General Chemistry
Assistant professor
•
Mervat Mohamed Hosny •
6-Quantum mechanical atom
(Schroedinger)
1-Democritus:
He theorized that all matter is composed of •
small indivisible particles called atoms •
2-Dalton’s atomic theory:
*each element is composed of minute •
indivisible particles called atoms •
*all atoms of a given element are chemically •
identical to each other ,atoms of one element
are different from the atoms of all other
element
3- during ordinary chemical reactions atoms of
one element cannot be changed into atoms of
different element.
4- atoms are not created or destroyed •
5- compound is formed when atoms of more •
than one element combine
3-J.J.Thomson-CRT
*he discovered the electron •
* in thomson‘s model ,electrons are •
embedded in a positive sphere of •
matter
4-Rutherford gold foil experiment:
*he established that the positive charged •
alpha particles emitted by certain radioactive
elements (helium) .
* he used these alpha particles to establish •
the nuclear nature of atoms .
*in these experiments ,he directed a stream of •
positive charged helium ions (alpha
particles)at a very thin sheet of gold foil
conclusion
*most of the mass and all of the positive •
charge of the atom are contained in a small
space called the nucleus •
*most of the volume of the atom is empty •
space occupied by tiny negatively charged
electrons
*negative charged electrons outside nucleus
=positive charge inside nucleus
*the atom is electrically neutral •
*protons:+vely charged subatomic particles •
found in nucleus
*neutrons : neutral (uncharged)subatomic •
particles found in nucleui
*electrons very small – vely charged •
subatomic particles
5- The Bohr model:
*electrons in an atom exist in specific regions •
at various distances from the nucleus .
*The electrons are rotating in orbits around •
the nucleus like planets rotating around the
sun.
*he describe hydrogen atom as a single •
electron rotating in an orbit about a relatively
heavy nucleus . •
*he applied the concept of energy quanta, •
proposed by the German physicist Planck
Planck stated that :
*energy is never emitted in a •
continuous stream but only in small
discrete packets
called quanta •
Bohr theorized that :
*There are several possible orbits for •
electrons at different distances from the
nucleus
*but electron had to be in one specific orbit or •
another .
*It could not exist between orbits
*when a hydrogen atom adsorbed one or •
more quanta of energy ,its electron jumped to
another orbit a greater distance from the
nucleus.
*when the electron fell back to lower orbits ,it •
emitted quanta of energy as light ,giving rise
to the spectrum of hydrogen .
*each orbit is at a different energy level
*an electron in the orbit closest to the nucleus •
•
is in the 1st energy level ,at greater distances it •
•
may be in the second ,3rd or fourth energy •
level
6-Quantum mechanical atom
)Schroedinger)
*They found that Bohr’s assumptions have to •
be modified
*Difficulty arise in applying the theory to •
atoms containing many electrons
*Bohr’s concept was replaced by quantum •
mechanics theory
One of the chief difference between
the 2 theories is that:
In the quantum mechanics theory electrons •
are not considered to be revolving around the •
nucleus in orbits but to occupy orbitals cloud •
like regions surrounding the nucleus and •
corresponding to energy levels •
Erwin Schrodinger introduced his
famous wave equation Quantum
mechanics or wave mechanics
He describe an electron as simultaneously •
having properties of:
1-a wave (like light) •
2-and a particle (have mass) •
The solution of the Schrodinger equation is
complex but as aconclusion :
There is four quantum numbers which •
define the location and properties of •
electrons in atoms: n,l,m,s •
n is the principle quantum no indicate
the energy levels of the electron
relative to their distance from the
nucleus
n=1,2,3,……… •
But always 1-7 •
‫• بعد االلكترون عن النواة‬
L=2nd quantum no explain the shape of
orbital
Electron exist in orbitals having specific •
shapes
S
P
d
f •
3-m magnetic quantum no
Orientation in space •
*electron orbitals have specific orientation in •
Space •
*This quantum number accounts for the •
number of s,p,d,f orbitals that can be present
in the principal energy level
4-Spine quantum no (s)
*an electron spins about its own axis in either •
a clockwise or counter clockwise direction
*S relates to the direction of spin of an •
electron
*when 2 electrons occupy the same •
orbital,they must have opposite spins
*when an orbital contain 2 electrons
the electrons are said to be paired
NO ELECTRONS IN AN ATOM CAN HAVE •
THE SAME 4 QUANTUM NUMBERS •
7- Energy levels of electrons
*all the electrons in an atom are not located •
the same distance from the nucleus
*as said in Bohr theory and quantum •
mechanics the probability of finding the
electrons is greatest at certain specified
distance s called energy levels, from the
nucleus
*energy levels are also referred to as electron
shell and may contain only a limited number of
electrons
* energy levels are numbered startly with •
n=1 to n=7
•
Or K,L,M,N,O,P,Q •
Where K=1st energy level,L=2nd energy level •
*the maximum number of
electrons that can occupy a
specific energy level =2(nxn(
n= number of the principle energy level •
E.g. for shell k or energy level 1=2x(1x1)=2 •
E.g. for shell l or energy level 2 =2x(2x2)=8 •
8- energy sublevels of electrons:
*the principle energy levels contain sublevels
designated by the letters s,p,d,f
*s sublevels consists of 1 orbital
*p sublevels consists of 3 orbitals
* d sublevels consists of 5 orbitals
* f sublevels consists of 7 orbitals
•
•
•
•
•
The maximum no of electrons that can
exist in these sublevels is :
S sublevel
P sublevel
d
10
f
2 electrons •
6 •
•
14
•
•
*No more than 2 electrons can
occupy an
orbital
*an electron will occupy the lowest •
possible sublevel
9-The atomic number of the element :
1- the elements are numbered consecutively •
from 1 to 106 coinciding with the number of
protons in the nucleus
2-H element number 1 has 1 proton in nucleus •
3- helium number 2 has 2 protons •
Mg 12 protons •
The atomic number of an element is the same as •
the number of protons in the nucleus ,the same
as positive charge and also number of electrons
in neutral atom.
1-Hydrogen atom atomic number of
the elements:
1- The H atom consisting of a nucleus containing •
one proton and an electron
2- orbital containing one electron ,is the simplest •
known atom
3- The electron occupies an S orbital in the 1st •
energy level
4-the electron doesn’t move in any definite path •
orbital but rather in a random motion within its
forming an electron cloud about the
nucleus •
11- Isotpes of the elements
*atoms of an element having the same atomic •
number but different atomic masses are
called
Isotopes of that element
•
*atoms of the isotopes of an element •
,therefore have the same number of protons
and elements but different numbers of
neutrons
12-atomic structure of the first twenty
elements :
* the structure of the atoms of the 1st 20 •
elements ,arranged in the order of increasing
atomic number (number of protons)
*the atoms of each succeeding element contain •
one more proton and one more electron than the
atoms of the proceeding element.
*the number of neutrons in an atom also •
increases as we progress from the simpler
elements to the more complex one
Periodic table page 23 •
Chemical bonding
2- bonding and molecular structure •
•
*chemical bond : •
The attractive force that hold atoms together •
in compounds are called chemical bonding
Bonding types
•
•
Bonding types
Ionic, covalent and metallic bonding •
1-ionic bond :term given to the electrostatic •
(charge-based)attractive forces which
Hold oppositely charged ions together •
•
2-Covalent bond : the sharing of electrons •
between two atoms that act s to hold the atoms
together
*metallic bond : is found in metals .Atoms of •
the metal are bound to several neighbors,
holding the atoms together but allowing
electrons to move freely
Ionic bonding
*The ionic bond is the electrostatic force •
which attracts particles with opposite
electrical charges
The formation of ions : •
*Atoms can gain or lose electrons to become •
charged particles called ions
Cations: Are positively charged ions formed •
when an atom loses
electrons
Anions: are negatively charged ions formed
when an atom gain electrons
An ion is formed when an atom gain or losses •
one or more electron
M →
X + e- →
M+
e– •
+
x-
•
*If electron lost by M is gained by x ,the overall
reaction will be:
M +
M+
X →M +
+ X- •
+ X-
→M + X - •
The ions attracted to each other because they have •
opposite charges ,the attraction is called an ionic bond
or electrovalent bond .
Lewis structure :
*Lewis discover a Lewis structure in which the •
chemical symbol for an atom is surrounded by
a number of dots corresponding to the
number of electrons in the valence shell of the
atom .
e.g Na atom has one valence –shell-electron •
so its Lewis structure is
Na .
•
. •
e.g. chlorine atom has 7 valence –shell electrons
so its lewis structure is:
.
:Cl:
..
•
•
•
The symbol = the nucleus plus all the inner
shell electrons ,It called the core
E.g Al=13 •
Electronic configration is 1S2 2S2 2P6 3S2 3P1 •
Lewis structure
..
Al .
Here the valences are shown as a pair (the 2
3S electrons) And a single electron (the 3
P)
•
•
•
Octet rule :
* The octet rule is a statement of the stability •
Of the nS2 –nP6 valence-shell configuration. •
Atoms which can achieve this configuration by •
the addition of only a few electrons that is,
tend to complete the octet . In adding
electrons the atom becomes a negative ion.
Thus the chloride ion is formed when one
electron adds to a chlorine atom.
..
.Cl :
..
+
e
-
..
→[ :Cl : ]
..
*Here the negative sign is written •
because the resulting particle is anion
In positive ions ,when has few valence electrons
and has an octet in the second shell from the
out side ,it tend to lose it
Valence electrons thereby exposing the octet. •
In this way the resulting positive ion ends up •
with an octet in what is now its outer shell.
Thus the sodium ion tends to lose its valence •
Electron to form a sodium ion : •
Na(1S2 2S2 2P6 3S 1)→Na+ (1S2 2S2 2P6 )+e - •
Na.
→ Na +
+
e- •
Lewis structure and ionic compounds
* To write the Lewis structure for an ionic •
compound ,we write structures for the
individual ions. Thus the Lewis structure for
NaCl is:
..
•
Na+ [ : Cl :] - •
.. •
Note that:
* The octet rule help us to predict •
stoichiometry that is , atomic combining ratio
in ionic compounds.In the NaCl example one
electron was transferred from one Na atom to
one Cl atom •
..
..
•
Na. + .Cl: → Na+ [ :Cl:] - •
..
..
•
In sodium oxide
Oxgyen has only six valence electrons so need •
to complete its octet
..
:O.
.
+
2e -
→
..
•
[ : O :]2 – •
..
•
Because Na atom has only one valence electron
to lose so 2 Na atom are requried to furnish
two electrons to a single electron
Na.
..
..
↘:O.} → { Na+
[ :O: ] 2Na. ↗ .
Na +
..
the Lewis structure for sodium oxide can
..
be written as 2 Na+ [:O:] 2..
•
•
•
•
•
•
•
Write the Lewis structure for
calcium chloride
Ca in group IIA of the periodic table,has 2
valence electrons,while chlorine in group VIIA
And has seven.A calcium atom can by losing
its 2 valences electrons,convert 2 Cl atoms to
ions
..
Ca 2+ 2[ : Cl:] ..
•
•
•
•
•
•
In this type of bonds
One atom has a low ionization energy the •
affinity other has a high electron
So one or two electron transfer from the first •
to the second .to form an ionic bond
Covalent bonding
*Covalent bonding occurs when 2 atoms are •
more nearly alike in their tendencies to gain
and lose electrons. •
So outright transfer of electron doesn’t occur. •
Instead, electrons are shared between the
atoms
Formation of covalent bond :
*In H2 molecule the H and H •
There are : •
Attractive force between electron of one atom •
and the nucleus of the other
And
•
Repulsion between the electron of one atom •
and the electron of the other atom
As 2 hydrogen atoms approach
each other
Each electron begins to ‘sense’ •
Electrostatically the presence of the nucleus •
Of the opposite atom. •
In terms of quantum mechanics this results in •
an increase in the probability of finding the 1st
atom’s electron near the second atom’s
nucleus and vice versa .
Eventually
*Each electron is equally influenced by the 2 •
nuclei ,and so the probability of finding each
electron is the same at each nucleus .
* so the 2 electrons occupy the same region of •
space.
*in any covalent bond the distance between •
the nulei of the bonded atoms is called the
bond distance or bond length.
Lewis structures and covalent
bonding :
Covalent bond: a bond formed between 2 •
Atoms by sharing of electrons •
Lewis structure for H2 and Cl2 •
H. +H. → H:H = H-H •
..
..
.. .. .. .. •
:Cl. + :Cl. → :Cl:Cl: =:Cl-Cl: •
..
..
.. ..
.. .. •
Lewis dot structure of hydrogen
flouride :
..
H. :F. →
..
..
•
H---F:
..
•
•
Drawing Lewis structure:
1 -Sum the valence electrons from all atoms in •
the species
2- write the atomic symbols for the atoms •
involved so as to show which atoms are
connected to which ,draw a single bond
between each pair of bonded atoms
3- Complete the octets of the atoms bonded •
to the central atom
4- Place leftover electrons on the central atom
even if it results in the central atom
having more than an octet
5- If there are not enough electrons to give •
the central atom an octet ,form multiple
bonds by pulling terminal electrons from a
peripheral atom and placing them into the
bond with the central atom
Draw the Lewis structure for
ammonia NH3
* Since each H can form only one covalent •
bond,the arrangement of atoms must be :
H
•
HNH
•
* From the periodic table ,N have 5 •
valence electrons .These ,plus one electron
from each H ,give a total of 8 .
Bonding the atoms in the molecule requires the
use of six valence electrons,as :
H
..
H:N:H
The remaining 2 valence electrons are then
assigned to N to complete its
octet
•
•
•
•
3H.
.
+. N . →
..
H
..
H:N:H
..
•
•
•
•
MOLECULAR STRUCTURE
1-Molecular structure and •
covalent bond theories
VALENCE SHELL ELECTRON PAIR
REPULSION (VSEPR)THEORY
*In a molecule composed of a central atom •
bonded covalently to several peripheral atoms
the bonding and lone pairs are oriented so
that electron-electrons are minimized while
electron nucleus attraction are maximized.
The method of determining this orientation is
called the valence-shell electron-pair
Repulsion or VSEPR method. The •
assumptionbehind the method are:
1- electron pairs in the valence shell tend to orient
themselves so that their total energy is minimized.
This means that they approach the nucleus as •
closely as possible,while at the same time •
staying as far away from each other as •
Possible,thus minimizing interelectronic •
repulsions. •
2-because lone pairs are spread out more
broadly than are bonding pairs
*repulsions are greatest between two lone •
pairs,intermediate between a lone pair and a
bonding pair, and weakest between two bonding
pairs.
Bonding pair
bonding pair lone pair
→
•
Increasing repulsion •
•
3-Repulsive forces decrease sharply with
increasing interpair angle
They are strong at 90 ◦ ,much weaker at •
120◦,and very weak at 180◦ •
Steric number and electron –pair
orientation:
*the first step in the VSEPR method for •
determining the shape of a molecule is to
draw its Lewis structure in order to find out
how many electron pairs are located around
the central atom.
* consider arsenic trichloride ,and sulfur •
tetraflouride as example .Their Lewis
structures are:
The steric number is defined as :
The total number of electron pairs (lone •
and bonding) around the central atom.
So arsenic has a steric number of 4 in arsenic •
trichloride
While in sulfur tetraflouride the steric •
number of sulfur is 5 (the valence shell of
sulfur has been expanded to 10 electrons.)
Special orientation of electrons
pairs around a central atom:
STERIC NUMBER
ORIENTATION
ANGLE
2
3
4
5
6
linear
180
120
109.5
90-120
90
Triangular planar
tetrahedral
Trigonal bipyramidal
octahedral
Valence – bond theory and orbital
overlap:
*Two approaches have been used for the •
purpose of describing the covalent bond and
the electronic structures of molecules.
1- Valence –bond(VB) theory ,consider that •
when a pair of atoms forms a bond ,the
atomic orbitals of each atom remain
essentially unchanged and that a pair of
electrons occupies an orbital in each of the
atoms .
2 -Molecular orbital (MO)
*this theory assume that the •
atomic orbitals of the original
unbonded atoms become replaced
by a new set of molecular energy
levels called molecular orbitals,
The hydrogen molecule:
* the hydrogen molecule formed from 2 •
isolated ,ground- state hydrogen atoms.
*each atom has at start a single electron in its •
atomic orbitals .
*If we call the two atoms A and B. •
*after the covalent bond has been •
formed,each electron now exists in the 1S
orbitals of both atoms.
According to valence-bond theory
Simultanious occupaucy of orbitals of 2 atoms by a •
pair of electrons is possible if the orbitals
overlap each other to an appreciable extent. •
The orbital overlap produces a region of enhanced •
electron probability denisty located directly
between the nuclei .
*the bond axis (the line connecting the 2 •
nuclie)passes through the middle of this region.
the bond in hydrogen is a sigma (Ơ) bond
* in which the charge –cloud of the chared •
pair is centered on and is symmetrical around •
the bond axis ..
•
Pi-bonding :
When p orbitals overlap sideways ,the results •
Are different .the resulting side to side overlap •
produces enhanced electron probability
density in two regions which are on opposite
sides of the bond axis .this is characteristic of
a pi ( π) bond.
Hybrid orbitals :
*Carbon forms countless compounds in which •
its atoms bond covalently to 4 other atoms .
E.g. methan CH4 •
How can we describe the 4 covalent bonds in •
this molecule in terms of orbital overlap?
The ground state electronic
configuration of C is
•
C
↓↑
1S
↑↓
↑
2S
2P
↑
•