|
|
|
|
|
|
|
|
ION | STOCK SYSTEM | OLD SYSTEM | ION | STOCK SYSTEM | OLD SYSTEM |
| Fe2+ | iron II | ferrous | Fe3+ | iron III | ferric |
| Cu+ | copper I | cuprous | Cu2+ | copper II | cupric |
| Au+ | gold I | aurous | Au3+ | gold III | auric |
| Sn2+ | tin II | stannous | Sn4+ | tin IV | stannic |
| Pb2+ | lead II | plumbous | Pb4+ | lead IV | plumbic |
| Hg+ | mercury I | mercurous | Hg2+ | mercury II | mercuric |
| Cr2+ | chromium II | chromous | Cr3+ | chromium III | chromic |
| Mn2+ | manganese II | manganous | Mn3+ | manganeseIII | manganic |
Cl +
e- +
Na+ Cl- + Na+
Na+Cl-
NaCl
Mg2+(SO4)2-
Mg(SO4) or MgSO4
Al3+(PO4)3-
Al(PO4) or AlPO4
FORMULA | COMMON NAME | SYSTEM NAME |
| N2O | nitrous oxide | dinitrogen monoxide |
| NO | nitric oxide | nitrogen monoxide |
| N2O3 | nitrous anhydride | dinitrogen trioxide |
| NO2 | nitrogen dioxide | nitrogen dioxide |
| N2O4 | nitrogen tetroxide | dinitrogen tetroxide |
| N2O5 | nitric anhydride | dinitrogen pentoxide |
| NO3 | nitrogen trioxide | nitrogen trioxide |
*FGP | number | *FGP | number | *FGP | number | *FGP | number |
| mono- | one | di- | two | tri- | three | tetra- | four |
| penta- | five | hexa- | six | hepta- | seven | octa- | eight |
| nona- | nine | deca- | ten | undeca- | eleven | dodeca- | twelve |
In an attempt to simplify, some books may seem to suggest that covalent and ionic bonds are two separate and completely different types of attachment. A covalent bond is a shared pair of electrons. The bond between the two atoms of any diatomic gas, such as chlorine gas, Cl2, is certainly equally shared. The two chlorine atoms have exactly the same pull on the pair of electrons, so the bond must be exactly equally shared. In cesium fluoride the cesium atom certainly donates an electron and the fluoride atom certainly craves an electron. Both the cesium ion (Cs+) and the fluoride (F-) ion can exist in solution independently of the other. The bond between a cesium and a fluoride ion to make cesium fluoride (CsF) would be clearly ionic because the difference in electronegativities (ΔEN) is so large.
The amount of pull on an atom has on a shared pair of electrons, called electronegativity, is what determines the type of bond between atoms. Considering the Periodic Table without the inert gases, electronegativity is greatest in the upper right of the Periodic Table and lowest at the bottom left. The bond in francium fluoride should be the most ionic. Some texts refer to a bond that is between covalent and ionic called a polar covalent bond. There is a range of bond between purely ionic and purely covalent that depends upon the electronegativity of the atoms around that bond. If there is a large difference in electronegativity, the bond has more ionic character. If the electronegativity of the atoms is more similar, the bond has more covalent character.
Back to the top of BONDS AND STRUCTURES
LEWIS STRUCTURES
Lewis structures are an opportunity to better visualize the valence electrons of elements. In the Lewis model, an element symbol is inside the valence electrons of the s and p subshells of the outer ring. It is not very convenient to show the Lewis structures of the Transition Elements, the Lanthanides, or Actinides. The inert gases are shown having the element symbol inside four groups of two electrons symbolized as dots. Two dots are above the symbol, two below, two on the right, and two on the left. The inert gases have a full shell of valence electrons, so all eight valence electrons appear. Halogens have one of the dots missing. It does not matter on which side of the symbol the dot is missing. Group 1 elements and hydrogen are shown with a single electron in the outer shell. Group 2 elements are shown with two electrons in the outer shell, but those electrons are not on the same side. Group 3 elements have three dots representing electrons, but the electrons are spread around to one per position, as in Group 2 elements. Group 4 elements, carbon, silicon, etc. are shown as having four electrons around the symbol, each in a different position.
Group 5 elements, nitrogen, phosphorus, etc. have five electrons in the outer shell. In only one position are there two electrons. So Group 5 elements such as nitrogen can either accept three electrons to become a triple negative ion or join in a covalent bond with three other items. When all three of the unpaired electrons are involved with a covalent bond, there is yet another pair of electrons in the outside shell of Group 5 elements.
Group 6 elements, oxygen, sulfur, etc., have six electrons around the symbol, again without any concern to position except that there are two electrons in two positions and one electron alone in the other two positions. Group 7 elements have all of the eight outside electrons spaces filled except for one. The Lewis structure of a Group 7 element will have two dots in all four places around the element symbol except for one.
Let's start with two atoms of the same type sharing a pair of electrons. Chlorine atoms have seven electrons each and would be a lot more stable with eight electrons in the outer shell. Single chlorine atoms just do not exist because they get together in pairs to share a pair of electrons. The shared pair of electrons make a bond between the atoms. In Lewis structures, the outside electrons are shown with dots and covalent bonds are shown by bars.
This covalent bond between chlorine is one of the most covalent bonds known. Why? A covalent bond is the sharing of a pair of electrons. The two atoms on ether side of the bond are exactly the same, so the amount of "pull" of each atom on the electrons is the same, and the electrons are shared equally.
Next let's consider a molecule in which the atoms bonded are not the same, but the bonds are balanced. Methane, CH4, is such a molecule. If there were just a carbon and a single hydrogen, the bond between them would not be perfectly covalent. In the CH4 molecule, the four hydrogen atoms exactly balance each other out. The Lewis structure of methane does not have any electrons left over. The carbon began with four electrons and each hydrogen began with two electrons. Only the bars representing the shared pairs of electrons remain. The carbon now shares four pairs of electrons, so this satisfies the carbon's need for eight electrons in the outside shell. Each hydrogen has a single shared pair in the outside shell, but the outside shell of the hydrogen only has two electrons, so the hydrogen has a full outer shell also.
Carbons and hydrogens are nice and easy to write in Lewis structures, because each carbon must have four attachments to it and each hydrogen atom must have one and only one attachment to it. When the bonds around a carbon atom go to four different atoms, the shape of the bonds around that carbon is roughly tetrahedral, depending upon what the materials are around the carbon. Carbons are also able to have more than one bond between the same two. Consider the series ethane (C2H6), ethene (C2H4),(common name is ethylene), and ethyne (C2H2), (common name is acetylene).
H3 – C – C – H3 ethane H2 – C = C – H2 ethylene H – C ≡ C – H acetylene
In writing the Lewis structure of compounds, the bars representing bonds are preferred to the dots representing individual electrons.
The double bars between the carbons in ethylene, C=C, represent a double bond between the two carbons, that is four shared electrons to make a stronger attachment between the two carbons. The triple bars between the carbons of acetylene represent a triple covalent bond between those two carbons, C≡C, three pairs of shared electrons between those carbons. Every carbon has four bonds to it showing a pair of electrons to make eight electrons (or four orbitals) in the outer shell. Each hydrogen atom has one and only one bond to it for two electrons in the outer shell that occupies the only orbital that hydrogen has. All of the outer shells are usually filled.
While we are doing this, notice that the Lewis structure of a molecule will show the shape of the molecule. All of the bonds in ethane are roughly the tetrahedral angle, so all of the hydrogen atoms are equivalent. This is true. The bonds in acetylene make it a linear molecule. The bonds in ethylene are somewhat trigonal around the carbons, and the carbons cannot twist around that bond as they can around a single bond, so that the molecule has a flat shape and the attachments to the carbons are not equivalent. This is also true. (You will see this in the study of organic chemistry. This type of difference between the positions of the hydrogen atoms is called cis - trans isomerism.)
The Lewis structure shows the shape of a molecule or polyatomic ion with the bonds to each atom drawn at 90 degrees (right, left, up, and down) from the atomic symbol and the non – bonded electrons as dots, usually in pairs, around the atomic symbol in the left, right, up, and down positions around the atom. We could set up a group of general guidelines for the drawing of Lewis structures for simple molecules or polyatomic ions.
Write all the atoms in the material in the form of the formula of the compound. CO2 can be an example.
TVE = 16 electrons
Electrons in bonds = - 4 electrons (two bonds)
Dots needed = 12 dots
. . . .
: O – C – O : This is the proposed shape for the CO2 molecule in the skeletal form.
. . . .
. . . .
: O = C = O :
This process of writing Lewis structures is very limited to small molecules. There are many exceptions to the process, for instance, there are some compounds in which one atom has only three orbitals around it. BF3, boron trifluoride is one in which the boron atom (central) is stuck with just three bonds to it. Some central atoms can have MORE than four orbitals around them. There is a phosphorus trichloride molecule (PCl3) that has the same shape as ammonia, but there is also a phosphorus pentachloride molecule (PCl5) that has five chlorine atoms attached to a central phosphorus. As you see, the scope of this tutorial goes only so far into the Lewis structure world.
With the warnings in mind, here are some general rules that will generally help you write Lewis structures.
The Lewis structures are usually good indicators of the actual shape of the molecule. We can tell that from the properties of the molecules. Rarely, but sometimes the best – looking Lewis structure is not the structure that predicts the properties of the material. In this case, the Lewis structure is wrong, and it probably makes some sense once the Lewis structure is written in the way that goes with the properties of the material.
Back to the top of BONDS AND STRUCTURES
SHAPES AROUND AN ATOM, VSEPR THEORY
There is no issue of shape around the Group 1 elements. There is only one attachment to them, so no angle is possible around them. But there are some molecular compounds with only two atoms, such as nitrogen monoxide, NO. The only feature of this molecule is the bond between the nitrogen atom and the oxygen atom. The small difference in electronegativity between the oxygen and the nitrogen give the molecule a small dipole, a small separation of charge, so a small amount of polarity. (Because there are an odd number of electrons in NO, this makes for an interesting Lewis structure. Try it.) Iodine fluoride, IF, is another diatomic compound that should have some polarity. Diatomic molecules like chlorine gas, Cl2, have no electronegativity difference (ΔEN) from side to side of the bond, so they are completely balanced and completely non – polar.
Group 2 elements have two electrons in the outer shell. Many of the compounds of Group 2 elements are ionic compounds, not really making an angle in a molecule. Molecules made with Group 2 elements that have two attached items to the Group 2 element have a linear shape, because the two attached materials will try to move as far from each other as possible. A linear shape means that a straight line could be made through all three atoms with the central element in the center. The shape of carbon dioxide is linear with the carbon in the center. O = C = O
VSEPR stands for Valence Shell Electron Pair Repulsion. The idea is a disarmingly simple one. Electrons are all negatively charged, so they repel each other. If an atom has two electron groups around it, the electrons, and the atoms they are bonded to, are likely to be found as far as they can be from each other. “As far as they can get from each other,” and still remain attached to the central atom means that the angle around the central atom is 180 degrees, a straight line. Molecules with two electron groups attached to a central atom have a linear electron group shape and a linear molecular shape. Unless there is a large difference in electronegativity from one side to the other of a linear compound, there is no separation of charge and no polar character of the molecule.
Covalent compounds with boron are good examples of trigonal shaped molecules. The trigonal shape is a flat molecule with 120 degree angles between the attached atoms. Again using the example of a boron atom in the center, the attached elements move as far away from each other as they can, forming a trigonal shape, also called triangular, or trigonal planar to distinguish it from the trigonal pyramidal shape of compounds like ammonia. BF3, boron trifluoride, is an example of a molecule with a trigonal planar shape. Each fluorine atom is attached to the central boron atom. There are three bonds to the boron, so the electron group shape is trigonal planar around boron. The molecular shape is also trigonal planar in boron trifluoride because each electron group has a fluorine atom attached to it.
But, what if the central atom has two other atoms and a lone pair of electrons attached to it? Nitrogen oxychloride is an example of that. NOCl, is a molecule with nitrogen in the center (See how to write Lewis structures above.) and an oxygen and a chlorine atom attached to the central nitrogen. When we go through the skeleton structure and distribute the electron dots, we find that there is a double bond between the nitrogen and the oxygen and a lone pair (unshared pair) of electrons on the nitrogen in addition to the single bond from the nitrogen to the chlorine. There are three electron groups around the nitrogen, making the electron group shape more or less trigonal planar. But only two of those electron groups have an atom attached, so the molecular shape of nitrogen oxychloride is bent or angular. NOCl is not a balanced shape, so it is likely that there is some separation of charge within the molecule, making it a somewhat polar compound.
Group 4 elements are not in the center of a flat molecule when they have four equivalent attachments to them. As with two or three attachments, the attached items move as far as they can away from each other. In the case of a central atom with four things attached to it, the greatest angle between the attached items does not produce a flat molecule. If you were to cut off the vertical portion of a standard three-legged music stand so that it was the same length as the three legs, the angles among all four directions would be roughly equal. Try this with a gumdrop or a marshmallow. Stick four different colored toothpicks into the center at approximately the same angle. If you have done it right, the general shape of the device will be the same no matter which one of the toothpicks is up. This shape is called tetrahedral. The shape of a tetrahedron appears with the attached atoms at the points of the figure and each triangle among any three of them makes a flat plane. A tetrahedron is a type of regular pyramid with a triangular base. Carbon is a group four element. Organic and biochemical compounds have carbon as a “backbone,” so this tetrahedral shape is very important. Methane, CH4, and carbon tetrachloride, CCl4, are good examples of tetrahedral shape. If you draw the Lewis structures of these compounds, you will see that there are four bonds to the central carbon atom, but no other electrons on the central atom. They have four electron groups (single bonds) around the central atom, so they have a tetrahedral electron group shape. Each bond to the central carbon has an atom attached, so they have a tetrahedral molecular shape. In both compounds, the four atoms attached to carbon are the same, so there is no separation of charge. All four atoms have the same electron pull in balanced directions, so these compounds are non – polar. Can a central carbon make molecules with other shapes around the central atom? Yes, you remember carbon dioxide, where there are two double bonds around the carbon.
O = C = O Each double bond is an electron group, so there are only two electron groups around the carbon in carbon dioxide.
See the “acid carbons,” the ones with the ionizable hydrogen (in blue) on it. The shape of around the acid carbons is trigonal planar because it has a double bond to it and only three electron groups, but the shape around the other carbons is tetrahedral. In the Lewis structures the atoms are drawn at ninety degrees from each other, but the real shape around those carbons exists in three – space.
Group 5 elements, for instance nitrogen or phosphorus, will become triple negative as they add three electrons in ionic reactions, but this is rare. Nitrides and phosphides do not survive in the presence of water. Covalent bonds with these elements do survive in water. From the Lewis structure of these elements in the previous section, you know that Group 5 elements have the capability of joining with three covalent bonds, but they don’t make the trigonal shape because the UNSHARED PAIR OF ELECTRONS ACTS LIKE ANOTHER BONDED ATTACHMENT. The shape of the bonds and the lone pair of electrons around nitrogen and phosphorus is tetrahedral, just like the bonds around Group 4 elements. The molecular shape is trigonal pyramidal.
Group 6 elements, oxygen and sulfur, have six electrons in the valence shell. The compounds they make usually have two pairs of unshared electrons. Just as in Group 5 elements, these pairs of unshared electrons serve as other attached atoms for the electron shape of the molecule. Group 6 elements make tetrahedral electron shapes, but now there are only two attached atoms. The angle between the hydrogens in water is about 105 degrees. This peculiar shape is one of the things that makes water so special.
Group 7 elements have only one chance of attachment, so there is not usually any shape around these atoms.
Back to the top of BONDS AND STRUCTURES
BONDING FORCES IN WATER
The alchemists of old had several other objectives aside from making gold. The thought of a fluid material that could dissolve anything, the universal solvent, was another alchemical project. No alchemist would say, though, what material would hold such a fluid. Surprisingly, the closest thing we have to a universal solvent is water. Water is not only a common material, but the range of materials it dissolves is enormous. The guiding principle for predicting which materials dissolve in which solvent is that 'like dissolves like.' Fluids in which the atoms are attached with covalent bonds will dissolve covalent molecules. Fluids with a separation of charge in the bonds will dissolve ionic materials.
The bonds that hold hydrogen atoms to oxygen atoms are closer to covalent than ionic, but the bond does have a great deal of ionic character. Oxygen atoms are more electronegative than hydrogen atoms, so the electron pair is held closer to the oxygen atom. Another way to look at it is that only a very small number of water molecules are ionized at any one time. The ionization of water, H2O → H+ + (OH)- , into hydrogen ions (actually, hydronium ions) and hydroxide ions happens in only a very small number of the water molecules, but the effect is quite important as the reason for the existence of acids and bases. Materials of a mildly covalent nature, such as small alcohols and sugars, are soluble in water due to the mostly covalent nature of the bonds in water.
The shape of the water molecule is bent at about a 105 degree angle due to the electron structure of oxygen. The two pairs of electrons that force the attached hydrogens into something close to a tetrahedral angle give the water molecule an unbalanced shape like a boomerang, with oxygen at the angle and the hydrogen atoms at the ends. We can think of the molecule has having an ‘oxygen side’ and a ‘hydrogen side’. Since the oxygen atom pulls the electrons closer to it, the oxygen side of the molecule has a slight negative charge. Cations (positive ions) are attracted to the partial positive charge on the oxygen side of water molecules. Likewise, the hydrogen side of the molecule has a slight positive charge, attracting anions. Polar materials such as salts, materials that have a separation of charge, dissolve in water due to the charge separation of water. The origin of the separation is called a dipole moment and the molecule itself can be called a dipole.
Molecules or atoms that have no center of asymmetry are non-polar. Atoms such as the inert gases have no center of asymmetry. Molecules such as methane, CH4, are likewise totally symmetrical. Very small forces, called London forces, can be developed within such materials by the momentary asymmetries of the material and induction forces on neighboring materials. These small forces account for the ability of non-polar particles to become liquids and solids. The larger the atom or molecule, the more potent the London forces, possibly due to the greater ability to separate charge within a larger particle. The larger the inert gas, the higher its melting point and boiling point. In alkanes, a series of non-polar hydrocarbon molecules, the larger the molecule, the higher the melting and boiling point.
There may be London forces in water molecules, but the enormous force of the dipole interaction completely hides the small London forces. The dipole forces within water are particularly strong for two additional reasons. Dipole forces that involve hydrogen atoms around a strongly electronegative material such as nitrogen, oxygen, fluorine, or chlorine are particularly strong due to the small size of the hydrogen atom compared to the size of the dipole force. Such dipoles have significantly stronger forces, and have been called hydrogen bonds. In water, this effect is even greater due to the small size of the oxygen atom, thus the whole water molecule. In a water molecule hydrogen bonding is a large intermolecular force in a small volume on a small mass that makes it particularly noticeable.
Compare methane, CH4, to water. They are similar in size and mass, but methane is non-polar and water is very highly polar due to the hydrogen bonding. The melting point for methane is -184 °C (89 K) and for water is 0 °C (273 K). The boiling point for methane is -161.5 °C (111.7 K) compared to water at 100 °C (373.2 K). The temperature range over which methane is a liquid is less than a quarter the range for water. Most of these differences are accountable from the hydrogen bonding of water.
The properties of water come directly from the molecular shape of it and the forces it has on it from that shape. Water is cohesive. It balls up with itself in zero gravity or on a non – polar surface like waxed paper. The surface tension of water is another product of the cohesive forces, mainly hydrogen bonding. Water is adhesive, that is, it clings to other things. It wets cotton or paper, it wets glass or ceramic, and it dissolves many compounds, to include polar compounds.
Water is a very important material for living things because:
There are three main types of bonding forces, forces that make compounds. Ionic bonding is just the attraction of a positive ion for a negative ion. Sodium chloride is a compound that is made of sodium ions, having lost an electron, with a positive charge, and negative chloride ions, negative because they attract another electron to fill the valence shell. Covalent bonds come about by a bonded pair of atoms sharing a pair (or more pairs) of electrons. Covalent bonds are usually stronger than ionic bonds. Ionic bonds can separate in water solution. Polar covalent bonds, such as the bonds between the hydrogen and oxygen atoms of water, happen when two atoms sharing a pair of electrons have a large difference in electronegativity.
Three main types of intermolecular forces, hydrogen bonding, dipole interactions, and dispersion forces, are forces that do not make compounds, but attract or repel on an atomic level. The name London forces (from Fritz London) is sometimes used for the small dipole interactions and even smaller dispersion forces. Dispersion forces are caused by the momentary unbalance of electrons around an atom. They are called “dispersion” forces for the uneven dispersion of electrons. Even noble gases can have these forces. In fact dispersion forces are the only forces that pull noble gases together. In atoms or small molecules, dispersion forces are very small. The melting and boiling points of noble gases are very low because it takes very little energy to overcome the dispersion forces. In macromolecules, though, the dispersion forces can develop to be much larger. In proteins and nucleic acids, dispersion forces rival the magnitude of the dipole forces and even hydrogen bonding.
Dipole forces, or dipole – dipole interactions are the forces from polar molecules pulling together by the difference in charge from one side of a molecule to another. Iodine fluoride, IF, is likely to have a small positive charge near the iodine and a small negative charge near the fluorine, because fluorine is by far the most electronegative. The IF molecules have a tendency to arrange themselves with the positive end of one molecule near the negative end of another molecule. The dipole forces of water are fairly large due to the highly polar nature of the water molecule.
In water, the most powerful intermolecular force is hydrogen bonding. Hydrogen bonding is the tendency of hydrogen atoms attached to highly electronegative atoms like fluorine, chlorine, or oxygen to seek other highly electronegative atoms in other molecules. The forces can make liquids viscous and cohesive. Water owes its cohesive properties mostly to hydrogen bonding. But hydrogen bonding is even more important in macromolecules. The secondary, tertiary, and quaternary structures of macromolecules are due in large part to hydrogen bonding. The association of opposing nucleotides in nucleic acids is due to hydrogen bonding.
|
Copyright © 1997-2009 Chemtutor, LLC. All rights reserved. |
Publication of any part is prohibited without prior written consent, except for instructor use publishing for instructor's own classroom students. All hard copy materials distributed under this exception must have on every page distributed reference to http://www.chemtutor.com as source. Under the same exception granted to classroom teachers, full recognition of Chemtutor must be given when all or any part is included in any other electronic representation, such as a web site, whether by direct inclusion or by hyperlink. |
