Atomic Structure and Interatomic Bonding

Frank venance,mcsematerial science and engineering

C h a p t e r 2 Atomic Structure andInteratomic Bonding

WHY STUDY ATOMIC STRUCTURE AND INTERATOMIC BONDING?

Reason is that, in some instances, the type of bond allows us to explain a material’s properties. For example, consider carbon, which may exist as both graphite and diamond. Whereas graphite is relatively soft and has a “greasy” feel to it, diamond is the hardest known material. This dramatic isparity in properties is directly attributable to a type of interatomic bonding found in graphite that does not existin diamond

INTRODUCTION

Some of the important properties of solid materials depend on geometrical atomic arrangements, and also the interactions that exist among constituent atoms or molecules.

Atomic StructureAll atom consists of a very small nucleus composed of protons and neutrons, whichis encircled by moving electrons. Both electrons and protons are electrically charged,the charge magnitude is 1.6x10-19c electrons is negative in sign Protons is positive sign neutrons are electrically neutral.

Masses for these subatomic particles are infinitesimally small; protons and neutrons have approximately the same mass, 1.67x10-27 kgMass of electron is approximately 9.11x10-31 kg

ATOMIC NUMBER (Z)

atomic number is the number of protons in the nucleus in element

If the number of protons is equals the number of electrons the atom is electrically neutral or complete atom, This atomic number ranges in integral units from 1 for hydrogen to 92 for uranium, the highest of the naturally occurring elements.

THE ATOMIC MASS (A)

The atomic mass is the sum of the masses of protons and neutrons within the nucleus. A=Z+NAlthough the number of protons is the same for all atoms of a given element, the number of neutrons (N) may be variable. Thus atoms of some elements have two or more different atomic masses, Isotopes is atoms of some elements have two or more different atomic masses,The atomic weight of an elementIs weighted average of the atomic masses of the atom’s naturally occurring isotopes

ELECTRONS IN ATOMS

ATOMIC MODELS1.QUANTUM MECHANICS is of set of principles and laws that govern systems of atomic and subatomic entities atom

. An understanding of the behavior of electrons in atoms and crystalline solids necessarily involves the discussion of quantum-mechanical concepts.BOHR ATOMIC MODEL In Bohr atomic model, in which electrons are assumed to revolve around the atomic nucleus in discrete orbitals, and the position of any particular electron is more or less well defined in terms of its orbital.

BOHR ATOMIC MODEL

With his model, Bohr explained how electrons could jump from one orbit to another only by emitting or absorbing energy in fixed quanta. For example, if an electron jumps one orbit closer to the nucleus, it must emit energy equal to the difference of the energies of the two orbits. Conversely, when the electron jumps to a larger orbit, it must absorb a quantum of light equal in energy to the difference in orbits

2;WAVE-MECHANICAL MODEL

The Bohr model was some significant limitations because of its inability to explain several phenomena involving electrons. A resolution was wave-mechanical model, in which the electron is considered to exhibit both wave-like and particle-like characteristics. With this model, an electron is no longer treated as a particle moving in a discrete orbital; rather, position is considered to be the probability of an electron’s being at various locations around the nucleus. In other words, position is described by a probability distribution or electron cloud.

QUANTUM NUMBERSQuantum number is used to describe the distribution of electrons in the atom. The only information that was important was the size of the orbit, which was described by the n quantum number. the electron to occupy three-dimensional space. It therefore required three coordinates, or three quantum numbers, to describe the orbitals in which electrons can be found. These quantum numbers describe the size, shape, and orientation in space of the orbitals on an atom.The three coordinates that come from wave equations areI. The principal quantum number (n) describes the size of the

orbital. Orbitals for which n = 2 are larger than those for which n = 1, Because they have opposite electrical charges, electrons are attracted to the nucleus of the atom. Energy must therefore be absorbed to excite an electron from an orbital in which the electron is close to the nucleus (n = 1) into an orbital in which it is further from the nucleus (n = 2). The principal quantum number therefore indirectly describes the energy of an orbital.

quantum numbers dictate the number of states within each subshell. Shells are specified by a principal quantumnumber n, which may take on integral values beginning with unity; sometimes these shells are designated by the letters K, L, M, N, O, and so on, which correspond, respectively, to as indicated

2;The angular quantum number or The second quantum number(l) describes the shape of the orbital.Orbitals have shapes that are best described as spherical (l = 0), polar (l = 1), or cloverleaf (l = 2). They can even take on more complex shapes as the value of the angular quantum number becomes larger.The second quantum number, l, signifies the subshell, which is denoted by a lowercase letter an s, p, d, or f; it is related to the shape of the electron subshell.

3,magnetic quantum number (m), Used to describe the orientation in space of a particular orbital. (It is called the magnetic quantum number because the effect of different orientations of orbitals was first observed in the presence of a magnetic field.)

Electron Configurations

electron states -values of energy that are permitted for electrons.To determine the manner in which these states are filled with electrons, we use the Pauli exclusion principle, this principle stipulates that each electron state can hold no more than two electrons, which must have opposite spins. Thus, s, p, d, and f subshells may each accommodate,respectively,

GROUND STATEWhen all the electrons occupy the lowest possible energies in accord with the foregoing restrictions, an atom is said to be in its ground state For most atoms, the electrons fill up the lowest possible energy states in the electron shells and sub shells, two electrons (having opposite spins) per state.

valence electrons are those electron that occupy the outermost shell.

ELECTRON CONFIGURATION OF AN ATOM

The electron configuration of an atom is a form of notation which shows how the electrons are distributed among the various atomic orbital and energy levels.The format consists of a series of numbers, letters and superscripts as shown below:1s2

      Here we see the electron configuration for the element helium.  This electron configuration provides us with the following information:

•The large number "1" refers to the principle quantum number "n" which stands for the energy level.  It tells us that the electrons of helium occupy the first energy level of the atom.•The letter "s" stands for the angular momentum quantum number "l".  It tells us that the two electrons of the helium electron occupy an "s" or spherical orbital.•The exponent "2" refers to the total number of electrons in that orbital or sub-shell.  In this case, we know that there are two electrons in the spherical orbital at the first energy level.

Order of Filling Sublevels with Electrons

The next thing that we need to recall is the fact that the energy sublevels are filled in a specific order that is shown by the arrow diagram seen below:

Table 2.2 A Listing of the Expected Electron Configurations for

Some of the Common Elements

Atomic Element Symbol Number Electron Configuration Hydrogen H 1 1s1 Helium He 2 1s2 Lithium Li 3 1s22s1 Beryllium Be 4 1s22s2 Boron B 5 1s22s22p1 Carbon C 6 1s22s22p2 Nitrogen N 7 1s22s22p3 Oxygen O 8 1s22s22p4 Fluorine F 9 1s22s22p5 Neon Ne 10 1s22s22p6 Sodium Na 11 1s22s22p63s1

Magnesium Mg 12 1s22s22p63s2

Aluminum Al 13 1s22s22p63s23p1

Silicon Si 14 1s22s22p63s23p2

THE PERIODIC TABLE

All the elements have been classified according to electron configuration in the periodic table.Here, the elements are situated, with increasing atomic number, in seven horizontal rows called periods. The arrangement is such that all elements arrayed in a given column or group have similar valence electron structures, as well as chemical and physical properties.These properties change gradually, moving horizontally across each period and vertically down each column.

The elements positioned in Group 0, the rightmost group, are the inert gases,which have filled electron shells and stable electron configurations. Group VIIA and VIA elements are one and two electrons deficient, respectively, from having stable structures. The Group VIIA elements (F, Cl, Br, I, and At) are sometimes termed the halogens. The alkali and the alkaline earth metals (Li, Na, K, Be, Mg, Ca, etc.) are labeled as Groups IA and IIA, having, respectively, one and two electrons in excess of stable structures.The elements in the three long periods, Groups IIIB through IIB, are termed the transition metals, which have partially filled d electron states and in some cases one or two electrons in the next higher energy shell. Groups IIIA,IVA, and VA (B, Si, Ge, As, etc.) display characteristics that are intermediate between the metals and nonmetals by virtue of their valence electron structures.

  1A IIA IIIB IVB VB VIB VIIB 8 9 10 1B IIB IIIA IVA VA VIA VIIA 0

1H He

1 2

2Li Be B C N O F Ne

2,1 2,2 2,3 2,4 2,5 2,6 2,7 2,8

3Na Mg Al Si P S Cl Ar

2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8

4K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

2,8,8,1 2,8,8,2 2,8,9,2 2,8,10,2 2,8,11,2 2,8,13,1 2,8,13,2 2,8,14,2 2,8,15,2 2,8,16,2 2,8,18,1 2,8,18,2 2,8,18,3 2,8,18,4 2,8,18,5 2,8,18,6 2,8,18,7 2,8,18,8

5Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

2,8,188,1

2,8,188,2

2,8,189,2

2,8,1810,2

2,8,1812,1

2,8,1813,1

2,8,1814,1

2,8,1815,1

2,8,1816,1

2,8,1818,0

2,8,1818,1

2,8,1818,2

2,8,1818,3

2,8,1818,4

2,8,1818,5

2,8,1818,6

2,8,1818,7

2,8,1818,8

6 Cs Ba * Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

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VII

Element Groups (Families)

Alkali Earth Alkaline Earth Transition Metals

Rare Earth Other Metals Metalloids

Non-Metals Halogens Noble Gases

electropositive elements,is the elements that they are capable of giving up their few valence electrons to become positively charged ions. Electronegative;is the elements that they readily accept electrons to form negatively charged ions, or sometimes they share electrons with other atoms. Furthermore, the elements situated on the right-hand side of the table are displays electronegativity values that have been assigned to the various elements arranged in the periodic table.Generally:electronegativity increases in moving from left to right and from bottom to top. Atoms are more likely to accept electrons if their outer shells are almost full,

Atomic Bonding in Solids

BONDING FORCES AND ENERGIES

physical properties of materials is predicated on a knowledge of the interatomic forces that bind the atoms together.Perhaps the principles of atomic bonding are best illustrated by considering the interaction between two isolated atoms as they are brought into close proximity from an infinite separation.At large distances, the interactions are negligible,but as the atoms approach, each exerts forces on the other.

These forces are of two types, attractive and repulsive, and the magnitude of each is a function of the separation or interatomic distance.The origin of an attractive force FA depends on the particular type of bonding that exists between the two atoms. The magnitude of the attractive force varies with the distance,But, the outer electron shells of the two atoms begin to overlap, and a strong repulsive force FR comes into play. The net force FN between the two atoms is just the sum of both attractive and repulsive components; that is, FN =FA +FR

Sometimes it is more convenient to work with the potential energies between two atoms instead of forces. Mathematically, energy (E) and force (F) are related as 

FdrE

r

R

r

NN drFdrFE

RAN EEE

in which EN, EA, and ER are respectively the net, attractive, and repulsive energies for two isolated and adjacent atoms.

Bonding energy between two atoms

• The interaction energy at equilibrium is called the bonding energy between the two atoms.

• To break the bond, this energy must be supplied from outside.• Breaking the bond means that the two atoms become infinitely separated.• In real materials, containing many atoms, bonding is studied by expressing the bonding energy of the entire materials in terms of the separation distances between all atoms,

Bonding

Primary bonding:1.Ionic (transfer of valence electrons)2.Covalent (sharing of valence electrons, directional)3.Metallic (delocalization of valence electrons)

Secondary or van der Waals Bonding:(Common, but weaker than primary bonding)1.Dipole-dipole2.H-bonds3.Polar molecule-induced dipole4.Fluctuating dipole (weakest)

Primary bonding 1.Ionic Bonds

Atoms like to have a filled outer shell of electrons. Sometimes, by transferring electrons from one atom to another, electron shells are filled. The donor atom will take a positive charge, and the acceptor will have a negative charge. The charged atoms or ions will be attracted to each other, and form bonds. The compound NaCl, or table salt, is the most common example.

The electron structure of atoms is relatively stable when the outer shells contain eight electrons (or two in the case of the first shell). An element like sodium with one excess electron will give it up so that it has a completely filled outer shell. It will then have more protons than electrons and become a positive ion (charged atom) with a +1 charge. An atom of chlorine, on the other hand, with seven electrons in its outer shell would like to accept one electron. When it does, it will have one more electron than protons and become a negative ion with a -1 charge. When sodium and chlorine atoms are placed together, there is a transfer of electrons from the sodium to the chlorine atoms, resulting in a strong electrostatic attraction between the positive sodium ions and the negative chlorine ions. This explains the strong attraction between paired ions typical of the gas or liquid state.

Figure A. Formation of ionic bond in NaCl.

Figure B. Na+ and Cl- ions formed by ionic bonding mechanism.

2.Covalent Bonds

Some atoms like to share electrons to complete their outer shells. Each pair of shared atoms is called a covalent bond. Covalent bonds are called directional because the atoms tend to remain in fixed positions with respect to each other. Covalent bonds are also very strong. Examples include diamond, and the O-O and N-N bonds in oxygen and nitrogen gases or H-H

“Sharing” of electrons• Why do some atoms want to share electrons?

Example1: CH4

C: has 4 valence e, needs 4 moreH: has 1 valence e,needs 1 more

3.Metallic (delocalization of valence electrons)

• Valence electrons are completely delocalized to form an electron cloud, in which positive ionic cores are embedded.• The remaining nonvalence electrons and atomic nuclei form “ion cores”, which posses a net positive charge equal in magnitude to the total valence electron charge per atom.• Metallic bonding is found in the periodic table for Group IA and IIA elements.• Electron delocalization is the origin of good electrical and thermal conductivities in metals. (Ionically BUT covalently bonded materials are typically electrical and thermal insulators, due to the absence of large numbers of free electrons) .

Figure 1. Metallic bond.

2.Secondary Bonds: Intermolecular Forces

• Secondary, Van der Waals, or physical bonds are weak in comparison to the primary bonds.• Secondary bonding exists between virtually all atoms or molecules, but its presence may be obscured if any of the three primary bonding types is present.• Secondary bonding forces arise from atomic or molecular dipoles. An electric dipole exists whenever there is some separation of positive and negative portions of an atom or molecule.• Dipole interactions occur between induced dipoles, between induced dipoles and polar molecules (which have permanent dipoles), and between polar molecules.• Hydrogen bonding, a special type of secondary bonding, is found to exist between some molecules that have hydrogen as one of the constituents.

Hydrogen Bonds

Hydrogen bonds are common in covalently bonded molecules which contain hydrogen, such as water (H2O). Since the bonds are primarily covalent, the electrons are shared between the hydrogen and oxygen atoms. However, the electrons tend to spend more time around the oxygen atom. This leads to a small positive charge around the hydrogen atoms, and a negative charge around the oxygen atom. When other molecules with this type of charge transfer are nearby, the negatively charged end of one molecule will be weakly attracted to the positively charged end of the other molecule. The attraction is weak because the charge transfer is small.

Polar molecule-induced dipole interaction:

Polar molecules (with asymmetric arrangement ofpositively and negatively charged regions) can inducedipoles in adjacent nonpolar molecules

-

+

PolarNonpolar(e.g. atom)

_

+

++

+

_

secondarybonding

FRANK VENANCE M032/T.11

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