Bonding in Minerals The Glue That Holds Minerals

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 H 2 • When 2 or more atoms of different elements combine together chemically, a compound is formed – example, water H 2 O • 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, Na. Cl, 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 • Na. Cl melts at 801°C, strong attraction between particles in solid lattice structure (Ionic bonding likely) • Al. Cl 3 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 CO 32 - ion into self-decomposition § Instead of Al 2(CO 32 -)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 H 2 S – hydrogen sulfide • Ending in –ate means oxygen is present § e. g. Ca. S = calcium sulfide § Ca. SO 4 = calcium sulfate 14

Ionic Bond Nomenclature II • Ending in –ite less oxygen present than in –ate compounds § e. g. Na. S = sodium sulfide § Na. SO 4 = sodium sulfate § Na. SO 3 = 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 C 2 H 6 • 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 s 2 p 6 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, s 2 p 6) is achieved when two fluorine atoms (s 2 p 5) 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 s 2 rather than s 2 p 6 • Thus we write Li. H 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 31

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, o. C Boiling Point, o. C H 2 O 0 100 H 2 S -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 NH 3, H 2 O 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, H 2, 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 - H 2, perhaps, or Br 2 • 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 53

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 57

Size and Dispersion Forces • As atomic size increases, so do the dispersion forces Element Helium Neon Argon Boiling Point, ºC -269 -246 -186 Krypton Xenon Radon -152 -108 -62 58

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 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, CH 3, and fluoromethane, CH 3 F, 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