Chapter 2 -
Chapter 2: Atomic Structure and Interatomic Bonding
• Atomic Structure • Electron Configuration • Periodic Table • Primary Bonding
– Ionic – Covalent – Metallic
• Secondary Bonding or van der Waals Bonding – Three types of Dipole Bonding
• Molecules
Chapter 2 -
Two Allotropes of CARBON
Graphite Diamond
- Relatively soft - Greasy feel to it - Reasonably good
conductor of electricity
- The hardest known material - Poor conductor of electricity
The disparities in properties are attributed to a type of interatomic bonding found in graphite that does not exist in
diamond
Why Study Atomic Structure and Interatomic Bonding?
Chapter 2 - 3
Atomic Structure (Freshman Chem.) • atom – electrons – 9.11 x 10-31 kg
protons neutrons
• atomic number = # of protons in nucleus of atom = # of electrons of neutral species
• A [=] atomic mass unit = amu = 1/12 mass of 12C Atomic wt = wt of 6.023 x 1023 molecules or atoms
1 amu/atom = 1g/mol
C 12.011 H 1.008 etc.
} 1.67 x 10-27 kg
Chapter 2 - 5
460 BC
Democritus’ atomos
Aristotle and Plato
360 BC
J.J. Thomson
1897
E. Rutherford
1909
N. Bohr
1913
Wave Mechanical
1923
Timeline of Atomic Theory
Dalton
1808
Chapter 2 -
~ 400 BC - Democritus
• Ancient Greek philosopher • Democritus coined the term átomos which
means "uncuttable" or "the smallest indivisible particle of matter".
Structure of Matter Physical world “VOID + BEING”
Chapter 2 -
• Democritus postulated that atoms
were completely solid, hard and small particles with no internal structure and has an infinite variety of shapes and sizes.
• This theory was ignored for more than 2000 years!
Democritus
7
Chapter 2 -
• There can be no ultimately
indivisible particles. • Believed that fire, earth, air and
water were the four main elements that world was made up of.
The Atomic Theory of Matter
Aristotle and Plato
8
Chapter 2 -
1803 – John Dalton
• English instructor and natural philosopher
• “Each element consists of atoms of single unique type and can join to form chemical compounds.”
• Beginning the modern atomic theory
Chapter 2 -
1. Elements are composed of
extremely small particles called atoms.
The Atomic Theory of Matter Dalton’s Atomic Theory Postulates
10
2. All atoms of a given element are
identical, having the same size, mass and chemical properties. The atoms of different elements are different. Different elements have different atomic properties such as atomic mass.
Chapter 2 -
The Atomic Theory of Matter
11
4. Compounds are formed from the combination of atoms of more than one element. A compound always has the same relative number and kind of atoms.
Dalton’s Atomic Theory
3. Atoms of an element are neither created nor destroyed by any chemical reactions. Chemical reactions only involve the combination, rearrangement or separation of atoms.
Assist. Prof. Dr. İlkay KALAY
Chapter 2 -
1869 - Mendeleev
• Building upon earlier discoveries by scientists, Mendeleev published the first functional PERIODIC TABLE.
• Certain chemical properties of elements repeat periodically when arranged by atomic number.
• The periodic table was first developed by Mendeleev and Meyer on the bases of similarity in chemical and physical properties exhibited on certain elements.
Chapter 2 -
1897 – Sir J. J. Thomson
• Although Dalton had postulated that atoms were indivisible, the studies have shown a more complex structure for an atom.
• J. J. Thomson conducted a series of
experiments which showed that the atoms were not indivisible
• Discovered the electron (1906 Nobel Prize
in Physics).
Chapter 2 - 17
J. J. Thomson proposed the atom consisted of a uniform positive sphere of matter in which the electrons were embedded.
J. J. Thomson's "plum-pudding" model of the atom
Atom is composed of even smaller particles but how the particles fit together?
Plum Pudding (1904): “The atom as being made up of electrons swarming in a sea of positive charge.
Chapter 2 -
• Results: – Majority of a particles transmitted (pass through) or
deflected through small angles – Tiny fraction deflected through large angles
1909 – E. Rutherford • Tested and disproved the Plum
Pudding Model.
Rutherford's experiment on the scattering of α particles by metal foil
Chapter 2 -
• Conclusion: – Disproved the Plum-Pudding Model – Large amount of the atom's charge and mass is
concentrated into a small region – Atom was mostly empty space
• Objections to Rutherford model – The laws of classical mechanics predict that the
electron will release electromagnetic radiation while orbiting a nucleus. Because the electron would lose energy, it would gradually spiral inwards, collapsing into the nucleus.
– This atom model is unsuccessful, because it predicts that all atoms are unstable.
1909 – E. Rutherford
Chapter 2 - 20
The Modern View of Atomic Structure
The structure of an atom
Electron
Proton Neutron
Nucleus Particle Charge Mass (amu)
Electron Negative (-1) 5.486 x 10-4
Proton Positive (+1) 1.0073 Neutron Neutral (0) 1.0087
Particle Charge (C) Electron -1.6022 x 10-19
Proton 1.6022 x 10-19
Neutron 0
Chapter 2 - 21
The Modern View of Atomic Structure The structure of an atom
The atoms are small and the atomic dimensions are expressed in terms of Angstrom (Å) unit.
1 Å = 10-10 m
Schematic view of an atom
~ 10-4 Å 1-5 Å
Nucleus
Assist. Prof. Dr. İlkay KALAY
Chapter 2 - Atomic number = Z
Mass number = A
22
The Modern View of Atomic Structure The structure of an atom
ü All atoms of an element have the same number of protons in the nucleus.
ü The number of protons in the nucleus of an atom is called atomic number (Z).
ü Because the atoms have no net electrical charge, # of protons = # of electrons.
ü The total number of protons and neutrons in a nucleus is called mass number (A).
612C
Assist. Prof. Dr. İlkay KALAY
Chapter 2 - 23
The Modern View of Atomic Structure ü All nuclei of the atoms of a particular element have the same
atomic number, but the the number of neutrons and so mass numbers may be different. Atoms of a given element with a different mass numbers are called isotopes.
ü Most elements occur in nature as mixtures of isotopes. Symbol # of
electrons # of protons # of
neutrons 6 6 5 6 6 6 6 6 7 6 6 8
613C612C
614C
611C
An atom of a specific isotope is called a nuclide.
Assist. Prof. Dr. İlkay KALAY
Chapter 2 - 24
The Periodic Table The periodic table, developed in 1869, shows the arrangement of elements with similar chemical and physical properties in order of increasing atomic number. Periodic table presents
atomic number, atomic symbol and atomic
weight for each element.
Zr91.224
40 Atomic number
Atomic weight Atomic symbol
Chapter 2 -
1912 – N. Bohr
• Many phenomena involving electrons in solids could not be explained in terms of CLASSICAL MECHANICS.
• We need QUANTUM MECHANICS…
Chapter 2 - 26
Line Spectra and The Bohr Model Bohr’s Model
Bohr proposed a model of the hydrogen atom that explains its line spectrum. Bohr’s postulates • Rutherford atom is correct • Classical EM theory not applicable to orbiting e- • Newtonian mechanics applicable to orbiting e- • Eelectron = Ekinetic + Epotential • Only orbits of certain radii, corresponding to certain definite energies are
permitted for electrons in an atom. • An electron in a permitted orbit has a specific energy and is in allowed
energy state. An electron in an allowed energy state will not radiate energy and therefore will not spiral into the nucleus.
• During the transition of an electron from one allowed energy state to another, energy is only emitted or absorbed by an electron. This energy is emitted or absorbed as a photon, ΔE = Ef-Ei= hν = hc/λ where c = νλ
Chapter 2 - 27
Line Spectra and The Bohr Model
Energy levels in the hydrogen atom from the Bohr model. -‐ The arrows refer to the transi7ons of the
electron from one allowed energy state to another.
-‐ The states shown are those for which n = 1 through n = 6, and the state for n = ∞, for which the energy, E, equals zero.
Energy States of the Hydrogen Atom
Chapter 2 - 28
Bohr’s Model
The lower the value of n, the smaller the radius of the orbit, and the lower
the energy level.
n=1 ground state (lowest energy state) (orbit closest to the nucleus)
n=2, 3, or higher excited state
Energy levels in the hydrogen atom from the Bohr model.
The arrows refer to the transi7ons of the electron from one allowed energy state to another.
Chapter 2 -
Nucleus: Z = # protons
2
orbital electrons: n = principal quantum number
n=3 2 1
N = # neutrons
Atomic mass A ≈ Z + N
Adapted from Fig. 2.1, Callister 6e.
BOHR ATOM
Chapter 2 -
1913 - Sommerfeld
• German theoretical physicist • Modified the Bohr Model • “suppose we have plurality of orbits” – a shell
containing multiple orbits: ORBITALS • How to capture these new ideas quantitatively? • We need new quantum numbers: n, l, m, s
n principal quantum number, distance of an electron from the nucleus l subshell, describes the shape of the subshell m number of energy states in a subshell s spin moment
Chapter 2 -
Wave mechanics to arrive at same place: E=E(n,l,m,s)
• The Bohr model – significant limitations • Resolution: Wave-mechanical model
(electron is considered to exhibit both wave-like and particle-like characteristics).
– De Broglie: “If a photon which has no mass, can behave as a particle, does an electron which has mass can behave as a wave (1920)?” λ = h/p = h/mv
– Heisenberg: Uncertainty Principle “I don’t know where any of one of electrons is, but I
can tell you an average where any of one of them is likely to be”
– Schrodinger
Chapter 2 -
Beyond Bohr’s Model De Broglie wavelength:
! =hm"
Heisenberg Uncertainty Principle: It is impossible to know simultaneously the exact posi7on and momentum of a par7cle. !x.!m! " h
4"
De Broglie wavelength states that electrons have wavelike mo7on
The more accurately we know the posi7on of the par7cle (smaller Δx), the less accurately we know its speed (larger Δu) and vice versa.
Wave mo>on of objects on the atomic scale
Schrödinger’s wave equaCon incorporates both the wave-‐like and par.cle-‐like behavior of the electron.
Wave function, ψ, Ψ2 provides information about an electron’s location when it s in an allowed energy state
Chapter 2 -
Quantum Mechanics and Atomic Orbitals
Electron-‐density diagram in the ground state of the hydrogen atom
Electron density: probability of the electron being at a point Higher density of dots region: larger values of Ψ2
Chapter 2 -
Bohr Model vs. Wave Mechanical Atom Model
Bohr WM
With WM model electron no longer treated as a particle moving in a discrete orbital,
rather position is described by a probability distribution
Chapter 2 -
Quantum Numbers
n…. The principal quantum number (K, L, M, N, O and so on that correspond to 1, 2, 3,4, 5…… As n increases: - The orbital becomes larger - The electron spends more time farther from the nucleus - The electron has a higher energy and is therefore less tightly bound to the nucleus. l…. The second quantum number • l can have integral values from 0 to n-1 for each value of n • defines the shape of the orbital • The value of l of a particular orbital is designated by s, p, d and f, corresponding l
values of 0, 1, 2 and 3 respectively s: sharp p:principal d:diffuse f: fundamental ml…. The magnetic quantum number • Have integral values between l and –l including zero. • Describes the orientation of the orbital in space
According to quantum mechanics, each electron in an atom is described by four different quantum numbers, three of which (n, l, ml) specify the wave function that gives the probability of finding the electron at various points in space.
Chapter 2 -
What is the filling sequence of electrons in orbitals by n, l, m, s is not adequate?
AUFBAU PRINCIPLE 3 principles:
1. Pauli Exclusion Principle:only one electron can have a given set of four quantum numbers.
2. Electrons -have discrete energy states -fill orbitals from lowest en. to highest en.
3. Hund’s rule
Chapter 2 - 37
Electron Configura>ons
Hund’s rule states that the lowest energy is attained by maximizing the number of electrons with the same electron spin. For example, for a carbon atom to achieve its lowest energy, the two 2p electrons will have the same spin.
Chapter 2 -
Quantum Numbers Electrons fill quantum levels in order of increasing energy ( only n and l make significant differences in energy configurations). 1s, 2s, 2p, 3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,…. When all electrons are at the lowest possible energy levels => ground state Excited states do exist such as in glow discharges etc…
Valence electrons occupy the outermost filled shell. Valence electrons are responsible for all bonding !
Chapter 2 - 40
Electronic Structure • Electrons have wavelike and particulate
properties. – This means that electrons are in orbitals defined by a
probability. – Each orbital at discrete energy level determined by
quantum numbers. Quantum # Designation
n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.) l = subsidiary (orbitals) s, p, d, f (0, 1, 2, 3,…, n -1) ml = magnetic 1, 3, 5, 7 (-l to +l) ms = spin ½, -½
Chapter 2 - 41
Electron Energy States
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
• have discrete energy states • tend to occupy lowest available energy state.
Electrons...
Adapted from Fig. 2.4, Callister 7e.
Chapter 2 - 42
• Why? Valence (outer) shell usually not filled completely.
• Most elements: Electron configuration not stable. SURVEY OF ELEMENTS
Electron configuration
(stable)
...
... 1s 2 2s 2 2p 6 3s 2 3p 6 (stable) ... 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable)
Atomic #
18 ... 36
Element 1s 1 1 Hydrogen 1s 2 2 Helium 1s 2 2s 1 3 Lithium 1s 2 2s 2 4 Beryllium 1s 2 2s 2 2p 1 5 Boron 1s 2 2s 2 2p 2 6 Carbon
... 1s 2 2s 2 2p 6 (stable) 10 Neon 1s 2 2s 2 2p 6 3s 1 11 Sodium 1s 2 2s 2 2p 6 3s 2 12 Magnesium 1s 2 2s 2 2p 6 3s 2 3p 1 13 Aluminum
... Argon ... Krypton
Adapted from Table 2.2, Callister 7e.
Chapter 2 - 43
Electron Configura>ons
Assist. Prof. Dr. İlkay KALAY
• Elements in any given group in the periodic table have the same type of electron arrangements in their outermost shells.
• The outer shell electrons those that lie outside the orbitals occupied in the next lowest noble gas element are called its valence electrons, whereas the electrons in the inner shells are called the core electrons.
Chapter 2 - 44
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
Chapter 2 - 46
Atomic Structure
• Valence electrons determine all of the following properties
1) Chemical 2) Electrical 3) Thermal 4) Optical
Chapter 2 - 47
Electronic Configurations ex: Fe - atomic # = 26
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
1s2 2s2 2p6 3s2 3p6 3d 6 4s2
Chapter 2 - 48
The Periodic Table • Columns: Similar Valence Structure
Adapted from Fig. 2.6, Callister 7e.
Electropositive elements: Readily give up electrons to become + ions.
Electronegative elements: Readily acquire electrons to become - ions.
give
up
1e
give
up
2e
give
up
3e
iner
t gas
es
acce
pt 1
e ac
cept
2e
O
Se Te Po At
I Br
He Ne Ar Kr Xe Rn
F Cl S
Li Be H
Na Mg
Ba Cs Ra Fr
Ca K Sc Sr Rb Y
Chapter 2 - 49 Smaller electronegativity Larger electronegativity
Electronegativity - Tells us whether a given bond will be nonpolar covalent, polar covalent or ionic. - The ability of an atom in a molecule to attract electrons to itself. - Ranges from 0.7 to 4.0.
Chapter 2 -
REVIEW OF ATOMIC STRUCTURE (FRESHMAN CHEMISTRY)
• Mass of an atom: – Proton and Neutron: ~ 1.67 x 10-27 kg – Electron: 9.11 x 10-31
kg • Charge:
– Electrons and protons: (±) 1.60 x 10-19 C
– Neutrons are neutral The atomic mass (A): total mass of protons + total mass of neutrons Atomic weight ~ Atomic mass # of protons are used to identify elements (Z) # of neutron are used to identify isotopes ( e.g. 14C6 and 12C6 ) Isotopes are written as follows: AXZ , i.e. 1H1, 2H1, 3H1
ATOMS = (PROTONS+NEUTRONS) + ELECTRONS NUCLEUS BONDING
Chapter 2 -
Atomic bonding in solids
Things are made of atoms—little particles that move around, attracting each other when they are a little distance apart, but repelling upon being squeezed into one another. In that one sentence ... there is an enormous amount of information about the world. — Richard P. Feynman
Chapter 2 -
Two Allotropes of CARBON
Graphite Diamond
- Relatively soft - Greasy feel to it - Reasonably good
conductor of electricity
- The hardest known material - Poor conductor of electricity
The disparities in properties are attributed to a type of interatomic bonding found in graphite that does not exist in
diamond
Why Study Atomic Structure and Interatomic Bonding?
Chapter 2 -
Atomic Bonding in Solids
• Attractive component is due to type of the bonding (minimize energy thru electronic configuration)
• Repulsive component is due to negatively charged electron clouds for two atoms and important only at small values of r
Interatomic separation, r
Interatomic forces that bind the atoms together are important to understand many properties of materials. • Start with two atoms infinitely separated • At large distances, interactions are negligible • At small distances, each atom exerts forces on the other. • Two types of forces: attractive, FA, and repulsive, FR
Chapter 2 -
Atomic Bonding
IONIC BONDING
COVALENT BONDING
METALLIC BONDING
Sulfur
Bromine Sucrose Magnesium
Gold
Copper
Magnesium oxide
Potassium dichromate
Nickel (II) oxide
Sharing of electrons between two atoms Bonding of metal atoms to
neighbor atoms
Electrostatic forces between ions
Essentially atoms either want to give up (transfer) or acquire (share) electrons to complete electron configurations; minimize their energy
Transfer of electrons => ionic bond Sharing of electrons => covalent Metallic bond => sea of electrons
Chapter 2 - 56
• Occurs between + and – ions (anion and cation). • Requires electron transfer. • Large difference in electronegativity required. • Example: NaCl
Ionic Bonding
Na (metal) unstable
Cl (nonmetal) unstable
electron
+ - Coulombic Attraction
Na (cation) stable
Cl (anion) stable
Chapter 2 - 57
Ionic bond – metal + nonmetal
donates accepts electrons electrons
Dissimilar electronegativities
ex: MgO Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4 [Ne] 3s2
Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6 [Ne] [Ne]
Chapter 2 - 8
IONIC BONDING
Oppositely charged ions attract, attractive force is coulombic. Ionic bond is non-directional, ions get attracted to one another in any direction. Bonding energies are high => 2 to 5 eV/atom,molecule,ion Hard materials, brittle, high melting temperature, electrically and thermally insulating
Chapter 2 - 59
Ionic Bonding • Energy – minimum energy most stable
– Energy balance of attractive and repulsive terms
Attractive energy EA
Net energy EN
Repulsive energy ER
Interatomic separation r
r A
n r B EN = EA + ER = - -
Adapted from Fig. 2.8(b), Callister 7e.
Chapter 2 - 60
• Predominant bonding in Ceramics
Adapted from Fig. 2.7, Callister 7e. (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.
Examples: Ionic Bonding
Give up electrons Acquire electrons
NaCl MgO CaF 2 CsCl
Chapter 2 - 61
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 7e.
Covalent Bonding • Requires shared electrons • Example: CH4
shared electrons from carbon atom
shared electrons from hydrogen atoms
H
H
H
H
C
CH 4
Chapter 2 - 10
COVALENT BONDING
Covalent bonds are formed by sharing of the valence electrons Covalent bonds are very directional Covalent bond model: an atom can have at most 8-N’ covalent bonds, where N’ = number of valence electrons Covalent bonds can be very strong, eg diamond, SiC, Si, etc, also can be very weak, eg Bismuth Polymeric materials do exhibit covalent type bonding.
Diamond, sp3
Chapter 2 - 11
• Molecules with nonmetals • Molecules with metals and nonmetals • Elemental solids (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
co
lum
n IV
A
Sn 1.8Pb 1.8
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.
EXAMPLES: COVALENT BONDING
Chapter 2 - 64
Primary Bonding • Ionic-Covalent Mixed Bonding
% ionic character =
where XA & XB are Pauling electronegativities of the
corresponding elements
%) 100 ( x
!
1"e" (XA"XB)2
4
#
$
% % %
&
'
( ( (
ionic 70.2% (100%) x e1 characterionic % 4)3.15.3( 2
=⎟⎟⎟
⎠
⎞
⎜⎜⎜
⎝
⎛
−=
−−
Ex: MgO XMg = 1.3 XO = 3.5
Chapter 2 - 12
• Arises from a sea of donated valence electrons (1, 2, or 3 from each atom).
• Primary bond for metals and their alloys
Adapted from Fig. 2.11, Callister 6e.
METALLIC BONDING
Non valence and atomic nuclei form ion cores. Ion cores in the “sea of electrons”. Valance electrons belong no one particular atom but drift throughout the entire metal. “Free electrons” shield +’ly charged ions from repelling each other…
Chapter 2 - 66
Arises from interaction between dipoles
• Permanent dipoles-molecule induced
• Fluctuating dipoles
-general case:
-ex: liquid HCl
-ex: polymer
Adapted from Fig. 2.13, Callister 7e.
Adapted from Fig. 2.14, Callister 7e.
SECONDARY BONDING
asymmetric electron clouds
+ - + - secondary bonding
H H H H
H 2 H 2
secondary bonding
ex: liquid H 2
H Cl H Cl secondary bonding
secondary bonding + - + -
secondary bonding secondary bonding
Chapter 2 - 67
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
Chapter 2 - 69
Bonding in Solids
The physical properties of crystalline solids, such as melting point and hardness, depend both on the arrangements of particles and on the attractive forces between them.
Chapter 2 - 70
• Bond length, r
• Bond energy, Eo
• Melting Temperature, Tm
Tm is larger if Eo is larger.
Properties From Bonding: Tm
r o r
Energy r
larger Tm
smaller Tm
Eo = “bond energy”
Energy
r o r unstretched length
Chapter 2 - 16
• Elastic modulus, E
• E ~ curvature at ro
ΔL F 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
Chapter 2 - 72
Ceramics (Ionic & covalent bonding):
Metals (Metallic bonding):
Polymers (Covalent & Secondary):
Large bond energy large Tm large E small α
Variable bond energy moderate Tm moderate E moderate α
Directional Properties Secondary bonding dominates
small Tm small E large α
Summary: Primary Bonds
secondary bonding
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