Cerium
1s22s22p63s23p64s23d104p65s24d105p66s24f2
(inner transition)
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Blocks of Elements
• The periodic table is divided into patterns of electron
configurations
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–
–
–
s block = groups 1A and 2A (including Helium)
p block = groups 3A, 4A, 5A, 6A, 7A, 8A
d block = transition metals
f block = inner transition metals
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Blocks in the periodic table
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Abbreviated electron configurations
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Trends in atomic size
• Atomic radius
– ½ the distance between the nuclei of two atoms of the
same element when the atoms are joined.
• 1 meter = 1012 picometers
• In general, atomic size tends to
– Increase from top to bottom within a group
– Decrease from left to right across a group
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Trends in atomic size
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Trends in Atomic Size
• Why does it increase down a group?
– Down a group you have 2 opposite effects:
• Nuclear attraction: more protons more positive charge; nucleus pulls
electrons in
• Shielding effect: increased # of energy levels shields the electrons from
the nucleus.
– The shielding effect is larger than the nuclear
attraction. The atom size gets bigger
• Why does it decrease across a period?
– Across a group 2 things are occurring:
• Nuclear attraction: more protons more positive charge; nucleus pulls
electrons in
• Shielding effect: CONSTANT… All electrons are in the same energy
level
– The nuclear attraction is stronger and pulls the
electrons in making the atom smaller.
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Ions
• An ION is an atom or group of atoms that has a
positive or negative charge.
• Positive and negative ions form when electrons are
transferred between atoms.
– Cation: ion with a positive charge
– Anion: ion with a negative charge
• Metals become cations by losing electrons
• Nonmetals become anions by gaining electrons.
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Trends in ionic size
• Cations are smaller than the atoms from which they
form.
– When an atom loses an electron, the remaining electrons
are drawn closer to the nucleus (increased attraction).
• Anions are always larger than the atoms from which
they form.
– When an atom gains an electron, the attraction of the
nucleus for any one electron decreases allowing the size to
get bigger.
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Trends in Ionic Size
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Trends in ionization energy
• If there is enough energy, an electron can be
removed from an atom.
• Ionization energy:
– The energy required to remove an electron from an
atom.
• First ionization energy tends
– to decrease from top to bottom within a group
– To increases from left to right across a group
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Ionization Energy Trend
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Trends in Ionization energy
• Why does it decrease down a group?
– It is easier to remove electrons because the shielding
effect is stronger than the nuclear attraction
hence less energy is required.
• Why does it increase across a period?
– Across a period, the shielding effect remains
constant but the nuclear attraction increases. It
takes more energy to remove an electron from the
atom.
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Trends in electronegativity
• Electronegativity: The ability of an atom of an
element to attract electrons when the atom is in a
compound.
– Electronegativity can help predict the type of bond
(ionic or covalent) that will form during a reaction.
– The noble gases do not form many compounds and
thus they do not have electronegativity values.
• In general, electronegativity values tend to
– Decrease from top to bottom within a group
– Increase from left to right across a period
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Trends in electronegativity
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Trends in electronegativity
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Summary of Trends
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