The previous chapter explained the structure of atoms through their electron configuration. This was possible through a process called Aufbau. Atoms can combine to form different substances, since they are the basic building blocks of matter. Through chemical bonds, atoms combine to form molecules or compounds. Particularly, for biochemistry, this allows biological structures such as membranes, chromosomes and muscle to be described in chemical detail by understanding the nature of these chemical bonds. As described in the previous chapter, all elements, with the exception of noble gases, are unstable and react chemically to achieve the outer-electron configuration of noble gases. The result of these chemical processes is the formation of compounds in which atoms are connected to one another by chemical bonds.
This chapter introduces the two major types of chemical bonds; ionic and covalent bonds. An ionic bond is a strong electrostatic attraction between a positive ion and a negative ion. It is formed when electrons are transferred from the valence shell of one atom into the valence shell of another. On the other hand, a covalent bond is a bond that is formed when electrons are shared between atoms. In general, ionic bonds are formed between metals and nonmetals and covalent bonds are formed between nonmetals.
3.1 - Ionic Vs. Covalent Bonds
The type of bond that will form as a result of a chemical reaction depends on the relative abilities of the reacting atoms to attract electrons. Remember, reacting atoms will always tend to form octet configurations on their valence shell. When one atom has a much greater ability to attract electrons that the other atoms does, electron transfer will occur to ensure octet formation. In this case, the bond will be ionic. When two atoms have similar abilities in attracting electrons, electron transfer is not possible. In this case, octet formation is accomplished through electron sharing and the resulting bond will be covalent.
The strength with which the positively charged nucleus interacts with the outer electrons often dictates how easy or how difficult it is to remove electrons from the valence shell. It takes work to overcome the attraction of the positively charged nucleus for the negatively charged electron. The energy required for this process is called the ionization energy.
Def 3.1 - Ionization Energy
Ionization energy is the work or energy necessary to overcome the force of attraction from a positively charged nucleus to remove a negatively charged electron from the valence shell of the atom.
In general, ionization energies increase across any given period. The reason for this behavior can be explained by the nuclei that grow in positive charge (increase in protons = increase in positive charge). Over a period however, the valence shell stays the same. When new electrons are added, they stay on this same valence shell until it is filled (and when the valence shell is filled, a new outer shell takes its place and becomes the new valence shell). This means that the distance between each nucleus and the electrons going into the outer shell remains the same over a period, while the force of attraction increases. Hence, it becomes more and more difficult to remove electrons from the valence shell, and thus requiring more energy.
On the other hand, ionization energy decreases from the top to the bottom of any group of the periodic table. This can again be explained by the fact that the number of electron shells increase down each group. The outer electrons therefor become increasingly insulated from the influence of the positively charged nucleus. For example, within group I the energy needed to form cations is greatest for lithium and decreases in the order lithium > sodium > potassium > rubidium > cesium.
3.2 - Ionic Bonds
As explained before, the chemical reactivity of elements depends on their different tendencies to gain, lose or share electrons in their valence shell. After the reaction, the valence shells have a configuration that is characteristic of a noble gas and they have acquired a complete octet (or duet in some special cases).
Lewis Symbols and Formulas of Ionic Compounds
In the previous chapter, Lewis symbols where introduced. These symbols can now be used to illustrate chemical reactions and octet formation. Remember that the group numbers of the elements describe the number of electrons in their valence shell. For example, lithium in group I has only one electron in its valence shell. Furthermore, the ionic charge of an atom is the relative charge it obtains after gaining (anion) or losing (cation) an electron. The convention used to describe ionic charge is to first write the number, followed by the sign of the charge.
Example 3.1 - Ionic Bonds using Lewis Symbols
In this example, Lewis symbols will be used to describe the reactions between the metals lithium, magnesium and aluminium with the nonmetal fluorine. The Lewis symbols depict the reactions, with the arrows indicating the transfer of electrons.
- In the reaction between lithium and fluorine, lithium only needs to lose one electron to achieve the helium electron formation (being a duet, rather than an octet). Electron loss can, however, only occur when the electron can be accepted by another atom. This is the reason why the process is described as electron transfer, rather than electron loss. Fluorine is an atom that only requires one electron for octet formation in its valence shell. The electron transfer ratio required is therefor 1:1 of reacting atoms. One lithium against one fluorine.
- In the reaction between magnesium and fluorine, magnesium needs to lose two electrons to obtain the neon electron configuration. Fluorine again only requires one electron to obtain an octet electron configuration. The transfer of two electrons from magnesium therefore requires acceptance of those two electrons by two fluorine atoms. This results in a transfer ratio of 1:2 of reacting atoms. One magnesium against two fluorine.
- Similarly, in the reaction between aluminium and fluorine, aluminium needs to lose three electrons to obtain the neon electron configuration. The transfer ratio required is therefor 1:3 of reacting atoms. One aluminium against three fluorine.
Def 3.2 - Oxidation & Reduction
- The loss of electrons is called oxidation.
- The gain in electrons is called reduction. The name reduction can be remembered by keeping in mind that when an atom gains electrons, its relative charge will decrease or reduce, hence the name.
Def 3.3 - Formula
A chemical compound is identified by its formula, which is a statement about the combining ratio of its constituent elements.
In the formula, the symbols of all elements in the compound are combined with subscripts that are the relative numbers of atoms of each element. For example, the compounds produced by the reactions described in the previous example are:
- ⇒
- ⇒
- ⇒
Compounds that consist of two elements such as the ones described above are called binary compounds.
Formulas of Ionic Compounds
Often it is the case where one wants to predict the formula of binary ionic compounds. Before this is possible, the ionic charges of all component elements needs to be known. The following table lists the ionic charges for Groups I to III and V to VII.
- The elements in Groups I, II and III are metals, they lose electrons from their valence shells to form cations, whose positive charge is equal to their group number.
- The elements in Groups V, VI and VII are nonmetals, they gain electrons in their valence shells to form anions, whose negative charge is equal to their group number minus eight.
Group Number | I | II | III | V | VI | VII |
Ion Charge | 1+ | 2+ | 3+ | 3- | 2- | 1- |
The next step in determining the formulas for binary ionic compounds is to satisfy the requirement that the net electrical charge of an ionic compound must be zero. This means that the magnitude of the total positive charge must equal the magnitude of the total negative charge. Ones the charges of the individual ions are known, this information can be used to calculate how many of each ion are needed to achieve a net charge of zero. This approach is called the net-charge approach. As will become clear in the next example, this is rather cumbersome and requires some computational or guesswork to find the right multipliers. The guesswork can be left out by using the cross-over approach, where the number of the cation charge can be used as the number of anions and the number of the anion charge can be used as the number of cations.
Example 3.2 - Net-charge Vs. Cross-over Approaches
Take the following reaction example where aluminium reacts with oxygen. Since one aluminium atom needs to loose three electrons, but one oxygen atom can only acquire two electrons, multiple atoms will be required to allow each of them to form an octet configuration. It can therefor be determined that aluminium will form cations with an ionic charge of 3+ and oxygen will for anions with an ionic charge of 2-.
Net-charge Approach
To find a combination of aluminium and oxygen ions in which the net charge is zero, we can multiply those charge numbers by factors that make the total positive and total negative charges equal.
The factors used to multiply each elemental ion can now be used to express the formula of the ionic compound. The compound formed is .
Cross-over Approach
Using the cross-over approach, it can be concluded that the charge of the cation is 3+ and the charge of the anion is 2-.
- The number of anions necessary is equal to the ionic charge number of the cation. Since aluminium is the cation and has an ionic charge of 3+, three oxygen ions are necessary.
- The number of cations necessary is equal to the ionic charge number of the anion. Since oxygen is the anion and has an ionic charge of 2-, two aluminium ions are necessary.
This results in the compound , which is the same as the compound obtained by the net-charge approach (as expected).
The following table summarizes the possible combining ratios for anions (abbreviated N for nonmetal) and cations (abbreviated M for metal) of various charges.
3.3 - Naming Binary Ionic Compounds
By convention, in the name and formula of every ionic compound, the cation (metal) is indicated first, followed by the anion (nonmetal). The name of the metal is included unchanged, while the name of the nonmetal is combined with the suffix -ide. Some examples to make this clear are
- ⇒ Potassium chloride
- ⇒ Aluminium oxide
This convention is easy to follow for elements within the first three periods of the PT. For higher-order periods, transition elements can form ions of more than one charge type (e.g. iron can form ions with charges 2+ and 3+). To name binary ionic compounds for these elements, the Stock system is used. In this system, the charge of the cation is indicated with a roman numeral, describing the ionic charge of the ion. Again, for the anion, the suffix -ide is used. Some examples include
- ⇒ Iron (II) chloride
- ⇒ Iron (III) chloride
3.4 - Polyatomic Ions
Def 3.4 - Polyatomic Ion
A polyatomic ion is an ion that consists of two or more atoms that are combined into a single charged unit.
The formulas and names for some important polyatomic ions can be found in the following table.
Ion Formula | Ion Name |
Ammonium | |
Nitrate | |
Nitrite | |
Hydroxide | |
Acetate | |
Cyanide | |
Carbonate | |
Hydrogen carbonate | |
Phosphate | |
Hydrogen phosphate | |
Dihydrogen phosphate | |
Sulfate | |
Hydrogen sulfate | |
Hydrogen sulfite |
As with all other ions, when a polyatomic ion is present in a compound, the sum of positive and negative charges must equal zero at all times. If two or more of a given polyatomic ion are present in a formula, the formula is written with the polyatomic ion enclosed in parentheses followed by a subscript indicating the number of those ions in the formula. For example:
- Aluminium carbonate ⇒
- Calcium phosphate ⇒
3.5 - Ionic Compound Formula And Structure
Def 3.5 - Ionic Solid
An ionic solid is a crystalline material composed of ions held together by strong electrostatic forces, forming a rigid lattice structure.
Ionic solids do not conduct electricity and neither does pure water (i.e. pure ). When an ionic solid such as sodium chloride () is dissolved in water or melted, the solutions and the melted solids do conduct electricity. Ionic compounds are therefor called electrolytes. Water and compounds which, when dissolved in water, do not conduct electricity are called nonelectrolytes. As the definition states, when in solid form, the ions of an ionic compound are locked into place and all their charge is neutralized. Dissolving or melting the solid frees individual ions, which in turn can carry an electric current. This reveals that an ionic bond only exists within a solid state. Once the solid is dissolved or melted, the constituent ions become independent of one another.
Finally, an ionic solid consists of a three-dimensional array of both positive ions surrounded by negative ions and negative ions surrounded by positive ions in a lattice-like structure. This arrangement makes it impossible to identify a discrete structural unit that can be represented by the formula of the solid. There is no such structural unit as a molecule for an ionic compound. In the next section, the nature of covalent bonds will be introduced. it will become clear that the formulas for compounds such as water , carbon dioxide and oxygen , whose chemical bonds are covalent, represent not only the combining ratios of component elements, but also the actual structural units that will be called molecules.
3.6 - Covalent Compounds & Nomenclature
Example 3.3 - Simple Covalent Bond
One of the simplest covalent bonds unites two hydrogen atoms to form the hydrogen molecule . Hydrogen atoms possess only one electron, and they can achieve the noble-gas configuration of helium (a duet rather than an octet) by gaining one electron more.
The helium configuration calls for a 1s orbital to contain two electrons. Hund’s rule tells us that such two electrons can only share an orbital if they have opposite spins. Therefore, a single covalent bond uniting two atoms consists of two spin-paired electrons shared by both atoms.
There a systematic method for naming simple, binary molecular compounds. This method is based on the fact that some pairs of elements can form more than one covalent covalently bonded compound. Examples of such pairs are (1) nitrogen and oxygen and (2) carbon and oxygen. The method uses Greek prefixes which can be found in the following table.
Number | Greek Prefix |
1 | mono- |
2 | di- |
3 | tri- |
4 | tetra- |
5 | penta- |
6 | hexa- |
7 | hepta- |
8 | octa- |
9 | nona- |
10 | deca- |
The first element in the formula is named first and the second element is named second, where the name of the second element is again suffixed by -ide. The Greek prefixes are added to the elements’ names to indicate the number of different kinds of atoms in each molecule of the compound. Note however that the prefix mono- is never used for naming the first element! The following table provides some examples.
Compound | Systematic Name |
dinitrogen oxide | |
nitrogen oxide | |
nitrogen dioxide | |
dinitrogen trioxide | |
dinitrogen tetroxide | |
dinitrogen pentoxide | |
carbon monoxide | |
dicarbon trioxide | |
carbon dioxide |
3.7 - Representation of Covalent Bonds
The electron-dot notation was introduced last chapter to represent the valence shell of elements. This representation can be easily adapted to represent covalent bonds as well. When the electron-dot notation is used to represent the structure of molecules, the resulting diagrams are called Lewis structures. To effectively build Lewis structures, the key concept is to satisfy the octet rule described in the previous chapter. When elements react to form covalent bonds, the number of bonds that are formed to complete the octet is called its combining power.
- Group IV elements are tetravalent, they require four additional electrons to fill an octet.
- Group V elements are trivalent, they require three additional electrons to fill an octet.
- Group VI elements are divalent, they require two additional electrons to fill an octet.
- Group VII elements are monovalent, they require one additional electron to fill an octet.
In general, lines can represent the covalent bonds between atoms, as will become clear in the following examples.
Example 3.3 - Predicting Formulas of Simple Molecular Compounds
Imagine we are given nitrogen and iodine and we are tasked with predicting the formula of the product of the reaction between these two elements.
- Nitrogen is an element in Group V and is trivalent.
- Iodine is an element in Group VII and is monovalent.
Therefore, one nitrogen atom will react with three iodine atoms and form the compound .
Example 3.4 - Lewis Method for Molecular Structure
Imagine two fluorine elements that react with each other to form a diatomic molecule. Fluorine is an element in Group VII and is monovalent, so two fluorine atoms will react and form a covalent bond between their free electrons. This is visualized in the following figure.
On the left side of the reaction, the reacting atoms are represented with the number of valence electrons (i.e. equal to their group number). On the right side of the reaction, the fluorine molecule can be seen and how each atom is surrounded by a complete octet. The blue overlapping circles represent the individuals atom structure. The electrons colored in red represent the electrons that are shared in the covalent bond. The Lewis structure can be further simplified by drawing a line between the shared electrons as follows.
The pairs of electrons that are not shared in a covalent bond are called nonbonded electrons or lone pairs. Lewis structures can be simplified even further by leaving out all nonbonded electrons. In this case, Lewis formulas are called structural formulas.
The molecular formula does tell the number of atoms for each element in a molecule of the compound. It does not, however, describe the molecule as a Lewis structure or structural formula. It the following paragraphs, a step-by-step guide will allow to construct Lewis structures from molecular formulas for different types of molecules through the form of examples.
Example 3.5 - Drawing Simple Lewis Structures
In this example, a step-by-step guide is given on how to construct the Lewis structure for ammonia .
Step 1
First, the symbols for the bonded atoms should be placed into an arrangement that will allow to start distributing electrons. In ammonia, there is only one atom of one type (in this case nitrogen ) and several others of another type (in this case hydrogen ). A good first approximation is to choose the singular atom as a central atom to which the rest will be bonded.
Step 2
Next, the total number of valence electrons in the molecule should be determined. To do so, the valence electrons that contribute to each atom can be added together.
Element | Valence Electrons | Atoms / Molecule | # Of Electrons |
Nitrogen | 5 | 1 | 5 |
Hydrogen | 1 | 3 | 3 |
This leads to a total of 8 available valence electrons.
Step 3
Shared pairs of electrons can be represented by drawing a line between bonded atoms as was done with structural formulas before. In the case of ammonia, each electron in the hydrogen atoms can be bonded with one electron in the valence shell of the nitrogen atom.
Step 4
Since the hydrogen atoms forms a duet (rather than an octet) with three of the five available valence electrons of the nitrogen atoms, there are still two electrons available. These remaining valence electrons are lone pairs that have to be positioned as to satisfy the octet rule (in this case, for nitrogen), resulting in the following Lewis structure.
Example 3.6 - Lewis Structures Without Central Atom
In some cases, the Lewis structure of a compound with no central atom has to be represented. In this example, the Lewis structure for the compound ethane will be constructed.
Step 1
Since hydrogen can only form a single bond, they cannot be inserted between two carbon atoms. This is why the carbon atoms can be placed such that they are “joined together” with the hydrogen atoms in the “terminal” positions.
Step 2
As in the previous example, the total number of valence electrons in the molecule should be determined.
Element | Valence Electrons | Atoms / Molecule | # Of Electrons |
Carbon | 4 | 2 | 8 |
Hydrogen | 1 | 6 | 6 |
This leads to a total of 14 available valence electrons.
Step 3
All the electrons are accounted for by drawing the seven bonds required to connect all the atoms of the molecule. These seven bonds use up all 14 free electrons and thus there are no more electrons that would represent lone pairs. The final Lewis structure can then be described as follows.
Example 3.7 - Lewis Structures With Double Bonds
Some interesting cases arise when there appear to be too few electrons to satisfy the octet rule. This “apparent” deficiency occurs when we try to write the Lewis structures for compounds such as ethylene or carbon dioxide .
Step 1
In the case of ethylene, steps similar to the previous example can be followed where hydrogen is placed at “terminal positions”.
Step 2
As in the previous example, the total number of valence electrons in the molecule should be determined.
Element | Valence Electrons | Atoms / Molecule | # Of Electrons |
Carbon | ㅤ | ㅤ | ㅤ |
ㅤ | ㅤ | ㅤ | ㅤ |