Transcript Periodic Trends Mrs.Kay

Periodic Trends
Mrs.Kay

Groups: vertical columns (18)
 Have
similar properties because have same
number of electrons in outer shell

Periods: horizontal columns (7)
Valence electrons:
electrons found in the
outer most shell or
valence shell
 Each energy
level/shell holds on a
certain number of
electrons

S- block
s1
s2
Alkali metals all end in s1
 Alkaline earth metals all end in s2
 really have to include He but it fits
better later.
 He has the properties of the noble
gases.

Transition Metals -d block
d1 d2 d3
s1
d5
s1
d5 d6 d7 d8 d10 d10
The P-block
p1 p2
p3
p4
p5
p6
F - block

inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
Metals:

To the left of the staircase line
Physical
Properties:
 Luster
(shiny)
 Good conductors
 High density
 High melting point
 Malleable

Chemical
Properties:
Easily lose electrons
 Corrode easily (ex:
rusting or
tarnishing)
 Low
electronegativity

Non metals:

to the right of the staircase line
Physical Properties:
Dull
 Poor conductor
 Brittle
 Not malleable
 Low density and
melting point

•
Chemical Property:


Tend to gain electrons
High electronegativity
Metalloids:
along the staircase line
Solids
 Shiny or dull
 Malleable
 Conduct heat and
electricity better than
non metal but not as
well as metals

Atomic Size
Atomic Size
The electron cloud doesn’t have a definite
edge.
 They get around this by measuring more
than 1 atom at a time.
 Summary: it is the volume that an atom
takes up

Atomic Size
}
Atomic
Radius
Radius = half the distance between two
nuclei of a diatomic molecule.
Group trends

As we go down a
group (each atom
has another energy
level) the atoms
get bigger, because
more protons and
neutrons in the
nucleus
H
Li
Na
K
Rb
Periodic Trends
atomic radius decreases as you go from left to right
across a period.
 Why? Stronger attractive forces in atoms (as you go
from left to right) between the opposite charges in
the nucleus and electron cloud cause the atom to be
'sucked' together a little tighter.
Na
Mg
Al
Si
P
S Cl Ar
Reactivity
Reactivity

Reactivity refers to how likely or vigorously
an atom is to react with other substances.
This is usually determined by how easily
electrons can be removed (ionization
energy) and how badly they want to take
other atom's electrons
For Metals:
Period - reactivity decreases as you go
from left to right across a period.
Group - reactivity increases as you go
down a group
 Why? The farther to the left and down the
periodic chart you go, the easier it is for
electrons to be given or taken away,
resulting in higher reactivity
For Non-metals
Period - reactivity increases as you go
from the left to the right across a period.
Group - reactivity decreases as you go
down the group.
 Why? The farther right and up you go on
the periodic table, the higher the
electronegativity, resulting in a more
vigorous exchange of electron.

Ionization Energy
Ionization Energy
The amount of energy required to
completely remove an electron from a
gaseous atom.
 An atom's 'desire' to grab another atom's
electrons.
 Removing one electron makes a +1 ion.
 The energy required is called the first
ionization energy.
X(g) + energy →X+ + e
Ionization Energy
The second and third ionization energies can
be represented as follows:
 X+ (g) + energy X2+ (g) + e X2+ (g) + energy X3+ (g) + e More energy required to remove 2nd
electron, and still more energy required to
remove 3rd electron
Group trends

Ionization energy decreases down the group.
Going from Mg to Be, IE decreases because:




Be outer electron is in the 3s sub-shell rather than the
2s. This is higher in energy
The 3s electron is further from the nucleus and
shielded by the inner electrons
So the 3s electron is more easily removed
A similar decrease occurs in every group in the
periodic table.
Period trends
IE generally increases from left to right.
 Why?
 From Na to Ar (11 protons to 18
protons), the nuclear charge in each
element increases.
 The electrons are attracted more
strongly to the nucleus – so it takes more
energy to remove one from the atom.
Why is there a fall from Mg to Al?
Al has configuration 1s2 2s2 2p6 3s2 3p1,
its outer electron is in a p sublevel
 Mg has electronic configuration 1s2 2s2
2p6 3s2.
 The p level is higher in energy and with
Mg the s sub level is full – this gives it a

slight stability advantage
Why is there a fall from P to S?





This can be explained in terms of electron
pairing.
As the p sublevel fills up, electrons fill up the
vacant sub levels and are unpaired.
This configuration is more energetically stable
than S as all the electrons are unpaired. It
requires more energy to pair up the electrons in
S so it has a lower Ionisation energy.
There is some repulsion between the paired
electrons which lessens their attraction to the
nucleus.
It becomes easier to remove!
Driving Force
Full Energy Levels are very low energy.
 Noble Gases have full energy levels.
 Atoms behave in ways to achieve noble
gas configuration.

2nd Ionization Energy
For elements that reach a filled or half
filled sublevel by removing 2 electrons 2nd
IE is lower than expected.
 Makes it easier to achieve a full outer shell
 True for s2
 Alkaline earth metals form +2 ions.

3rd IE
Using the same logic s2p1 atoms have an
low 3rd IE.
 Atoms in the aluminum family form +3
ions.
 2nd IE and 3rd IE are always higher than 1st
IE!!!

Electron Affinity
Electron Affinity
The energy change associated with adding
an electron to a gaseous atom.
 Easiest to add to group 7A.
 Gets them to full energy level.
 Increase from left to right atoms become
smaller, with greater nuclear charge.
 Decrease as we go down a group.

Ionization energy, electronegativity
Electron affinity INCREASE
Atomic size increases,
shielding constant
Ionic size increases
Ionic Size
Cations form by losing electrons.
 Cations are smaller than the atom they
come from.
 Metals form cations.
 Cations of representative elements have
noble gas configuration.

Ionic size
Anions form by gaining electrons.
 Anions are bigger than the atom they
come from.
 Nonmetals form anions.
 Anions of representative elements have
noble gas configuration.

Configuration of Ions
Ions always have noble gas configuration.
 Na is 1s22s22p63s1
 Forms a +1 ion : 1s22s22p6
 Same configuration as neon.
 Metals form ions with the configuration of
the noble gas before them - they lose
electrons.

Configuration of Ions
Non-metals form ions by gaining electrons
to achieve noble gas configuration.
 They end up with the configuration of the
noble gas after them.

Periodic Trends
Across the period nuclear charge increases
so they get smaller.
 Energy level changes between anions and
cations.

Li+1
B+3
Be+2
C+4
N-3
O-2
F-1
Size of Isoelectronic ions
Iso - same
 Iso electronic ions have the same # of
electrons
 Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
 all have 10 electrons
 all have the configuration 1s22s22p6

Size of Isoelectronic ions

Positive ions have more protons so they
are smaller.
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
Electronegativity
(optional coverage)
Electronegativity
The tendency for an atom to attract
electrons to itself when it is chemically
combined with another element.
 How fair it shares.
 Big electronegativity means it pulls the
electron toward it.
 Atoms with large negative electron affinity
have larger electronegativity.

Group Trend
The further down a group the farther the
electron is away and the more electrons an atom
has.
 So as you go from fluorine to chlorine to
bromine and so on down the periodic table, the
electrons are further away from the
nucleus and better shielded from the nuclear
charge and thus not as attracted to the nucleus.
For that reason the electronegativity
decreases as you go down the periodic table.

Period Trend
Electronegativity increases from left to
right across a period
 When the nuclear charge increases, so
will the attraction that the atom has for
electrons in its outermost energy level and
that means the electronegativity will
increase

Period trend
Electronegativity increases as you go from left to
right across a period.
 Why? Elements on the left of the period table
have 1 -2 valence electrons and would rather
give those few valence electrons away (to
achieve the octet in a lower energy level) than
grab another atom's electrons. As a result, they
have low electronegativity. Elements on the right
side of the period table only need a few
electrons to complete the octet, so they have
strong desire to grab another atom's electrons.
Group Trend
electronegativity decreases as you go down a group.
 Why? Elements near the top of the period table have
few electrons to begin with; every electron is a big deal.
They have a stronger desire to acquire more electrons.
Elements near the bottom of the chart have so many
electrons that loosing or acquiring an electron is not as
big a deal.
 This is due to the shielding affect where electrons in
lower energy levels shield the positive charge of the
nucleus from outer electrons resulting in those outer
electrons not being as tightly bound to the atom.
Shielding
Shielded slightly from the
pull of the nucleus by the
electrons that are in the
closer orbitals.
 Look at this analogy to
help understand
