Download Chapter 6 lecture 2013

Chapter 6: The Periodic Table
Elements
 Science has come a
long way since
Aristotle’s theory of Air,
Water, Fire, and Earth
 Scientists have
identified 90 naturally
occurring elements, and
created about 28 others
A Little Bit of the History..
 He originally
organized the periodic
table by atomic mass
Dmitri Mendeleev
Could not make a
complete table, only had
63 elements leaving
many spaces between
elements
Used properties of other
elements to predict
undiscovered elements
properties
A Little Bit of the History..
Mendeleev’s original Periodic Table
The Periodic Law
• In the modern periodic table, elements are arranged in
order of increasing atomic number.
• Organized into columns and rows.
• 7 horizontal rows correspond to the energy levels
found outside an atom’s nucleus.
The Periodic Law
• Horizontal rows called periods
• Maximum of 7 periods that represent the energy
levels of the atom
• Every element in the first period has ONE energy level
• Second period – TWO energy levels
• Third period – THREE energy levels
The Periodic Law
•
•
•
•
Columns known Groups or Families
Elements in the same column have similar properties
Group 1 Alkali Metals, 1 valence electron, very reactive
Group 2 Alkaline Earth Metals, 2 valence electrons
The Periodic Law
• Farthest column to right: Noble Gases
• Each element has 8 valence electrons
• The elements that are truly inert: Helium and Neon
The Periodic Law
•
•
•
•
Column to the left of noble gases: Halogens
Each element has 7 valence electrons
Transition metals: in the “d block”
Rare Earth metals: in the “f block”
How are Elements Classified?
Three Regions
How are Elements Classified?
Three Regions: Representative, Transition, and Rare Earth
Properties of Metals
Metals are:
Good conductors of heat
and electricity
Shiny
Ductile (can be stretched
into thin wires)
Malleable (can be pounded
into thin sheets)
Mostly solid at room temp
(except Hg)
Properties of Metals
Metals are not found in
their pure form in nature
Found in ore, mixed with
other elements.
Must be refined to get the
pure metal.
Let’s Look at the Families
•
•
•
•
Alkali
Alkali Earth
Halogen
Noble Gases
Alkali Metals
 Group 1A (Not including hydrogen)
 Very reactive metals
 Always combined with something else in
nature (like salt)
 1 valence electron
 Malleable, ductile, good conductors of
electricity
 Soft enough to cut with a butter knife
 Two most reactive elements: Cs and Fr
Explodes in water
http://www.youtube.com/watch?v=92Mfric7JUc
Li
Na
K
Rb
Cs
Fr
Alkaline Earth Metals
Be
 Group 2A
Mg
 Reactive metals that are
always combined with
nonmetals in nature
Ca
 2 valence electrons
Ba
 Several of these are
important mineral
nutrients (Mg and Ca)
Ra
Sr
Transition Metals
 38 elements
 Valence electrons: The
electrons used to combine
with other elements.
 Valence electrons present
in more than one shell.
We’ll deal with oxidation states next chapter.
 Iron, Cobalt and Nickel
produce magnetic fields
d block
Other Metals
(We won’t include this as a separate category)
 7 elements (Al, Ga, In Sn,
Th, Pb, Bi)
 Similar to transition
metals – Ductile and
malleable
 Different than transition
metals – valence
electrons only present in
outer shell
 3 physical properties –
solid, high density,
opaque
Metalloids
Have properties of both
metals and non-metals
Semi-conductor: Can carry
an electric charge better
than non-metals but not as
well as metals
 Silicon and Germanium
 Useful in computers and calculators
Silicon
Properties of Non-Metals
 Poor conductors of heat
and electricity, are not
ductile or malleable and
are brittle
 3 states of matter at room
temperature: Gas, solid or
liquid
 Have no luster and do not
reflect light
Sulfur
Halogens
F
Cl
 Are non-metals
 “Halogen” means salt
former. Compounds
containing halogens are
called salts
 NaCl – table salt
 7 valence electrons
 Liquid (Br), Gas (F, Cl),
Solid (I, At)
Br
I
At
The Noble Gases
He
Ne
Ar
VERY stable gases
because they have the
maximum number of
electrons in their outer
shell
 Helium – 2 valence electrons
 All others – 8 valence electrons
Used in lighted “neon”
signs
Used in blimps to fix the
Hindenberg problem
Kr
Xe
Rn
Rare Earth Metals
 30 elements in the
f- block
 Trans-uranium
means Man-Made or
Synthetic
 One element in
Lanthanides and
most of Actinides
are man-made
 In periods 6 and 7
f block
lanthanides
actinides
OK, Let’s Review…
 What are some examples of metals?
Gold, silver, magnesium, lead, aluminum
 What is a metal that is liquid at room temp?
Mercury
 What are some examples of non-metals?
Oxygen, fluorine, nitrogen, sulfur
 What are some of the characteristics of the
metalloids ?
 Better conductors than non-metals, are shiny or dull, etc.
Periodic Trends
• So – we’ve learned about how the periodic table is
organized by….
– PERIODS (Rows):
• Each period represents elements with electrons in a different energy
level.
• Properties of elements change as you move across a period
– GROUPS/FAMILIES (Columns):
• Elements in the same group have similar properties
• For representative elements, group number (1A – 8A) = the number of
valence electrons (electrons in the s & p sublevels in the highest energy
level)
Periodic Trends
• Now we’re going to look at how element properties
change across a period or down a group:
“PERIODIC TRENDS”
• The periodic trends of properties that we will look at
are:
–
–
–
–
atomic size or atomic radius
ionic size or ionic radius
ionization energy
electronegativity
Trends in Atomic Size
• The atomic radius is the distance between the
nucleus and the outer edge of the electron cloud.
Trends in Atomic Size
• In general, atomic size:
– increases from top to bottom within a group and
– increases from right to left across a period.
Atomic radius vs. atomic number
Atomic Radius (pm)
250
K
200
Na
Li
150
Mg
Al Si
Be
100
Ca
P S Cl
B C N
O F
Ar
Ne
50
H
0
0
He
2
4
6
8
10
12
Element
14
16
18
20
Trends in Atomic Size
Ions
• Ions form when electrons are transferred between atoms.
• If an atom loses an electron: cation (positive charge)
• METALS ALWAYS LOSE (electrons)
Ions
• If an atom gains an electron: anion (negative charge)
• Non-metals always gain electrons
Trends in Ionic Size
• The size of ions:
– Cations (LOSE ELECTRONS) are always smaller than the
original atoms.
– Anions (GAIN ELECTRONS) are always larger than the original
atoms.
Trends in Ionic Size
•
Trends in Ionic Size: Size of anions and cations increase from right to left
Size generally increases
Positive Ions (cations)
Negative Ions (anions)
Trends in Ionization Energy
• Atoms hold onto its electrons with different “strengths”
• The energy required to remove an electron from an atom is
called ionization energy.
– Energy required to remove the first electron is called the first
ionization energy.
– Energy required to remove an electron from an ion with a 1+ charge
is called the second ionization energy.
Trends in Ionization Energy
• First ionization energy:
– increases from bottom to top within a group, and
– increases from left to right across a period.
Ionization energy vs. atomic number
He
Ionization energy (kJ/mol)
2500
Ne
2000
Ar
F
1500
N
H
Cl
C
Be
1000
O
P S
B
500
Mg Si
Al
Li
Ca
Na
K
0
0
2
4
6
8
10
12
Element
14
16
18
20
Trends in Electronegativity
Electronegativity is the ability of an atom to attract
electrons when the atom is in a compound
*
* Fluorine has the highest electronegativity
Trends in Electronegativity
• Electronegativity is the ability of
an atom to attract electrons when
the atom is in a compound.
• Fluorine has the highest
electronegativity.
• Noble gases have zero
electronegativity (they already
have a full valence shell)
Periodic Trends
1. How is the atomic radius measured?
2. What is the ionization energy?
3. Do cations gain or lose electrons?
Periodic Trends
4.
What is the ionic size?
5. What is electronegativity?
6. Which element has the highest electronegativity?