Transcript Slide 1

Orbital filling table
•(1869) Dmitri Mendeleev(Russian
chemist) shows a first version of the
periodic table.
•He noticed that
classifying the
elements by
their atomic mass
a periodicity in certain properties
could be seen. The first table consisted
of 63 elements.
•Periodicity: the regular repeating of
properties according to the
arrangement of elements in the PT.
*Henry Mosely(British) discovered
nuclear charges of all known
elements and
that chemical properties of elements
are related to their atomic numbers
but not atomic weights.
He stated that elements should be
arranged in order of increasing
atomic numbers.
So, today’s periodic table was
formed.
In modern periodic table, elements
are listed in order of increasing
atomic numbers.
Elements with similar chemical
properties are placed in the same
vertical columns.
1A:Alkali metals
2A:Alkaline earth metals
3A:Earth metals
4A:Carbon Family
5A:Nitrogen Family
6A:Oxygen Family
7A:Halogens
8A:Noble(Inert)gases
GROUPS/FAMILIES:
The vertical columns
 Elements in the same group have
-similar chemical properties(exception:
H in 1A group)
-Same number of valence electrons
and orbitals.(exception:He in 8A group)
- The effective nuclear charge (the
charge acting on valence electrons) is
the same.(exception:He in 8A group)

*
*For A groups; # of valence electrons=
# of the group(except He in 8A)
in 7A;number of valence electrons=7
However, we can’t say the same thing
for B groups.
*Lanthanides&Actinides belong to 3B
group(the biggest group including 32
elements)
PERIODS:





The horizontal rows
There are 7 periods
Valence shell determines the period
number
Each of them starts with a metal and ends
with a noble gas.(except first and seventh
ones)
Elements in the same period have the
same # of energy levels or shells or
principle quantum numbers.
 PERIODIC
TRENDS:
1)ATOMIC&MASS NUMBER:
AN,MN increases
AN,MN
increases
2)ATOMIC RADIUS:
Radius of an atom: Half the
distance between two nuclei.
2r
Br
Br
Covalent radius: Half the distance between two
nuclei in a covalent molecule consisting of identical
atoms.
Van der Waals radius: This is for 0 group gases.
Ionic radius: This for ions in an ionic compound.
2)ATOMIC RADIUS:

Atomic size (volume,radius) is affected by
mainly two factors in the periodic table:

1)The # of shells (as it increases, atomic
volume also increases)

2)Nuclear charge(as the p+ # increases,
atomic volume decreases)
1A
Li
Na
K
Atomic volume decreases
2A
Be
3A
B
4A
5A
6A
7A
C
N
O
F
WHY?
Within the same period;All the
elements have the same # of shells
but the p+ # increases from left to
right.Therefore,atomic radius
decreases from left to right.
Proceeding down a group; The # of shells of atoms
increases but the p+ # of atoms also
increases.However,the increase in shells makes a
bigger effect on the radius than the nuclear
charge.Therefore, the atomic volume or radius
increases down a group.
IONIC VOLUME:
X-2
X-
X
X+
X+2
Less e-e repulsion
More e-e repulsion
For isoelectronic species;
The greater the nuclear charge,the smaller the
radius(or volume)
-3, O-2, Na+, Al+3 are isoelectronics.The
N
7
8
11
13
relationship between their radii;
-3 > O-2 > Na+ > Al+3
N
7
8
11
13
3)IONIZATION ENERGY:
The minimum amount of energy
required to remove the most losely
bound e-s from a mole of gaseous
atoms is called “the ionization
energy(I).”
X(g) + I1
X+(g) + e-
WARNING!!! X2(s)+I1
X+(s) + e-
!!!It’s not the
IE(bec. X is a solid
and molecular)
3)IONIZATION ENERGY:
 The
energy required to remove (1 mole
of) the first electron from 1 mole of
gaseous atom is called the first
ionization energy.

2
X (g )  I 2    X (g )

X
(g )
 I1    X

(g )
Ionization Energy

The second ionization energy is the energy
required to remove (1 mole of) the second
electron(s).
X

(g )
 I2    X
2
(g )
Always greater than first IE.
 The third IE is the energy required to
remove a third electron.
 Greater than 1st or 2nd IE.

What determines IE
The greater the nuclear charge, the
greater IE.
 The greater the effective nuclear
charge, the greater IE.
 Greater distance from nucleus (atomic
radius) decreases IE
 Shielding of electrons in filled inner
orbitals

***Because I of d block elements are
irregular, rules that we talk about the IE belong
to A group elements.
The variation of first ionization energies
within the same period:
As the atomic volume increases, the
attraction of the nucleus on the electrons
decreases.
*Ionization energies in the same period:
Noble gases > nonmetals > metals
Atomic volume decreases& IE
generally increases
Ionization energy
All the atoms in the same period have
the same energy level.
 Same shielding.
 But, increasing nuclear charge and
effective nuclear charge, ENC.
 So IE generally increases from left to
right.

***In
the same period from left to right
the ionization energies:
1A < 3A < 2A < 4A < 6A < 5A < 7A < 8A
irregularities
The variation of IE within the same group:
Down the group, atomic volumes of
elements increase and more shielding
effect, same ENC.Therefore, the IE of
elements decrease in the same group from
top to bottom.
First Ionization energy
He
He has a greater IE
than H.
 same shielding
 greater nuclear charge

H
Atomic number
First Ionization energy
He
 Li
H
Li
has lower IE
than H
 Outer electron
further away
 outweighs 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
B
First Ionization energy
He
H
Be
B
Li
has lower IE than
Be
 B has greater
shielding
 greater nuclear
charge

p orbital is slightly more
diffuse and its electron
easier to remove
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
Breaks the pattern,
because the outer
electron is paired in a
p orbital and
experiences interelectron repulsion.
B
Li
Atomic number
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
Ne
First Ionization energy
He
Ne has a lower IE
than He
 Both are full,
 Ne has more
shielding
 Greater distance

N F
H
C O
Be
B
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
Why the drop between groups
IIA and IIIA(Be-B)?
The explanation lies with the structures of
Boron and Aluminium. The outer electron is
removed more easily from these atoms than
the general trend in their period would
suggest.
Be
1s22s2
1st I.E. = 900 kJ mol-1
B
1s22s22px1
1st I.E. = 799 kJ mol-1
You might expect the Boron value to be
more than the Beryllium value because of
the extra proton. Offsetting that is the fact
that Boron's outer electron is in a 2p orbital
rather than a 2s. 2p orbitals have a slightly
higher energy than the 2s orbital, and the
electron is, on average, to be found further
from the nucleus. This has two effects.
The increased distance results in a reduced
attraction and so a reduced ionisation
energy.
The 2p orbital is screened not only by the
1s2 electrons but, to some extent, by the
2s2 electrons as well. That also reduces the
pull from the nucleus and so lowers the
ionisation energy.
Why the drop between groups
IIA and IIIA(Mg-Al)?
The explanation for the drop between
Magnesium and Aluminium is the same,
except that everything is happening at
the 3-level rather than the 2-level.
12Mg
13Al
1s22s22p63s2
1st I.E. = 736 kJ mo
1s22s22p63s23px1
1st I.E. = 577 kJ mo
The 3p electron in Aluminium is slightly
more distant from the nucleus than the 3s,
and partially screened by the 3s2 electrons
as well as the inner electrons. Both of these
factors offset the effect of the extra proton.
If the outer electron looks in towards the
nucleus, it doesn't see the nucleus sharply.
Between it and the nucleus there are the
two layers of electrons in the first and
second levels. The 11 protons in the
sodium's nucleus have their effect cut
down by the 10 inner electrons. The outer
electron therefore only feels a net pull of
approximately 1+ from the centre.
This lessening of the pull of the nucleus by
inner electrons is known as screening or
shielding.
Why the drop between groupsVA
and VIA (N-O and P-S)?
Once again, you might expect the
ionisation energy of the group VIA element
to be higher than that of group VA because
of the extra proton. What is offsetting it this
time?
•
22s22p 12p 12p 1 1st I.E. = 1400 kJ mol-1
N:
1s
7
x
y
z
22s22p 22p 12p 1 1st I.E. = 1310 kJ mol-1
O:
1s
8
x
y
z
Two electrons in the same orbital
experience a bit of repulsion from each
other. This offsets the attraction of the
nucleus, so that paired electrons are
removed rather more easily than you
might expect.
The screening is identical (from the 1s2 and, to
some extent, from the 2s2 electrons), and the
electron is being removed from an identical orbital.
The difference is that in the Oxygen case the
electron being removed is one of the 2px2 pair. The
repulsion between the two electrons in the same
orbital means that the electron is easier to remove
than it would otherwise be.
The drop in ionisation energy at Sulphur is
accounted for in the same way.
Increases
Increases
***We can decide about the group number of A
group elements by considering their ionization
energy values.
Mg(g)(1s22s22p63s2)+I1(176Kcal/mole)
Mg+(g) + eMg+(g)(1s22s22p63s1)+ I2(348Kcal/mole) Mg+2(g)+ eMg+2(g) (1s22s22p6)+ I3(1847Kcal/mole) Mg+3(g)+ e1st&2nd IE for Mg atom belong to the removal of
valence electrons which are bounded very weakly to
the nucleus .However,3rd erequires considerably more energy than the removal
of valence e-s as it will experience higher ENC.
WARNING!!!There is always needed much
more energy to remove inner electrons than
the outer electrons since the inner electrons
will experience a higher ENC.
The sharp jumps between ionization energies
help us to find out group numbers of A
group elements.
We can say that Mg atom is in 2A group by
considering the sharp increase between its
second and third ionization
energies.(because,there is very sharp
increase between its 2nd&3rd ionization
energies)
electron being lost:
1st
2nd
3rd
4th
(2A)
(3A)
(4A)
(5A)
(6A)
5th
6th
7th
The rare gases (He, Ne, Ar, Kr, Xe, Rn) appear at
peak values of ionization energy, which reflect
their chemical inertness, while the alkali metals
(Li, Na, K, Rb, Cs) appear at minimum values of
ionization energy, in keeping with their reactivity
and ease of cation formation.
4)ELECTRON AFFINITY:

The electron affinity is the energy that is
released when an atom in the gas phase gains
an electron and is thus converted to an anion,
also in the gas phase:
Electron affinities are difficult to
measure and there is no reliable data
available for most elements. However,
the larger the atom, the lower its electron
affinity, as shown with Group VII
elements:
For reasons outside the scope of this
discussion, the electron affinity of Fluorine is
an exception to this trend.
5) ELECTRONEGATIVITY:
 Electronegativity is the power of an
atom to attract electron density in a
covalent bond (Linus Pauling)
Electronegativity
 Electronegativity
describes
how electrons are shared in
a compound
 Consider the compound HCl
+
–
H
Cl
• The electron clouds represent where the two
electrons in the HCl bond spend their time
(sizes of atoms are not being shown)
• The shared electrons spend more time
around Cl than H. In other words Cl is more
electronegative than H.
Electronegativity
+

–

H Cl
0

0

H H
Electronegativity
Pauling’s electronegativity scale
H
2.1
He
-
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Ne
-
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
-



These numbers are derived from
several factors including EA, IE, atomic
radius
You do not need to understand where
the numbers come from
You need to know that a high number
means the element has a greater pull
on electrons
Calculating EN differences
 The
first step in defining the polarity of a
bond is to calculate electronegativity
difference ( EN)
  EN = EN large - EN small
 E.g. for NaCl,  EN = 3.2 –0.9 = 2.3, ionic.
Bigger electronegativity,
-Bigger tendency in gaining electrons.

Nuclear charge increases
Shielding increases
Atomic radius increases
Ionic size increases
Ionization energy decreases
Electronegativity decreases
Summary
Shielding is constant
Atomic Radius decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases
OXIDATION STATES OF
ELEMENTS
Group
IA
IIA
IIIA
IVA
VA
VIA
VIIA VIII
A
MAX +1
(+)
+2
+3
+4
+5
+6
+7
NA
MIN NA
(-)
NA
NA
-4
-3
-2
-1
NA
Important Groups
Groups of Elements
• Alkali Metals(IA)(except Fr):
– Group 1A metals
– Soft, silvery colored metals
that react violently with H2O
to form basic solutions
- They have low melting points and
densities due to being the largest
atom in their period of the PT.
• Alkali Metals(IA)(except Fr):
 Going
down the group, the metals
get softer and mp decreases due to
increase in atomic size.
-
-
They are the most reactive of all
the metals on the periodic table
since they can easily lose their one
e.
Their rectivity also increases as
you move down their
family.Therefore; most reactive
ones are: Cesium & Francium


Because the elements are very
reactive, they easily combine with
water and Oxygen.Therefore, most of
these elements are not found freely
in nature.That is why they are stored
in oil in a bottle in the laboratory in
order to prevent any reaction with
Oxygen.
They tarnish (lose lustre) rapidly when
exposed to the air.
Chemical Reactions of Alkali Metals

Reaction with water (Their rectivity increases as you move down
the group)

2Na(s) + H2O(l)




2Na+(aq) + 2OH- (aq) + H2(g)
With lithium, the reaction occurs slowly and steadily
In the case of sodium, the reaction is vigorous, producing enough
heat to melt the sodium which fizzes around the surface quite
vigorously
With potassium the reaction is violent and the heat produced is
enough to ignite the hydrogen gas evolved, which burns with a
purple flame.
They produce alkaline solution and hydrogen gas as a result of the
rxn w/ water.
Alkali Metal Family
Li
Na
K
•They give caharacteristic colour in the
flame(chemical peoperty).
Chemical Reactions of Alkali Metals

Reaction with oxygen,

React with oxygen to produces oxides.

4Li(s)
+ O (g)
2
2Li2O(s)
Chemical Reactions of Alkali Metals

Reaction with halogens

Reaction with halogens produces salts

2Na(s) + Cl (g)
2K(s) + Br (g)
2Cs(s) + l (g)


2
2
2
2NaCl(s)
2KBr(s)
2CsI(s)
•HALOGENS(VIIA) except At:


Exist as diatomic molecules in which atoms
are joined by a single cov. bond.
The elements in this group are referred to
as Halogens because they produce salts
when combined with alkali metals
(e.g.NaCl).these salts are usually white &
soluble in water.
•HALOGENS(VIIA) except At:



They are all very reactive and quite
electronegative nonmetals. The ease w/
which they gain electrons decreases going
down the group.
Halogens tend to be less reactive as
you move down the group.Flourine is
the most reactive Halogen and
combines with other elements very
readily.
Oxidizing power decreases down the
group.




Most are Poisonous .
When Fluorine combines with Na to form
NaF,it is an effective cavity fighter
that is added to toothpastes.
Chlorine is a great bacteria fighter so it
is used in swimming pools and household
cleaning agents.
Iodine is also useful for eliminating
bacteria.Since it is not as reactive as
Chlorine, it can be used on humans.
•HALOGENS(VIIA) except At:





Going down the group, their physical state
varies at room temp & pressure depending
on the Van Der Waals force strength present
between molecules since the molecules have
different molecular mass values.
F2, Cl2 -- gas
Br2 --- liquid
I2 -- solid, forms a purple gas on heating.
They all need an electron to become stable,
thus form negative ions
•HALOGENS(VIIA) except At:





Are slightly soluble in water as they are non-polar
molecules.
Concentrated solutions of chlorine- green tinge
Solutions of bromine- darken from yellow
through orange to brown as the concentration
increases.
Iodine dissolved in non-polar solvents like
hexane--violet solution.
Iodine dissolved in polar solvents like water &
ethanol--brown solution.
•HALOGENS(VIIA) except At:

Halogens dissociate slightly in aqueous solutions,
forming an acidic solution:

Cl2(aq) + H2O(l)  H+ (aq)+ Cl-(aq) + HOCl(aq)

HOCl(hypochlorous acid): weak acid, reacts as an
oxidant since it donates its one oxygen.It oxidises
colored dyes to colorless products.
HOCl turns blue litmus paper into red, and then
make it colorless. Therefore, HOCl and OCl- are
used in bleaches.They are also toxic to microbes.

Displacement reactions of Halogens
A more reactive halogen is
capable of replacing less
reactive one from its
solution
Cl2 reacts with Br- and I-
Cl2(aq) + 2Br-(aq)  2Cl-(aq) + Br2(l)
Cl2(aq) + 2I-(aq)  2Cl-(aq) + I2(s)
Br2 reacts with I-
Br2(aq) + 2I-(aq)  2Br-(aq) + I2(s)
I2 non-reactive with halide ions
84
•HALOGENS(VIIA) except At:

The common insoluble halides (ions of halogens)
are those of Pb and Ag.
 PbI2 -- is a bright yellow colored, can be used
as a test for iodide ion.
 Look
at P. 77
& 78 for the colors of
halogens and tests
of halide ions.
Oxides of period 3 elements
Metallic Oxides in Period 3
Sodium oxide: Na2O
Magnesium oxide: MgO
Aluminum oxide: Al2O3
ionic
ionic
ionic
Metalloid oxide in Period 3
Silicon dioxide: SiO2
covalent
Nonmetallic oxides in Period 3
Tetraphosphorus decoxide: P4O10
Sulfur trioxide: SO3
Dichlorine heptoxide: Cl2O7
covalent
covalent
covalent
86
Discuss the changes in nature, from ionic
to covalent and from basic to acidic, of
the oxides across period 3.
Acidic/Basic
Metallic oxides in Period 3 are basic
Sodium oxide:
Na2O + H2O  2 NaOH
Magnesium oxide:
MgO + H2O  Mg(OH)2
basic
basic
Net ionic eqn: O2- + H2O  2 OHAluminum oxide:
Al2O3 + 6HCl  2 AlCl3+ 3H2O amphoteric
Al2O3 +2 NaOH + 3H2O  2NaAl(OH)4
87
Oxides of period 3 elements
Metalloid oxide in Period 3 is acidic
Silicon dioxide:
SiO2 + H2O  H2SiO3
acidic
(silicic acid)
Nonmetallic oxides in Period 3 are acidic
Tetraphosphorus decoxide: P4O10 + 6H2O  4H3PO4
Sulfur trioxide:
SO3 + H2O  H2SO4
Dichlorine heptoxide: Cl2O7 + H2O  2HClO4
Argon does not form an oxide
acidic
acidic
acidic
Look at!!
 P.80