Transcript AP Chemistry Chapter 7

Chemistry Chapter 5
The Periodic Law
Mendeleev’s Periodic Table
Dmitri Mendeleev
Mendeleev – organized periodic
table
• Vertical columns in atomic mass order
• Made some exceptions to place elements in rows
with similar properties (Te and I)
• Horizontal rows have similar chemical properties
• Gaps for “yet to be discovered” elements
• Left questions: why didn’t some elements fit in
order of increasing mass? Why did some elements
exhibit periodic behavior?
Moseley
• Discovered that periodic table was in
atomic number order, not atomic mass order
• Explained the Te-I anomaly
Periodic Law
• Physical and chemical properties of the
elements are periodic functions of their
atomic numbers
Modern Periodic Table
• Discovery of noble gases yields new family
(Group 18 – aka inert gases)
• Lanthanides (#58 - #71)
• Actinides (#90 – #103)
Periods and Blocks of the Periodic
Table
• Periods – horizontal rows
• Groups/Families – vertical columns; these
elements share similar chemical properties
(they have the same number of valence
electrons)
• Blocks – periodic table can be broken into
blocks corresponding to s, p, d, f sublevels
Orbital filling table
Blocks and Groups – s block
• Group1 – “The alkali metals”
• One s electron in outer shell
• Soft, silvery metals of low density and low
melting points
• Highly reactive, never found pure in nature
Blocks and Groups – s block
• Group 2 – “Alkaline Earth Metals”
• 2 s electrons in outer shell
• Denser, harder, stronger, less reactive than
Group 1
• Too reactive to be found pure in nature
Periodic Table with Group Names
The Properties of a Group:
the Alkali Metals
Easily lose valence electron
(Reducing agents)
React violently with water
React with halogens to form salts
Blocks and Groups – d block
•
•
•
•
Groups 3 -12
Metals with typical metallic properties
Referred to as transition metals
Group number = sum of outermost s and d
electrons
Properties of Metals
 Metals are good
conductors of heat and
electricity
 Metals are malleable
 Metals are ductile
 Metals have high tensile
strength
 Metals have luster
Examples of Metals
Potassium, K
reacts with
water and
must be
stored in
kerosene
Copper, Cu, is a relatively soft
metal, and a very good electrical
conductor.
Zinc, Zn, is
more stable
than potassium
Mercury, Hg, is the only
metal that exists as a
liquid at room temperature
Blocks and Groups – p block
• Groups 13-18
• Properties vary greatly – metals, metalloids,
and nonmetals
• Group 17 – halogens are most reactive of
non metals
• Group 18 – noble gases are NOT reactive
Properties of Nonmetals
Carbon, the graphite in “pencil lead” is a great
example of a nonmetallic element.
 Nonmetals are poor conductors of heat and
electricity
 Nonmetals tend to be brittle
 Many nonmetals are gases at room
temperature
Examples of Nonmetals
Sulfur, S, was
once known as
“brimstone”
Graphite is not the only
pure form of carbon, C.
Diamond is also carbon;
the color comes from
impurities caught within
the crystal structure
Microspheres
of phosphorus,
P, a reactive
nonmetal
Properties of Metalloids
Metalloids straddle the
border between metals
and nonmetals on the
periodic table.
 They have properties of both metals and
nonmetals.
Metalloids are more brittle than metals, less
brittle than most nonmetallic solids
 Metalloids are semiconductors of electricity
 Some metalloids possess metallic luster
Silicon, Si – A Metalloid
 Silicon has metallic luster
 Silicon is brittle like a nonmetal
 Silicon is a semiconductor of
electricity
Other metalloids include:





Boron, B
Germanium, Ge
Arsenic, As
Antimony, Sb
Tellurium, Te
Blocks and Groups – f block
• Lanthanides – shiny metals similar to group
2
• Actindes – all are radioactive; plutonium –
lawrencium are man-made
Determination of Atomic Radius:
Half of the distance between nucli in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to decreased shielding
Radius increases down a group
Addition of principal quantum levels
Table of
Atomic
Radii
Ionization Energy - the energy required to remove an
electron from an atom
Increases for successive electrons taken from
the same atom
Tends to increase across a period
Electrons in the same quantum level do
not shield as effectively as electrons in
inner levels
Irregularities at half filled and filled
sublevels due to extra repulsion of
electrons paired in orbitals, making them
easier to remove
Tends to decrease down a group
Outer electrons are farther from the
nucleus
Ionization of Magnesium
Mg + 738 kJ  Mg+ + eMg+ + 1451 kJ  Mg2+ + eMg2+ + 7733 kJ  Mg3+ + e-
Table of 1st Ionization Energies
Another Way to Look at Ionization
Energy
Ionic Radii
Cations
Anions
Positively charged ions
Smaller than the corresponding
atom
Negatively charged ions
Larger than the corresponding
atom
Table of Ion Sizes
Electronegativity
A measure of the ability of an atom in a chemical
compound to attract electrons
Electronegativities tend to increase across
a period
* more nuclear charge, more power to
attract electrons
Electronegativities tend to decrease down
a group or remain the same
* additional energy levels result in less
attraction to the nucleus
Periodic Table of Electronegativities