Chapter 2.

Atomic Structure
• Atomic Structure
• Ionic Bond
• Covalent Bond
• Metallic Bond
• Secondary, or van der Waals, Bond
• Materials Bonding Classification

1
 Atomic Structure
• Materials available for engineering applications are
– Metals, Ceramics, and Polymers
• One basis for classification of materials is the chemical
bonding
– Primary bonding involves sharing of electrons
• Ionic, covalent, metallic
– Secondary bonding involves relatively weak attraction of atoms
with no sharing or exchange of electrons
• Van der waals, hydrogen

2
 Atomic Structure sp3
1s

• Atomic structure consists of a nucleus with orbiting electrons

• Atoms are made up of complex subatomic particles.
– Protons are positively charged identical particles. Mass of 1.66 x 10-24 g
– Neutrons remain at the nucleus or center of the atom and are neutral identical particles. Mass of 1.66 x 10-24 g each
– Electrons are negatively charged particles that orbit the nucleus at speeds approaching the speed of light. Mass of an
electron is 1/1000 that of a proton.
– Mass of an atom consists of the number of protons and neutrons times the mass of each. Atomic mass unit is the
mass of the protons, neutrons, and electrons
• Carbon 12 has 6 neutrons and 6 protons.
• Amadeo Avogadri (1776-1856) Italian physicist coined the word molecule and hypothesized…
• All gases (at a given temp and pressure) contained the same number of molecules per unit volume.
• Recognized after his death

3
 Atomic Structure
• Avogadro’s number represents the number of protons or neutrons necessary to produce a mass of 1 gram.
– Avogadro number is 0.6023 x 1024 amu per gram
– Avogadro’s number of atoms of a given element is termed gram-atom, for a compund is in mole.
• 1 mole of NaCl contains Avogadro’s number of Na atoms and Avogadro’s number of Cl atoms.
• Avogadro’s number of C12 atoms
– Would have a mass of 12.00g
– Naturally occurring Carbon actually has atomic mass of 12.011 amu because not all carbon atoms contains six neutrons in the
nuclei. Some have seven and are called isotropes.
• Variations in combinations of protons, neutrons, and electrons make up difference between atom types.
– The most fundamental difference between atoms is the number of protons.
– Atoms which differ in the number of neutrons are called isotopes.
• e.g., H has 3 isotropic forms: (1) one proton and one neutron, atomic wt 1; (2) one proton and two neutrons, atomic wt 2; (3) one proton
and three neutrons, atomic wt 3. Final atomic weight = weighted average of the 3.

4
 Quantum Mechanics
• Electrons orbit the nucleus of the atom. The tiny solar system
behaves in ways that cannot be predicted by common laws of
physics. Need quantum mechanics.
• Quantum Mechanics is a field of study that uses energy levels,
motion analysis, and probability theories to study atoms.
• In Quantum Mechanics electrons behave in a wavelike fashion
rather than individual particles. Waves can be diverted by reflection
or diffraction.
• The location of the electron is described by energy levels rather than
by individual positions. The higher the energy level the further
away from the nucleus.

5
 Quantum Levels
• Electrons exist in orbits around the nucleus. More than one electron can be
in each orbit due to alternating spins.
• Energy levels appear at predictable intervals in disctict orbits or shells,
e.g., 1s (2 electrons), 2s (2 electrons) and 2p (6 electrons), 3s, 3p, etc.
• s, p, f, and d are Quantum Levels 1, 2, 3, and 4
• Vertical groupings in periodic table are based upon similar electron
configuration and similarities in both chemical and physical properties
– Group 1A alklai metals; Group IIA are alkaline-earth metals
– Group B subgroups- transition elements
– Group IIIA through VA and VIIA are mostly non-metals
– Group VIII- Noble or Inert Gases
– Octet rule- accept/give electrons to fill s and p orbitals (2+6=8)

6
 Periodic Table
• Invented by Dmitri Medeleyev in the late 1800’s
• Many of the elements in the table were not discovered until long after
the table was invented
• All elements are in their most basic form and cannot be simplified
• Table lists the atomic number and atomic mass
• The atomic number is the number of protons in the nucleus or the
center of the atom
• The atomic mass is the sum of the masses of the protons and neutrons.
Electrons weigh about 1/2000 as a proton.
• Carbon (C) has atomic number 6 because there are 6 protons in the
nucleus.

7
 Periodic Table
IA IIA ….Groups… IIIA IVA VA VIA VIIA VIIIA
Atm # 1
Symbol H
Wgt 1.01
B Groups

Lanthanides

Actinides
8
Excellent Reference: http://www-tech.mit.edu/Chemicool/
 Atomic Structure
• The periodic table is a sequential listing of the number of protons from 1 to 110. H has 1 proton, He has 2, Li
has 3…
– The number of protons is the atomic number.
– The sum of protons and neutrons is the atomic weight.
– Atomic weight is not a whole number due to the isotopes that have a different number of neutrons.
• Periodic table has Groups in columns and Periods in Rows
– Group I elements are most likely to give up electrons
– Group II elements are 2nd most likely to give up electrons
– Group VIII are stable and do not accept electrons
– Group VII are very reactive and accept one electron
• Valence
– the number that reflects the number of electrons the atom will usually give up (+1 for H, Li, Na, etc.) or receive (-1 for F, Cl,etc)

9
 Periodic Chart Organization
• Elements are divided into two groups- metals and nonmetals
– Of the 120 known elements (including synthesize), approximately 81 are metals. 92 occur
naturally in the earth.
• Metals have the following characteristics to varying degrees
– High electrical conductivity and High thermal conductivity
– Luster- ability to reflext light. Ductility, maleability
– Loose electrons (low ionization energy) readily when they react with nonmetals
– Metallic character should decrease as we move across the Periodic Table and increase as
we move down.
• Nonmetals tend to be insulators (solid, liquid, or gas)
– Gain electrons in chemical reactions
• Noble gases are inert
• Metalloids are semiconductors

10
 Periodic Table
• Groups
– I: Group 1 is also called the alkali metal group. These are strong
metals that are unusually soft and very reactive toward Oxygen
forming Oxides and water forming hydroxides of the metal.
– II: Group 2 is called the alkaline earth metals. These metals are not
as soft as Group 1 metals. Theyalso react more mildly with
Oxygen to produce oxides of the metals and only react with water
at temperatures where the water is steam.
– Groups 3-12 are referred to as the transition metal groups. These
metals are not as predictable

11
 Periodic Table
• Groups 3B- are referred to as the transition metal groups.
– These metals are not as predictable because of the shielding effect
of the inner electrons. As for the "shielding effect" this refers to
the inner electrons found in the transition state elements and the
inner transition (rare earth)elements. These electrons had a
tendency to block the electrical effect of the positive nucleus upon
the outer valence electrons of those atoms. This shielding effect
helps to partially explain the erratic placement of the electrons in
the d and f orbitals relative to the s and p orbitals.
• Groups 1A-2A and 3A-8A are referred to as the
representative elements
• Group 7A is referred to as the halogen group
• Group 8A is referred to as the Noble gas group previously
known as the inert gas group.
12
 Periodic Table
• The metals which tend to have their atoms losing electrons
during a chemical change are roughly found to the left
Group 3
• Non-metals which tend to have their atoms gaining electrons
during chemical change are roughly found in Group 6A-7A
with some elements in the lower parts of Groups 5A.
• Metalloids which tend to have their atoms sometimes losing
and sometimes gaining electrons during chemical change are
generally found in Groups 4A-6A
• The Noble gases really belong to their own category since
their atoms tend neither to lose or gain electrons. There are
only a handful of compounds involving the Noble Gases
(mostly involving Xenon).
13
 Periodic Table
Properties

• As you proceed to the left in a period or as you proceed down within a group:
– The metallic strengths increase (non-metallic strengths decrease).
– The atomic radius of atoms (distance from the nucleus to the outermost occupied region)
increases. Atomic radii are the distance between the outermost occupied probability
region of an atom and its nucleus.
– The ionization potential (energy required to remove an electron from an atom) decreases.
Ionization Potential is energy required to remove electron from atom.
– The electron affinity (energy released as electron is picked up by an atom) decreases.
– The electronegativity (the electron attracting ability of an atom) decreases.

14
 Bonding
• Bonding occurs when two or more atoms come into close
proximity resulting in an attraction between atoms.
• If atoms are large distance from each other, there is little
interaction.
– As the distance between the atoms decreases, the energy begins to
decrease and the system becomes more stable.
– Eventually, the atoms reach an optimal separation which is the
bottom of the energy curve. The separation is the bond length.
Bond energy is depth of well.
– Atoms too close have repulsive energy.
• Bonding is attributed to interactions between the electrons.
unstable Repulsive energy
Figure 2.3
System
Energy Separated atoms
Bond
Stable Energy
15
Distance between atoms
 Bonding
• Ionic bonding (ceramics, e.g., salt and clay)
– Nondirectional
– Forms when an atoms that has a strong tendency to give up electrons (a metal)
is in close proximity to an atom that has a strong tendency to accept electrons
(nonmetal).
• Transfer of one or more electrons from the outer shell of one atom to the outer shell
of the other atom depending on the valence of the atoms.
• Results in an electron arrangement when many ions (+ and -) are in close proximity,
e.g., NaCl, that has a polar arrangement of the ions similar to a magnet.
• Forms crystalline structure
Metal cation Non-metal anion

e- e- e- e- Na+ Cl- Na+ Cl- Na+ Cl-

e-
e- Cl e- e- Cl- e-
e- e- e- e- Cl- Na+ Cl- Na+ Cl- Na+
e- e-
e- Na+
Na Na+ Cl- Na+ Cl- Na+16 Cl-
atoms ions
 Bonding
• Coordination Number
– Number of adjacent ions (or atoms) surrounding a reference ion
(or atom)
• For each ion in Fig 2-5, the CN is 6 (each has 6 nearest neighbors)
• For ionic compounds, CN of the smaller ion can be calculated by
considering the greatest number of larger ions of opposite charge that can
be in contact with the smaller ion.
• Radius ratio (r/R) of smaller ion (r in nm) and larger ion (R in nm).
• Table 2.1
Coordination Coordination
Number r/R Geometry
2 0 < r/R < 0.155
3 0.155 < r/R < 0.225
4 0.225 < r/R < 0.414

6 0.414 < r/R < 0.732

8 0.732 < r/R < 0.1 17

12 1
 Bonding
• Covalent bonding (most important for plastics)
– Occurs when two nonmetal atoms are in close proximity.
– Both atoms have a tendency to accept electrons, which results in
shared outer electron shells of the two atoms.
– Number of shared electrons is usually to satisfy the octet rule.
– Resulting structure is substantially different that the individual
atoms, e.g., C and H4 make CH4, a new and distinct molecule.
– Atoms is covalent bonds are not ions since the electrons are shared
rather than transferred as in ionic or metallic bonds.
e- e-
H H H
e- e- e-
e-
H H
e- e- C e-
H e- e- e- H
e- C e- e-
18
e- H
 Bonding
• Covalent bonding (most important for plastics)
– Highly directional sharing of valence electrons.
– Ethylene monomer goes to polyethylene polymer with 1,000 to
10,000 repeating units

H H H H H H H H
C C C C C C …
C C
Heat, Pressure,
Initiator H H H H H H
H X

– Table 2.2. Bond Energies of Representative of the C-C, C-H, C-O

bonds. These can be measured in the FTIR lab.
19
 Carbon Chain Polymers
• Homopolymers
– Simplest plastic containing one basic structure
– If X = H then Polyethylene
– If X = Cl the PVC
– If X = CH3 then PP
– If X = Benzene Ring then Polystyrene
• Through Addition Polymerization from monomer

H H H H

C C Heat, Pressure, C C
Initiator
H X H X 20
n
• Metallic bonding
Bonding
– Occurs when two metal atoms are in close proximity.Both atoms have tendency
to give up electrons. Electrons are free to move about entire atoms structure
– Releasing electrons yields a lower energy state.
– The metal atoms approach each other and give up electrons when in close
proximity to a sea of electrons.
– Charged metal ions cancel the repulsive forces due to the electron movement.
– Crystal structures can form in some atoms but the forces are not as strong as
ionic bonds in ceramics.
– Metallic alloys can form when each gives up electrons and form a positively
charged ion.
Metal cations Sea of electrons
Fe atom
e- e- Fe atom Fe++ e- Fe++ e- Fe++ e- Fe++ e-
Fe Fe e- e- e- e-
e- e- Fe++ e-Fe++ e- Fe++ e- Fe++ e-
e- e- e- e-
e- Fe++ e- Fe++ e-Fe++ e- Fe++ e-
Fe++ Fe++ e-
e- e- e- e- e- e- 21
Electrons (free to move)
 Metallic Bonding
• Bond energies for metals and ceramics are best
characterized by energies of the bulk metal are more useful
than for individual atoms as with covalent
• Heat of sublimation
– Represents the amount of thermal energy necessary to turn 1 mold of solid
directly into vapor at a fixed temperature.
– Indicates the relative strength of the metallic bonding in the solid
– Table 2.3
Metal/oxide Heat of Sublimation, KJ/mol
Al 326
Cu 338
Fe 416
Mg 148
Ti 473
FeO 509
MgO 605
TiO 597
22
 Secondary Bonding
• Secondary bonding: weaker than ionic, metallic, covalent
– Hydrogen bonding
• Occurs between the positive end of a bond and the negative end of another
bond.
• Example, water the positive end is the H and the negative end is O.
– van der Waals
• Occurs due to the attraction of all molecules have for each other, e.g.
gravitational. Forces are weak since masses are small
– induced dipole
• Occurs when one end of a polar bond approaches a non-polar portion of
another molecule.

23
 Bonding Classification
• Relative Bond Energies
– Can be obtained with comparison of melting point
• Sufficient energy to break cohesive bonds.
Material Bonding Type Melting Point
NaCl Ionic 801 C
C diamond Covalent 3550 C
HDPE Covalent and Seondary 120 C
Cu Metallic 1085 C
Ar Secondary Induced dipole -189 C
Water Secondary permanent dipole 0 C

Material Bonding Character Example

Metal Metallic Iron and ferrous alloys
Ceramic and Glasses Ionic/covalent Silica (SiO2): crystalline and amorphous)
Polymers Covalent and secondary HDPE, LDPE, PP, PS, etc..
Semiconductors Covalent or covalent/ionic Si, cadmium sulfide (CdS)

24