COMPOUNDS

A bond is an attachment among atoms. Atoms may be held together for any of several reasons, but all bonds have to do with the electrons, particularly the outside electrons, of atoms. There are bonds that occur due to sharing electrons. There are bonds that occur due to a full electrical charge difference attraction. There are bonds that come about from partial charges or the position or shape of electrons about an atom. But all bonds have to do with electrons. Since chemistry is the study of elements, compounds, and how they change, it might be said that chemistry is the study of electrons. If we study the changes brought about by moving protons or neutrons, we would be studying nuclear physics. In chemical reactions the elements do not change from one element to another, but are only rearranged in their attachments.

A compound is a group of atoms with an exact number and type of atoms in it arranged in a specific way. Every bit of that material is exactly the same. Exactly the same elements in exactly the same proportions are in every bit of the compound. Water is an example of a compound. One oxygen atom and two hydrogen atoms make up water. Each hydrogen atom is attached to an oxygen atom by a bond. Any other arrangement is not water. If any other elements are attached, it is not water. H2O is the formula for that compound. This formula indicates that there are two hydrogen atoms and one oxygen atom in the compound. H2S is hydrogen sulfide. Hydrogen sulfide does not have the same types of atoms as water. It is a different compound. H2O2 is the formula for hydrogen peroxide. It might have the right elements in it to be water, but it does not have them in the right proportion. It is still not water. The word formula is also used to mean the smallest bit of any compound. A molecule is a single formula of a compound joined by covalent bonds. The Law of Constant Proportions states that a given compound always contains the same proportion by weight of the same elements.

Some atoms, such as metals tend to lose electrons to make the outside ring or rings of electrons more stable and other atoms tend to gain electrons to complete the outside ring. An ion is a charged particle. Electrons are negative. The negative charge of the electrons can be offset by the positive charge of the protons, but the number of protons does not change in a chemical reaction. When an atom loses electrons it becomes a positive ion because the number of protons exceeds the number of electrons. Non-metal ions and most of the polyatomic ions have a negative charge. The non-metal ions tend to gain electrons to fill out the outer shell. When the number of electrons exceeds the number of protons, the ion is negative. The attraction between a positive ion and a negative ion is an ionic bond. Any positive ion will bond with any negative ion. They are not fussy. An ionic compound is a group of atoms attached by an ionic bond that is a major unifying portion of the compound. A positive ion, whether it is a single atom or a group of atoms all with the same charge, is called a cation, pronounced as if a cat were an ion. A negative ion is called an anion, pronounced as if Ann were an ion. The name of an ionic compound is the name of the positive ion (cation) first and the negative (anion) ion second.

The valence of an atom is the likely charge it will take on as an ion. The names of the ions of metal elements with only one valence, such as the Group 1 or Group 2 elements, is the same as the name of the element. The names of the ions of nonmetal elements (anions) develop an -ide on the end of the name of the element. For instance, fluorine ion is fluoride, oxygen ion is oxide, and iodine ion is iodide. There are a number of elements, usually transition elements that having more than one valence, that have a name for each ion, for instance ferric ion is an iron ion with a positive three charge. Ferrous ion is an iron ion with a charge of plus two. There are a number of common groups of atoms that have a charge for the whole group. Such a group is called a polyatomic ion or radical. Chemtutor suggests it is best to learn by rote the list of polyatomic ions with their names, formulas and charges. Chemtutor provides a Quickquiz on the common ions and a quiz on reading and writing ionic compounds.

NOTE THERE ARE TWO COMMON NAMES FOR THE IONS. YOU SHOULD KNOW BOTH THE STOCK SYSTEM AND THE OLD SYSTEM NAMES.

The ions by the Stock system are pronounced, ?copper one?, ?copper two?, etc. Notice that the two most likely ions of an atom that has multiple valences have suffixes in the old system to identify them. The smallest of the two charges gets the ?-ous? suffix, and the largest of the two charges has the ?-ic? suffix. This leads to the amusing possibility of Saint Nickelous coming down your chimney. (Boo! Hiss!)

The following radicals or polyatomic ions are groups of atoms of more than one kind of element attached by covalent bonds. They do not often come apart in ionic reactions. The charge on the radical is for the whole group of atoms as a unit. These are common radicals you should learn WITH THEIR CHARGE AND NAME.

These are written here with the parentheses around the polyatomic ions to show their origin. Usually these compounds are written without the parentheses, such as HNO3 or H2SO4. Note that the polyatomic ions with a single negative charge only have one hydrogen. Polyatomic ions with two negative charges have two hydrogens.

In the lists above, the radicals and compounds have a small number after and below an element if there is more than one of that type of that atom. For instance, ammonium ions have one nitrogen atom and four hydrogen atoms in them. Sulfuric acid has two hydrogens, one sulfur, and four oxygens.

Knowing the ions is the best way to identify ionic compounds and to predict how materials would join. People who do not know of the ammonium ion and the nitrate ion would have a difficult time seeing that NH4NO3 is ammonium nitrate. Chemtutor very highly recommends that you know all the above ions, complete with the valence or charge.

Let?s consider what happens in an ionic bond using electron configuration, the octet rule, and some creative visualization. A sodium atom has eleven electrons around it. The first shell has two electrons in an s subshell. The second shell is also full with eight electrons in an s and a p subshell. The outer shell has one lonely electron, as do the other elements in Group 1. This outside electron can be detached from the sodium atom, leaving a sodium ion with a single positive charge and an electron. A chlorine atom has seventeen electrons. Two are in the first shell, eight are in the second shell, and seven are in the outside shell. The outside shell is lacking one electron to make a full shell, as are all the elements of Group 7. When the chlorine atom collects another electron, the atom becomes a negative ion. The positive sodium ion missing an electron is attracted to the negative chloride ion with an extra electron. The symbol for a single unattached electron is e^-.

Any compound should have a net zero charge. The single positive charge of the sodium ion cancels the single negative charge of the chloride ion. The same idea would be for an ionic compound made of ions of plus and minus two or plus and minus three, such as magnesium sulfate or aluminum phosphate

But what happens if the amount of charge does not match? Aluminum bromide has a cation that is triple positive and an anion that is single negative. The compound must be written with one aluminum and three bromide ions. AlBr3. Calcium phosphate has a double positive cation and a triple negative anion. If you like to think of it this way, the number of the charges must be switched to the other ion. Ca3(PO4)2. Note that there must be two phosphates in each calcium phosphate, so the parentheses must be included in the formula to indicate that. Each calcium phosphate formula (Ionic compounds do not make molecules.) has three calcium atoms, two phosphate atoms, and eight oxygen atoms.

There are a small number of ionic compounds that do not fit into the system for one reason or other. A good example of this is magnetite, an ore of iron, Fe3O4. The calculated charge on each iron atom would be +8/3, not a likely actual charge. The deviance from the system in the case of magnetite could be accounted for by a mixture of the common ferric and ferrous ions.

The word binary means that there are two types of atom in a compound. Covalent compounds are groups of atoms joined by covalent bonds. Binary covalent compounds are some of the very smallest compounds attached by covalent bonds. A covalent bond is the result of the sharing of a pair of electrons between two atoms. The chlorine molecule is a good example of the bond, even if it has only one type of atom. Chlorine gas, Cl2, has two chlorine atoms, each of which has seven electrons in the outside ring. Each atom contributes an electron to an electron pair that make the covalent bond. Each atom shares the pair of electrons. In the case of chlorine gas, the two elements in the bond have exactly the same pull on the electron pair, so the electrons are exactly evenly shared. The covalent bond can be represented by a pair of dots between the atoms, Cl:Cl, or a line between them, Cl-Cl. Sharing the pair of electrons makes each chlorine atom feel as if it has a completed outer shell of eight electrons. The covalent bond is much harder to break than an ionic bond. The ionic bonds of soluble ionic compounds come apart in water, but covalent bonds do not usually come apart in water. Covalent bonds make real molecules, groups of atoms that are genuinely attached to each other. Binary covalent compounds have two types of atom in them, usually non-metal atoms. Covalent bonds can come in double (sharing of two pairs of electrons) and triple (three pairs of electrons) bonds.

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

With the compounds of nitrogen and oxygen to use as examples, we see that there are often more ways for any two elements to combine with each other by covalent bonds than by ionic bonds. Many of the frequently seen compounds already have names that have been in use for a long time. These names, called common names, may or may not have anything to do with the makeup of the material, but more of the common names of covalent compounds are used than of the ionic compounds.

The system names include numbers that indicate how many of each type of atom are in a covalent molecule. The Fake Greek Prefixes (FGP?s above in the chart) are used to indicate the number. It would be wise of you to know the FGP?s.

In saying or writing the name of a binary covalent the FGP of the first element is said, then the name of the first element is said, then the FGP of the second element is said, and the name of the second element is said, usually with the ending ?-ide? on it. The only notable exception for the rule is if the first mentioned element only has one atom in the molecule, in which case the ?mono-? prefix is omitted. CO is carbon monoxide. CO2 is carbon dioxide. In both cases there is only one carbon in the molecule, and the ?mono-? prefix is not mentioned. For oxygen the last vowel of the FGP is omitted, as in the oxides of nitrogen in the above table.

Here?s a checklist of the things you need to know to be able to correctly write the formulas for materials.

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 and the fluoride ion can exist independently of the other. The bond between a cesium and a fluoride ion is clearly ionic.

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 Chart without the inert gases, electronegativity is greatest in the upper right of the Periodic Chart 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.

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 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.

There is a more extensive and better illustrated section on the Lewis structures of elements in the atomic structure chapter in Chemtutor. In this section there will be emphasis on the Lewis structure of small compounds and polyatomic ions.

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 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).

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, three pairs of shared electrons between those carbons. Every carbon has four bonds to it showing a pair of electrons to make eight electrons in the outer shell. Each hydrogen has one and only one bond to it for two electrons in the outer shell. All of the outer shells are 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 hydrogens 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 can not twist around that bond as they can around a single bond, so that the molecule has a flat shape and the hydrogens 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 hydrogens is called cis - trans isomerism.)

We could set up a group of general guidelines for the drawing of Lewis structures for more complex molecules or polyatomic ions.

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. 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 Group 2 element in the center.

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.

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.

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 around nitrogen and phosphorus is tetrahedral, just like the bonds around Group 4 elements.

Group 6 elements, oxygen and sulfur, have two pairs of unshared electrons. Just as in Group 5 elements, these two pairs of unshared electrons serve as another attached atom for the shape of the molecule. Group 6 elements make tetrahedral molecules also, but now the items making the points of the tetrahedron are now limited to two. 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 no shape around these atoms.

The alchemists of old had several other objectives aside from making gold. The though 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 there are only a very small number of water molecules ionized. The ionization of water, H2O

H⁺ + (OH)^-, into hydrogen 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. More about water later.

Write chemical formula as requested. Show subscript numbers where needed. Show valences for all ions.

1. hydrochloric acid _________________ 2. sodium chloride ________________ 1. HCl 2. NaCl

3. sodium hexafluoride _____________ 4. strontium nitrate ________________ 3. NaF6 4. Sr(NO3)2

5. calcium chloride _________________ 6. acetic acid ___________________ 5.CaCl2 6.HC2H3O2

7. phosphoric acid __________________ 8. ammonia ______________________ 7. H3PO4 8. NH3

9. chlorine ______________________ 10. lithium sulfate ___________________ 9. Cl2 10. Li2SO4

11. potassium chromate ____________ 12. calcium hydroxide ____________ 11.K2CrO4 12.Ca(OH)2

13. aluminum foil _________________ 14. ammonium sulfate ______________ 13.Al 14.(NH4)2SO4

15. sulfuric acid __________________ 16. ammonium iodide ______________ 15. H2SO4 16. NH4I

17. acetylene _____________________ 18. rubidium nitrite _______________ 17. C2H2 18. RbNO2

19. lead II sulfite __________________ 20. copper I sulfide ________________ 19. PbSO3 20.Cu2S

21. aluminum oxide _______________ 22. magnesium bromide _____________ 21.Al2O3 22.MgBr2

23. sodium chlorate ________________ 24. iron II chloride ________________ 23.NaClO3 24.FeCl2

25. hydrogen gas __________________ 26. silver chromate ________________ 25. H2 26. Ag2CrO4

27. zinc bicarbonate _______________ 28. barium oxide ________________ 27.Zn(HCO3)2 28.BaO

29. aluminum nitrate ______________ 30. diphosphorus pentoxide __________ 29.Al(NO3)3 30.P2O5

31. aluminum hydroxide ___________ 32. chromium III oxide _____________ 31.Al(OH)3 32.Cr2O3

33. lithium phosphate ________________ 34. ice ________________________ 33. Li3PO4 34. H2O

35. nitrogen dioxide _________________ 36. iron III oxide _________________ 35. NO2 36. Fe2O3

37. sodium peroxide ________________ 38. copper II oxide ________________ 37.Na2O3 38.CuO2

39. liquid nitrogen _______________ 40. lead II acetate _________________ 39.N2 40.Pb(C2H3O2)2

41. lead IV fluoride ________________ 42. ferrous bromide ________________ 41. PbF4 42. FeBr2

43. carbonic acid _________________ 44. silver bisulfite ________________ 43.H2CO3 44.AgHSO3

45. cupric hydroxide ________________ 46. nitric acid __________________ 45.Cu(OH)2 46.HNO3

47. mercury II bromide _______________ 48. stannic sulfide ________________ 47. HgBr2 48. SnS2

49. hydrofluoric acid _______________ 50. potassium phosphate _____________ 49. HF 50. K3PO4

51. iodine tribromide _______________ 52. phosphorus pentafluoride __________ 51. IBr3 52. PF5

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

*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