Transcript Atoms and Bonding

ENG2000 Chapter 2
Atoms and Bonding
ENG2000: R.I. Hornsey
Atom: 1
Overview
• Atomic structure
 fundamentals
 electrons and atoms
• Atomic bonding
 bonding mechanisms and forces
 bonding types
 molecule
• Atomic bonding is determined by the electronic
configurations of the atoms
• Atomic bonding determines all the fundamental
physical and electronic, magnetic, optical etc
properties
ENG2000: R.I. Hornsey
Atom: 2
Atoms
• For our purposes, atoms are made from three
fundamental particles




proton (charge = +q, m = 1.67 x 10-27kg)
neutron (charge - 0, m = 1.67 x 10-27kg)
electron (charge = -q, m = 9.1 X 10-31kg)
q = 1.6 x 10-19C
• An element is defined by its atomic number, Z
 Z = number of protons in the atomic nucleus
 1 (H) ≤ Z ≤ 92 (U) for naturally occurring elements
• The atomic mass (A) is the sum of proton and
neutron masses in the nucleus
 # neutrons (N) can vary to give different isotopes of the
same element
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Atom: 3
Masses
• The atomic weight (really a mass) is typically
given in units of grams per mole (g/mol.)
 1 mole of a substance contains 6.023 x 1023 particles –
Avogadro’s Number
 e.g. iron has an atomic weight of A = 55.85 g/mol.
• Where several isotopes of a substance are
present, the atomic weight is calculated from the
appropriate fractions of the weights of the
individual isotopes
ENG2000: R.I. Hornsey
Atom: 4
Bohr Atom
• In the early years of the 20th century, atomic
spectroscopy indicated that electron energies are
quantised
 the Bohr planetary model of the atom is an early attempt to
visualise a system that would yield quantised energies
 it is incomplete because it does not explain why the orbiting
electrons do not emit electromagnetic radiation
nucleus
orbiting
electrons
ENG2000: R.I. Hornsey
http://www.marxists.org/reference/subject/philosophy/images/bohr.jpg
http://csep10.phys.utk.edu/astr162/lect/light/bohrframe/bohr2.gif
Atom: 5
Energy levels
0
-1.5eV
-3.4eV
-13.6eV
n=3
n=2
n=1
potential
energy
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• These are the first three energy
levels for an isolated H atom
• 1eV = 1.6 x 10-19J
 the energy gained by an electron
accelerated through a potential
difference of 1V
• To move between energy levels
requires a ‘quantum jump’
• More refined measurements
showed that each ‘n’ level was in
fact composed of several
discrete energies
• Better models needed
Atom: 6
Other energy levels
• Due to electrostatic (and
other) interactions
between electrons, each
primary energy is in fact
several closely spaced
levels
• These are named s, p, d, f
 after the shapes of the
spectroscopic lines in the early
experiments
 sharp, principal, diffuse, fine
• Energy levels are identified
by four quantum numbers
ENG2000: R.I. Hornsey
0
-1.5eV
-3.4eV
-13.6eV
n=3
n=2
n=1
3d
3p
3s
2p
2s
1s
potential
energy
Atom: 7
Wave mechanics
• Numerous pieces of evidence suggest that all
particles can be thought of as both particles and
waves
 interference effects – quintessentially wave-like phenomena
– can be seen with electrons
 quantum-mechanical tunnelling (see later)
 called wave-particle duality
• A particle’s wavelength is calculated from the de
Broglie formula (1924)
h

mv
 where h is Planck’s constant (1901); h = 6.62 x 10-34 Js
 m is the mass, v is the velocity
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Atom: 8
• The spatial properties of the
wave (x, y, t, intensity) are
closely related to the
probability of finding the
particle at a particular
location
 the important part here is that the
wave mechanical nature of an
electron implies that we do not
know the precise position
 only a probability function giving
the likelihood of an electron’s
position
Atom: 9
ENG2000: R.I. Hornsey
Callister
Quantum numbers
• Principal quantum numbers are n = 1, 2, 3, 4 …
 they correspond to energy shells K, L, M, N, …
• Second quantum number, l, is [s, p, d, f]
 related to the spatial shape of the energy level
 the number of sub-shells is limited to the ‘n’ for the level
• A third number, ml (the magnetic quantum
number), describes the number of available
energy states per sub-shell
 1 for s, 3 for p, 5 for d, 7 for f
 the energies of these states are identical in the absence of a
magnetic field, but split when a field is applied
• The last quantum number is the spin, ms
 ms = ± 1/2
ENG2000: R.I. Hornsey
Atom: 10
Planetary picture
• Very, very roughly these for quantum numbers
can be visualised in terms of a planetary orbit




n corresponds to the radius of the orbit
l corresponds to the shape of the orbit
ml corresponds to the tilt (or inclination) of the orbit
ms represents the two directions the ‘planet’ can spin
l
n
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ml
ms
Atom: 11
Maximum number of states
n
1
2
3
4
sub-shell
K
L
M
N
# states
max # electrons
sub-shell*
shell
2
s
1
2
s
1
2
p
3
6
s
1
2
p
3
6
d
5
10
s
1
2
p
3
6
d
5
10
f
7
14
8
18
32
* # states x 2, because two electrons (with ± spin) can exist in each state
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Atom: 12
Notation
• The conventional notation is: n [s,p,d,f]#
 where # is the number of available states that actually
contain electrons
• For example:








H = 1s1
He = 1s2
Li = 1s22s1
Be = 1s22s2
B = 1s22s22p1
Ne = 1s22s22p6
Na = 1s22s22p63s1
Al = 1s22s22p63s23p1
ENG2000: R.I. Hornsey
Atom: 13
Filling the energy levels
• Electrons occupy the lowest energy state
available
 note that e.g. 4s < 3d, so fills first
http://www.webelements.com/webelements/elements/media/e-config/H.gif
http://www.chemtutor.com/scheme.gif
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Atom: 14
Valence electrons
• The number of electrons occupying the
outermost shell of an atom – the valence
electrons – is important for determining the
chemical properties of the atom
 because these electrons will be involved with the bonding of
atoms
• Atoms with one electron too many (e.g. Na) or
one too few (e.g. F) are highly reactive
• Atoms with full shells (e.g. Ne, Ar) tend to be inert
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Atom: 15
Periodic table
• The periodic table of the elements was originally
drawn up according to the chemical properties of
the elements
 as we have seen these properties are closely related to the
atomic electron configurations
 the seven horizontal rows are called ‘periods’
 chemical properties vary from one end of the period to the
other
 each column – a ‘group’ – displays similar chemical
properties and similar valence structures
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Atom: 16
http://helios.augustana.edu/physics/301/periodic-table-fix.jpg
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Groups
• Group 0 contains the inert (noble) gases
• Group IV includes Si
 important materials in Si chip manufacture are B (III) and P
(V) – as we will see later
 together, these materials are between metals and nonmetals
• Group VII are the ‘halogens’ and are one electron
deficient in the valence shell
• Groups IA and IIA are the alkali and alkaline earth
metals
• Groups IIIB – IIB are the transition metals, which
have partially filled lower (d) energy states
 includes ‘real’ metals and magnetic materials
ENG2000: R.I. Hornsey
Atom: 18
Bonding
• Atomic bonding determines many of the physical
properties of a material
• If two isolated atoms are brought closer together
the net force varies with distance
 there is a mechanism-specific attractive force (FA)
 and a repulsive force (FB), which increases when the atoms
are sufficiently close for the outer shells to overlap
 equilibrium is reached when FA + FB = 0
 this is at r0 on the following page
• The potential energy at r0 is the bonding energy,
E0
 and represents the energy required to separate the atoms to
an infinite distance
 e.g. thermal energy to melt the material
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Atom: 19
Callister
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Atom: 20
Ionic bonding
• Ionic bonding occurs in materials composed of a
metallic and a non-metallic element
 the metallic element easily donates its electron to the nonmetallic element
 the metal becomes a positive ion, while the non-metal is
negatively ionised
ENG2000: R.I. Hornsey
http://www.agen.ufl.edu/~chyn/age4660/lect/lect_02/2_11a.gif
http://www.astro.lsa.umich.edu/users/cowley/lecture11/images/NaCl.jpg
Atom: 21
• Here, the attractive forces are coulombic, arising
from the attraction of oppositely charged ions
 E0 ≈ 600 – 1500 kJ/mol., or 3 – 8 eV/atom
 this relatively large bonding energy is reflected in typically
high melting temperatures for ionically bonded materials
 including ceramics
ENG2000: R.I. Hornsey
Atom: 22
Covalent bonding
• As the name suggests, covalent bonds are
formed by sharing valence electrons between the
constituent atoms
 thereby causing all atoms to achieve a full – and stable –
outer shell
 the classic example is methane, CH4
ENG2000: R.I. Hornsey
http://www.mse.cornell.edu/courses/engri111/images/covalent.gif
Atom: 23
• Covalent bonds are also common in elements
from the right-hand side of the periodic table
 notably the semiconductors silicon and germanium, as well
as carbon
 also compound semiconductors, e.g. GaAs and InP
• The number of atoms participating in the bond is
determined by the number of valence electrons
 Si is in group IV, so has 4 valence electrons, and therefore
bonds with 4 neighbouring atoms
ENG2000: R.I. Hornsey
Atom: 24
Metallic bonding
• Metallic elements have one or two (possibly
three) ‘loose’ valence electrons
 which are relatively freely donated by all atoms
• The result is a structure in which ionised atoms
(because they have donated their electron) are
‘suspended’ in a ‘sea’ of electrons
 the ions are fixed in place because the negatively charged
electron sea exerts an equal attraction in all directions
+ve ion cores
ENG2000: R.I. Hornsey
-ve electron sea
Atom: 25
Metallic bonding
• Because the donated
electrons are freely mobile,
the electrical conductivity
of metals is high
 heat can also be transmitted by
electrons, so metals are good
thermal conductors
• Ionically and covalently
materials are typically good
electrical insulators
 there is another mechanism for
thermal transport which means
that e.g. ceramics can be good
thermal conductors
ENG2000: R.I. Hornsey
http://207.10.97.102/chemzone/lessons/03bonding/mleebonding/metallicblue.gif
Atom: 26
Other bonding types
• Ionic, covalent and metallic are the primary
bonding types
• Secondary bonds are those that exist between all
atoms, but are relatively weak and may be
obscured by the primary bonds
• van der Waals bonds are typically only
0.1eV/atom (c.f. 8eV/atom for ionic)
 and results from atomic or molecular dipoles
+
–
+
–
• Dipoles can result from molecular bonds –
especially those involving H atoms – atomic
vibrations, or external electric fields
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Atom: 27
Melting temperatures
Type
Ionic
Covalent
Metallic
van der Waals
Hydrogen
ENG2000: R.I. Hornsey
Substance
Energy (eV/atom) Melt. Temp (°C)
NaCl
3.3
801
MgO
5.2
2800
Si
4.7
1410
C
7.4
>3550
Hg
0.7
-39
Al
3.4
660
Fe
4.2
1538
W
8.8
3410
Ar
0.08
-189
Cl2
0.32
-101
NH3
0.36
-78
H 20
0.52
0
Atom: 28
Summary
• Atomic structure is determined by quantum
mechanics
 four quantum numbers determine energy states
 states may or may not be occupied by electrons
• Atomic structure determines chemical and
physical properties of the elements
 periodic table
• Structure also determines how atoms bond
 primary – ionic, covalent, metallic
 secondary – van der Waals, hydrogen
ENG2000: R.I. Hornsey
Atom: 29