Download The Periodic Table

Periodic Table
Chapter 5
History
Organization
Introduction to Bonding
Trends
Terms
• The Periodic Law
States that the physical & chemical properties
of the elements are periodic functions of
their atomic numbers.
• The Periodic Table
An arrangement of the elements in order of
their atomic numbers so that elements with
similar properties fall in the same columns.
• The Modern Periodic Table
Is organized by electron configuration.
Periodic Table
& electron configurations
Blocks of the Periodic Table
s
p
d
f
Video Clip: Periodic Table by
cassiopeiaproject
• https://youtu.be/5MMWpeJ5dn4
Extended Periodic Table
History
1817-Dobereiner’s Triads
Groups of 3 elements with similar
properties, the atomic mass of second was
an average of all three.
EX:
Ca = 40
Sr = 88
Ba = 137
History
1863- Newland’s Law of Octaves
If elements are arranged by increasing
atomic mass, their properties repeat
every 8 elements.
Chapter 5 – Section 1: History of the Periodic Table
Mendeleev and Periodicity
• The first periodic table of the
elements was published by
Russian chemist Dmitri Mendeleev.
• Mendeleev left empty spaces in his
table and predicted
elements that would fill
3 of the spaces.
• By 1886, all 3 of these
elements had been
discovered.
History
1869- Mendeleev’s First Periodic Table
He arranged all the known elements in rows
by increasing mass AND he grouped
elements with similar properties in
columns.
Brilliant in spite of discrepancies: Co & Ni, Te & I
Brilliant because of his predictions:
ekasilicon = Ge
Video Clip: Periodic Table of
Elements - Chemistry: A Volatile
History - BBC
• https://youtu.be/nsbXp64YPRQ
Chapter 5 – Section 1: History of the Periodic Table
Mosley and the Periodic Law
• In 1911, the English scientist
Henry Moseley discovered that
the elements fit into patterns
better when they were arranged
according to atomic number,
rather than atomic weight.
• The Periodic Law states that the physical
and chemical properties of the elements are
periodic functions of their atomic numbers.
History
1913- Moseley discovers protons
Nuclear charge (protons) increases by one
for each element!
Atomic number, not atomic mass, is the
basis for organization!
Organization
Groups, Periods and Blocks
• Elements in the periodic table are arranged
into vertical columns, called groups or
families, that share similar chemical
properties.
• Elements are
also organized
horizontally
in rows,
or periods.
GROUPS
• 18 vertical columns (families) with family
names
• same outer e- configuration.
• similar chemical properties.
• labeling columns has changed over time…
Labeling Groups… then & now
PERIODS
• 7 horizontal rows
• corresponds to main Energy level that is
filled
• length of row determined by sublevels
that are filling
Video Clip: The (truly) Periodic
Table by scienceoffice.org
• https://youtu.be/xd4-Uy2FLWc
BLOCKS
The representative elements of the
Main Groups: s & p blocks only
Group 1: Alkali Metals
• Group 1 elements are called alkali metals.
• Alkali metals have a silvery appearance
and are soft enough to cut with a knife.
• They are extremely reactive and are not
found in nature as free elements.
• They must be stored under oil or kerosene.
Group 2: Alkaline Earth Metals
• Elements in group 2 are known
as the alkaline earth metals.
• Group 2 metals are harder, denser
and stronger than alkali metals,
and have higher melting points.
• Less reactive than
group 1, but still too
reactive to be found
in nature as free
elements.
Group 17: Halogens
• Elements in group 17 are
known as the halogens.
• Halogens are the most
reactive nonmetals,
reacting vigorously with
metals to form salts
•Most halogens exist in
nature as diatomic molecules
(i.e. F2, Cl2, Br2 and I2.)
Group 18: Noble Gases
• Elements in group 18 are
known as noble gases.
• They are completely nonreactive and don’t form
compounds under normal
conditions.
• A new group was added to
the periodic table in 1898
for the noble gases.
d-block: Transition Metals
• Elements in the d-block are
called transition metals.
• They have typical metallic
properties such as conduction
of electricity and high luster.
• Less reactive than group 1 and 2 elements.
• Some (i.e. platinum & gold) are so unreactive
they usually don’t form compounds.
f-block: Lanthanides & Actinides
• Elements in period 6
of the f-block are called
lanthanides (or rare-earth).
• Lanthanides are shiny
metals similar in reactivity
to alkaline earth metals.
• Elements in period 7 of the
f-block are called actinides.
• Actinides are all radioactive,
and many of them are known
only as man-made elements.
Classification & Location
of the elements
Metals: left of staircase
Nonmetals: right of staircase
Semimetals: touch the staircase (except Al)
Why are H & He
sometimes
separated from
Periodic Table?
• Hydrogen has one valence
electron so it is normally
placed in group 1.
• Hydrogen has a second
home at the top of group
17 because it can gain an
electron as the halogens
do.
• Helium has two valence
electrons like group 2 but
is clearly a noble gas
because its outer energy
level is full.
Introduction to Bonding
PHYSICAL PROPERTIES OF METALS
luster, conductivity, solid, malleable,
ductile
3 or fewer valence e- = defines a metal
Introduction to Bonding
CHEMICAL PROPERTIES OF
METALS….HOW THEY FORM BONDS
Lose valence e- to nonmetals → Ionic
compounds
Bond with each other → Metallic bonds
Mix with other metals → Alloys
Introduction to Bonding
PHYSICAL PROPERTIES OF NONMETALS
No luster, poor conductors, not malleable
or ductile
5 -8 valence e- = defines a nonmetal
Can be soft solids: C, P, S, Se, I2
Can be liquids: Br2
Can be gases: N2, Cl2, H2, O2, F2, & the
noble gases
 All gases are nonmetals but not all nonmetals are gases!
Introduction to Bonding
CHEMICAL PROPERTIES OF NONMETALS …. HOW
THEY FORM BONDS
Gain valence e- from metals → Ionic
compounds (EX: salt: NaCl)
Share valence with themselves→ Diatomic
Molecules (EX: N2, H2 )
Share valence with other nonmetals →
Molecules (EX: H2O, CO2)
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Atomic Radii
• Atomic radius – one-half the distance
between the nuclei of identical atoms
that are bonded together.
Group 1
•Atomic radii tend to increase
as you go down a group
because electrons occupy
successively higher energy
levels farther away from the
nucleus.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Atomic Radii (continued)
• Atomic radii tend to decrease as you go
across a period because as more electrons
are added they are pulled closer to the more
highly charged nucleus with more protons.
Period 2
Atomic Radius (Size)
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Atomic Radii
Sample Problem
Of the elements Mg, Cl, Na, and P, which has
the largest atomic radius? Explain.
Solution:
Na has the
largest radius.
All of the
elements are
in the 3rd period,
and atomic radii
decrease across
a period.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Ionization Energy
• An ion is an atom of group of bonded atoms
that has a positive or negative charge.
• The energy required to
remove an electron from
a neutral atom of an
element is called the ionization energy (IE).
• Ionization energy tends to increase across
each period because a higher nuclear charge
more strongly attracts electrons in the same
energy level.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Ionization Energy (continued)
• Ionization energy tends to decrease down
each group because electrons farther from
the nucleus are removed more easily.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Ionization Energy
Sample Problem
Consider two elements, A and B. A has an IE of 419
kJ/mol. B has an IE of 1000 kJ/mol. Which element is
more likely to be in the s block? Which will be in the
p block? Which is more likely to form a positive ion?
Solution:
Element A is most likely to be in the s-block since IE
increases across the periods.
Element B would most likely lie at the end of a period
in the p block.
Element A is more likely to form a positive ion since it
has a much lower IE than B.
Ionization Energy
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Electron Affinity and Electronegativity
• Electron affinity is the energy change that
occurs when an electron is acquired by a
neutral atom.
• Electronegativity is a measure of the ability
of an atom in a chemical compound to
attract electrons from another atom in the
compound.
• Electronegativity applies to atoms in a
compound, while electron affinity is a
property of isolated atoms.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Electron Affinity and Electronegativity
(continued)
• Electron affinity and electronegativity both
tend to increase across periods, and
decrease (or stay the same) down a group.
Electron Affinity
ELECTRONEGATIVITY
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Electronegativity
Sample Problem
Of the elements Ga, Br, and Ca, which has
the highest electronegativity? Explain .
Solution:
All of these elements are in the fourth period.
Br has the highest atomic number and is
farthest to the right in the period.
Br would have the highest electronegativity
since electronegativity increases across
a period.
Periodic Trends
• Atomic Radius, and
Metallic Character
• Factors affecting size:
Distance of outer e- from
nucleus, shielding effect of
inner e- and size of the ecloud  with each sublevel
makes atoms larger down a
group.
Nuclear charge  across the
period pulls e- cloud in and
makes the atoms smaller.
Periodic Trends
• Ionic Radius
• Factors affecting size:
Metals lose valence e- the
ion is smaller =
previous noble gas!
Nonmetals gain valence eion is larger =
next noble gas!
Periodic Trends
First Ionization Energy:
E req’d to remove outer eElectron Affinity:
Atom’s attraction for e-
Metals
low IE
low e- affinity
Nonmetals
high IE
high e- affinity
Periodic Trends-Bonded
Atoms
Electronegativity:
The ability of an atom to
attract e- in a chemical
bond.
Fluorine is assigned
highest value = 4.0
Nonmetals
high EN
high e- affinity
Metals
low EN
low e- affinity
START CH.7
valence electrons
• Occupy highest principal energy level
• Responsible for chemical properties =
bonding
• Elements in a group have similar
properties b/c valence e- same.
• Outer E level has e- in s & p sublevels
only
• Outer E level full when s & p sublevels
are full = noble gas configuration
Ion formation in atoms
 Ion:
an atom that has a charge b/c it has lost or
gained e The Octet (Duet) Rule:
atoms will gain, lose or share e- in order to
acquire a full set of valence e- = 8
 The Noble Gas Rule:
atoms attain the nearest noble gas configuration
when they become ions.
Types of Ions
 CATIONS are positive ions
 ANIONS are negative ions
 Metals form monatomic cations
 To name: use name of element
Na+1
Mg+2
magnesium ion…)
Al+3
Si+4
(sodium ion,
Types of Ions
 Nonmetals form monatomic anions
 To name: use name of element + ide
N-3
O-2
F-1
(nitride ion, oxide ion,
P-3
S-2
Cl-1
(phosphide, sulfide, chloride)
fluoride ion)
Transition Metals: d block
• 1 or 2 valence e• Several of these metals form more than one type
of ion. (MULTIPLE OXIDATION STATES)
• How? Will lose the valence s e- then the d e- one
at a time.
• To name, use roman numerals
Cu+1
ion
Cu+2
ion
ion
Fe+2
iron II
copper II ion
Fe+3
iron III
copper I
Lower left of p block = metals
• 3 to 5 valence e• Several of these metals form more than one type
of ion.
• How? Will lose all the valence s & p electrons
OR just the p electrons.
• To name, use roman numerals
Tl+1
ion
Tl+3
thallium I ion
Sn+2 tin II
thallium III ion
Sn+4 tin IV
Types of Ions
 Polyatomic ions: groups of covalently
bonded atoms that carry a charge.
NH4+1
ammonium ion
SO4-2
sulfate ion
OH-1
hydroxide ion
CN-1
cyanide ion
PO4-3
phosphate ion
CO3-2
carbonate ion
NO -1
nitrate ion