Download The Periodic Table

The
Periodic
Table
Why is the Periodic Table important
to me?
• The periodic table
is the most useful
tool to a chemist.
• You get to use it on
every test.
• It organizes lots of
information about
all the known
elements.
Pre-Periodic Table Chemistry …
• …was a mess!!!
• No organization of
elements.
• Imagine going to a grocery
store with no organization!!
• Difficult to find information.
• Chemistry didn’t make
sense.
Dmitri Mendeleev:
Father of the Periodic Table
SOME PROBLEMS…
HOW HIS WORKED…
• Put elements in rows by • He left blank spaces for
what he said were
increasing atomic weight.
undiscovered elements.
• Put elements in columns
(Turned out he was
by the way they reacted.
right!)
• He broke the pattern of
increasing atomic weight
to keep similar reacting
elements together.
Periodic Table
• This table is a remarkable way to show the
manifold relationships between differing kinds of
elements
• The modern table was devised in 1869 by Dimitri
Mendeleyev
• He arranged the elements by weight and by their
chemical properties
Periodic Table
• "...if all the elements be arranged in order of their
atomic weights a periodic repetition of properties
is obtained." - Mendeleyev
The Current Periodic Table
• Mendeleev wasn’t too far off.
• Now the elements are put in rows by increasing
•
ATOMIC NUMBER!!
This is called “periodic law”.
• The horizontal rows are called periods and are
labeled from 1 to 7.
• The vertical columns are called groups are
labeled from 1 to 18.
Metals vs Non Metals
Metals
•
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•
•
Have Luster (shiny / reflect light)
Solid at room temperature (except mercury)
Gray or silver color (except gold and copper)
Malleable (Bend, dent, deform when hit with mallet)
Ductile (Can be drawn into thin wire)
Good conductor of electricity
Good conductor of heat
METALS
Metals vs Non Metals
NON Metals
• Are dull (do NOT reflect light)
• Some are solid, liquid or gas at room temperature
• Come in many colors including colorless, black,
purple, yellow, green, gray, etc.
• Brittle (shatter when struck with mallet)
• Poor conductor of electricity
• Poor conductor of heat
NONMETALS
Metalloids or semi-metals
Border the staircase
Have blends of properties:
some metal characteristics
some non metal characteristics
Si, Ge, As, Sb, Te, sometimes B
METALLOIDS
B
Si
Ge
As
Sb
Te
Periods and Groups
• Horizontal rows in the periodic table
are called periods.
• Vertical columns are called groups.
• We know that the electron configuration
of the elements matches the way they
are arranged. (see p362-363 in your textbook)
The relation between orbital filling
and the periodic table
Groups…Here’s Where the Periodic
Table Gets Useful!!
• Elements in the
same group
have similar
chemical and
increasing
physical
properties!!
•
(Mendeleev did that on purpose.)
Why??
• They have the same
number of valence
electrons.
• They will form the same
kinds of ions.
• They increase in size
from smallest at top to
largest at bottom.
Families on the Periodic Table
• Columns are also grouped into families.
• Families may be one column, or several
columns put together.
• Families have names rather than
numbers.
• Most noted families of the periodic
table are
A separate slide for each follows
• Alkali Metals
• Alkaline Earth Metals
• Halogens
• Noble Gases
Hydrogen
• Hydrogen is an exception, as it
1
s
sometimes is shown with the
alkali metals family because of its
electron configuration and
sometimes with the halogens
because it needs one more to fill
its valence shell; it belongs to a
family of its own.
• Hydrogen is a diatomic element
and a reactive gas.
• In water, hydrogen will usually
make the +1 ion, but can make
the -1 ion.
Alkali Metals s1
• 1st column on the
periodic table (Group 1)
not including hydrogen.
• Very reactive metals,
always combined with
something else in nature
(like in salt).
• Soft enough to cut with a
butter knife.
• Makes the +1 ion.
Alkaline Earth Metals s2
• Second column on the
periodic table. (Group 2)
• Reactive metals that are
always combined with
nonmetals in nature.
• Several of these
elements are important
mineral nutrients (such
as Mg and Ca.
• Make the +2 ion.
Transition Metals s2 d1through10
• Elements in groups 3-12,
the d-block.
• Less reactive harder metals
• Includes metals used in jewelry
and construction.
• Metals used “as metal.”
• f-block is called the inner
transition metals. s2 f1through14
• Sometimes can make more
than one ion: called polyvalent.
Boron Family s2 p1
• Elements in group 13
• Aluminum metal was
once rare and expensive,
not a “disposable metal.”
• This family crosses the
metal /nonmetal line,
its family pattern is
not always apparent.
Carbon Family s2 p2
• Elements in group 14
• Contains elements
important to life and
computers.
• Carbon is the basis for
an entire branch of
chemistry.
• Silicon and Germanium
are important
semiconductors.
Nitrogen Family s2 p3
• Elements in group 15
• Nitrogen makes up over
¾ of the atmosphere.
• Nitrogen and phosphorus
are both important in
living things.
• The red stuff on the tip of
matches is phosphorus.
• Typically forms the -3 ion
in ionic compounds.
Oxygen Family or Chalcogens s2 p4
• Elements in group 16
• Oxygen is necessary for
respiration.
• Many things that stink,
contain sulfur (rotten eggs,
garlic, skunks,etc.)
• When in an ionic compound,
the -2 ion usually forms.
Halogens s2 p5
• Elements in group 17
• Very reactive, volatile,
diatomic, nonmetals
• Always found combined with
other element in nature .
• Used as disinfectants and to
strengthen teeth.
• When in an ionic compound,
halogens form the -1 ion.
The Noble Gases s2 p6
• Elements in group 18
• VERY unreactive, inert,
monatomic gases
• Used in lighted “neon” signs
• Used in blimps to fix the
Hindenberg problem.
• Have a full valence shell.
• Does not bond.
• Does not form an ion.
• Group 1 is called the alkali metals
and does not include hydrogen.
• Group 2 is called the alkaline earth metals.
• Groups 3-12 are called the transition metals
• Groups 13-15 are not usually named
• Group 16 are the chalcogens or oxygen family.
• Group 17 are the halogens
• Group 18 are the noble gases & are basically inert
To review:
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•
•
•
•
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Valence Shell by Group #:
Group 1
Group 2
Groups 13
Groups 14
Groups 15
Group 16
Group 17
Group 18
s1
s2
s2p1
s2p2
s2p3
s2p4
s2p5
s2p6
These groups are called the
“representative elements”.
The representative elements in
the s-block and p-block are
more predictable and have less
exceptions than the
d-block and f-block elements.
You must know these to be successful in
Unit 4 (e- configuration and periodic table) and Unit 5 (bonding)
•
•
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•
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Group 1
Group 2
Groups 13
Groups 14
Groups 15
Group 16
Group 17
Group 18
s1
s2
s2p1
s2p2
s2p3
s2p4
s2p5
s2p6
These are the valence shell
electron configurations that
we will be using repeatedly.
Write them down. Study them.
Know how they are used.
Periods and Groups
• Sizes of the atoms increase as we move
from top to bottom down a group.
• This is due to the increasing distance the
electrons are from the nucleus. Each time a new
principle energy level (shell) is added the orbiting
electrons are at a higher energy and are further
away from the nucleus.
Periods and Groups
• Sizes of the atoms decrease as we move
from left to right across a period.
• This is due to the increasing number of protons
in the nucleus, so the electrical attraction
between the nucleus and the orbiting electrons
gets stronger and pulls the electrons closer to
the nucleus.
• Trends within periods and groups – There is
usually a trend within a period (across, horizontal)
and a trend within a group (down, vertical)
• Atomic Radius trend – DECREASES ACROSS the
period, INCREASES DOWN the group…
Atomic Radius
• Atomic size is a periodic trend influenced by
electron configuration
• Electron clouds do NOT have a clearly definded
edge.
• The outer limit of an electron cloud is defined as a
spherical surface within which there is a 90%
probability of finding an electron.
• Atomic size is defined by how closely an atom lies
to a neighboring atom.
Electronegativity
• We can display the table to demonstrate other
properties as well
• As you move from left to right across a period,
the ability of the atom to attract electrons in a
shared bond increases
• This property is called electronegativity
Electronegativity
• The Electronegativity of an element is the relative
ability of it’s atoms to attract electrons in a chemical
bond.
• Trends – INCREASES as you go ACROSS THE
PERIOD, INCREASES as you go UP THE GROUP.
Ionization Energy
• Ionization energy is defined as the energy
required to remove an electron from a gaseous
atom…making the element a positive ion
• Ionization energy can give us an indication of
how strongly an atom’s nucleus holds onto its
valence electrons. More about valence electrons
later
• High Ionization energy values = nuclei’s with
strong holds on its electrons, which, in return,
Ionization Energy’s Periodic Trend
• Within the PERIOD, ionization energy
INCREASES as you ACROSS the row.
• Within the GROUP, ionization energy
INCREASES as you go UP the column.
Periodicity
• 3 Kinds of Periodic Trends…
• General trends are called periodicity. The
properties repeat in a periodic pattern.
• Atomic Radius, Electronegativity, Ionization
Energy, Metallic Character, Ionic Radius
and other properties, are periodic trends.
Know the reasons as well as the trends:
properties related to electron configuration
TRENDS, TRENDS, TRENDS!
Periodicity
• Left to right horizontal/across a period:
– greater positive charge of nucleus:
• Holds/attracts electrons in
the SAME energy level more tightly.
• Top to bottom vertical in a group:
– more energy levels, farther distance from nucleus:
• Attractive force is reduced,
electrons at HIGHER energy levels
more loosely held.
EXCEPTIONS: bumps and blips in
• trends
Metals to nonmetal trend often dictates type of
ion formed: positive cation vs. negative anion.
• Full sublevels are slightly more stable than
partially filled sublevels: s2 more stable than s1
• Half filled sublevels are slightly more stable than
partially filled sublevels: p3 more stable than p4
Both full s and p sublevels (s2p6 most stable, like noble gases)
have a “sheilding effect” dropping attractive force on
valence electrons. Abbreviated noble gas electron
configuration helps in recognizing/remembering this effect.
To Review
• Group 1 is called the alkali metals
and does not include hydrogen.
• Group 2 is called the alkaline earth metals.
• Groups 3-12 are called the transition metals
• Groups 13-15 are not usually named
• Group 16 are the chalcogens
• Group 17 are the halogens
• Group 18 are the noble gases & are basically inert
Lots of Questions
• Why are there only two elements in the first
period?
• Why are there eight elements in the second and
third periods?
• Why do we jump to eighteen elements across in
the fourth period?
• Why are the chemical properties of a group so
similar?
Lots of Questions
• What you learned about electron configuration
should answer the first three questions.
• Our next unit on bonding will address the last
question.
• This chapter will involve discovering the periodic
trends, making predictions based on the trends
and deducing the reasons behind this periodictity
• Read Chapter 14 for next week!