Download first ionization energy.

Hydrogen and Helium
Hydrogen does not share the same
properties as the elements of group 1.
Helium has the electron configuration of
group 2 elements however it behaves like
group 18 (Noble Gases)
f-block elements
Lanthanides & Actinides
Lanthanides – shiny metals
Actinides – radioactive; only the 1st four
are found naturally
The Octet Rule
• All atoms want their outer shell (highest
energy level) to have 8 electrons.
• If the outer shell has eight, the atom is
stable (non-reactive)
Valence Electrons
• e- available to be lost, gained, or shared
in forming chemical compounds
• these e- are in highest energy level
# of valence electrons?
Li
C
Ca
Ne
As
At
Ions
Ions are formed when:
• a metal gives up valence electrons
• a nonmetal gains valence electrons
They do this so their highest energy level
has 8 valence electrons = octet
IONS
• This transfer of electrons creates an
element (ion) that is no longer neutral
• Now the protons ≠ electrons
• Cations are ions with a positive charge
• Anions are ions with a negative charge
Cations
(The “t” in “cation” looks like a plus sign)
• All metals form positive cations when they give
electrons from their valence shell.
• Alkali metals (1 valence e- to give away), so they
become a ion with +1 charge b/c they have 1
more p+ than e-)
ex. Li+, Na+, K+
• Alkaline Earth Metals (2 valence e- to give
away), so…
ex. Ca+2, Ba+2
Common Ionic Charges
Anions (“A-negative-ion”)
• All non-metals form anions when they gain
electrons from metals to fill their outer
shell with a total of eight.
• More e- than p+ means they have a
- charge
All halogens (7 valence electrons); so they
form an ion with a -1 charge b/c gained 1eex. Cl-, F-, Br-
Common Ionic Charges
Your Turn!
• potassium ion?
K+1
• oxygen ion?
O-2
• magnesium ion?
Mg+2
• noble gases?
Already stable!
Ionic Radii (Trend)
Cations (+ ions) are smaller b/c they lose eAnions (- ions) are bigger b/c they gain e• Metals (left side p.table) form cations
• Nonmetals (right side p.table) form anions
Ionic Radii INCREASES as you go across
and down the p.table
Atomic Size - (Trend)
As we increase the
atomic number (or
go down a group). .
.
each atom has
another energy
level,
so the atoms get
bigger.
H
Li
Na
K
Rb
Atomic Size - Period Trends
Going from left to right across a period, the
size gets smaller.
Electrons are in the same energy level.
But, there is more nuclear charge.
Outermost electrons are pulled closer.
Na
Mg
Al
Si
P
S Cl Ar
Rb
K
Atomic Radius (pm)
Period 2
Na
Li
Kr
Ar
Ne
H
3
10
Atomic Number
Ionization Energy (Trend )
Ionization energy is the amount of
energy required to completely remove
an electron (from a gaseous atom).
Removing one electron makes a 1+
ion.
The energy required to remove only the
first electron is called the first
ionization energy.
Ionization Energy
The second ionization energy is the
energy required to remove the
second electron.
Always greater than first IE.
The third IE is the energy required to
remove a third electron.
Greater than 1st or 2nd IE.
Symbol First
Second
1312
2731
520
900
800
1086
1402
1314
1681
2080
5247
7297
1757
2430
2352
2857
3391
3375
3963
H
He
Li
Be
B
C
N
O
F
Ne
Third
11810
14840
3569
4619
4577
5301
6045
6276
Ionization Energy - Group trends
As you go down a group, the
first IE decreases because...
The electron is further away
from the attraction of the
nucleus, and
There is more shielding.
Ionization Energy - Period trends
All the atoms in the same period have
the same energy level.
Same shielding.
But, increasing nuclear charge
So IE generally increases from left to
right.
Exceptions at full and 1/2 full orbitals.
First Ionization energy
He
H
He has a greater IE
than H.
Both elements have the
same shielding since
electrons are only in the
first level
But He has a greater
nuclear charge
Atomic number
First Ionization energy
He
 Li
H
Li
has lower IE
than H
 more shielding
 further away
 These outweigh
the greater
nuclear charge
Atomic number
First Ionization energy
He
 Be
H
Be
has higher IE
than Li
 same shielding
 greater nuclear
charge
Li
Atomic number
First Ionization energy
He
B
H
Be
B
Li
has lower IE
than Be
 same shielding
 greater nuclear
charge
 By removing an
electron we make
s orbital filled
Atomic number
First Ionization energy
He
H
Be
C
B
Li
Atomic number
First Ionization energy
He
N
H
C
Be
B
Li
Atomic number
First Ionization energy
He
N
H
C O
Be
B
Oxygen breaks
the pattern,
because removing
an electron leaves
it with a 1/2 filled p
orbital
Li
Atomic number
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
First Ionization energy
He
Ne
N F
H
C O
Be
B
Ne has a lower IE
than He
Both are full,
Ne has more
shielding
Greater distance
Li
Atomic number
Ne
First Ionization energy
He
N F
H
C O
Be
B
Li
 Na
has a lower
IE than Li
 Both are s1
 Na has more
shielding
 Greater
distance
Na
Atomic number
Atomic number
First Ionization energy
Electronegativity
• Electronegativity refers to the ability of an
atom to attract the electrons from
another atom bonded together in a
compound.
• EA (e- affinity) deals with an isolated atom
and Electronegativity deals with 2 atoms
bonded together
Electronegativity Trends
(same trend as EA)
• Electronegativity generally increases
as you move across the periodic table
• Electronegativity decreases as you
move down a group because of the
shielding effect
Electronegativity
Sample Problems
Which has the greatest electronegativity?
Cl, S, P, Al
Ra, Ba, Ca, Be
Fr, Ba, Cs, Al
Element with Highest
Electronegativity?
Fluorine, has the highest electronegativity.
•
•
•
•
•
Electron Affinity
EA
The change in energy (always negative)
when you add an e- to an atom.
Affinity = “you like it”
EA increases (more negative #) as you
move across because those atoms want
to gain e- to gain an octet
EA decreases as you move down
because atoms want to gain e- less
because of the shielding effect
Noble gases have an EA of ZERO, b/c
they already have an octet
Atomic Radii Decrease as You
Move Across a Period
Why??????
• The atomic radius decreases because the
greater positive charge (more p+) in the
nucleus pulls the electrons closer..
(opposite charges attract)
Atomic Radii Increase as you Move
Down a Group
• WHY?????
• As you move down you add energy levels
which increases the size and radius of the
atom.
Sample Problems
Ca, Be, Ba, and Sr: largest atomic radius?
Al, Mg, Si, and Na: smallest radius?
Largest atomic radius?
Li
O
C
F
Ionization Energy
• The IE is the minimum amount of energy
required to remove an electron from the
outer shell of a neutral atom.
• energy to remove the first e- is called the
“first ionization energy” IE1
Ionization Energy Trend
• IE (energy to take an e-) increases as you
move across because non-metals want to gain
e– , not lose their e• nuclear charge increases due to adding more
# p+, so it takes more energy to remove an
electron.
• IE decreases as you move down the periodic
table because the electrons are farther away
from the nucleus and easier to remove.
• Why? all the electrons in the many energy
levels between repel the valence e- = shielding
effect