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Unit 3: The Periodic Table
Elements
 Science has come
along 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
• On the periodic table, elements are:
– In order of increasing atomic number.
– In groups (columns) with similar properties
– In periods (rows) with valence electrons in the
same energy level
How are Elements Classified?
Three Regions
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.
Metalloids
The elements touching the
diagonal, but not Al or Po.
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
Let’s Look at the Groups
•
•
•
•
Alkali
Alkali Earth
Halogen
Noble Gases
GROUPS or FAMILIES
•
•
•
•
Elements in the same column (groups) have similar properties
Group 1 Alkali Metals, 1 valence electron, very reactive
Group 2 Alkaline Earth Metals, 2 valence electrons
Group 8: Noble Gases, 8 valence electrons, inert
The Periodic Law
• Group 7: Halogens, 7 valence electrons
• Transition metals: in the “d block”
• Rare Earth metals: in the “f block”
Alkali Metals
 Group 1A (Not including hydrogen)
Li
Na
K
Rb
 Very reactive metals
Cs
 Two most reactive elements: Cs and Fr
Fr
Explodes in water
https://www.youtube.com/watch?v=bze3hN9j9Cw
Alkaline Earth Metals
Be
 Group 2A
Mg
 2 valence electrons
Ca
 Several are important
mineral nutrients (Mg
and Ca)
Sr
Ba
Ra
Halogens
F
Cl
 Are non-metals
 “Halogen” means salt
former.
 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 nonmetal
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
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
– ionization energy
– electronegativity
Definition: Atomic Size
• Atomic Radius: the distance between the nucleus
and the outer edge of the electron cloud.
– Measured by the distance between two nuclei divided by 2
Definition: Ionization Energy
• Ionization Energy:
the energy required
to remove an
electron from an
atom.
Definition: Electronegativity
Electronegativity:
the ability of an
atom to attract
electrons in
another atom to fill
its valence shell.
Periodic Trends
• As you move from left to right along a period,
the number of electrons and protons increase.
• This leads to a stronger pull on the electrons
by the protons.
e
e
e
e
e
+++
++
e
e
e
+++
+++
e
e
e
+++
+++
+
e
e
e
e
e
e
e
Boron: 5 p+, 5 e
Carbon: 6 p+, 6 e
Nitrogen: 7 p+, 7 e
Periodic Trends
Atomic Radius Trend:
• Along a period, the atomic radius decreases
as you move from left to right.
• Stronger pull by protons = smaller atom.
e
e
e
e
e
+++
++
e
e
e
+++
+++
e
e
e
+++
+++
+
e
e
e
e
e
e
e
Boron: 5 p+, 5 e
Carbon: 6 p+, 6 e
Nitrogen: 7 p+, 7 e
Periodic Trends
Ionization Energy Trend:
• Along a period, the ionization energy increases as
you move from left to right.
• Stronger pull by protons = more difficult to remove
electron.
e
e
e
e
e
+++
++
e
e
e
+++
+++
e
e
e
+++
+++
+
e
e
e
e
e
e
e
Boron: 5 p+, 5 e
Carbon: 6 p+, 6 e
Nitrogen: 7 p+, 7 e
Periodic Trends
Electronegativity Trend:
• Along a period, the electronegativity increases as
you move from left to right.
• Stronger pull by protons = stronger pull on
electrons in another atom.
e
e
e
e
e
+++
++
e
e
e
+++
+++
e
e
e
+++
+++
+
e
e
e
e
e
e
e
Boron: 5 p+, 5 e
Carbon: 6 p+, 6 e
Nitrogen: 7 p+, 7 e
Trends in Electronegativity
• Fluorine has the highest
electronegativity.
• Noble gases have zero
electronegativity (they already
have a full valence shell).
Group Trends
• As you move from top to bottom along a group, the
number of energy levels increases.
• This leads to a weaker pull on the electrons by the
protons.
e
e
e
e
e
e
3 p+
e
e
11 p+
e
e
Lithium: 3 p+, 3 e
e
e
e
e
e
e
e e
e
Sodium : 11 p+, 11 e
e
e
e
e
19 p+
e
e
e
e
e
e
e
e e e
Potassium: 19 p+, 19 e
Group Trends
Atomic Radius Trend:
• As you move down a group, atomic radius increases.
• More energy levels = larger atomic size.
e
e
e
e
e
e
3 p+
e
e
11 p+
e
e
Lithium: 3 p+, 3 e
e
e
e
e
e
e
e e
e
Sodium : 11 p+, 11 e
e
e
e
e
19 p+
e
e
e
e
e
e
e
e e e
Potassium: 19 p+, 19 e
Group Trends
Ionization Energy Trend:
• As you move down a group, ionization energy decreases.
• More energy levels = less pull on electrons & easier to
remove.
e
e
e
e
e
e
3 p+
e
e
11 p+
e
e
Lithium: 3 p+, 3 e
e
e
e
e
e
e
e e
e
Sodium : 11 p+, 11 e
e
e
e
e
19 p+
e
e
e
e
e
e
e
e e e
Potassium: 19 p+, 19 e
Group Trends
Electronegativity Trend:
• As you move down a group, electronegativity decreases.
• More energy levels = less pull on electrons in another atom
and “shielding” by inner electrons.
e
e
e
e
e
e
3 p+
e
e
11 p+
e
e
Lithium: 3 p+, 3 e
e
e
e
e
e
e
e e
e
Sodium : 11 p+, 11 e
e
e
e
e
19 p+
e
e
e
e
e
e
e
e e e
Potassium: 19 p+, 19 e
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.
Trends in Atomic Size
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
• Which has the largest atomic radius?
• Be or B?
Be
• Be or Mg?
Mg
• Ar or Na?
Na
• K or Li?
K
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
• increases from bottom to top within a group, and
• increases from left to right across a period.
*
• Fluorine has the highest electronegativity,
• Noble gases have ZERO electronegativity
Ion Size
• Metals lose electrons: cation (positive charge)
• Cations are always smaller than the original atoms.
Na+ has valence electrons in energy level 2 more pull on
valence electrons – smaller ion.
Ions
• Non-metals gain electrons: anion (negative charge)
• Anions are always larger than the original atoms.
More electrons, more repulsion, so electrons “spread out”:
larger atom.
ISOELECTRIC SERIES
• “Isoelectronic” means ions/atoms with the
same number of electrons.
N3- O2- F- He
Na+
Mg2+ Al3+
Each has 10 electrons.
ISOELECTRIC SERIES
What is the trend of atomic size for these isoelectric
atoms/ions?
N3- O2- F- He
Na+
Mg2+ Al3+
Each has 10 electrons.
Decreasing radius with increasing # of protons.
(Due to stronger pull from MORE protons.)
Periodic Trends
1. What is the ionization energy?
The energy required to remove one electron from the atom
2. What is electronegativity?
The ability of atoms to attract electrons.
3. Which element has the highest electronegativity?
Fluorine
Periodic Trends
Which has the higher electronegativity, Na or Cl?
Cl
• Which has the higher ionization energy, F or Br?
F
• What is the larger atom, Be or N?
Be
• What is larger, Mg or Mg2+?
Mg