atomic structure

Atomic Structure

The Structure of the Atom

Atoms are basic building blocks of matter, and cannot be chemically subdivided by ordinary means.

The Atom

in 1808 John Dalton (an English scientist) states a theory about the nature of the elements known as Dalton’s atomic theory , the main ideas of this theory can be stated as follows:

All atoms of a given element are identical.

Elements are made of a tiny particle called an atom.

The atoms of a given element are different from those of any other element.

Atoms of one element can combine with atoms of other elements to form compounds.

In chemical reactions atoms are neither created nor destroyed they simply change the way they are grouped together.

Later on Thomson & Rutherformed worked on the structure of the atom & they discovered that the atom is

consist of a tiny nucleus ( about 10-13 cm in diameter) and electrons that move around the nucleus . The nucleus contains protons which have a positive charge equal in magnitude to the electron's negative charge , and neutrons , which have almost the same mass as a proton but no charge.

Atoms are composed of three type of particles:

Protons

Neutrons

Electron

The mass & charge of the electron, proton & neutron are given below:

particleparticleRelative Relative massmass

Relative Relative chargecharge

electronelectron 11- - 11

protonproton 18361836+ + 11

neutronneutron 18391839No chargeNo charge

Modern atomic theory:

In 1911 Niels Bohr construct a model of the hydrogen atom with quantized energy levels ,

Bohr picture the electrons moving in circular orbits (like planet orbiting the sun) , corresponding to the various allowed energy levels. He suggested that the electron could jump to different orbit by absorbing or emitting energy.

Bohr’s atomic orbital:

• Is a specific path on which the electrons travel about the nucleus.

+

--

-

Although Bohr's model opened the way for the later theories it is

important to realize that electrons do not move around the nucleus in circular orbits like the planet orbiting the sun.

Later on Schrodinger found that it is not precisely to describe the

electrons path , he could only predict the probability of finding the electron at a given point in space around the nucleus .

The probability map, or orbital that describes the hydrogen electron in it's lowest possible energy state . The more intense the color of a

given dot the more likely it is that the electron will be found at that point.

In its ground state the hydrogen electron has a probability map

Shrodinger showed that the orbitals of electrons are regions of

electron density with the location and routs of electrons described as probabilities.

Atoms with the same number of

protons but different number of

neutrons. In nature elements are

usually found as a mixture of

isotopes.

Isotopes:

Hydrogen 1 (hydrogen)

1 proton, 0 neutronsMass number = 1

  Hydrogen 2 (deuterium)

1 proton, 1 neutronMass number = 2

  Hydrogen 3 (tritium)

2 neutronsMass number = 3

Example

Three isotopes of elemental carbon are C6

12 , C613, C6

14 . Determine the number of each of the three types sub atomic particles in each of these carbon atoms.

Mass number: The sum of the number of neutrons & number of protons in a given nucleus.

Atomic number: The number of protons in the nucleus.

Electrons may be added to a certain atoms to form a negatively charged particle (anion). These charged particles whether positive or negative are called ions.

Ions:

Under certain circumstances it is possible to remove electrons from a neutral atom leaving a positively

charged particle (cation) .

Quantum numbers:

The various orbitals available to an atom are

described by four quantum numbers, which can

take certain values to create differently sized and

shaped orbital of various energies:

Atomic Orbital

describes the:

Shells:(an electron shell is collection of orbital's)

Size of orbital

Numbered 1, 2, 3, 4, 5, etc

are often lettered (K, L, M, etc.).

The principal quantum number (n)

describes the:

Subshells : are groups of orbitals with in an electron

Shape of orbital (number of lobes).

Given letters s, p, d, f, g, h, i, etc.

The values of ( l ) run from 0 to n − 1

The subsidiary quantum number (l)

describes:

Orientation of the orbitals in space.

Named after the directions they point in (x, y, z, etc.)

Can also be given numbers ranging from 0, ±1, ±2, ±3 … ±l.

The magnetic quantum number (m)

The spin quantum number (s).

describes:

spin of an electron on its own axis

May have the values +½ or -½.

Two electrons in the same orbital have paired

(opposite) spins

The Quantum Numbers

namesymbolvalues

Principal Quantum Numbernany integer from 1 to infinity

Subsidiary Quantum Numberlany integer from 0 to n-1

Magnetic Quantum Numbermany integer from - l to + l

Spin Quantum Numbers -/+1/2

L-valueorbital type

No. of orbitalsMax no. of electrons

0s12

1p36

2d510

3f714

Using symbols, the valid quantum states can be listed in the following manner: 1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f 5g

6s 6p 6d 6f 6g 7h

7s 7p 7d 7f 7g 7h 8i

Atomic Orbitals

The S Orbital

The simplest orbital in the atom is the 1s orbital.

The 1s orbital is simply a sphere of electron density.

There is only one s orbital per shell

The s orbital can hold two electrons have different

spin quantum numbers l = 0

m = 0

S = +1/2 , -1/2

1s Shell

The P Orbitals

Starting from the 2nd shell, there is a set of p

orbitals

There are 3 choices for the magnetic quantum

number, which indicates 3 differently orientated

p orbitals px, py, and pz

each orbital can accommodate two electrons,

giving a total capacity of 6 electrons.

p orbitals are very often involved in bonding

l = 1

m = -1, 0, 1

S = +1/2 , -1/2

2 px

2py

2pz

px py pz

Electronic configuration of elements

Rules for electronic configuration:

The Pauli Exclusion principle:

Only two electrons with opposite spin can occupy an atomic orbital (or no two electrons have the same (4) quantum numbers n, l, m and s)

1s

Hund's Rule:

Electrons prefer parallel spins in separate orbitals of subshells

The Aufbau Principle:

Explains the order in which the electrons fill the various orbitals in an atom.

Filling begins with the orbitals in the lowest- energy shells and continues through the higher-energy shells

The energy relationships among the first three levels of orbitals;

1s 2s 3s 4s 5s 6s 7s

2p 3p 4p 5p 6p 7p

3d 4d 5d 6d

Example :

Write a complete electronic configuration for the noble gases; He, Ne, Ar, Kr, Xe,

Solution:

This can be written in the normal configuration representation As shown below:

He2 1s2

Ne10 1s2 2s2 2p6

Ar18 1s2 2s2 2p6 3s2 3p6

Kr36 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

Xe54 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

Or can be shown in the box representation as shown below :

He2

1s

Ne10

Ar18

1s

1s

2s

2s

2p

2p 3s 3p

Example :

Write a complete electronic configuration for each of the eight elements; from Na to Ar:

Na11 1s2 2s2 2p6 3s1 or [Ne] 3s1

Mg 12 1s2 2s2 2p6 3s2 [Ne] 3s2

Al13 1s2 2s2 2p6 3s2 3p1 [Ne]3s2 3p1

Si14 1s2 2s2 2p6 3s2 3p2 [Ne] 3s2 3p2

P15 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3

S16 1s2 2s2 2p6 3s2 3p4 [Ne] 3s2 3p4

Cl17 1s2 2s2 2p6 3s2 3p5 [Ne] 3s2 3p5

Ar18 1s2 2s2 2p6 3s2 3p6 [Ne] 3s2 3p6

Special electronic configuration: The pairing of electrons raise the orbital energy

slightly.

Half-filled and full-filled subshell low the energy.

For example the electronic configuration of Cr and

Cu

Cr24

1s 2s 2p 3s 3p 4s 3d

Cu29

1s 2s 3s 4s 3d2p 3p

1 -which of the following pair of atoms contain the same

number of neutrons?

A- C614 & C6

12

B- F919 & Ne10

22

C- S1632 &Al13

29

Tutorial (A)

2 -Which of the following particles does not contain

the same number of electrons as fluoride ion (F -)

A- Ne

B- Li +

C- Na +

3 -the electronic configuration of helium atom ,boron

atom , carbon atom & element X are given below.

Which one could be the electronic configuration of

element X?A- 1s2 2s22P1

B- 1s2

C- 1s2 2s1

D- 1s2 2s2 2p2

4- atoms which have the same electronic configuration are said to be isoelectronic which of the following is not isoelectronic with O2-?

A- N 3-

B- Al 3+

C- Na +

D- Na

5 -the magnetic quantum number (m) gives

A- the sub shells

B- the orbitals

C- the spin of the electron .

Tutorial (B)

1- an orbital can take a maximum of…………..electrons.

2- an( s ) sub level can take a maximum of……electrons.

3- the (p) sub level can take a maximum of…….electrons.

4- how many electrons can fit in to a set of (5d) orbitals?

5- what is the maximum numbers of electrons in levels,

2, 3, 4 ?

6 -which of the following orbitals could not be exist?

1s,1p,1d,2s,2p,2d,3p,3d,3f,4s,4f.

7- write the four quantum numbers for energy

level(4)?

8- For the H-like atom, which subshell has the highest

energy level?

4f, 3d, 2p, 1s

9- How many electrons are required to fill all the

following subshells?

1s 2s 2p 3s 3p 4s 3d 4p

10- Which of the following two electronic configuration

is more stable?

a- [Ar]4s1 3d5

b- [Ar]4s2 3d4

11- Which of the following two electronic configurations is more stable?

a [Ar]4s2 3d9

b [Ar]4s1 3d10

12- Choose the electronic configuration for palladium, Pd (Z = 46).

a- [Kr]5s1 4d7

b- [Kr]5s1 4d8

c- [Kr]5s0 4d10

d- [Kr]5s1 4d10

By the late 1800's many elements had already been discovered.

The scientist Dmitri Mendeleev, a Russian chemist, proposed an arrangement of known elements based on their atomic mass

The modern arrangement of the elements is known as the Periodic Table of Elements and is arranged according to the atomic number of elements.

The Periodic Table of Elements

Periodic law:

The modern periodic law states that ; when elements

are arranged by atomic number; their physical and

chemical properties vary periodically .

The periodic table displays the elements in rows (periods ) and columns in order of increasing atomic number.

Elements that have similar chemical properties fall into vertical columns called groups or families

Most of the elements are metals and located on the left hand side of the periodic table

The nonmetals appear on the right hand side of the periodic table

From the periodic table we can know many information

directly like symbol, atomic number, atomic mass & you

can also know whether the element is metal, non metal

or metalloid.

Today's periodic table consist of seven horizontal rows

called periods & a number of vertical columns called

groups.

All elements in each group have the same number of

electrons in their outer most shells so the behave similarly.

We have two types of groups:

Group (A): ( representative or main elements)

Group (B):( transition elements)

Metals : are found on the left-hand and at the centre of the periodic table

Non metals: are relatively few they are in the upper-right hand corner the table

Metalloids : those are few elements exhibit both metallic and nonmetallic behavior (also known as semimetals)

Group 1A metals (Li, Na, K, Rb, Cs, and Fr) are called the alkali metals , and they are the most reactive metals in the periodic table.

Group 2A metals (Be, Mg, Ca, and Ra) are called the alkaline earth metals.

Group 1B (Cu, Ag and Au) are called the coinage metals

Although these elements have outer electronic configuration similar to those of the alkali metals and the alkaline earth metal the are much less reactive

Group 7A (F, Cl, Br and I) are called halogens and they are the most reactive nonmetals in the periodic table

Group 8A (He, Ne, Ar, Kr, Xe, and Rn) are called the rare or noble gases,

these gases are characteristic by their completely filled shells, for this reason they are chemically inert.

Physical properties of elements

Ionization Energy (I):Is the energy required to remove an electron from an isolated atom in its ground state.X(g)

Chemical Bonds:

In nature elements or compounds exist due to combination of similar or different atoms

Atoms and molecules are electrically neutral, according to this fact in atom combination each tries to exhibit 8 electrons on the outer most shell that by loosing, gaining or sharing electrons.

Metals: have three or less electrons e.g. Na, Ca, Fe

Nonmetals: Have five or more electrons e.g. H, N, S

Carbon atom: have four electrons in the outer most shell liable to loose them or to gain more electrons

The Bonds By which atoms can combine are :

Ionic Bond

Covalent Bond

Co-ordinate Bond

Ionic BondIonic Bond

This is Characteristic of metallic and non metallic

combination forming a neutral molecule.

It is achieved via two steps:

Ionization of atoms

Formation of the molecule

Ionization of atoms:

Na Na+

+ e-

Ca Ca+ +

+ 2e- Cations

Cl Cl-

+ e-

S + 2e S- - Anions

Formation of the molecule

In formation of molecules, loss or gain of electron

occurs at the same time.

e.g. Formation of NaCl:

Na Na+

+ e -

Chlorine atom attacks this electron to its outermost

shell and become an anion.

Cl Cl+ e- -

Thus two ions of different charges of equal number are

formed, attraction between them taken place, leading

to the formation of the neutral molecule NaCl.

Na+

Cl-

+ NaCl

NaCl is said to be ionic compound

Example:

Formation of CaCl2 ;

Ca Ca+ +

+ 2e-

Cl Cl-

+ e-

Ca+ +

Cl-

Cl-

+CaCl2

Covalent BondCovalent Bond

The covalent bond is the chemical bond in which

two or more non metal atoms share electrons Both atoms are unable to loose or gain electrons

By sharing electrons both atoms reach octet state

e.g. Fluorine atom has 7 electrons in its outermost

shell. It needs one electron to reach its octet.

To achieve this F atom shares an electron

belonging to another F atom

By this F2 molecule is formed

F + F F F

F F F2

Single Covalent Bond

Examples:

Water molecule:

OH H ++ OH H

OH H H2O

Single Covalent Bond

Ammonia molecule:

N + 3H N

H

HH

NH H

H

NH3

Single Covalent Bond

Oxygen molecule:

O O+ O O

O O O2

Double Covalent Bond

Nitrogen molecule:

N N+ N N

N N N2

Triple Covalent Bond

Co-ordinate BondCo-ordinate Bond

Coordinate covalent bonding is a special type of

bonding, in which the bonding electrons originate

solely from another atom.

e.g. SO2 molecule:

O + O + S SO O

O OS SO2

Co-ordinate Covalent Bond