atomic structure
DESCRIPTION
TRANSCRIPT
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