1. Chapter 2: Atomic Structure and Interatomic Bonding
These notes have been prepared by Jorge Seminario from the textbook material

2. ISSUES TO ADDRESS... • What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding?

3. Basic idea … properties of materials are a consequence of
Identity of the atoms
Spatial arrangement of the atoms
Interactions between the atoms
Thus, we need to study atomic structure/bonding!
End of lecture 1

4. Atomic Structure Changes as you move around the periodic table!
Atomic number
Number of electrons/electronic configuration
e.g. you might expect electronic properties to depend on the number of electrons an atom has (i.e. open v closed shells)
Type of bonding (ionic versus covalent)
Changes as you move around the periodic table!
End of lecture 1

5. 2.2 Fundamental Concepts Atoms consist of a small nucleus containing
Protons
+1.60 x C = e
1.67 x kg
Neutrons
0 C (neutral)
Electrons (which circle the nucleus)
-1.60 x C = -e
9.11 x kg

6. 2.2 Fundamental Concepts Atomic Number (Z) Atomic Mass (A) Isotopes
Number of protons in the nucleus
Electrically neutral or complete atom: Z = # electrons
Atomic Mass (A)
Sum of the masses of protons and neutrons; atomic mass unit = amu = 1/12 mass of 12C
Isotopes
Atoms of the same element with different atomic masses due to varying number of neutrons (e.g. 12C, 13C, 14C

7. Atomic Structure 2.2 Fundamental Concepts
Atomic weight is the weighted average of the element based on the relative amounts of its isotopes (e.g. 1 mol/carbon = g/mol, NOT 12 g/mol!)   Atomic wt = wt of NA molecules, atoms, etc. and
NA = x 1023   1 amu/atom = 1g/mol
C g/mol H g/mol, etc.
Mole
Quantity of a substance corresponding to 6.022X1023 atoms, molecules, …
1 amu/ atom (or molecule) = 1g/mol

8. Basic concepts Atoms are made of protons, neutrons and electrons
me = x = x10-31 kg = MeV
mp = x kg = MeV
mn = x kg = MeV = mp MeV
Charge of a proton and electron are the same: x10-19 C
However p are +’ve and e are –’ve
Since J = C x V (1 joule = 1 coulomb x 1 volt), 1 eV = x10-19 J
mass is related to energy by E = mc2
End of lecture 1

9. Examples How many grams are there in one amu of a material?
The two major isotopes of carbon:
98.93% of 12C with an atomic weight of amu, and
1.07% of 13C with an atomic weight of amu.
Confirm that the average atomic weight of C is amu.
Sum the product of the isotope atomic weight and the percent abundance.
(12 amu)*(.9893)+( amu)*(.0107) = amu

10. 2.3 Electrons In Atoms Bohr Atomic Model
Early outgrowth of quantum mechanics
Electrons revolve around nucleus in discrete orbitals
Electrons closer to nucleus travel faster then outer orbitals
Principal quantum number (n); 1st shell, n=1; 2nd shell, n=2; 3rd shell, n=3

11. c02f02 Quantum Numbers For the H atom Scaled for hydrogen-like atoms
Degenerate states
Same energy
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12. c02f03
Bohr Atom
Wave-mechanical atom
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13. Atomic Models Wave-Mechanical Model
Electron exhibits both wave-like and particle-like characteristics
Position is now considered to be the probability of an electron being at various locations around the nucleus, forming an electron cloud

14. Electron Configuration
Pauli Exclusion Principle
Stipulates that electron states (orbital) can have no more than two electrons, must have opposite spins
Ground state
All electrons occupy the lowest energies
Electrons can move to higher states
Filled shells are more stable

15. Electronic Structure
Electrons have wave-like and particle-like properties (old view)
We can better say that the wave-particle nature is the real thing; individual wave and particle states are limiting cases, observed in measurements (collapse of the wave function)
To better understand electronic structure, we assume
Electrons “reside” in orbitals.
Each orbital, at a discrete energy level, is determined by quantum numbers.
c Quantum numbers Designation
n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.)
l = angular (orbitals) s, p, d, f (0, 1, 2, 3,…, n -1)
ml = magnetic 1, 3, 5, 7 (-l to +l)
ms = spin ½, -½

16. Quantum Numbers
c02f04
Relative electrons energies (E) for shells and subshells
if n↓ then E↓
Within each shell E↑ with quantum number
Overlapping in energy of a state in one shell with states in adjacent shells, true of d and f states
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17. c02tf01
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18. Electron Configurations
Valence electrons – those in unfilled shells
Filled shells more stable
Valence electrons are most available for bonding and tend to control the chemical properties
example: C (atomic number = 6)
1s2 2s2 2p2
valence electrons

19. Atomic Models Quantum numbers
Principal quantum number n, represents a shell
K, L, M, N, O correspond to n=1, 2, 3, 4, 5....
Quantum number l, signifies the subshell
Lowercase italics letter s, p, d, f; related to the shape of the subshell
Quantum number ml , represents the number of energy states
s, p, d, f have 1, 3, 5, 7 states respectively
Quantum number ms, is the spin moment
Each electron is a spin moment
(+1/2) and (-1/2)

20. Electron Configuration Silicon (Si)
Electron configuration represents the manner in which the states are occupied
Valence electrons
Occupy the outermost shell
Available for bonding
Tend to control chemical properties

21. Electron Configurations - Pauli Exclusion Principle
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Na Atom
Z = 11
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22. c02tf02
When some elements covalently bond, they form sp hybrid bonds, e.g., C, Si, Ge
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23. Examples Give the electron configurations for the following: C
1s2 2s2 2p2 Br
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
Mn+2
1s2 2s2 2p6 3s2 3p6 3d5 F-
1s2 2s2 2p6 Cr
1s2 2s2 2p6 3s2 3p6 4s1 3d5

24. Electronic Configurations
ex: Fe - atomic # = 26
1s2 2s2 2p6 3s2 3p6
3d 6 4s2
valence
electrons 1s 2s 2p
K-shell n = 1
L-shell n = 2 3s 3p
M-shell n = 3 3d 4s 4p 4d
Energy
N-shell n = 4
Adapted from Fig. 2.4, Callister & Rethwisch 3e.
Notice that 2s and 2p do not have the same energy

25. Electron Configuration
“Stable electron configurations”
States within the outermost or valence electron shell are completely filled
Some atoms of elements with unfilled shells assume stable electron configurations by gaining or losing electrons to form charged ions
Sometimes s and p orbitals form hybrid spn orbitals
3A, 4A, and 5A group elements typically
Lower energy state for the valence electrons

26. SURVEY OF ELEMENTS • Most elements: Electron configuration not stable.
... 1s 2 2s 2p 6 3s 3p 3d 10 4s 4p
Atomic # 18 36
Element 1
Hydrogen
Helium 3
Lithium 4
Beryllium 5
Boron
Carbon
Neon 11
Sodium 12
Magnesium 13
Aluminum
Argon
Krypton
Adapted from Table 2.2, Callister & Rethwisch 3e.

27. STABLE ELECTRON CONFIGURATIONS
• have complete s and p subshells
• tend to be unreactive.
Adapted from Table 2.2, Callister 6e. 4

28. 2.4 Periodic Table
Elements classified according to electron configuration
Elements in a given column or group have similar valence electron structures as well as chemical and physical properties
Group 0 – inert gases, filled shells and stable
Group VIIA – halogen
Group IA and IIA - alkali and alkaline earth metals
Groups IIIB---IIB – transition metals
Groups IIIA, IVA and VA – characteristics between the metals and nonmetals

29. 2.4

31. c02f07 Electropositive: Electronegativity Values Electronegative:
Capable of giving up their valence electrons to become positively charged
Electronegative:
Readily accept electrons to form negatively charged ions
Sometimes share electrons with other atoms

32. Atomic Bonding
Valence electrons determine all of the following properties
Chemical
Electrical
Thermal
Optical
Deteriorative
etc.

33. Atomic Bonding in Solids

34. 2.5 Bonding Forces and Energies
When 0 = FA + FR, equilibrium exists. The centers of the atoms will remain separated by the equilibrium spacing ro.
This spacing also corresponds to the minimum of the potential energy curve. The energy that would be required to separate two atoms to an infinite separation is Eo
FN = FA + FR
EN = EA + ER
Figure 2.8

35. 2.5 Bonding Forces and Energies
A number of material properties depend on Eo, the curve shape, and bonding type
Material with large Eo typically have higher melting points
Mechanical stiffness is dependent on the shape of its force vs. interatomic separation curve
A material’s linear coefficient of thermal expansion is related to the shape of its Eo vs. ro curve

36. Bonding in Solids 2.5 Bonding forces and energies
Far apart: atoms don’t know about each other
As they approach one another, start to exert force on one another
two types of forces
Attractive (FA) – slowly changing with distance
Repulsive (FR) – typically short-range
Net force is the sum of these
FN = FA + FR
At some point the net force is zero; at that position a state of equilibrium exists
End of lecture 1

37. Bonding in Solids
Bonding forces and energies
& setting our ZERO ENERGY reference at ∞
The point where the forces are zero also corresponds to the minimum potential energy for the two atoms, which makes sense because -dE/dr = F = 0 at a minimum.
End of lecture 1
The interatomic separation at that point (ro) corresponds to the potential energy at that minimum
Eo, it is also the bonding energy
Eo is the energy needed to separate the atoms

38. Examples (assume the book wrong sign of the F)
Calculate the force of attraction between ions X+ and an Y-, the centers of which are separated by a distance of 2.01 nm.
&

39. 2.6 Primary Interatomic Bonds
Types of chemical bonds found in solids
Ionic
Covalent
Metallic
As you might imagine, the type of bonding influences properties – why?
End of lecture 1
Bonding involves the valence electrons!!!

40. 2.6 Primary Interatomic Bonds
Ionic Bonding
Compounds composed of metallic and nonmetallic elements
Coulombic Attractive Forces: positive and negative ions, by virtue of their net electrical charge, attract one another
EA = -A/r
ER = B/rn
Bonding is nondirectional: the magnitude of the bond is equal in all directions around an ion
Properties: generally large bonding energies ( kJ/mol) and thus high melting temperatures, hard, brittle, and electrically and thermally insulative
A, B, n are constants
Coulombic bonding Force
Cl-
Na+

41. 2.6 Primary Interatomic Bonds
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42. 2.6 Primary Interatomic Bonds
Ionic bonding
Prototype example – sodium chloride (NaCl)
Sodium gives up one its electrons to chlorine – sodium becomes positively charged, chlorine becomes negatively charged
The attraction energy is electrostatic in nature in ionic solids (opposite charges attract)
The attractive component of the potential energy (for 2 point charges) is given by
End of lecture 1
The repulsive term is given by

43. Ionic bond: metal + nonmetal
donates accepts electrons electrons  
Dissimilar electronegativities  
ex: MgO Mg 1s2 2s2 2p6 3s O 1s2 2s2 2p4
[Ne] 3s2 
Mg2+ 1s2 2s2 2p O2- 1s2 2s2 2p6
[Ne] [Ne]

44. Ionic Bonding - + • Occurs between + and - ions.
• Requires electron transfer.
• Large difference in electronegativity required.
• Example: NaCl
Na (metal)
unstable
Cl (nonmetal)
electron + -
Coulombic
Attraction
Na (cation)
stable
Cl (anion)

45. Ionic Bonding Energy – minimum energy most stable A B EN = EA + ER = r
Energy balance of attractive and repulsive terms A B
EN = EA + ER = - + r r n
Attractive energy EA
Net energy EN
Repulsive energy ER
Interatomic separation r
Adapted from Fig. 2.8(b), Callister & Rethwisch 3e.

46. Examples: Ionic Bonding
• Predominant bonding in Ceramics
NaCl
MgO
Give up electrons
Acquire electrons
CaF 2
CsCl
Adapted from Fig. 2.7, Callister & Rethwisch 3e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

47. 2.6 Primary Interatomic Bonds
Covalent Bonding
Stable electron configurations are assumed by the sharing of electrons between adjacent atoms
Bonding is directional: between specific atoms and may exist only in the direction between one atom and another that participates in electron sharing
Number of covalent bonds for a particular molecule is determined by the number of valence electrons
Bond strength ranges from strong to weak
Rarely are compounds purely ionic or covalent but are a percentage of both.
Sharing 4 electrons
Sharing 2 electrons
%ionic character = {1 – exp[-(0.25)(XA-XB)2]} x 100
XA and XB are electronegatives

48. Covalent bonding Sharing of electrons between adjacent atoms
Most nonmetallic elements and molecules containing dissimilar elements have covalent bonds
Polymers!
Bonding is highly directional!
Number of covalent bonds possible is guessed by the number of valence electrons
Typically is 8 – N, where N is the number of valence electrons
Carbon has 4 valence e’s – 4 bonds (ok!)
End of lecture 1

49. EXAMPLES: COVALENT BONDING
Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is
adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.
• Molecules with nonmetals
• Molecules with metals and nonmetals
• Elemental solids (RHS of Periodic Table)
• Compound solids (about column IVA) 11

50. Covalent Bonding similar electronegativity  share electrons
bonds determined by valence – s & p orbitals dominate bonding
Example: CH4
shared electrons
from carbon atom
from hydrogen
atoms H C CH 4
C: has 4 valence e-,
needs 4 more
H: has 1 valence e-,
needs 1 more
Electronegativities
are comparable.
Adapted from Fig. 2.10, Callister & Rethwisch 3e.

51. Bonding in Solids
Many materials have bonding that is both ionic and covalent in nature (very few materials actually exhibit pure ionic or covalent bonding)
Easy (empirical) way to estimate % of ionic bonding character:
End of lecture 1
XA, XB are the electronegativities of atoms A and B involved
Notice: this is a very very very empirical formula

52. Primary Bonding Ionic-Covalent Mixed Bonding % ionic character =
where XA & XB are Pauling electronegativities %)
100 ( x
Ex: MgO XMg = XO = 3.5

53. Example
Compute the percentage ionic character of the interatomic bond for zinc oxide (ZnO). Refer to the periodic Table for electronegativity values. Note: Electronegativity values in slides differ slightly from those in book.
% ionic character = {1-exp[-(0.25)*(XZn-XO)2]}*100
= {1-exp[-(0.25)*( )2]}*100
= 55.51%

54. 2.6 Primary Interatomic Bonds
Metallic Bonding
Found in metals and their alloys
1 to 3 valence electrons that form a “sea of electrons” or an “electron cloud” because they are more or less free to drift through the entire metal
Nonvalence electrons and atomic nuclei form ion cores
Bonding energies range from weak to strong
Good conductor of both electricity and heat
Most metals and their alloys fail in a ductile manner
Ion Cores + + + - - + + + - - + + +
Sea of Valence Electrons

55. METALLIC BONDING • Arises from a sea of donated valence electrons
(1, 2, or 3 from each atom).
Adapted from Fig. 2.11, Callister 6e.
• Primary bond for metals and their alloys 12

56. Bonding in Solids Metallic bonding
Most metals have one, two, or at most three valence electrons
These electrons are highly delocalized from a specific atom – have a “sea of valence electrons”
End of lecture 1
Free electrons shield positive core of ions from one another (reduce ER)
Metallic bonding is also non-directional
Free electrons also act to hold structure together
Wide range of bonding energies, typically good conductors (why?)

57. 2.7 Secondary Bonding or van der Walls Bonding
Also known as physical bonds
Weak in comparison to primary or chemical bonds
Exist between virtually all atoms and molecules
Arise from atomic or molecular dipoles
bonding that results from the coulombic attraction between the positive end of one dipole and the negative region of an adjacent one
a dipole may be created or induced in an atom or molecule that is normally electrically symmetric

58. 2.7 Secondary Bonding or van der Waals Bonding
Fluctuating Induced Dipole Bonds
A dipole (whether induced or instantaneous) produces a displacement of the electron distribution of an adjacent molecule or atom and continues as a chain effect
Liquefaction and solidification of inert gases
Weakest Bonds
Extremely low boiling and melting point
Atomic nucleus
Atomic nucleus
Instantaneous
Fluctuation
Electron
cloud
Electron
cloud

59. 2.7 Secondary Bonding or van der Waals Bonding
Polar Molecule-Induced Dipole Bonds
Permanent dipole moments exist by virtue of an asymmetrical arrangement of positively and negatively charged regions
Polar molecules can induce dipoles in adjacent nonpolar molecules
Magnitude of bond greater than for fluctuating induced dipoles
Atomic nucleus -
Electron Cloud +
Polar Molecule
Induced Dipole

60. 2.7 Secondary Bonding or van der Waals Bonding
Permanent Dipole Bonds
Stronger than any secondary bonding with induced dipoles
A special case of this is hydrogen bonding: exists between molecules that have hydrogen as one of the constituents
Hydrogen Bond H Cl H Cl

61. van der Waals interactions between polar molecules
Permanent dipoles
Hydrogen-bonds
These interactions are fairly strong, very complex, and surprisingly not well understood!
2.82 Å
109.47°
van der Waals interactions between polar molecules
Best known example hydrogen bonding

62. c02tf03
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63. MATERIAL OF IMPORTANCE
Water
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Many molecules do not have a symmetric distribution/arrangement of positive and negative charges (e.g. H2O, HCl)
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64. c02uf01
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65. Properties From Bonding: Tm
• Bond length, r
• Melting Temperature, Tm r o
Energy r
• Bond energy, Eo Eo =
“bond energy”
Energy r o
unstretched length
smaller Tm
larger Tm
Tm is larger if Eo is larger.

66. Properties From Bonding : a
• Coefficient of thermal expansion, a
coeff. thermal expansion D L
length, o
unheated, T 1
heated, T 2 D L = a ( T - T ) 2 1 L o
• a ~ symmetric at ro r o
smaller a
larger a
Energy
unstretched length Eo
a is larger if Eo is smaller.

67. PROPERTIES FROM BONDING: E
• Elastic modulus, E
E ~ dF/dr|ro elastic modulus 16

68. Summary: Primary Bonds
Ceramics
Large bond energy
large Tm
large E
small a
(Ionic & covalent bonding):
Metals
(Metallic bonding):
Variable bond energy
moderate Tm
moderate E
moderate a
Polymers
(Covalent & Secondary):
Directional Properties
Secondary bonding dominates
small Tm
small E
large a
secondary bonding

69. Summary: Bonding Type Bond Energy Comments Ionic Large!
Nondirectional (ceramics)
Covalent
Variable
Directional
(semiconductors, ceramics
polymer chains)
large-Diamond
small-Bismuth
Metallic
Variable
large-Tungsten
Nondirectional (metals)
small-Mercury
Secondary
smallest
Directional
inter-chain (polymer)
inter-molecular

70. JIC, additional vg’s

71. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2

72. Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
8.2

73. Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
8.2

74. Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.
8.3

75. 8.3

76. I1 first ionization energy
Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.
I1 + X (g) X+(g) + e-
I1 first ionization energy
I2 + X (g) X2+(g) + e-
I2 second ionization energy
I3 + X (g) X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3
8.4

77. 8.4 Filled n=1 shell Filled n=2 shell Filled n=3 shell

78. Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
X (g) + e X-(g)
F (g) + e X-(g)
DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e O-(g)
DH = -141 kJ/mol
EA = +141 kJ/mol
8.5

79. Group 2A Elements (ns2, n  2)
M M+2 + 2e-
Be(s) + 2H2O(l) No Reaction
Mg(s) + 2H2O(g) Mg(OH)2(aq) + H2(g)
M(s) + 2H2O(l) M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
Increasing reactivity
8.6

80. Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g) Al2O3(s)
2Al(s) + 6H+(aq) Al3+(aq) + 3H2(g)
8.6

81. Group 7A Elements (ns1np5, n  2)
X + 1e X-1
X2(g) + H2(g) HX(g)
Increasing reactivity
8.6