Transcript chemistry chapter 11 & 12

CHEMISTRY CHAPTER 11
& 12
THE ELEMENTS OF THE
PERIODIC TABLE
1
Organization of the Elements
•
•
•
•
Metals vs. non-metals.
Listing according to mass.
Listing according to properties
Listing according to Chemistry
– MO ratio; reactions with Hydrogen,
halogens etc.
Organization of the Elements
• Dmitri Mendeleev
(1860)
– Established the Periodic Table
– Stated the Periodic Law:
– “Elements were placed on the table
according to their atomic masses, with
columns established by periodic
recurring similar properties.”
Organization of the Elements
Mendeleev's Periodic Table
•Vertical columns in atomic weight
order
•Mendeleev made some exceptions
to place elements in rows with
similar properties
•Tellurium and iodine's places were
switched
Organization of the Elements
•Horizontal rows have similar
chemical properties
•Missing Elements
-Gaps existed in Mendeleev’s
table
•Mendeleev predicted properties of
the “yet to be discovered”
elements
•Scandium, Germanium and Gallium
agreed with predictions
Unanswered Questions
1. Why didn't some elements fit in
order of increasing atomic mass?
2. Why did elements exhibit
periodic behavior?
Moseley and the Periodic Table (1911)
•Protons and Atomic Number
•X ray experiments revealed a way to
determine the number of protons
in the nucleus of an atom
•The periodic table was found to be
in atomic number order, not
atomic mass order
• The tellurium-iodine anomaly was
explained
The Periodic Law
•The physical and chemical
properties of the elements are
periodic functions of their
atomic numbers
The Modern Periodic Table
•Discovery of the Noble Gases
1868 - Helium discovered as a
component of the sun, based on the
emission spectrum of sunlight
1894 - William Ramsay discovers argon
1895 - Ramsay finds helium on Earth
1898 - Ramsay discovers krypton and
xenon
1900 - Freidrich Dorn discovers
radon
The Lanthanides
Early 1900's the elements from
cerium (#58) to lutetium (#71) are
separated and identified
Organization of the Elements
The Modern Periodic Table
– Listed according to atomic
numbers.
Organization of the Elements
• The Modern Periodic Table
– Vertical columns = Families
• Group 1 (1A) = Alkali Metals
• Group 2 (2A) = Alkaline Earth
Metals
Organization of the Elements
•Group 17 (7A) Halogens
•Group 18 (8A) Noble Gases
•Group 3-12: Transitional
Metals
•Horizontal Row: Periods
Periodic Table and econfiguration
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•
•
•
•
Alkali - s1
Alkaline Earth - s2
Halogens - p5
Noble gases - p6
Valence Electrons= electrons in outer
most shell.
– The outermost shell electrons are the
most easily removed.
Periodic Table and econfiguration
– The outermost (valence) electrons.
– In the outermost s and p orbitals.
• What is the maximum number of
valence orbitals?
FOUR
• What is the maximum number of
valence electrons?
EIGHT
Periodic Trends
• Metallic properties:
– Decrease across a period.(left to right)
– Increase going down a column.
• Ionization Energies:
– Increase across a period. .(left to
right)
– Decrease going down a column.
Ionization Energy
•The energy required to remove
one electron from a neutral atom
of an element, measured in
kilojoules/mole (kJ/mol)
A + energy -> A+ + e-
Ionization Energy Trends
•Ionization
energy of main-group elements
tends to increase across each period (Left to
Right)
•Atoms are getting smaller, electrons are closer
to the nucleus
•Ionization energy of main-group elements
tends to decrease as atomic number increases in
a group
Ionization Energy Trends
•Atoms are getting larger, electrons are
farther from the nucleus
•Outer electrons become increasingly more
shielded from the nucleus by inner
electrons
•Metals have a characteristic low ionization
energy
•Nonmetals have a high ionization energy
•Noble gases have a very high ionization
energy
Ionization
Energy
Decreases
Metallic
Properties
Increase
Metallic Ionization
Properties Energy
Decrease Increases
Periodic Trends
• Atomic radii:
– Decrease across a period.
– Increase going down a column.
•Atomic Radius
-One half the distance between nuclei
of identical atoms that are bonded
together
Ionic Radii
Ionic radii:
–Positive: Decrease across a period.
Predominately Metals
–Negative: Decrease across a period.
Predominately non-metals
Cations
-Positively charged ions
-Smaller than the corresponding atom
-Less shielding of electrons
Ionic Radii
Anions
-Negatively charged ions
-Larger than the corresponding atom
-Greater electron-electron repulsion
-Ion size tends to increase downward within
a group
This table
shows the
relative
sizes of
atoms
based on
increasing
ATOMIC
RADIUS
Size is dependent on energy level and
electrons filling level. As you fill energy
level across the period, the atomic radius
decreases
Argon fills the 3p orbitals and
Potassium starts the 4s
(larger sized) orbital. Calcium
has more electrons filling level
four shell which shrinks the
atom size (negative shell
attracts to positive nucleus
As Atoms lose electrons the radius of
the ion gets smaller
As elements receive electrons the
radius of the ion gets larger
Atomic Radii
Increase
Atomic
Radii
Decrease
Electronegativity
• A measure of the ability of an atom in a
chemical compound to attract electrons
• Electronegativities tend to increase across
a period
• Electronegativities tend to decrease down
a group or remain the same
Example: F(4.0) vs. Cl(3.0), O(3.5) vs. Ba(0.9)
• Elements that do not form compounds are
not assigned electronegativities
Table of Electronegativities
Electronegativity Trends
•Nonmetals have characteristically
high electronegativity
•Highest in the upper right corner
•Metals have characteristically low
electronegativity
•Electron
Electron Affinity
Affinity: The energy change that
occurs when an electron is acquired by a neutral
atom, measured in kJ/mol
•Most atoms release energy when they acquire
an electron
A + e- -> A- + energy (exothermic)
•Some atoms must be forced to gain an electron
A + e- + energy -> A - (endothermic)
Electron Affinity Trends
1. Halogens have the highest electron
affinities
2. Metals have characteristically low
electron affinities
3. Affinity tends to increase across a
period
a. Irregularities are due to the extra
stability of half-filled and filled
sublevels
Electron Affinity Trends
4. Electron affinity tends to
decrease down a group
5. Second electron affinities are
always positive (endothermic)
Chemical Families
• Alkali Metals:
– Very reactive metals
– All form 1+ ions and are found in
nature as ions. Never found pure in
nature
– Sodium and Potassium are the most
common in the family
- Soft, silvery metals of low density and
low melting points
- Highly reactive,
The Properties of a Group:
the Alkali Metals
• Easily lose valence electron
(Reducing agents)
•React violently with water
Large hydration energy
•React with halogens to form
salts
Chemical Families
– React readily with halogens to form
common salts
• Example: NaCl (table salt)
– React readily with water to form
basic solutions (alkali), hydrogen and
Energy.
Chemical Families
• Alkaline Earth Metals
– 2nd most reactive metallic family.
– All form 2+ ions.
- Denser, harder, stronger, less
reactive than Group 1
- Too reactive; not found pure in
nature
– Mg and Ca are the most common
family members.
Chemical Families
• Halogens Group 17
– Very reactive non-metals, each having
SEVEN valence electrons.
– Valence Electrons: The electrons
available to be lost, gained, or shared
in the formation of chemical
compounds (outer most)
– Question: In what orbitals are the
valence electrons?
s and p
Chemical Families
• Halogens React readily with
Alkali metals to form salts
(salt formers).
• React with hydrogen to form
hydrogen halides (strong acids).
Chemical Families
• Halogens Group 17 are:
– Natural occurring as diatomic
molecules.
– all COLORED.
– all (1-) ions.
Chemical Families
• Noble Gases: Group 18
–Also known as “inert gases,”
they are characteristically
unreactive.
Noble gases have EIGHT valence
electrons.
Chemical Families
• Noble Gases: Group 18
– First found to react in 1962,
but formed unstable
compounds.
Chemical Families
• Transitional Metals
– Columns 3-12
– Have more than one Oxidation state.
– Most common: Iron
– Includes “precious metals”
– Compound often highly colored.
Chemical Families
– Gemstones most often are transitional
metal compounds.
• Metaloids: In between metals and
non-metals
– Include semiconductors
• Non-Metals: Carbon, Oxygen,
Nitrogen families
Chemical Families
– Allotropic forms are known of both
metaloids and non-metals: same
element having different structures.
• Inner transitional elements: (Rare
Earth Elements)
– Lanthanide series (Atomic # 57-70)
– Actinide series (#89-102)
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
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
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
Determination of Atomic Radius:
Ionization of Magnesium
Mg + 738 kJ  Mg+ + eMg+ + 1451 kJ  Mg2+ + eMg2+ + 7733 kJ  Mg3+ + e-
Affinity tends to increase across a period
Affinity tends to decrease as you go down
in a period
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals
Electron Affinity - the energy change associated
with the addition of an electron
Ionization
Energy
Decreases
Metallic
Properties
Increase
Metallic Ionization
Properties Energy
Decrease Increases
5-2 Electron Configuration and the
Periodic Table
I. Periods and the Blocks of the
Periodic Table
A. Periods
1. Horizontal rows on the periodic table
2. Period number corresponds to the
highest principal quantum number of the
elements in the period
B. Sublevel Blocks
1. Periodic table can be broken into
blocks corresponding to s, p, d, f
II. Blocks and Groups
A. s-Block, Groups 1 and 2
1. Group 1 - The alkali metals
a. One s electron in outer shell
b. Soft, silvery metals of low density and lo
c. Highly reactive, never found pure in natur
2. Group 2 - The alkaline earth metals
a. Two s electrons in outer shell
b. Denser, harder, stronger, less reactive t
c. Too reactive to be found pure in nature
B. d-Block, Groups 3 - 12
1. Metals with typical metallic properties
2. Referred to as "transition" metals
3. Group number = sum of outermost s and
C. p-Block elements, Groups 13 - 18
1. Properties vary greatly
a. Metals
(1) softer and less dense than d-block meta
(2) harder and more dense than s-block me
b. Metalloids
(1) Brittle solids with some metallic and
properties
(2) Semiconductors
c. Nonmetals
(1) Halogens (Group 17) are most react
D. f-Block, Lanthanides and Actinides
1. Lanthanides are shiny metals similar i
2. Actinides
a. All are radioactive
b. Plutonium (94) through Lawrencium (10
5-3 Electron Configuration and Periodic P
V. Valence Electrons
A. Valence Electrons
1.
2. Main group element valence electrons are outermost
energy level s and p
sublevels
Group # 1 2 13 14 15 16 17 18
Number of valence Electrons 1 2 3 4 5 6 7 8
Periodic Trends
– Positive: Increase down a column.
– Negative: Increase down a
column.
B.
1.
C.
1.
a.
2.
Ionization Energy
Tends to increase across d- and f-B
Ion Formation and Ionic Radii
Electrons are removed from the out
Most d-block elements form 2+ ions
Ions of d- and f-Blocks are cations
D. Electronegativity
1. Characteristically low
electronegativity of metals
2. Electronegativity increases as
atomic radius decreases
B. Trends
1. Atomic radius tends to decrease acr
2. Atomic radii tend to increase down a
VII. Periodic Properties of the d- and fBlock Elements
A. Atomic Radii
1. Smaller decrease in radius across a
period within the d- Block than within
the main-group elements
a. Added electrons are partially shielded
from the increasing positive
nuclear charge
b. Slight increase at the end of the dBlock is due to electron-electron
repulsion
2. Little change occurs in radius across