![Na (metal)
unstable
Cl (nonmetal)
unstable
electron
+ -
Coulombic
Attraction
Na (cation)
stable
Cl (anion)
stable
• Occurs between + and - ions. Non-directional
• Requires/powered by electron transfer.
• Larger difference in electronegativity required.
• Example: NaCl
IONIC BONDING
• High Melting Temperature Tm , hard, brittle, insulator
• Predominant bonding in Ceramics
Give up electrons Acquire electrons
He
-
Ne
-
Ar
-
Kr
-
Xe
-
Rn
-
F
4.0
Cl
3.0
Br
2.8
I
2.5
At
2.2
Li
1.0
Na
0.9
K
0.8
Rb
0.8
Cs
0.7
Fr
0.7
H
2.1
Be
1.5
Mg
1.2
Ca
1.0
Sr
1.0
Ba
0.9
Ra
0.9
Ti
1.5
Cr
1.6
Fe
1.8
Ni
1.8
Zn
1.8
As
2.0
CsCl
MgO
CaF2
NaCl
O
3.5
Adapted from Fig. 2.7, Callister 6e.
EXAMPLES: IONIC BONDING
• Requires shared electrons
• Example: Methane CH4
C: has 4 valence e,
needs 4 more
(to complete outer shell)
H: has 1 valence e,
needs 1 more
Bonding is DIRECTIONAL
• Electronegativities
are comparable.
shared electron
from carbon ato
shared electron
from hydrogen
atoms
H
H
H
H
C
CH4
COVALENT BONDING
• Most bonds are partially ionic partially covalent
%ionic bonding={1-exp[-(.25)(XA-XB)2]}x100
X-electronegativity of A or B elements. See Table 2.7 p-24
• Molecules of nonmetals
• Molecules with metals and nonmetals
• Elemental solid molecules (RHS of Periodic Table)
• Compound solids (about column IVA)
He
-
Ne
-
Ar
-
Kr
-
Xe
-
Rn
-
F
4.0
Cl
3.0
Br
2.8
I
2.5
At
2.2
Li
1.0
Na
0.9
K
0.8
Rb
0.8
Cs
0.7
Fr
0.7
H
2.1
Be
1.5
Mg
1.2
Ca
1.0
Sr
1.0
Ba
0.9
Ra
0.9
Ti
1.5
Cr
1.6
Fe
1.8
Ni
1.8
Zn
1.8
As
2.0
SiC
C(diamond)
H2O
C
2.5
H2
Cl2
F2
Si
1.8
Ga
1.6
GaAs
Ge
1.8
O
2.0
columnIVA
Sn
1.8
Pb
1.8
EXAMPLES: COVALENT BONDING
Usual
examples
Bonds diagonally farther = higher ionic bond, closer = higher covalent bond](https://image.slidesharecdn.com/ch02-atomicstructbonding-fall2016-sent-161102134154/85/Ch02-atomic-structbonding-fall2016-sent-5-320.jpg)
![14
Type
Ionic
Covalent
Metallic
Secondary
Bond Energy
Large!
Variable
large-Diamond
small-Bismuth
Variable
large-Tungsten
small-Mercury
smallest
Comments
Nondirectional (ceramics)
Directional
(semiconductors, ceramics
polymer chains)
Nondirectional (metals)
Directional
inter-chain (polymer)
inter-molecular
SUMMARY- BONDING
15
• Defn: Bond length, r
• Defn: Bond energy, Eo
F
F
r
• Melting Temperature, Tm
Eo=
“bond energy”
Energy (r)
ro
r
unstretched length
r
larger Tm
smaller Tm
Energy (r)
ro
Tm is larger if Eo is larger.
PROPERTIES FROM BONDING: TM
Melting Temperature of some Compounds
IonicDistance [Å] MeltingTemperature [°C]
NaF 2.31 988
NaCl 2.81 801 Na Halides
NaBr 2.98 755
NaI 3.23 651
MgO 2.1 2800
CaO 2.4 2580 Oxides
SrO 2.57 2430
BaO 2.76 1923
LiF 2.01 824
NaF 2.311 988 Florides
KF 2.67 846
RbF 2.82 775
16
• Elastic modulus, E
• E ~ curvature at ro
cross
sectional
area Ao
L
length, Lo
F
undeformed
deformed
LF
Ao
= E
Lo
Elastic modulus
r
larger Elastic Modulus
smaller Elastic Modulus
Energy
ro
unstretched length
E is larger if
Eo is larger.
PROPERTIES FROM BONDING: E](https://image.slidesharecdn.com/ch02-atomicstructbonding-fall2016-sent-161102134154/85/Ch02-atomic-structbonding-fall2016-sent-7-320.jpg)
Ch02 atomic structbonding-fall2016-sent
- 1. Main Topics this chapter... • What promotes bonding? • What types of bonds are there? What properties are inferred from bonding? ME3101-Materials Science and Engineering Ch2: BONDING AND PROPERTIES Nucleus: Z = # protons- responsible for primary classification of the material orbital electrons: n = principal quantum number n=3 2 1 = 1 for hydrogen to 94 for plutonium N = # neutrons Atomic mass A ≈ mass of nucleus = Z + N Adapted from Fig. 2.1, Callister 6e. BOHR ATOMIC MODEL Negatively charged electrons spin around positively charged nucleus in specified orbits. WAVE MECHANICAL MODEL • Every Electron- characterized by 4 quantum numbers (ZIP CODES !) • They affect shape, size and orientation of electron’s probability density in the orbit- where and when electrons may be found/exist • Bohr levels – divide into number of electron shells- depends on quantum number, which are: 1. Principal quantum number-1, 2, 3,.. Main energy states. ~ Bohr shell # - designated K L M N 2. Subsidiary (Angular Momentum) Subshells within the main energy states Standard designations… s, p, d, f (sharp, principal, diffuse, and fundamental. Original appearance of spectral lines) 3. Magnetic quantum number- number of energy states within subshells (energy levels) in s:1 in p:3 in d:5 in f:7 4. Spin moment (similar to Clockwise/counterclockwise) designated as + ½ - ½ Possible in every energy state ELECTRONS IN SHELLS/SUBSHELLS summary
- 2. • have discrete energy states • tend to occupy lowest available energy state. • notice overlap: 3d higher than 4s ! Increasingenergy n=1 n=2 n=3 n=4 1s 2s 3s 2p 3p 4s 4p 3d Electrons... (in practice) ELECTRON (relative) ENERGY STATES SeeFiguresintext • have complete s and p subshells • tend to be inert/non-reactive. • known as noble gas/elements. Elements with stable electron configurations... Z Element Configuration 2 He 1s2 10 Ne 1s22s22p6 18 Ar 1s22s22p63s23p6 36 Kr 1s22s22p63s23p63d104s24p6 Adapted from Table 2.2, Callister STABLE ELECTRON CONFIGURATIONS Other elements will have different outer sub-shell configurations… Review 1. How many quantum numbers? 1 2 3 4 6 8 12 2. What are they called? Principal, sub…, mag.., sp 3. What are the various levels of materials structure? Macroscopic, mi… ? 4. What is the common characteristics of the atomic structures of noble elements? sp sub-shells are …? A SIMPLE RULE FOR DETERMINING ELECTRON CONFIGURATION 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Example: 1s2 2s2 2p6 3s2 3p6 4s2 3….. Special circumstances: spn hybrid orbital, n= 1,2, or 3, that causes an exception to this rule. P-22 see foot note table 2.2 which elements more prone to hybridization? Reason: lower energy states. Let’s do a class exercise: Look at the periodic table (inside back cover) and select a material, say Chromium (Symbol= ?, Z= ?, A= ) Now start the electron configuration process using the rule/chart above: 1s2 2s2 ……………………Now look at Table 2.2 for Cr, p-25
- 3. • Why? Valence (outer) shell usually not filled completely. • Most elements: Electron configuration not stable. Element Hydrogen Helium Lithium Beryllium Boron Carbon ... Neon Sodium Magnesium Aluminum ... Argon ... Krypton Atomic # 1 2 3 4 5 6 10 11 12 13 18 ... 36 Electron configuration 1s1 1s2 (stable) 1s22s1 1s22s2 1s22s22p1 1s22s22p2 ... 1s22s22p6 (stable) 1s22s22p63s1 1s22s22p63s2 1s22s22p63s23p1 ... 1s22s22p63s23p6 (stable) ... 1s22s22p63s23p63d104s246 (stable) SURVEY OF ELEMENTS (See p-25) So ? • Rows: Periods ---- Columns: Groups Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions. He Ne Ar Kr Xe Rn inertgases accept1e accept2e giveup1e giveup2e giveup3e FLi Be Metal Nonmetal Intermediate H Na Cl Br I At O SMg Ca Sr Ba Ra K Rb Cs Fr Sc Y Se Te Po THE PERIODIC TABLE Analyze Periodic Table Characteristics • Horizontal rows – Period – repeats structure/characteristics • Vertical Columns = Groups: similar valence/bonding electron, properties- changes gradually horizontally and vertically down • Group 0- inert gas- • VIIA (halogens), VIA – 1 , 2 electron short of complete shell • IA, IIA – Alkali/Alkaline Earth Metals- 1, 2 electron excess in orbit than complete shell – • IIIB through IIB in 3 long periods – Transition metals. d- subshells are partially filled, sometimes with 1,2 electrons in higher subshells • IIIA, IVA, VA – intermediate between metal and nonmetal because of their valence electron structure. ELECTRONEGATIVITY • Electronegativity values assigned: from 0.7 to 4.0 • Large values: more tendency to acquire electrons. • Electropositive Elements- readily give up electrons to become electropositive ions. E.g. metals • Electronegativity: readiness to accept electrons to form negatively charged ions. See Figure in Text
- 4. 7 Smaller electronegativity Larger electronegativity He - Ne - Ar - Kr - Xe - Rn - F 4.0 Cl 3.0 Br 2.8 I 2.5 At 2.2 Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Fr 0.7 H 2.1 Be 1.5 Mg 1.2 Ca 1.0 Sr 1.0 Ba 0.9 Ra 0.9 Ti 1.5 Cr 1.6 Fe 1.8 Ni 1.8 Zn 1.8 As 2.0 ELECTRONEGATIVITY (contd.) • How about noble elements! Elements are more likely to attract electrons if their outer shell is almost full, and closer to the nucleus. Which means diagonally from B-L to T-R. Atomic Bonding in Solids Bonding- determines many physical properties such as elasticity, strength, melting temp. etc. Interatomic separation vs repulsive, attractive and net (a) forces and (b) (b) energies. At distance r0, net force = ? Net Energy = ? Bonding Types Bonding Primary/chemical Secondary Note: Both Primary and secondary bonding can and do exist simultaneously Review • How many quantum numbers ? • Hybrid orbital spn - effect/reason ? • Electronegativity - ? Willingness to accept ?? • When bonded, Net Force = ? » Net Energy = ? • Halogens in Group ? • Alkali Metals in Group ?
- 5. Na (metal) unstable Cl (nonmetal) unstable electron + - Coulombic Attraction Na (cation) stable Cl (anion) stable • Occurs between + and - ions. Non-directional • Requires/powered by electron transfer. • Larger difference in electronegativity required. • Example: NaCl IONIC BONDING • High Melting Temperature Tm , hard, brittle, insulator • Predominant bonding in Ceramics Give up electrons Acquire electrons He - Ne - Ar - Kr - Xe - Rn - F 4.0 Cl 3.0 Br 2.8 I 2.5 At 2.2 Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Fr 0.7 H 2.1 Be 1.5 Mg 1.2 Ca 1.0 Sr 1.0 Ba 0.9 Ra 0.9 Ti 1.5 Cr 1.6 Fe 1.8 Ni 1.8 Zn 1.8 As 2.0 CsCl MgO CaF2 NaCl O 3.5 Adapted from Fig. 2.7, Callister 6e. EXAMPLES: IONIC BONDING • Requires shared electrons • Example: Methane CH4 C: has 4 valence e, needs 4 more (to complete outer shell) H: has 1 valence e, needs 1 more Bonding is DIRECTIONAL • Electronegativities are comparable. shared electron from carbon ato shared electron from hydrogen atoms H H H H C CH4 COVALENT BONDING • Most bonds are partially ionic partially covalent %ionic bonding={1-exp[-(.25)(XA-XB)2]}x100 X-electronegativity of A or B elements. See Table 2.7 p-24 • Molecules of nonmetals • Molecules with metals and nonmetals • Elemental solid molecules (RHS of Periodic Table) • Compound solids (about column IVA) He - Ne - Ar - Kr - Xe - Rn - F 4.0 Cl 3.0 Br 2.8 I 2.5 At 2.2 Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Fr 0.7 H 2.1 Be 1.5 Mg 1.2 Ca 1.0 Sr 1.0 Ba 0.9 Ra 0.9 Ti 1.5 Cr 1.6 Fe 1.8 Ni 1.8 Zn 1.8 As 2.0 SiC C(diamond) H2O C 2.5 H2 Cl2 F2 Si 1.8 Ga 1.6 GaAs Ge 1.8 O 2.0 columnIVA Sn 1.8 Pb 1.8 EXAMPLES: COVALENT BONDING Usual examples Bonds diagonally farther = higher ionic bond, closer = higher covalent bond
- 6. • Arises from a sea of donated valence electrons (1, 2, or 3 from each atom). • The electron fluid/cloud acts as ‘glue’ for the ‘ion cores’ • Primary bond for metals and their alloys • Atoms stack orderly- forms crystal structure. • Floating electrons = carriers of energy (heat, electricity, ) + + + + + + + + + METALLIC BONDING Electron cloud (glue) Nucleus (ion core) Arises from interaction between dipoles (2 charged poles) • Fluctuating dipoles – temporary • Weak bond HH HH H2 H2 secondary bonding ex: liquid H2asymmetric electron clouds + - + -secondary bonding SECONDARY BONDING Examples- liquids Bonding between noble element molecules SECONDARY BONDING (Contd.) • Permanent dipoles-molecule induced + - secondary bonding + - H Cl H Clsecondary bonding secondary bonding -general case: -example: liquid HCl -example: between polymer chains Polar Molecules SECONDARY BONDING (Contd.) • Hydrogen-bond -A special case of polar molecule bonding -Strongest of secondary bonding -Occurs between molecules where H is covalently bonded to F (e.g. HF) ,O (e.g. H2O), or N (e.g. NH3). -Directional ? Water O HH
- 7. 14 Type Ionic Covalent Metallic Secondary Bond Energy Large! Variable large-Diamond small-Bismuth Variable large-Tungsten small-Mercury smallest Comments Nondirectional (ceramics) Directional (semiconductors, ceramics polymer chains) Nondirectional (metals) Directional inter-chain (polymer) inter-molecular SUMMARY- BONDING 15 • Defn: Bond length, r • Defn: Bond energy, Eo F F r • Melting Temperature, Tm Eo= “bond energy” Energy (r) ro r unstretched length r larger Tm smaller Tm Energy (r) ro Tm is larger if Eo is larger. PROPERTIES FROM BONDING: TM Melting Temperature of some Compounds IonicDistance [Å] MeltingTemperature [°C] NaF 2.31 988 NaCl 2.81 801 Na Halides NaBr 2.98 755 NaI 3.23 651 MgO 2.1 2800 CaO 2.4 2580 Oxides SrO 2.57 2430 BaO 2.76 1923 LiF 2.01 824 NaF 2.311 988 Florides KF 2.67 846 RbF 2.82 775 16 • Elastic modulus, E • E ~ curvature at ro cross sectional area Ao L length, Lo F undeformed deformed LF Ao = E Lo Elastic modulus r larger Elastic Modulus smaller Elastic Modulus Energy ro unstretched length E is larger if Eo is larger. PROPERTIES FROM BONDING: E
- 8. 16 • E ~ slope (blue line) of Net Force curve (red) at zero F More vertical slope means stiffer (high E) material More horizontal slope means more flexible material. PROPERTIES FROM BONDING: E(contd.) Another way is to examine the bonding force curve. 17 • Coefficient of thermal expansion, • ~ symmetry at ro is larger if Eo is smaller. L length, Lo unheated, T1 heated, T2 = (T2-T1) L Lo coeff. thermal expansion r smaller larger Energy ro PROPERTIES FROM BONDING: 18 Ceramics (Ionic & covalent bonding): Metals (Metallic bonding): Polymers (Covalent & Secondary): secondary bonding Large bond energy large Tm large E small Variable bond energy moderate Tm moderate E moderate Directional Properties Secondary bonding dominates SUMMARY: Properties Tm E Guess these 3 properties small small large