Download 5-1 Development of the Periodic Table

5-1 Development of the Periodic
Table
Why have a table?
• Chemists developed the Periodic Table to
help organize and classify the elements.
J.W. Dobereiner
• Early 1800’s
• Observed elements could be classified into
sets of three
–
–
–
–
Called Triads
Ex. Li, Na, K
Ex. Ca, Sr, Ba
Cl, Br, I
• Elements within a triad have similar
chemical properties.
– In each triad, the properties of the middle
element are close to the averages of the
properties of the 1st and 3rd element.
– Ca Sr Ba (40 + 137) ÷ 2 = 88
40 88 137
J.A.R. Newlands
• 1865
• By this time, 62 elements were known.
• He observed that when elements were
arranged in order of increasing atomic mass,
the properties of the 8th element were like
the 1st, 9th
– Law of Octaves
Dmitri Mendeleev
• 1869
• Like Newlands, he saw a periodic repetition
of properties when elements were arranged
by increasing atomic mass
– He developed the first periodic table
• Elements in the same column have similar
properties
• He also knew that not all elements were
known and left gaps in this table for these
undiscovered elements
– He predicted the properties of the undiscovered
elements, and it turns out he was correct.
• Example: Ekasilicon  Germanium
H.G.J. Moseley
• Remember Mendeleev organized his
periodic table in order of increasing atomic
mass.
– This was not quite correct, the elements are
arranged in order of increasing atomic number.
• Periodic Law – When elements are arranged
in order of increasing atomic number, their
physical and chemical properties show in a
periodic pattern.
• 5-1 Section Review
• 1. What is the periodic law?
• 2. How did Dobereiner and Newlands help
to develop the periodic table
• 3. How did Mendeleev demonstrate that his
table was valid? What would have helped
him develop the current PT?
• 4. What did Moseley contribute to the
development of the PT?
• Do # 5 on p.163
5-2 Reading the Periodic Table
Organizing Squares / Labeling
and Naming Groups
• Each square represents an element. It
contains its symbol, atomic number and
atomic mass. The table is read like a book,
from left to right, top to bottom, in order of
increasing atomic number. (Protons)
• http://viewpure.com/DYW50F42ss8
• The stability of the atom depends on the
electron configuration
• The outermost electrons determine the
chemical reactivity
– These are called Valence Electrons
Groups or Families
• vertical columns with elements that have
similar properties
• 18 total.
• The number with the “A” groups gives
the number of valence shell electrons for
all elements in the group.
• As you move down a group, the boiling
points and densities increase.
Groups or Families
Periods
• Horizontal rows 1-7.
• What does the row number tell you?
Metals, Nonmetals and
Semimetals (metalloids)
• Metals-left side of PT. Characteristicsluster or shine, conductors of heat and
electricity, usually solid at room temp,
malleable, and ductile. Group 1 and 2.
• Nonmetals-right side of PT. no luster,
non conductors, non malleable or ductile,
many are gases at room temp. Diverse
properties
• Metalloids (semimetals)-stair step from
boron. Some properties of metals and
nonmetals. Silicon is a metalloid, it is a
semiconductor and therefore used in
computer chips so you don't fry your
unit.
There are different “Blocks” on
the Table
S-Block
• Groups IA and IIA
• Outermost electrons of these elements are in sorbitals
• Group IA  1 valence electron (Alkali Metals)
• Group IIA  2 valence electrons (Alkaline
Earth Metals)
P- Block
• Groups IIIA through VIIIA
• Outermost electrons of these elements are in
the p – orbitals
• How many elements can each sublevel or
block hold?
• s=
p=
d=
f=
• How many electrons?
• The shape of the periodic table is a result of
the way the electrons fill the s, p, d, and f
orbitals.
5-2 Section Review
1. Why do elements within a group have similar
properties?
2. Color in a periodic table by orbitals (blocks),
indicate the s, p, d, and f orbitals.
3. Describe the general difference between the
elements on the right and the left side of the PT.
4. What information is present in each square of the
PT?
5. Circle the ions most like the ions of Br?
Be, Mg, F, O, C, B, I, Cr, F, K
6. Why would you expect the properties of those ions
circled to be similar to Br ion?
5-3 Periodic Trends
Atomic Radii
one-half the distance between the nuclei of
identical atoms that are bonded together.
• Atom size increases going down a group
because electrons are being added to higher
energy sublevels. (don’t worry about
memorizing exceptions).
Atomic Radii
• o Atom size decreases going across a periods left
to right.
– This is opposite what you would intuitively think.
• Why?
• As you move across a period, the atom gains
electrons as well as protons. Because the effective
nuclear charge has increased, the electrons will
therefore be pulled closer to the nucleus, causing
the atom to become smaller.
Ionic Radii – the size of an ion.
• Cation – a positive ion:Na+,Mg2+,Fe2+, Fe3+,
– Formed by the loss of an electron
– Cations are smaller than their neutral atom.
• As an electron is lost, the effective nuclear charge is
greater, causing the electron cloud to become pulled
closer to the nucleus, making it smaller
• Cations are pawsitive
• Anion – a negative ion: Cl-, F-, N3-, …
– Formed by the gain of an electron(s).
– As an electron is added, the nucleus remains
the same so the electrons are not drawn as
close to the nucleus. Also, added electron
causes repulsion among other electrons.
Ionization Energy
• Energy required to remove one electron
from an atom
• 1st ionization energy, 2nd ionization energy,
3rd ionization energy, …
A + energy  A+ + e-
• Ion – an atom or group of bonded atoms
that has a positive or negative charge.
– Ex. Na+, Cl-, O2-, Fe3+, NH4+, CO32-, …
– Monatomic ion – ion containing one atom
– Polyatomic ion – ion containing atoms bonded
together
Ionization Energy
• Noble Gases have the highest ionization
energies because they have a full octet
• Group IA have the lowest ionization
energies
– Ionization energy increases going across a
period left to right and decreases down a group
• Decreases down a group because electrons in outer
energy levels are easier to remove
• Increases across period because more electrons in
an energy level make it more difficult to get to the
magic number of 8
• It is possible to remove electrons from ions
– Ex. Cu+ + energy  Cu2+ + e– Each successive electron removal from an ion
feels an increasingly stronger effective nuclear
charge
• Look at trend on fig 5-20 pg. 180
Electron Affinity
• the energy change that occurs when an electron is
acquired by a neutral atom
• Most atoms release energy when they
acquire an electron
– A + e- → A- + energy
• If energy is released, use negative sign
• If energy is absorbed, use positive sign
• Common unit is kJ/mol
• Generally, the halogens gain electrons
more readily; therefore, their values are
most negative
– This fact helps to explain their great reactivity
• Notice electron affinities become more
negative going across period
• Generally, electron affinities become more
negative going up a group
– There are exceptions to this rule
• It is also more difficult to add an electron to
an already negative ion, so the second
electron affinities are all positive.
– Ex. Cl- takes less energy to form than Cl2-
Electronegativity
• a measure of the ability of an atom in a
chemical compound to attract electrons
– In a chemical compound, the more
electronegative atom will attract electrons
more
– Ex. HCl : H-Cl : electrons are pulled closer to
Cl
• The most electronegative atom on the
Periodic Table is Fluorine
• Electronegativity tends to increase across a
period
• Electronegativity tends to decrease down a
group