Transcript Bonding in Minerals - Florida Atlantic University

Bonding in Minerals
The Glue That Holds Minerals Together
GLY 4200 –Fall, 2012
1
Types of Bonds
• Intramolecular
 Ionic
 Covalent
 Metallic
• Intermolecular
 Hydrogen
 Van der Waals
2
Definition of Bonding
• A chemical bond is an attraction between
atoms brought about by:
 A sharing of electrons between two atoms or,
 A complete transfer of electrons
• When a chemical bond is formed, energy is
released
• Breaking chemical bonds requires energy
3
Substances Formed by Bonding
• When two or more atoms of the same
element bond together, a molecule is
formed – example, hydrogen H2
• When 2 or more atoms of different elements
combine together chemically, a compound
is formed – example, water H2O
• Most minerals are compounds
4
Ionic Bonding
• Ionic bonding is the result of electrostatic
attraction between two oppositely charged
ions
• Positive ions are formed from metals
(usually) and negative ions are usually
formed from non-metals
5
Halite
• Halite, NaCl, is a
classical example of
an ionically bonded
substance
• The sodium donates
an electron to chlorine
to complete the eightelectron subshell on
chlorine
6
Physical Properties of
Ionically Bonded Crystals
• Ionic bonding is non-directional
• Ionically bonded minerals may yield ions to
solution
• Moderate hardness
• Fairly high to very high melting points &
boiling points
• Poor thermal & electrical conductors except
near the melting points
7
Polarization
• Polarity is the distortion of the electron cloud of one
atom by another.
• A standard example is often hydrogen chloride
(HCl)
8
Does Size Affect Polarizing Power?
• Yes, and so does electronegativity
• The greater the electronegativity, the greater the
polarizing power
• So for hydrogen halogen compounds:
• Bond polarity has a huge hand in determining
chemistry
9
Relative Size of Ions
• The size mismatch of the anions and cations
is of importance also
• If two ions are similar in size, then they
exist quite happily
• If there is a size mismatch, then is it quite
likely that covalent bonding will occur
10
Size Mismatch
• NaCl melts at 801°C, strong attraction between
particles in solid lattice structure (Ionic bonding
likely)
• AlCl3 sublimes (goes from solid to gas not via the
liquid phase) at 180°C, so there are no strong
attractions present (Covalent bonding likely)
11
Polarizing Cations
• If the cation is small and highly charged, it has a
large polarizing power
• If the anion is large and has a relatively low
charge, then it is said to have a large polarizability
• In the first case, the anion is being polarized by
the cation
 There will be a significant degree of covalent character
to the bond
12
Non-Existent Compounds
• There are some ionic compounds that do not
exist at all
• Aluminum carbonate is an example
 The aluminum 3+ cation is so small and highly
polarizing that is completely distorts the large
CO32- ion into self-decomposition
 Instead of Al2(CO32-)3, carbon dioxide is driven
off, leaving aluminum oxide
13
Ionic Bond Nomenclature
• Compounds ending in –ide are simple
binary compounds containing 2 elements even if there is no metal
 e.g H2S – hydrogen sulfide
• Ending in –ate means oxygen is present
 e.g. CaS = calcium sulfide
 CaSO4 = calcium sulfate
14
Ionic Bond Nomenclature II
• Ending in –ite less oxygen present than in
–ate compounds
 e.g. NaS = sodium sulfide
 NaSO4 = sodium sulfate
 NaSO3 = sodium sulfite
15
Covalent Bonding
• Covalent bonds involve a complete sharing
of electrons and occur most commonly
between atoms that have partially filled
outer shells or energy levels
• Thus, if the atoms are similar in
electronegativity then the electrons will be
shared
16
Carbon
C2H6
• Carbon forms covalent
bonds
• The electrons are in hybrid
orbitals formed by the
atoms involved as in this
example: ethane
• Diamond is strong
because it involves a vast
network of covalent bonds
between the carbon atoms
in the diamond
17
Physical Properties of Covalently
Bonded Crystals
• Covalent bonds are directional and
molecules are often formed.
• Covalently bonded crystals do not yield
ions to solutions, as ionically bond crystals
sometimes do
• Covalent crystals have very high melting
points & boiling points
18
Octet Rule
• The idea that the noble-gas configuration is a
particularly favorable one which can be achieved
through formation of electron-pair bonds with
other atoms is known as the octet rule
• Present-day shared electron-pair theory is based
on the premise that the s2p6 octet in the outermost
shells of the noble gas elements above helium
represents a particularly favorable configuration
19
Basis of Octet Rule
• By allowing each nucleus to claim halfownership of a shared electron, more
electrons are effectively “seeing” more
nuclei, leading to increased electrostatic
attractions and a lowering of the potential
energy
20
Fluorine
• Noble gas configuration (in this case, that of neon,
s2p6) is achieved when two fluorine atoms (s2p5) are
able to share an electron pair,which becomes the
covalent bond
• Only the outer (valence shell) electrons are
involved
21
Covalent Bonds Between
Different Elements
• Hydrogen chloride (aka
hydrochloric acid)
• The hydrogen has a helium
structure, and the chlorine an
argon structure
22
Octet Limitations –
Light Elements
• For the lightest atoms the octet rule must be
modified, since the noble-gas configuration
will be that of helium, which is simply s2
rather than s2p6
• Thus we write LiH as Li:H, where the
electrons represented by the two dots come
from the s orbitals of the separate atoms
23
Octet Limitations –
Heavy Elements
• The octet rule applies quite well to the first full row of the
periodic table (Li through F), but beyond this it is generally
applicable only to the non-transition elements, and even in
many of these it cannot explain many of the bonding
patterns that are observed
• The principal difficulty is that a central atom that is bonded
to more than four peripheral atoms must have more than
eight electrons around it if each bond is assumed to consist
of an electron pair
• In these cases, we hedge the rule a bit, and euphemistically
refer to the larger number of electrons as an “expanded
octet”
24
Metallic Bonding
• A metallic bond occurs when positive metal
ions like Cu+2 or Fe+3 are surrounded by a
"sea of electrons" or freely-moving valence
electrons.
• The valence electrons are not bound to any
particular cation, but are free to move
throughout the metallic crystal
25
Sea of Electrons
• In the picture, the red
circles are metal
cations packed in a
crystal lattice
• The black dots
represent the "sea" of
freely moving valence
electrons
26
Minerals with Metallic Bonding
• Only native metals display metallic bonding
• Alkaline metals are for too reactive to be
found uncombined in nature
• Only a few minerals, such as gold, silver,
copper and the platinum group are
metallically bound
27
Conductivity Properties
• Metals are good conductors of electricity
 Electric current is a movement of free electrons
 Substances with partial metallic bonding may
be semiconductors
• Metals are good conductors of heat
 Heat is transferred by the increased speed of
electrons
28
Flexibility Properties of Metallic
Bonding
• The model of metallic bonding explains the
flexibility properties of metals
 Metals are ductile - They can be drawn into
wires because electrons are mobile.
 Metals are malleable - They can be hammered
into sheets due to mobility of electrons
 Metals are tenacious – they do not break easily
29
Electronic Forces in Metals
• Strong attraction between positive nuclei
and the electrons
• The positive ions repel as do the negative
electrons
• The electrons move constantly, but some
electrons will always be between the layers
creating an attraction and keeping them
attracted to one another
30
Explanation of
Metallic Properties
• An impact will allow a shearing effect as
there is a degree of repulsion between layers
• The sea of electrons allows movement of
ions, therefore pure metals are not brittle
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Other Physical Properties
• Low hardness
• Low melting point & boiling point
32
Optical Properties
• Metallically bonded minerals
are opaque –
 This is often true at very
small thicknesses, such as
the 30 micron thickness of a
thin section
• Metallically bonded
substance usually show
metallic luster
Thin section of llanite, a
hypabyssal rhyolite porphyry
dike – opaque mineral grains
are magnetite
 Weathering may make this
luster dull
33
Intermolecular Bonds
• Bonds which hold molecules together are
called intermolecular bonds
• In minerals, the concept of a “molecule” is
often inapplicable, but the term is still used
34
Hydrogen Bonding
• In some substances, hydrogen is bonded to
elements which are quite electronegative,
and which possess “lone pairs” of electrons
• Examples include water and ammonia
• Hydrogen bonding leads to the many
anomalous properties of water and ammonia
35
Hydrogen Bond Image
• The δ+ hydrogen is so
strongly attracted to the lone
pair that it is almost as if you
were beginning to form a coordinate bond
• It doesn't go that far, but the
attraction is significantly
stronger than an ordinary
dipole-dipole interaction
36
Relative Bond Strength
• Hydrogen bonds have about a tenth of the strength
of an average covalent bond, and are being
constantly broken and reformed in liquid water
• If you liken the covalent bond between the oxygen
and hydrogen to a stable marriage, the hydrogen
bond has "just good friends" status
• On the same scale, van der Waals attractions
represent mere passing acquaintances!
37
Relative Boiling Points
Compound
Melting Point, oC Boiling Point, oC
H2O
0
100
H2S
-85.5
-60.7
38
Relative Boiling Points
• The boiling point of the hydride of the first element
in each group is abnormally high
• In the cases of NH3, H2O and HF there must be some
additional intermolecular forces of attraction,
requiring significantly more heat energy to break
• These relatively powerful intermolecular forces are
39
described as hydrogen bonds
Water
• Each water molecule can potentially form
four hydrogen bonds with surrounding
water molecules
• There are exactly the right numbers of δ+
hydrogens and lone pairs so that every one
of them can be involved in hydrogen
bonding
40
Ammonia and Hydrogen Fluoride
• In the case of ammonia, the amount of hydrogen
bonding is limited by the fact that each nitrogen
only has one lone pair
 In a group of ammonia molecules, there aren't enough
lone pairs to go around to satisfy all the hydrogens
• In hydrogen fluoride, the problem is a shortage of
hydrogens
• In water, there are exactly the right number of
each
• Water could be considered as the "perfect"
hydrogen bonded system
41
Hydrogen Bonding in Biology
• Hydrogen bonding also holds the DNA
double helix together
• During sexual reproduction, the hydrogen
bonds break, allowing each parent to pass
on a strand of DNA
 The strands recombine to form a new double
helix, a combination of genetic material from
each parent
42
Residual Bonding Forces
• All molecules experience intermolecular
attractions, although in some cases those
attractions are very weak
• Even in a gas like hydrogen, H2, if you slow
the molecules down by cooling the gas, the
attractions are large enough for the
molecules to stick together eventually to
form a liquid and then a solid
43
Hydrogen and Helium
• In hydrogen's case the attractions are so
weak that the molecules have to be cooled
to 21 K (-252°C) before the attractions are
enough to condense the hydrogen as a liquid
• Helium's intermolecular attractions are even
weaker - the molecules won't stick together
to form a liquid until the temperature drops
to 4 K (-269°C)
44
Van der Waals Bonding
• There are two types of Van der Waals
forces
 Dispersion forces are also known as "London
forces" (named after Fritz London who first
suggested how they might arise)
 Dipole-dipole interactions
45
Electrical Attractions
• Attractions are electrical in nature
• In a symmetrical molecule like hydrogen,
however, there doesn't seem to be any
electrical distortion to produce positive or
negative parts
• But that's only true on average
46
Distortion of Electron Cloud
• The lozenge-shaped diagram represents a small
symmetrical molecule - H2, perhaps, or Br2
• The even shading shows that on average there is no
electrical distortion
47
Mobile Electrons
• But the electrons are mobile, and at any one
instant they might find themselves towards
one end of the molecule, making that end δ• The other end will be temporarily short of
electrons and so becomes δ +
48
Temporary Fluctuating Dipoles
• An instant later the
electrons may well
have moved up to the
other end, reversing
the polarity of the
molecule
49
Momentary Dipoles
• This constant "sloshing around" of the
electrons in the molecule causes rapidly
fluctuating dipoles even in the most
symmetrical molecule
• It even happens in monatomic molecules molecules of noble gases, like helium,
which consist of a single atom
50
Helium
• If both the helium electrons
happen to be on one side of
the atom at the same time,
the nucleus is no longer
properly covered by
electrons for that instant
51
Temporary Dipoles and
Intermolecular Attractions
• Imagine a molecule which has a temporary polarity being
approached by one which happens to be entirely non-polar
just at that moment
 A pretty unlikely event, but it makes the diagrams much easier
to draw!
 In reality, one of the molecules is likely to have a greater
polarity than the other at that time - and so will be the dominant
one
52
Induced Dipoles
• This sets up an induced dipole in the approaching
molecule, which is orientated in such a way that
the δ+ end of one is attracted to the δ- end of the
other
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Fluctuating Induced Dipoles
• An instant later the electrons in the left hand molecule may
well have moved up the other end
• In doing so, they will repel the electrons in the right hand
one
• The polarity of both molecules reverses, but you still have
δ+ attracting δ• As long as the molecules stay close to each other the
polarities will continue to fluctuate in synchronization so
that the attraction is always maintained
54
Synchronization in a Lattice
• There is no reason why this has to be
restricted to two molecules
• As long as the molecules are close together
this synchronized movement of the
electrons can occur over huge numbers of
molecules
55
Lattice Diagram
• Diagram shows how a whole lattice of molecules
could be held together in a solid using Van der
Waals dispersion forces
• An instant later, of course, you would have to draw
a quite different arrangement of the distribution of
the electrons as they shifted around - but always in
56
synchronization
Strength of Dispersion Forces
• Dispersion forces between molecules are
much weaker than the covalent bonds
within molecules
• It isn't possible to give any exact value,
because the size of the attraction varies
considerably with the size of the molecule
and its shape
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Size and Dispersion Forces
• As atomic size increases, so do the
dispersion forces
Element
Helium
Neon
Argon
Krypton
Xenon
Radon
Boiling Point, ºC
-269
-246
-186
-152
-108
-62
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Larger Temporary Dipoles
• The reason that the boiling points increase
as you go down the group is that the number
of electrons increases, and so also does the
radius of the atom
• The more electrons you have, and the more
distance over which they can move, the
larger the possible temporary dipoles, and
therefore the larger the dispersion forces
59
Increased “Stickiness”
• Because of the greater
temporary dipoles, xenon
molecules are "stickier"
than neon molecules
• Neon molecules will
break away from each
other at much lower
temperatures than xenon
molecules - hence neon
has the lower boiling
point
60
Bigger Molecules, Higher B.P.’s
• This is the reason that (all other things
being equal) bigger molecules have higher
boiling points than small ones
• Bigger molecules have more electrons and
more distance over which temporary dipoles
can develop - and so the bigger molecules
are "stickier"
61
Dipole-Dipole Interactions
• A molecule like HCl has a permanent dipole
because chlorine is more electronegative
than hydrogen
• These permanent, in-built dipoles will cause
the molecules to attract each other rather
more than they otherwise would if they had
to rely only on dispersion forces
62
Addition of Forces
• It's important to realize that all molecules
experience dispersion forces
• Dipole-dipole interactions are not an alternative to
dispersion forces - they occur in addition to them
• Molecules which have permanent dipoles will
therefore have boiling points higher than
molecules which only have temporary fluctuating
dipoles
63
Relative Strength of Dipoledipole vs. Dispersion Forces
• Surprisingly dipole-dipole attractions are
fairly minor compared with dispersion
forces, and their effect can only really be
seen if you compare two molecules with the
same number of electrons and the same size
64
Molecular Comparison
• For example, the boiling points of ethane,
CH3CH3, and fluoromethane, CH3F, are:
65
Why Ethane and Fluoromethane?
• Both have identical numbers of electrons,
and if you made models you would find that
the sizes were similar - as you can see in the
diagrams
• That means that the dispersion forces in
both molecules should be much the same
66
Importance of Permanent Dipole
• The higher boiling point of fluoromethane is
due to the large permanent dipole on the
molecule because of the high
electronegativity of fluorine
• However, even given the large permanent
polarity of the molecule, the boiling point
has only been increased by some 10°
67
Resonant Bonds
• When a bond has elements of more than one
type of ideal bond, i.e. partial ionic, partial
covalent, it is said to be a resonant bond
• Many carbon compounds exhibit this
behavior
68
Presence of Multiple Bond Types
• If a crystal has more than one type of bond
the weakest bonds present determine the
physical properties which may be very
directional
69
Graphite
• Covalent bonding
within a sheet
• The sheets are held
together by Van de
Waals bonds – very
easy to break in one
direction
• Thus soft with perfect
cleavage
70