1a. Lewis Dot
The Basics (review)
Lewis Dot Diagrams, sometimes referred to as “Electron-Dot Notation,” give you the ability to figure out the types of covalent bonds that an element may make in certain situations. Lewis Dot Diagrams can also be used to predict the type of ion that an atom might make (when it forms an ion). Each dot diagram of a single element consists of: an elemental symbol, which goes in the center of the dots, and a group of 1-8 dots which shows the configuration of the outer-most electron shell of the atom, which we also call the “valence shell.”

Below shows the order of filling in the dots on LDD’s for an element that has eight valence electrons. The first two dots can be put on any side, but the rest of the dots should be placed in either a clockwise or counter clockwise manner, with no side (other than the first) receiving two dots until each side gets one.
Note: When working with a specific element, that element’s symbol will replace the “X” used here.


For a more specific example of basic lewis dot diagrams (still a review), see below. Chlorine has 7 valence electrons.

external image LewisCl.gif
Molecules
Electron-dot notation is not only used for elements, but can also be used with molecules. Luckily, this is not as difficult as it sounds. Lewis dot structures for molecules are made by combining the structures for the individual elements.
Key terms to remember: Long pair (also called an unshared pair): a pair of electrons that is NOT involved in bonding and that belongs only to one atom. Single bonds: as you already know, a single bond is a covalent bond produced by the sharing of one pair of electrons between two atoms. In Lewis Dot Structures, this is represented by just two dots in between the “shells.”
See below for an example. Double bonds are represented in lewis dot structures with a double pair of dots (four dots).

The Octet Rule: see below in the “octet rule section.” This is important to keep in mind with Lewis Dot structures.

Lewis Structures specifically for Ionic Compounds
The overall charge on the compound must be equal to zero. In other words, the number of electrons that are lost by one atom has to be equal to the number of electrons gained by the other atom.

The Lewis Dot Structure (also known as the electron dot diagram) of each ion is then used to construct the Lewis Structure for the ionic compound.

Example #1
Lithium fluoride (LiF)

  • Lithium atom loses one electron to form the cation Li+
  • Fluorine atom gains one electron to form the anion F-
  • The Lithium fluoride compound can be represented as either:
Li+ OR

Example #2
Lithium oxide (Li2O)
  • Each lithium atom loses one electron to form 2 cations Li+ (2 electrons in total are lost)
  • Oxygen atom gains two electrons to form the anion O2-
  • Lithium oxide compound can be represented as
2Li+ OR Li+ Li+ OR


1b. Ions

An ion is an atom or group of bonded atoms that has a positive or negative charge. Whether the atom has a positive or negative charge depends on the differing numbers of protons and electrons in the atom. If there are more protons than electrons in the atom, it is called a cation. A cation is created when an atom loses electrons. Conversely, if there are more electrons than protons in the atom, it is called an anion. An anion is created when an atom gains electrons.
To see some examples of the creation of cations and anions, check out the Electron Transfer section on this page.


1c. Formation of Ions

a. Formation of monatomic ions


A monatomic ion is one that contains, as the name implies, only one ion (for example, Na+, but not NH4+). They are formed by the addition of electrons to the valence shell of the atom, or the losing of electrons from this shell. Ionization is the process wherein a neutral atom or molecule acquires or loses electrons. Though atoms can go undergo ionization through radiation, the the transfer of electrons, effected by the attaining of closed shell electronic configurations, is much more common in chemistry.

Example:
A sodium atom, Na, has a single electron in its valence shell, surrounding a closed inner shell of 10 electrons. Since the 10-electron configuration is very stable, Na wants to lose its extra electron so it can achieve this stable configuration, becoming the sodium cation (remember, cation = +) in the process:
Na → Na+ + e−
Conversely, a chlorine atom, Cl, has 7 electrons in its valence shell, which is one short of the filled shell with 8 electrons. Thus, chlorine wants to gain an extra electron in order to attain to the stable 8-electron configuration, becoming the chloride anion in the process:
Cl + e− → Cl−
This driving force is what causes sodium and chlorine to undergo a chemical reaction, where the extra electron is transferred from sodium to chlorine, forming sodium cations and chloride anions. Being oppositely charged, these cations and anions combine together to form sodium chloride, NaCl.

If you need more help, check out this animation.

b. Formation of polyatomic and molecular ions

Polyatomic and molecular ions are often formed by the gaining or losing of elemental ions such as H+ in neutral molecules. For example, when ammonia, NH3, accepts a proton, H+, it forms the ammonium ion, NH4+. Ammonia and ammonium have the same number of electrons in essentially the same electronic configuration, but ammonium has an extra proton that gives it a net positive charge.
Ammonia can also lose an electron to gain a positive charge, forming the ion NH3·+. However, this ion is unstable, because it has an incomplete valence shell around the nitrogen atom, making it a very reactive radical ion.
Due to the instability of radical ions, polyatomic and molecular ions are usually formed by gaining or losing elemental ions such as H+, rather than gaining or losing electrons. This allows the molecule to preserve its stable electronic configuration while acquiring an electrical charge.

c. Formation of positive ions


Metals usually have 1-4 electrons in the outer energy level. The electron arrangement of a rare gas is most easily achieved by losing the few electrons in the newly started energy level. The number of electrons lost must bring the electron number "down to" that of a prior rare gas.

Let's take a look at this example.

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What is the octet rule?

It is an elementary principle which declares that atoms combine in such a way that they each have eight electrons in their valence shells, thus making their electronic configuration the same as a noble gas.

How will sodium complete its octet?

Our first task is to look at the atom's electronic arrangement. The atomic number is eleven, which means there are eleven electrons and eleven protons. We also see that Na only possesses a single electron in its outer level. Since we want only eight electrons in the outer level, the single electron is dropped and its level ceases to exist. We now have a new outer level with eight electrons, thus satisfying the octet rule.

What is the ionic charge?

What is the charge on sodium ion as a result of losing one electron? Let's compare the atom and the ion.

Sodium Atom

Sodium Ion

11 p+
to revert to
11 p +
Protons are identical in the atom and ion. Positive charge is caused by lack of electrons.
12 n
an octet
12 n
11 e-
lose 1 electron
10 e-
0 charge

+ 1 charge

What are the basic principles of ionic compounds?

An ionic compound is born out of the total transfer of electrons from a metal to a non-metal and the subsequent fulfilling of the octet rule.

1d. Electron Configuration of Ions and Octet Rule

Electron configuration is the arrangements of an atom and how they are distributed throughout the different orbitals.

Physicists and chemists use a standard notation to describe the electron configurations of atoms and molecules. For atoms, the notation consists of a string of atomic orbital labels (eg, 1s, 2p, 3d, 4f) with the number of electrons assigned to each orbital (or set of orbitals sharing the same label) placed as a superscript. For example, hydrogen has one electron in the s-orbital of the first shell, so its configuration is written 1s. Lithium has two electrons in the 1s-subshell and one in the (higher-energy) 2s-subshell, so its configuration is written 1s2 2s1. Phosphorus (atomic number 15), is as follows: 1s2 2s2 2p6 3s2 3p3.

We use the Aufbau Principle to help predict the electron configuration of an atom. The Aufbau Principle states the order in which the different orbital levels will be filled.
The figure below shows the relationship between the periodic table and the orbitals being filled during the aufbau process. The two columns on the left side of the periodic table correspond to the filling of an s orbital. The next 10 columns include elements in which the five orbitals in a d subshell are filled. The six columns on the right represent the filling of the three orbitals in a pf subshell subshell. Finally, the 14 columns at the bottom of the table correspond to the filling of the seven orbitals in an .
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Electrons fill orbitals according to rules first stated by Friedrich Hund. Hund's rules can be summarized as follows.
  • One electron is added to each of the orbitals in a subshell before two electrons are added to any orbital in the subshell.
  • Electrons are added to a subshell with the same value of the spin quantum number until each orbital in the subshell has at least one electron.
The outer shell of an atom is called the valence shell. Ions are atoms that do not have a full valence shell and when they combine to form molecules they combine in ways so that they have 8 electrons in their valence shell. This is called the octet rule.

Below is a carbon dioxide molecule. Since there are 8 electrons surrounding each atom, CO2 is a stable atom.
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There are some exceptions to the octet rule, for example, hydrogen follows the duet rule. Since hydrogen is so small, it only needs two electrons to fill its valence shell.

For more information on electron configuration and the octet rule watch these videos.





2a. Electron Transfer

Electron transfer is the process by which an electron moves from one atom or molecule to another atom or molecule

Example #1:
A lithium atom (Li) has one electron in its valence shell which surrounds a stable, closed inner shell of 10 electrons. Since the 10-electron configuration is very stable, lithium wants to lose its extra electron so that it can attain its stable configuration, becoming a lithium cation in the process: Li → Li+ + e−

Example #2
A Fluorine atom (F) has 7 electrons in its valence shell. Because of this, it wants to gain an electron in order to attain the stable 8-electron configuration, becoming a fluoride anion in the process: F + e− → F−

If you are confused or would like more examples, watch this video:




Making Compounds (Main Group Elements)

Before we begin discussing making compounds for ionic bonding, we should first define ionic bonding. Ionic bonding is the chemical bonding that result from the electrical attraction between positively-charged cations and negatively-charged anions. As seen by the diagram below, an outer valence electron has transferred from the first atom to the second atom, resulting in the formation of a cation in the first molecule, and an anion in the second molecule.

external image ionic_bonding.gif

Now the atom on the left has a full outer shell, making it stable. The atom on the right has a complete octet (which was defined in earlier sections). The two opposite charges strongly attract each other, producing an ionic bond. The strength of this bond can be observed by measuring the melting point, and noting that compounds with ionic bonds have much higher melting points than do covalent bonds, because of the fact that the ionic bonds are stronger.

Compounds that contain ionic bonds are produced when cations and anions are brought on contact with each other. A simple example would be the synthesis of Sodium Chloride (NaCl) from Sodium metal and Chlorine gas. This violent reaction is described in the below scheme:
external image Formation-Sodium-Chloride.gif
The sodium electron, represented as an “X”, is removed from the outer shell and transferred to the Chlorine atom outer shell, resulting in a cation and an anion, which then form an ionic bond.

Similarly, ionic bonds can form between multivalent atoms as well. For example, the reaction of Magnesium metal with Chlorine gas results in the transfer of two electrons from magnesium to two different Chlorine atoms.
external image Magnesium-Chloride-Formation.gif

In this example, Magnesium has a full outer shell, and both Chlorines have satisfied the octet rule.


In contrast to Magnesium Chloride, which has a +2 charged cation and a -1 anion, Sodium Sulfate represents a pair where there is a +1 cation and a -2 anion, as seen below:

external image Sodium_sulfate.png
For more help on making compounds with ionic bonds for main group elements, please see the video on the link below:

Ionic Compounds


Finding the Charges of Transition Metals


Finding the charges of transition metals can be tricky, because transition metals don’t have a set charge. The transition metal charge can be found from the chemical formula if it is accompanied by a Roman numeral. For example, Lead (II) is the metal Lead with a charge of +2. Similarly, Iron (III) is the metal Iron with a charge of +3. Once you understand this, finding the charges of transition metals is much easier.

If this information is not given, you will be able to find the charge of the transition metal by the following procedure. In these cases, you find the charges of everything else that metal is bonded to and subtract this number from the total charge of the compound. The most important thing to remember is that the sum of the charges must add up to zero in order to form a neutral compound. Here are some examples:

(1) CuSO4

Sulfate has a charge of -2, and since there is only one Sulfate in the compound, the overall charge is -2. To balance this out, the Copper has to have a charge of +2 for the compound to be neutral because (-2) + (+2) = 0
Therefore, this is Copper (II) Sulfate.

(2) Fe2O3

Each oxygen atom has a charge of -2, but there are three oxygen atoms in this compound. So, three oxygen would add up to (-2) + (-2) + (-2) = -6. The charge of all of the Iron atoms is +6, but to find the charge of the Fe, you must divide by two, because there are two Iron atoms present. Therefore, (+6) / 2 = +3. So you can now figure out that this is Iron (III) Oxide.

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This graph shows the possible charges, or oxidation numbers, of some transition metals. For example, the possible charges for Nickel are +4, +3, +2, and +1.

*Note that all transition metals have a positive charge.

2d. Making Compounds-transition metals

Before we discuss transition metal compounds, we should make sure we remember what ionic bonding means. As mentioned above, ionic bonding is defined as the chemical bonding that results from the electrical attraction between positively-charged cations and negatively-charged anions. In one sentence, transition metals form compounds by losing electrons to form positive ions.
  • Transition metals are the elements found in group D of the period table and include elements such as copper (Cu), iron (Fe), and platinum (Pr). Transition metals serve as a bridge between the two sides of the periodic table.
  • Transition metals can also be defined as elements whose atoms have an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell. Transition metals serve as a bridge between the two sides of the periodic table.
  • Transition metals are commonly made to combine with other elements and to form catalysts.
  • Catalysts are chemical substances that can either increase or decrease the speed of a reaction.
The other elements combined with transition metals to form compounds are usually nonmetals. Some common examples are oxygen (O), carbon (C), and nitrogen (N).
Transition metals/compounds are similar to but not the same as those described in the “Making Compounds (Main Group Elements)” section above.
The fact that the two best conductors of electricity are a transition metal (copper) and a main group metal (aluminum) shows the extent to which the physical properties of main group metals and transition metals overlap.
However, there are major differences between these metals. The transition metals are more electronegative than the main group metals, for example, and are therefore more likely to form covalent compounds.

Another difference between the main group metals and transition metals can be seen in the formulas of the compounds they form.
As seen above in the general explanation of making compounds, transition metals form compounds similarly to main group metals. However, transition metals are much more likely to form “complexes,” which have an excess of negative ions.

CrCl3(s) + 6 NH3(l) -----> CrCl3 6 NH3(s)

The example above shows the behavior of the formation of the CrCl3 6 NH2(s) compound. This is significant because main group elements form aqueous solutions, and they also form stable compounds with neutral molecules with less ease than transition elements.



2e. Oxidation States

An oxidation state is what indicates the degree of oxidation of an atom in a chemical compound. The formal oxidation state is the hypothetical charge that an atom would have if all of the bonds to atoms of different elements were completely (100%) ionic.
Basically, an oxidation number tells how many electrons an atom has shared, lost, or gained in order to become stable.
Positive, negative, or “0” integers are what typically represent oxidation states.
Oxidation is the rise (increase) in the oxidation state of a given atom through some chemical reaction.

Further, an oxidation state is defined as the charge an atom might have when electrons are counted according to a certain set of rules listed as follows:
1) The oxidation state of a free element (uncombined element) is zero
2) For a simple (also known as monatomic) ion, the oxidation state is the same as the net charge on that ion
3) (This just a fact): Hydrogen has an oxidation state of 1 and oxygen has an oxidation state of -2 when present in most compounds. (Some exceptions to this are that hydrogen has an oxidation state of -1 in hydrides of active metals, e.g. LiH, and oxygen has an oxidation state of -1 in peroxides, e.g. H2O2)
4) The sum of the oxidation states of all of the atoms in a neutral molecule has to be zero, while in ions the sum of the oxidation states of the “constituent” atoms must be equal to the ion’s charge

When there is a Lewis Structure (see above section for basic information) of a molecule available, the oxidation states may be found simply by computing the difference between the number of valence electrons that a neutral atom of that element would have and the number of electrons that "belong" to it in the Lewis structure.
Let’s use acetic acid as an example.

Without a Lewis structure:
As previously stated, the algebraic sum of the oxidation states of all atoms in a neutral molecule has to be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. This fact, combined with the fact that some elements almost always have certain oxidation states, allows one to compute the oxidation states for atoms in simple compounds. Some typical rules that are used for assigning oxidation states of simple compounds follow (this is just for a general intro to oxidation states):

1. Fluorine has an oxidation state of -1 in all of its possible compounds because it’s electronegativity is higher than that of all other reactive elements.
2. Hydrogen usually has an oxidation state of +1 except when it is bonded to more electropositive elements (such as sodium, boron, and aluminum, as in NaH, LiAlH4, NaBH4, where each H has an oxidation state of -1).
3. Oxygen usually has an oxidation state of -2.
4. The Alkali metals have an oxidation state of +1 in almost all of their compounds.
5. The Alkaline earth metals have an oxidation state of +2 in basically all of their compounds.
6. Halogens other than fluorine have an oxidation state of -1 except when bonded to oxygen, nitrogen or with another halogen.

However, these general rules actually appear to over-complicate things: again, in ions the oxidation state calculation is fairly simple: the algebraic sum of the oxidation states of the constituent atoms MUST be equal to the charge on the ion.

One example is H2S.
H2S has 2 hydrogens. H has a +1 (the first element in an ionic compound always has the positive charge), so S is -2 (remember to make sure everything is equal).



3a. Naming of Compounds Without a Transition Metal

Ionic compounds, in which ion are held together by ionic bonds in a lattice structure (see below), are formed when a metal gives up its electrons to a non-metal. In layman's terms, it is ionic if the compound contains a metal. However, there are different sets of rules for transition metals. A transition metal is an element with an atomic number of 21 to 30, 39 to 48 or 57 to 80, or found in the middle region of the periodic table.

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a. For a compound with any other metal, use these rules:

    • The metal ion's name does not change regardless of charge
    • The non-metal's name ends in ide.

For example:
AlCl3 = aluminum chloride
Na2S = sodium sulfide
K2O = potassium oxide
MgH2 = magnesium hydride

Remember, the subscripts play no part in naming the compounds.

b. Going the opposite way (from name to formula) is a bit harder. In this case, the total charge of the (+) and (-) ions (called cations and anions, respectively) must equal zero.

Example: What is the correct formula for calcium phosphide?

Here, we have to consider the common charges for calcium and phosphorus, which are +2 and -3, respectively. Calcium, usually loses two electrons to get a noble gas electron arrangement, and phosphorus needs three more electrons.
Ca+2 and P-3
One of each would create a sum of 2 + (-3) = -1. To get a sum of zero, we need three Ca+2 ions and two P-3 for a total of 3(2)+ 2(-3) = 0.

Thus, the answer is Ca3P2.

c. Polyatomic Ions

When metals are bonded to polyatomic ions, which consist of two or more atoms with one overall charge, the same rules apply, but it is absolutely imperative that you learn all the polyatomic ions. (You will fail several tests if you don't.) I recommend using some sort of mnemonic device or code to remember these.

Some examples:

Polyatomic Ion
Name
SO4-2
sulfate
PO4-3
phosphate
NO3-1
nitrate
CO-2
carbonate
HCO3-1
hydrogen carbonate/bicarbonate
ClO3-1
chlorate
NH4+1
ammonium

external image ammonium2a.jpg
The structure of ammonium.


To go the opposite way (from name to formula):

Aluminum sulfate
This has Al+3 and SO4-2. To get a total charge of zero, we need two aluminum ions and three sulfates, so the formula becomes Al2(SO4)3. Remember, when there is more than one polyatomic group, we make use of parenthesis.


If you are still puzzled, watch this video:




3b. Naming of Compounds With a Transition Metal


This is the exact same procedure as the above, but remember we have to specify the transition metal ion's charge by using a Roman numeral. If you don't know the basic Roman numerals (which would be very sad), you should probably learn them.

Roman Numeral
Charge
I
+1
II
+2
III
+3
IV
+4
V
+5
VI
+6

The use of Roman numerals in ionic compounds is to indicate which charge to use, since transition metals often can host more than one charge.

Example: Manganese (II) oxide contains Mn+2 and O-2. So we just need one of each and the formula becomes MnO.


To find the proper Roman numeral, we must discern the charge of the transition metal in a particular situation.

Example: What is the correct name of CrCl3?

We do not know the charge of Cr, so it becomes x in our equation, but we do know that the charge of Cl is (-1). The sum of the charges has top be zero, therefore (we are going to do some basic algebra here):

x +3(-1) = 0.
x = 3.
Answer: CrCl3 = chromium (III) chloride.


Polyatomic Ions

Polyatomic ions are charged ions made up of two or more atoms covalently bonded that act as a single unit. Polyatomic ions can either be positively charged or negatively charged. When a polyatomic ion is positively charged it is called a cation, and when it is negatively charged it is called an anion.

Below is a list of polyatomic ions you should know.


Arsenate
AsO4 ^-3
Borate
BO3 ^-3
Tetraborate
B4O7 ^-2
Bromate
BrO3 ^-1
Carbonate
CO3 ^-2
Cyanide
CN^-1
Oxalate
C2O4^-2
Acetate
C2H3O2^-1
Tartrate
C4H4O6^-2
Chlorate
ClO3^-1
Chromate
CrO4^-2
Dichromate
Cr2O7^-2
Iodate
IO3^-1
Bicarbonate
HCO3^-1
Bisulfate
HSO4^-1
Binoxalate
HC2O4^-1
Hydrogen Phosphate
HPO4^-2
Dihydrogen Phosphate
H2PO4^-1
Hydrogen Sulfide
HS^-1
Permaganate
MnO4^-1
Amide
NH2^-1
Ammonium
NH4^+1
Nitrate
NO3^-1
Hydroxide
OH^-1
Peroxide
O2^-2
Phosphate
PO4^-3
Thiocyanate
SCN^-1
Thiosulfate
S2O3^-2
Sulfate
SO4^-2
Selenate
SeO4^-2
Hexafluorosilicate
SiF6^-2
Silicate
SiO3^-2


Polyatomic ions are the building blocks of many ionic compounds and they should be memorized.

How to name polyatomic ions
First, think of the -ate ion as being the "base" name, in which case the addition of a per- prefix adds an oxygen. Changing the -ate suffix to -ite will reduce the oxygens by one, and keeping the suffix -ite and adding the prefix hypo- reduces the number of oxygens by two. In all situations, the charge is not affected.

modify stem name with:
meaning
examples
-ate
a common form, containing oxygen
chlorate, ClO3-
nitrate, NO3-
sulfate, SO42-
-ite
one less oxygen than -ate form
chlorite, ClO2-
sulfite, SO32-
nitrite, NO2-
per-, -ate
same charge, but contains one more oxygen than -ate form
perchlorate, ClO4-
perbromate, BrO4-
hypo-, -ite
same charge, but contains one less oxygen than the -ite form
hypochlorite, ClO- hypobromite, BrO-
thio-
replace an O with an S
thiosulfate, S2O32-
thiosulfite, S2O22-

For more help, watch this video









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