Download Chapter 5 ppt

Chapter 5



Arranges elements according to similar
periodic properties
Developed by Dmitri Mendeleev (atomic
mass), refined by Moseley (atomic number)
Elements share similar properties with
elements in the same group
•
After Mendeleev placed all the known elements
in his periodic table, several empty spaces were
left.
•
In 1871 Mendeleev predicted the existence and
properties of elements that would fill three of the
spaces.

Horizontal rows = periods
◦ There are 7 periods

Vertical column = group (or family)
◦ Similar physical & chemical properties
◦ Identified by number at top



When elements are placed in proper order,
there is a regular, repeating pattern of
properties and trends
Metals – more reactive going down a group
Nonmetals – more reactive going up a group
(Alkali metals video)
Group 1A are the alkali metals
(but NOT H)
Group 2A are the alkaline earth metals
H
•
The elements of Group 1 of the periodic table
are known as the alkali metals.
• lithium, sodium, potassium, rubidium, cesium, and
francium
• In their pure state, all of the alkali metals have a silvery
appearance and are soft enough to cut with a knife.
•
The elements of Group 2 of the periodic table
are called the alkaline-earth metals.
• beryllium, magnesium, calcium, strontium, barium, and
radium
• Group 2 metals are less reactive than the alkali metals,
but are still too reactive to be found in nature in pure
form.
Group 8A are the noble gases
 Group 7A is called the halogens


Halogens – Very reactive
◦ Readily steal electrons from Group 1 and 2 metals

Noble Gases – completely stable
◦ Do not react with other elements
◦ Have a full outer shell of electrons
1.
2.
Periodic trends – what happens as you go
across a period
Group trends – what happens as you go
down a group
Influences to look for this chapter…
1. Energy Level
Higher energy levels are further away
from the nucleus.
2. Charge on nucleus (# protons)
More charge pulls electrons in closer.
1
2
3
4
5
6
7
2
3
n=
4
n=
3
n=
2
n=
1
nucleus



Opposites attract
More protons = greater nuclear charge
Greater nuclear charge = greater e- attraction
}
Radius
• Atomic Radius - half the distance between the
two nuclei of a diatomic molecule.
• Measured this way because a single atom
doesn’t have a definite edge
Group Trends



As we move
down a group…
each atom has
another energy
level,
so the atoms get
H
Li
Periodic Trends

Na

K
bigger.
Rb

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 (+ from
nucleus)
Outermost electrons
are pulled closer.
Na Mg Al
Si P S
Cl Ar
Atomic Radius
Increases
Atomic Radius Decreases
Cations


Anions
Form when metals LOSE
electron(s) from their
outermost energy level
Na  Na+1 + 1e-


Na


Cl
Na+1
Cation size is always
smaller than size of
parent atom
Why? Same energy level
and nuclear charge
attracting fewer electrons
Form when nonmetals
GAIN electron(s) into
their outermost energy
level
Cl + 1e-  Cl-1


Cl-1
Anion size is always
bigger than size of the
parent atom
Why? Same energy level
nuclear charge attracting
more electrons
Group Trends



Generally, ionic size
increases as you go
down a group
Why?
You add energy levels
as you move down a
group
nucle
us
Periodic Trends



Generally, ionic size
DECREASES as you go
left to right across a
period
Why?
Increased nuclear
charge in the same
energy level
+7
+8
+9
Ionic Radius
Increases
Ionic Radius Decreases
• Ionization energy is the amount of
energy required to completely
remove an electron (from a gaseous
atom).
• Easier to remove = lower IE
• Harder to remove =higher IE
Na  Na+1 + e- IE = 495.8 kJ/mol
K  K+1 + eIE = 418.8 kJ/mol
Group trends
As we move
down a group…
 Outer level
electrons are
further from the
nucleus and are
shielded
 so the 1st
ionization
energy
decreases
(easier to remove
e- )

Periodic Trends
H

Li
Na

K
Rb
Going left to right
across a period, nuclear
charge increases (+
from nucleus).
So the 1st ionization
energy increases (it
becomes harder to
remove an electron)
Na Mg Al
Si P S
Cl Ar
Ionization Energy
Decreases
Ionization Energy Increases
• The electron on the
outermost energy
level has to look
through all the other
energy levels to see
the nucleus.
• Filled and half-filled orbitals
have lower energy, so achieving
them is easier, lower IE.


Energy required to remove a second electron
from an ion
Generally, it requires more energy to remove
the second electron than the first
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
Why did these values
5247
7297
1757
2430
2352
2857
3391
3375
3963
increase so
much?
11810
14840
3569
4619
4577
5301
6045
6276
• Electron affinity is a measure of
an atom’s affinity or attraction
for electrons
• More negative = more affinity
Group trends
As we move down a
group…
 The pull from the
nucleus is less
because it is farther
away
 so the electron affinity
GENERALLY decreases
going down a group
(harder to add an e-)

Periodic Trends



Going left to right
across a period, nuclear
charge increases (+
from nucleus).
So the electron affinity
GENERALLY increases
(easier to add an e-)
Halogens generally have
the most negative
electron affinities. Why?
Electron Affinity – decreases
Electron Affinity – increases


Electronegativity is the ability of an atom to
attract bonded electrons
Bonded electrons are valence electrons
(electrons in the outermost energy level –
given by main group number)
Group trends
As we move down a
group…
 The pull from the
nucleus (+) is less
because it is farther
away
 so the ability of the
atom to attract
bonded electrons
decreases
(lower
electronegativity)

Periodic Trends


Going left to right
across a period, nuclear
charge increases (+
from nucleus).
So the ability of an atom
to attract bonded
electrons increases
(higher
electronegativity)
Electronegativity
Decreases
Electronegativity Increases
 Pg
166
 #1-5, 7-9, 13, 18, 20, 24,
27, 28, 30
 Due tomorrow!