Transcript Periodic_Table

The Periodic Table
History of the Periodic Table
1) Doberiner - Doberiner’s triads
• Grouped together elements in groups of 3’s with
similar chemical properties.
(ex) Cl, Br, I
Ca, Ba, Sr
• Discovered that the atomic mass of one element
from the triad is close to the ave atomic mass of the
the other two elements from the triad.
• Found 4-5 triads
• First attempt to classify the elements in any way.
2) John Newlands
• First to lay things out horizontally by increasing at.
mass.
• When the properties of a series (row) began to
repeat themselves he started a new row.
• Noted that every 8th element was similar to the
first
• Law of Octaves: That when the elements were
placed in order by increasing atomic mass every 8th
element had similar properties to the 1st element.
• He obtained 7 groups of 7.
3) Sir William Ramsay - discovered the Noble
(Inert) gases in the 1890’s.
4) Demitri Mendeleev (1869)
• Consisted of rows & columns.
• Elements placed in order by increasing at. Mass
• Left to right and top to bottom.
• Placed elements with similar properties in the same
column.
• 8 columns wide.
• Sometimes the elements properties did not line up
properly. As a result he believed that not all the
elements had been discovered.
• Therefore, left gaps in his P. Table for the
undiscovered elements.
• Predicted the chemical properties of these
undiscovered elements and was very accurate in
doing so.
• Mendeleev’s Periodic Law: Properties of
elements are a periodic function of their at. mass.
(If you order the elements by increasing at. mass
the properties of the elements go thru a cycle and
repeat themselves).
• Also discovered a problem with the P. Table. If
he ordered the elements by increasing at. Mass
some of the properties of the elements were out of
order. (ex) Co and Ni
• But if he lined the elements up by properties not
all of the atomic masses were in order.
• Discovered the problem, but did not know what
to do about it.
5) Mosely
• 45 years later Mosely solved Mendeleev’s
problem with the properties not lining up when you
ordered the elements by increasing at. mass.
• Mosely was performing x-ray experiments that
showed the # of protons per nucleus varied
progressively from element to element.
• X-rays are a form of electromagnetic radiation.
• X-rays have high frequency and therefore short
wavelengths.
• X-rays are produced when high speed electrons
hit a metal target in an evacuated tube.
• Mosely discovered that the higher the atomic #
(Z) = #p, the shorter the wavelength of the x-ray.
• He found in some cases the wavelength was 2x
shorter than he expected - proved Mendeleev correct
in that some elements had not yet been discovered.
• When he listed elements by increasing at. # (Z)
instead of increasing at . mass the properties lined up.
• The Periodic Law - the properties of the elements
are a periodic function of their at. #.
• When you list the elements by increasing at. # the
properties of the elements go thru a cycle and repeat
themselves.
Arrangement of the Modern Periodic Table
• Horizontal row = series or a period
• Vertical column = Group or family.
Draw diagram on the board.
• Elements with similar properties have the same
number of outer shell e- (valence e-) and are in the
same Group on the P. Table.
•Name of elements on the P. Table.
•Group IA = 1 valence e- = alkali metals
•Group IIA = 2 valence e- = alkaline earths
Group IIIA = 3 valence e- = B family
Group IVA = 4 valence e- = C family
Group VA = 5 valence e- = N family
Group VIA = 6 valence e- = O family
• (Groups IIIA - VIA are named according to
the top element in the family.)
Group VIIA = 7 valence e- = halogens
Group VIIIA = 8 valence e- = Nobel or Inert
gases (octet = stable)
• No compounds containing He, Ne, and Ar are
known.
• A few compounds contain Kr, Xe, and Rn.
• Transition Elements = d sect. of P. Table
• Transition elements and the bottom of the p
sect. = heavy metals.
• 4f series = lanthanide series (La - Yb).
• 5f series = actinide series (Ac - No).
• f sect. = rare earths
Metals, nonmetals, and metalloids
•Metals - left of the P. steps
•Nonmetals - right of the P. steps.
•Metalloids - border the P. steps on 2 sides
(except Al)
• Most active metal is left and bottom. Also the
most basic.
• Most active nonmetal is right and top. Also the
most acidic.
• Metals tend to have 3 or fewer valance e-.
• Nonmetals tend to have 5 or more valence e-.
• Determining the Outer Most Electrons by
Position on the Periodic Table
(ex) 51 Sb
Steps:
1) Locate sect. = sublevel
2) count down = en level
3) count over = # of valence e(ex) 46 Pd
(ex) 74 W
(ex) 95 Am
Locating the element from the last subshell to
receive electrons
Steps:
1) sect.
2) en level
3) # valence e(ex) 4p4
(ex) 5d6
(ex) 5f4
Periodic Properties
•An element’s position on the periodic table and its
properties are a result of e- configuration.
•Atoms of elements in the same column have
similar outer e- configurations.
•The change in structure from one column to the
next as we scan the p. table varies in a set way.
•Properties can be predicted by e- configuration as
well as the position on the P. table.
Trends in Properties
1) Types of Elements
A) Series
metals -- metalloids -- nonmetals -- Noble gases
(border the steps)
B) Family
IA
VA
VIA
Metals
nonmetals
nonmetals
metalloids
metals
metalloids
Summary
nonmetals
metals -- metalloids -- nonmetals -- Noble gases
metals
2) Atomic Radius
A) Family (draw graph)
Atomic size increases as you go down a family.
Show examples down a family.
• Generally as you go down a family, each element
has one more shell or en level than the element
above it. An increase in the number of en levels
means an increase in the distance from the nucleus
making the atoms larger.
• Due to an increase in the number of en levels
there is also an increase in the screening or
shielding effect (when inner en levels block the
nuclear pull on outer en level e-). Since the e- are
not held as tightly the size of the atom increases.
• Atoms are becoming more metallic and therefore
want to lose e- and do not hold them as tightly
making the atom larger.
•B) Series (draw graph)
• Generally there is a decrease in size across a
series. Each element has a greater (+) charge
which results in a greater total force of attraction
between the (+) nucleus and the (-) electrons.
• There are more e- added also but no new en
levels. The increase in p overcompensates for the
increase in e-.
• Atoms become more nonmetallic want to gain
e- so hold onto their existing e- very tightly.
• Becoming more stable so hold the e- they
already have very tightly.
• Group VIIIA gets bigger, but not as big as Group
IA.
• It is believed that Group VIIIA gets bigger due to
the fact that it is stable.
•Summarize!
Radi of Ions
• Which is bigger a Na atom or an Na ion and give
all reasons why? Explain on the board.
• Which is bigger a Cl atom or a Cl ion and give all
reasons why? Explain on the board.
• Metallic atoms lose e- to from smaller ions.
• Nonmetallic atoms gain e- to form larger ions.
3) First Ionization Energy (I.E.)
• The amount of energy needed to remove 1 efrom an atom.
Atom + en -----> cation + 1e-
• 1st I.E. - en needed to remove the 1st e-.
• 2nd I.E. - en needed to remove the 2nd e-.
• 3rd I.E. - en needed to remove the 3rd e-.
• Write the I.E. for Na (do this on the board)
A) Series (draw graph)
How? I.
B) Family (draw graph)
Summary
Ionization energies to remove successive e• I.E. for success e- increases because as you
move more e- there will be less e- than p resulting
in a greater nuclear pull (the e- are held more
tightly) therefore, requiring more en to rip e- off.
(ex) write the 3 I.E’s for Al (put this on the
board)
4) Electron Affinity (E.A.)
• An element’s desire for additional electrons.
• The amount of en released when you add an
e-.
Atom + 1 e- ------> en + Anion
A) series
How?
E. A. increases as you go across a series.
Why?
•The atoms are becoming more nonmetallic,
therefore they want to gain e-.
•The atoms are becoming smaller - strong nuclear
pull.
•The atoms are becoming closer to being stable and
gaining e- will make them even closer.
B) Family
How?
E.A. decreases down a family
Why?
• Atoms are becoming more metallic - want to lose
e- not gain.
• Atoms are becoming bigger, decrease in nuclear
pull - can’t hold their own e- very tightly so do not
want more.
• More en levels therefore atoms are larger and have
a larger screening/shielding effect which decreases
the nuclear pull - own e- are not being held tightly -
5) Electronegativity
An atom’s desire to share e-.
A) series
How?
Electronegativity increases across a series.
Why?
The same reasons as for E.A.
B) Family
Electronegativity decreases down a family.
Why?
The same reasons as for E.A.
Summary of all Trends and Properties:
Summarize