The Glue That Holds Minerals Together Bonding in Minerals The Glue That Holds Minerals Together GLY 4200 – Fall, 2019 © D. L. Warburton 2019 So far the discussion has centered on individual atoms and ions – but minerals are made of tremendous numbers of atoms and ions – (one mole = one formula weight = 6.02 x 1023 atoms or molecules). What holds the crystal structure together? Are all the links, called bonds, between ions the same? There are five bond types recognized at present – the type of chemical bonding is largely responsible for the physical properties of crystal, which are much easier to observe and thus more familiar – these properties include: Hardness, cleavage, fusibility, electrical or thermal conductivity, and the coefficient of thermal expansion. Some of these properties will be discussed in more detail later, but in general the stronger the bond the harder the crystal, the higher its melting point, and the smaller the thermal expansion coefficient.

Types of Bonds Intramolecular Intermolecular Ionic Covalent Metallic Hydrogen Van der Waals These are the extreme cases – all bonds have various degrees of these principal types present at the same time.

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

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 Chemical elements combine when it is energetically favorable for them to do so – the types of bonds formed depend on the electronic configuration for the ions, and are energetically the most stable – energy is released when elements combine.

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 Ionic bonds – in an ideal ionic bond the cation completely loses an electron, which is transferred to the union – the resulting electrostatic attraction holds the ions together.

Halite Halite, NaCl, is a classic example of an ionically bonded substance The sodium donates an electron to chlorine to complete the eight-electron subshell on chlorine

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 Because the attraction is electrostatic each cation tries to surround itself with anions and each anion with cations – result is that the ionic bond is non-directional. Crystals with ionic bonding do not have discrete molecules. The physical properties of ionically bound crystals are: (as above)

Polarization Polarity is the distortion of the electron cloud of one atom by another. A standard example is often hydrogen chloride (HCl) Ionic bond produces polarization, a distortion of the electron cloud of one atom by another.

Does Size Affect Polarizing Power? Yes, and so does electronegativity The greater the electronegativity, the greater the polarizing power So for hydrogen halogen compounds: Ionic size and electronegativity both affect polarization. As the ionic size increases, polarizing ability decreases, so fluorine is much more polarizing than iodine. An increase in electronegativity increases polarization. Bond polarity has a huge hand in determining chemistry

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 The relative size of ions also is of importance. Anions and cations of similar sizes form stronger, more stable bonds than ions of greatly differing sizes

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) If the ions are of greatly differing size, the bonding is more covalent and less ionic.

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 Some cations are small, with high charges. Such cations tend to be have a large polarizing power. Large anions with low charge have a large polarizability. Cations with high polarizing power tend to polarize anions, and the bonding will be more covalent.

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 Indeed, some highly polarizing cations may destroy anionic groups, making the compounds so unstable that they do not exist.

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 There is a nomenclature associated with ionic bonds. Suffixes added to a compound name describe the compound.

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

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 Covalent bond – atoms involved share electrons in order to achieve inert gas configuration. Covalent bonding (or electron – sharing) is the strongest type of bond. Covalent bonds are formed between similar atoms, from the same part of the periodic table – in contrast with ionic bonds which form only between widely separated elements. The electronegativity of these atoms will be similar. Indeed, covalent bonds may form between elements of the same element, so that the electronegativity difference is zero.

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 This may lead to the formation of hybrid orbitals, such as is seen in many carbon compounds. C2H6

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

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

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

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 Source: http://www.chem1.com/acad/webtext/chembond/cb03.html Only the outermost, “valance” electrons are involved in bonding.

Covalent Bonds Between Different Elements Hydrogen chloride (aka hydrochloric acid) The hydrogen has a helium structure, and the chlorine an argon structure Covalent bonding between different elements follows the octet rule.

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 Like all rules, the octet rule has limitations. For light elements, we are dealing with s electrons, so the bonding is s2.

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”

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 Metallic bonding – metals are characterized a different type of bond than there salts are.

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 

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

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

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

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

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

Other Physical Properties Low hardness Low melting point & boiling point

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 dull Image source: http://uts.cc.utexas.edu/~rmr/llanite.html Thin section of llanite, a hypabyssal rhyolite porphyry dike – opaque mineral grains are magnetite

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 Although often overlooked in discussions of bonding, they are extremely important in the physical and chemical properties of materials.

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 Hydrogen bonding – Hydrogen can donate or share only one electron so it should form only one hand. However, if the H bonds to an anion of high electronegativity, both electrons will be primarily on the anion. The H ion then carries a δ+ charge large enough to attract electrons from other anions.

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 Hydrogen bonds are weak ionic bonds but are stronger than Van der Waals – they don’t play a large role in mineralogy, but nevertheless are important. They are extremely important in the properties of water, and in minerals containing hydroxyl groups.

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! The importance in water is shown by a comparison of the melting and boiling points of various hydrogen- oxygen group compounds.

Relative Boiling Points Compound Melting Point, oC Boiling Point, oC H2O 100 H2S -85.5 -60.7 H-bonding stabilizes water – thus it remains a liquid rather than being a gas.

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 The boiling point of the hydride of the first element in the nitrogen, oxygen, and halogen groups is abnormally high.

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

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

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

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

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) Two of the best examples of this are the two most common elements in the universe, hydrogen and helium. The weakness of the intermolecular forces accounts for the low boiling points of these elements, 21K and 4K, respectively. The forces involved are the fifth type of bonding, known as Van der Waals.

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 Much of the information for this section, as well as the images, comes from: http://www.chemguide.co.uk/atoms/bondingmenu.html#top Van der Waals attractions occur when atoms or molecules come into close proximity, which may allow a dipole to be induced

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

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

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 +

Temporary Fluctuating Dipoles An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule This produces fluctuating dipoles.

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 The polarized atoms or ions form very weak electrostatic bonds known as Van der Waals bonds. This is the London or dispersion case.

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 In helium, we are dealing with a single atom, with two electrons. By chance, both electrons may be on the same side of the atom at times. This creates a momentary dipole.

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 Interactions between the momentary dipole and an unpolarized atom or molecule create induced dipoles.

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 The induced dipoles may fluctuate, which again allows a weak attraction.

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 Such interactions are not restricted to pairs of atoms or molecules. They may be synchronized throughout a lattice.

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 In mineralogy these bonds occur between neutral molecules or uncharged structure units held in the crystal lattice.

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

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 The strength of the dispersive forces is essentially impossible to measure, because the size of the attraction varies considerably with the size of the molecule and its shape.

Size and Dispersion Forces Fritz London 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 This leads to an increase in properties like boiling point. Why?

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

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 As atoms become larger, the number of electrons increases, as does atomic radius. The greater number of electrons allows for a greater number of potential dipoles. In addition, the outermost electrons are much more weakly held, and therefore are easier to polarize. This accounts for differences in boiling points between the noble gas atoms.

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"

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 Some molecules, like hydrogen chloride, have permanent dipoles. Interactions between the permanent dipoles of these molecules produces the dipole-dipole type of Van der Waal bonding.

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

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

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

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

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°

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

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

Graphite Covalent bonding within a sheet The sheets are held together by Van der Waals bonds – very easy to break in one direction Thus soft with perfect cleavage Source: http://www.benbest.com/cryonics/graphite.gif