Notes

Ionic, Covalent, and Metallic Bonds:

So, we have three kinds of atoms: atoms that are burdened with "extra" electrons, atoms that are in "need" of more electrons, and atoms that have "enough" electrons. Depending on the types of atoms involved, three different types of bonds can
occur: ionic bonds, metallic bonds, and covalent bonds.

In this lesson, you will explore how elements participate in the different types of bonds, how each type of bond is formed, and the properties that each type of bond confers upon the compounds formed.

The Octet Rule:
As you know, the arrangement of electrons in the atom is called its electron configuration. The outermost energy level of each atom is the valence shell, and the electrons present in the valence shell are called valence electrons. Valence electrons play a vital role in the formation of compounds because each element tries to attain stability by following the octet rule.

The octet rule refers to the tendency of elements to gain, lose, or share electrons so that the atom's valence shell is filled. It is called the octet rule because a filled valence shell contains eight electrons (octet means a "set of eight"). Hydrogen and helium are excluded from this rule because their valence shell is in energy level one and requires only two electrons to fill.
The octet rule can be illustrated with the help of a Lewis dot diagram. Take a look at this graphic showing the Lewis structures of aluminum (Al) and chlorine (Cl). The Lewis dot diagram for aluminum is drawn with three dots, which represent its three valence electrons. The diagram for chlorine is drawn with seven dots representing its valence electrons.

Each dot in a Lewis dot diagram is placed sequentially around each side of the symbol. The maximum number of electrons allowed per side is two. Lewis dot diagrams are useful because they illustrate the rearrangement of electrons in atoms during chemical reactions to achieve a full set of valence electrons, or an octet.

Elements that have the required number of electrons in the valence shell of their atoms are called noble elements. With the exception of helium, which has two electrons in its valence shell, noble elements have eight electrons in the the valance shell. As these elements all happen to be gases under normal conditions, they are called noble gases. They remain stable (or inert) because they almost never react with other elements. All of the other elements have either more or fewer than eight electrons in the valence shell, which makes them unstable.
Elements follow the octet rule by striving to have exactly eight electrons in the valence shell and thus attain the stability level of a noble gas. The only way an element can do so is by combining with other elements. In this process, an atom may gain electrons, lose electrons, or share electrons with other atoms.

During the bonding process it is common for metals to lose electrons and most nonmetals to gain or share electrons to attain stability. The tendency of atoms to follow the octet rule often leads to the formation of ions.

Ions:
An ion is an electrically charged atom (or group of atoms) formed when the atom gains or loses electrons. An atom is always electrically neutral, as the number of protons within its nucleus equals the number of electrons surrounding the nucleus. When an atom gains or loses electrons while combining with other atoms, it is no longer neutral because the number of electrons becomes either less or more than the number of protons. So, the atom converts into an electrically charged ion. Depending on its charge, it's either an anion (negatively charged) or a cation (positively charged).
Anions:
Anions are negatively charged ions. When an atom gains electrons, it ends up with more electrons than protons. So, the net charge of the atom, which is now called an anion, is negative. The suffix –ide is used to name an anion; for example, chlorine forms the chloride ion.

Nonmetals have more than three valence electrons in their outer shell. It's easier for a nonmetal to gain electrons than to lose electrons, so nonmetals tend to form anions. For example, chlorine (Cl) has seven valence electrons in its neutral state. It gains one electron to complete its octet and forms a chloride ion (Cl-). Chlorine can also form an octet by losing seven electrons, but since it's easier for it to gain one electron than to lose seven, it's more likely to form an anion.

The gain of this one electron gives the chloride ion the same number of electrons as argon, its nearest noble gas.
Cations:
Cations, on the other hand, are positively charged ions. When an atom loses electrons, it ends up with fewer electrons than protons. So, the net charge of the atom, now called a cation, is positive. A cation takes the same name as the element; sodium, for example, forms the sodium ion.

Most metals have one, two, or three valence electrons in the outer shell. Losing those electrons requires less energy than gaining enough electrons to reach noble gas stability. So, metals tend to lose electrons and form cations. For example, sodium (Na) has one valence electron in its neutral state. It loses one electron to attain an octet and forms a sodium ion (Na+). Sodium can also form an octet by gaining seven electrons, but since it's easier for it to lose one electron than to gain seven, it's more likely to form cations.

The loss of this one electron gives the sodium ion the same number of electrons as neon, its nearest noble gas.
cations form when an atom loses one or more electrons

Ionic Bonding:
Ionic bonds are formed by transferring electrons between two atoms. Chlorine has seven valence electrons, so it needs only one more to complete its octet. Sodium, on the other hand, has an extra electron to give up to complete its octet. Do they seem like a perfect match? You bet!

When sodium (Na) is heated and exposed to chlorine gas (Cl2), the ionic compound sodium chloride is formed. The sodium atom gives away one electron to the chlorine atom, completing the octet of electrons. Consequently, the sodium atom acquires a positive charge and the chlorine atom acquires a negative charge. The two ions, Na+ and Cl-, are held together by an ionic bond. Watch the video, which shows how sodium chloride (NaCl) is formed.
In an ionic bond, anions and cations are bound by an electrostatic force, which is the force of attraction or repulsion between charged bodies. A compound thus formed is called an ionic compound, an ionic solid, or a salt.

Another example of an ionic compound is potassium oxide. Potassium has one valence electron, and oxygen has six valence electrons. So, two potassium atoms can lose one electron each as they bond with an oxygen atom to complete their octets. Thus, two potassium atoms and one oxygen atom participate in the formation of potassium oxide, as shown above.

Properties of Ionic Compounds:
The crystal lattice structure of ionic compounds is responsible for their characteristic properties. The crystal lattice of NaCl is shown here. Let's explore what a crystal lattice is and how this structure is formed.

In an ionic bond, the ions are tightly packed together because the oppositely charged ions strongly attract each other. Consequently, the ions in an ionic compound align in an alternating pattern of cation, anion, cation, and so on, to form a three–dimensional structure called a crystal lattice. This alternating arrangement of oppositely charged ions maximizes the attraction between opposite charges and minimizes the repulsion between like charges.

When cations and anions bond, their union creates a neutral,
non–charged compound because the cations and anions have opposite charges of the same magnitude, which cancel each other. In the case of NaCl, for example, the charge on the Na ion is 1+ and the charge on the Cl ion is 1-. This one–to–one ratio of the two ions neutralizes their opposite charges, thus producing a neutral compound.
Because of the geometric arrangements of the crystal lattice, compounds created by ionic bonds are solid at room temperature, have a high melting point, and are hard, brittle, and electrically neutral unless dissolved in water.

Ionic compounds are solid at room temperature and have high melting points because the bonds between cations and anions are very strong. It requires a lot of energy to break those bonds and allow the ions to flow freely in a liquid state.

The crystal lattice is also responsible for making most ionic compounds brittle solids. When struck, the ions in the crystal shift, which can cause cations to align with cations and anions to align with anions. The repulsion between ions with the same charge becomes so intense that the crystal shatters along the fissure where the ions shifted.

Ionic compounds are electrically neutral because the crystal lattice structure does not allow an electrical charge to move through it. So, these compounds conduct electricity only when melted or dissolved in water, as these conditions allow the ions to move.
these propertys are the result of the crystal lattice structure of ionic compound crystalline solids

Covalent Bonding:
Unlike ionic bonds, where atoms transfer electrons while forming bonds, covalent bonds are formed by the sharing of electrons between atoms of nonmetals. Two or more atoms linked together by covalent bonds form a molecule; the compounds formed by covalent bonding are called molecular compounds. The number and types of atoms that constitute a molecule are described by a molecular formula. Here again, the atoms try to attain the electron configuration of a noble gas.

Chemists often use Lewis structures to show how atoms are covalently bonded to one another. Lewis structures resemble the Lewis dot diagrams used to show ionic bonding, but in covalent bonds, the dots are often replaced by dashes to show that electrons are shared instead of given or taken. Each dash represents two shared electrons in a bond and counts as two valence electrons in each atom's valence shell.
A methane gas molecule (CH4) is covalently bonded, containing one carbon atom and four hydrogen atoms. The carbon atom has four valence electrons and requires four more to fill its valence shell. The hydrogen atom has one valence electron and requires one more to fill its valence shell.

Look at the 3–D structure and the Lewis structure of methane. Two electrons are shared between the carbon atom and each of the four hydrogen atoms. This provides the two electrons to fill each hydrogen atom's valence shell and the eight electrons needed to fill the carbon atom's valence shell.

Not all electrons in the valence shell are required to bond to form molecules. Look at the 3–D structure and the Lewis structure of water (H2O). Oxygen has six valence electrons and hydrogen has only one. The Lewis structure shows how these electrons are shared between the hydrogen and oxygen atoms.

The octet requirements for all three atoms are fulfilled. Two pairs of electrons on the oxygen atom remain unshared. These are referred to as unshared electrons, nonbonding pairs, or lone pairs.

Exception to the Octet Rule:
As with most rules, there are exceptions to the octet rule. For example, carbon monoxide (CO) does not follow the simplest form of the octet rule: equal donation of shared electrons.

In a carbon monoxide molecule, carbon has four valence electrons and oxygen has six. The carbon and oxygen atoms share six electrons between them, forming a triple covalent bond. Four of the six shared electrons come from oxygen while only two come from carbon. Still, no matter where they originated from, both atoms now have eight valence electrons: three shared pairs and one unshared pair.

A covalent bond is formed when two atoms contribute one electron each to the formation of the bond. But a coordinate covalent bond is formed when two electrons of the same atom contribute to the formation of the bond. Carbon monoxide has two covalent bonds and one coordinate covalent bond.
Covalent compounds such as methane, water, and carbon monoxide owe their characteristic properties to their covalent bonds. Unlike ionic compounds, which are solid at room temperature, covalent compounds may be gaseous, liquid, or solid at room temperature. Carbon monoxide and methane are gases, water is a liquid, and sugar is a solid at room temperature.

Both the melting and boiling points of most covalent compounds are lower than those of ionic compounds. Covalently bonded molecules do not typically form covalent bonds with each other. The weak intermolecular forces that hold neutral molecules together in a liquid or a solid are not very strong.

Network solids are exceptions to this rule. In a network solid, each atom is covalently bonded to many other atoms, forming a sturdy structure. For example, in a diamond, each carbon atom covalently bonds with four carbon atoms to form a stable network solid as shown in the graphic. These covalent compounds melt only at extremely high temperatures. A diamond molecule, for example, is so tough that it remains solid even at 1,064ºC, which is the melting point of gold.
ammonia (NH3) molecule is formed with covalent bonds when A nitrogen atom shares one electron each with three hydrogen atoms

Metallic Bonding:
If ionic and covalent bonding represent the two extremes of the bonding spectrum, metallic bonding falls somewhere in between. Metallic bonds occur only between metal atoms. Metals are made up of cations, and electrons drift freely among these cations.

The simplest model of metallic bonding is the electron sea model. This model suggests that mobile valence electrons surround closely packed metal cations, forming a structure resembling marbles packed in a jar with ball bearings. The electrons essentially act as ball bearings, allowing the cations, represented by the marbles, to slide past one another while reducing the repulsive force between them. Thus the electrons are free to move around the cations instead of being anchored to them.
metal atom may have one, two, or three electrons in its valence shell. The fewer the electrons in the valence shell, the weaker the bond between the valence electrons and the nucleus of the atom.

This means that metal atoms can easily part with their valence electrons to form metal ions. The electrons that separate from their metal atoms during the formation of cations are known as delocalized electrons. These electrons surround the cations in the metal. The attraction between these free–flowing electrons and the cations is the metallic bond, which gives the metal its structure and properties.

For example, in aluminum, aluminum ions are surrounded by an electron cloud that moves throughout the metal. These mobile electrons are responsible for conducting heat and electricity through the metal, as well as its malleability.
The electron sea model explains many properties of metals, such as their electrical and thermal conductivity, ductility, and malleability. An electric current occurs due to the free movement of electrons. Metals are good electrical conductors because electrons constantly move through a metal. Metals are also good conductors of heat because the moving electrons transfer thermal energy.

Ductility and malleability are properties referring to the ability to shape metals. Ductile objects can be easily drawn into wires. Copper, for example, is both ductile and a good conductor of electricity, so copper wires are extensively used in electrical appliances.

Malleable objects can be pounded and bent into shape. The malleability of metals, such as gold, differs from the brittleness of ionic crystal lattices, which break along fissures. The reason metals do not break like ionic crystals is that the electrons essentially act as ball bearings, allowing the cations to slide past one another while reducing the repulsive force between the cations.

An ionic bond is formed by the give and take of electrons between two atoms. It binds the anions and cations together, forming an ionic compound. Fluorite and pyrite are examples of ionic compounds.

Covalent bonds are formed by the sharing of electrons between atoms of nonmetals. Compounds formed by covalent bonding are called molecular compounds. Water and carbon dioxide are examples of molecular compounds

A metallic bond is fomed between metal cations and free-flowing electrons in a metal. Iron and platinum, for example, have metallic bonds.
ionic bond=crystalline
metallic bond=ductile
covalent bond=poor thermal conductivity
covalent bond=molecular compounds
ionic bond=hard and brittle
metallic bond=electron sea model