1

Bonding in Minerals

The Glue That Holds Minerals Together

GLY 4200 –Fall, 2012

2

Types of Bonds

• Intramolecular Ionic Covalent Metallic

• Intermolecular Hydrogen Van der Waals

3

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

4

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

5

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

6

Halite

• Halite, NaCl, is a classical example of an ionically bonded substance

• The sodium donates an electron to chlorine to complete the eight-electron subshell on chlorine

7

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

8

Polarization

• Polarity is the distortion of the electron cloud of one atom by another.

• A standard example is often hydrogen chloride (HCl)

9

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

10

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

11

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)

12

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

13

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 CO3

2- ion into self-decomposition

Instead of Al2(CO32-)3, carbon dioxide is driven

off, leaving aluminum oxide

14

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

15

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

16

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

17

Carbon• 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

C2H6

18

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

19

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

20

Basis of Octet Rule

• By allowing each nucleus to claim half-ownership of a shared electron, more electrons are effectively “seeing” more nuclei, leading to increased electrostatic attractions and a lowering of the potential energy

21

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

22

Covalent Bonds Between Different Elements

• Hydrogen chloride (aka hydrochloric acid)

• The hydrogen has a helium structure, and the chlorine an argon structure

23

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

24

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”

25

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

26

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 

27

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

28

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

29

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

30

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

31

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

32

Other Physical Properties

• Low hardness

• Low melting point & boiling point

33

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 Weathering may make this

luster dullThin section of llanite, a hypabyssal rhyolite porphyry dike – opaque mineral grains are magnetite

34

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

35

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

36

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 co-ordinate bond

• It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction

37

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!

38

Relative Boiling Points

Compound Melting Point, oC Boiling Point, oC

H2O 0 100

H2S -85.5 -60.7

39

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 described as hydrogen bonds

40

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

41

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

42

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

43

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

44

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)

45

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

46

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

47

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

48

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 +

49

Temporary Fluctuating Dipoles

• An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule

50

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

51

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

52

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

53

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

54

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

55

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

56

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 synchronization

57

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

58

Size and Dispersion Forces

• As atomic size increases, so do the dispersion forces

Element Boiling Point, ºC

Helium -269

Neon -246

Argon -186

Krypton -152

Xenon -108

Radon -62

59

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

60

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

61

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"

62

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

63

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

64

Relative Strength of Dipole-dipole 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

65

Molecular Comparison

• For example, the boiling points of ethane, CH3CH3, and fluoromethane, CH3F, are:

66

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

67

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°

68

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

69

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

70

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