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General Chemistry Assistant professor Mervat Mohamed Hosny

General Chemistry Assistant professor Mervat Mohamed Hosny

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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 atomSchroedinger ))

•*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 1-7النواة • عن االلكترون بعد

L=2nd quantum no explain the shape oforbital

•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 electronsthe 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 2 electrons

• P sublevel 6•d 10

•f 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 but rather in a random motion within its orbital 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 → M+ + e –

•X + e- → x-

*If electron lost by M is gained by x ,the overall :reaction will be

•M + X →M + + X -

• M + + 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 .

•*Here the negative sign is written because the resulting particle is anion

.. ...Cl : + e - →[ :Cl : ] -

.. ..

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: ] 2-

• .. + Na . ↗ 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

44-- Place leftover electrons on the central atom Place leftover electrons on the central atom even if it results in the central atom even if it results in the central atom octet octet having more than an having more than an

•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

•H N H

• * 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

• H• .. .

•3H. + . N . → 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 NUMBERORIENTATIONANGLE

2linear180

3Triangular planar120

4tetrahedral109.5

5Trigonal bipyramidal90-120

6octahedral90

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

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