Periodic Trends

The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements. Elements tend to gain or lose valence electrons to achieve stable octet formation. Stable octets are seen in the inert gases, or noble gases, of Group VIII of the periodic table. In addition to this activity, there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group. These trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity

 

Atomic Radius

The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching each other. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the largest atomic radii are located in Group I and at the bottom of groups.

Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Electrons within a shell cannot shield each other from the attraction to protons. Since the number of protons is also increasing, the effective nuclear charge increases across a period. This causes the atomic radius to decrease.

Moving down a group in the periodic table, the number of electrons and filled electron shells increases, but the number of valence electrons remains the same. The outermost electrons in a group are exposed to the same effective nuclear charge, but electrons are found farther from the nucleus as the number of filled energy shells increases. Therefore, the atomic radii increase.

 

Ionization Energy

The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.

 

Electron Affinity

Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affinities of certain groups in the periodic table. The Group IIA elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high electron affinities because the addition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases, have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other groups have low electron affinities.

 

Electronegativity

Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus. In a group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine.

 

Summary of Periodic Table Trends

Moving Left --> Right

• Atomic Radius Decreases
• Ionization Energy Increases
• Electronegativity Increases

Moving Top --> Bottom

• Atomic Radius Increases
• Ionization Energy Decreases
• Electronegativity Decreases

Trends Across a Period Quiz

Periodic Trends Quiz

The modern periodic table shows the elements arranged in order of increasing atomic number.

v  Arranged in this manner, the elements show a periodic pattern in their properties.

v  This is why it’s called a Periodic table!

What this means is that certain groups of elements share similar physical & chemical properties. 

v  Periodically, as you run through the elements by increasing atomic number, you come across an element with similar properties to any element of your interest. 

v  This is similar to a calendar; as you run through the days by increasing number, you periodically come across a Monday, Wednesday, etc. 

These periodic trends are due to the atoms’ composition and their desire to have a full outer energy level.  Remember the octet rule: atoms desire to gain or lose e^- to achieve 8 valence e^- (to be like noble gases).

v  Metals typically give up their e^- (form positive ions) to be like the previous noble gas.

v  Nonmetals typically gain more e^- (form negative ions) to be like following noble gas.

Periodic Trends

1.      Atomic Radius: Radius of atom without regard to surrounding atoms.  (Size of the atom)

v  Left to right, atomic radius ______________.  The reason for this is _______________________

______________________________________________________________________________

v  Top to bottom, atomic radius ______________. The reason for this is _____________________

______________________________________________________________________________

2.      Ionic Radius: Radius of ion formed from atom.  (Did it gain or lose electrons?)

v  Left to right and top to bottom, ionic radius mostly follows the same trend as atomic radius (which should make sense – think about from where the ion comes).

v  Positive ions are ______________ than the uncharged atom. The reason for this is ___________

_____________________________________________________________________________

v  Negative ions are ______________ than the uncharged atom. The reason for this is __________

_____________________________________________________________________________

3.      Electron Affinity: Attraction an atom has for an extra e^-. (How much it wants to form a [-] ion.)

v  Left to right, electron affinity ______________.  The reason for this is ____________________

_____________________________________________________________________________

v  Top to bottom, electron affinity ______________. The reason for this is ___________________

_____________________________________________________________________________

4.      Ionization Energy: Energy needed to remove an e^- from atom. (Energy needed to form a [+] ion.)

v  Left to right, ionization energy ______________.  The reason for this is ___________________

_____________________________________________________________________________

v  Top to bottom, electron affinity ______________. The reason for this is ___________________

_____________________________________________________________________________

5.      Electronegativity: Tendency of an element to pull e^- toward itself in a bond. (Follows e^- affinity trend)

v  Left to right, electronegativity _______________.  The reason for this is ___________________

_____________________________________________________________________________

v  Top to bottom, electronegativity _______________. The reason for this is _________________

_____________________________________________________________________________

The Enthalpies

What is Enthalpy?

Usually there is energy change involved in a chemical reaction somehow. In chemistry this change in heat energy is represented as enthalpy change using the following symbol:

So that this figure is the same everywhere and fair, it is taken at standard conditions which are 298K and 100 kPa pressure. Here are two reactions, each with a different enthalpy change.

For the first rection, the enthalpy change is negative this is because the reaction is exothermic meaning it gives out heat. The reason it is negative is energy is leaving the reaction.
The second reaction has a positive enthalypy change, this is because it is endothermic and heat is entering the system.

Enthalpy of Reaction, Combustion, and Formation

These are three different types of situation that an energy change may take place. On the top image x is used to show where a letter is written to indicate what type of enthalpy is shown: r for reaction and so on.
The easiest one to get your head around is standard enthalpy of reaction this is merely the energy change of a reaction in the quantities that are expressed.
The standard enthalpy of combustion is the enthalpy change when 1 mole of substance is completely burned in excess oxygen under standard conditions, all reactants in their standard states. The enthalpies of combustion can be determined using a calorimeter which is explained below; it uses the symbol c.
The standard enthalpy of formation is the enthalpy change when 1 mole of compound is produced from its elements in standard conditions, all products and reactants in the standard state. The standard state is important, it means the way that the element is, at standard conditions (see above). So you would have H₂(g) NOT H on its own. Therefore you sometimes need to do fractions of elements, as with the standard enthalpy of formation of ethanol:

C (s) + 2H₂ (g) + ½O₂ (g) ® CH₃OH (l)

other enthalpies....

Enthalpy	Definition
Atomisation ΔHθa	The standard enthalpy change when 1mol of gaseous atoms are formed. Depending on the standard state of the element, this could also be called the enthalpy of sublimation if going from solid to gas.
Bond dissociation, ΔHθdiss	The standard enthalpy change when 1mol of gaseous covalently bonded molecule is broken to form 2 radicals. So the enthalpy change for A (g) ® X. (g) + Y..
Electron Affinity, ΔHθea	The standard molar enthalpy change when an electron is added to an atom in the gas phase. For example: Cl (g) + e- ® Cl- (g)
Lattice Enthalpy, ΔHθL	The standard enthalpy change accompanying dissociation: the seperation of 1 mol of solid ionic lattice into gaseous ions; formation the creation of 1 mol of solid from gas ions. Note that formation enthalpies are always negative and dissociation positive.
Hydration, ΔHθhyd	The standard molar enthalpy change for the process: X+/- (g)® X+/- (aq).
Solution, ΔHθsol	The standard enthalpy change when 1mole of ionic solid dissolves in enough water that none of the dissolved ions interact with each other

Chemical Bonding Quiz

Chemical Bonds (Ionic and Covalent) Quiz

Test your knowledge of ionic and covalent bonds.

The chemical formula of the product formed from the reaction between Mg and O2 is _______. (1 point) MgO2 MgO Mg2O

How is the bond in F2 different from the bond in KCl ? (1 point) F2 is covalent and KCl is ionic F2 is ionic and KCl is ionic F2 is ionic and KCl is covalent

Non-metals tend to _______ electrons to be come ________ ions. (1 point) lose, positive gain, negative lose, neutral

From the list of elements given, select 2 elements that would likely form an ionic bond. K, C, Br, Ar. (1 point) K, Br Ar, C K, Ar

Metals tend to _______ electrons to become _____ ions. (1 point) lose, positive gain, negative lose, neutral

In an ionic bond, electrons are _________ or _______. (1 point) lost or gained shared none of the above

Charged atoms are called as _____. (1 point) protons electrons ions

An atom that has lost or gained electrons becomes a (an) _______. (1 point) proton electron ion

Ionic bond is formed between atoms of ________ and _______. (1 point) metals and non-metals metals non-metals

When a Na atom, loses one electron, it gets a charge of _______. (1 point) -1 +1 0

Write the chemical formula for a compound that has one Calcium atom and 2 Chlorine atoms. Predict the bond between them. (1 point) CaCl2, ionic CaCl2, covalent Ca2Cl, ionic

Oxygen has 6 valence electrons and it is diatomic. How many covalent bonds are there in an O2 molecule ? (1 point) single double triple

When a Cl atom gains an electron, it gets a charge of ________. (1 point) -1 +1 0

The bond formed when Mg combines with O2 is _____. (1 point) ionic covalent metallic

N has 5 valence electrons and it is diatomic. How many covalent bonds are there in a N2 molecule ? (1 point) single double triple

Calcium Oxide is a (an) _____ compound and the formula for Calcium Oxide is ________. (1 point) covalent, CaO2 ionic, CaO ionic, CaO2

The process of becoming an ion from an atom is called as ________. (1 point) ionization electrolysis electron affinity

Covalent bond is formed between atoms of ________. (1 point) metals non-metals metals and non-metals

Hydrogen Fluoride has a (an) ________ bond and its chemical formula is ________. (1 point) covalent, HF ionic, HF covalent, H2F

Magnesium Bromide is a (an) ________ compound. (1 point) metallic covalent ionic

In a covalent bond, electrons are _______. (1 point) lost or gained shared none of the above

Intermolecular Forces

INTERMOLECULAR FORCES Introduction: The physical properties of melting point, boiling point, vapor pressure, evaporation, viscosity, surface tension, and solubility are related to the strength of attractive forces between molecules. These attractive forces are called Intermolecular Forces. The amount of "stick togetherness" is important in the interpretation of the various properties listed above. There are four types of intermolecular forces. Most of the intermolecular forces are identical to bonding between atoms in a single molecule. Intermolecular forces just extend the thinking to forces between molecules and follows the patterns already set by the bonding within molecules. 1. IONIC FORCES: The forces holding ions together in ionic solids are electrostatic forces. Opposite charges attract each other. These are the strongest intermolecular forces. Ionic forces hold many ions in a crystal lattice structure. Review - Ionic Bonds
	2. DIPOLE FORCES: Polar covalent molecules are sometimes described as "dipoles", meaning that the molecule has two "poles". One end (pole) of the molecule has a partial positive charge while the other end has a partial negative charge. The molecules will orientate themselves so that the opposite charges attract principle operates effectively. In the example on the left, hydrochloric acid is a polar molecule with the partial positive charge on the hydrogen and the partial negative charge on the chlorine. A network of partial + and - charges attract molecules to each other. Review - Polar Bonds
	3. HYDROGEN BONDING: Link to more extensive discussion: Hydrogen Bonding The hydrogen bond is really a special case of dipole forces. A hydrogen bond is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine. In other words - The hydrogen on one molecule attached to O or N that is attracted to an O or N of a different molecule. In the graphhic on the left, the hydrogen is partially positive and attracted to the partially negative charge on the oxygen or nitrogen. Because oxygen has two lone pairs, two different hydrogen bonds can be made to each oxygen. This is a very specific bond as indicated. Some combinations which are not hydrogen bonds include: hydrogen to another hydrogen or hydrogen to a carbon.
	4. INDUCED DIPOLE FORCES: Forces between essentially non-polar molecules are the weakest of all intermolecular forces. "Temporary dipoles" are formed by the shifting of electron clouds within molecules. These temporary dipoles attract or repel the electron clouds of nearby non-polar molecules. The temporary dipoles may exist for only a fraction of a second but a force of attraction also exist for that fraction of time. The strength of induced dipole forces depends on how easily electron clouds can be distorted. Large atoms or molecules with many electrons far removed from the nucleus are more easily distorted. Review - Non-Polar Bonds

Metallic Bonding

More than 80 elements in the periodic table are metals. Metals are solids at ordinary temperature and pressure, with the exception (of mercury and gallium). Metals have characteristic properties such as:

• High thermal and electrical conductivity.

• Luster and high reflectivity.

• Malleability and ductility. They can be beaten or shaped without fracture.

• Variability of mechanical strengths (ranging from soft alkali metals to Tungsten, which is hard).

The force that binds together the atoms of metals is called metallic bond. The properties of metals cannot be explained in terms of common types of bonds such as ionic and covalent bonds. The inadequacy of these two type of bonds for metal formation can be explained as under.
The atoms of metals are all alike therefore they cannot form ionic bonds. Moreover, ionic compounds do not conduct electricity in the solid state and ionic compounds are brittle as opposed to properties of metals. The atom of metallic elements contain only 1 to 3 valence electrons, therefore these atoms cannot form covalent bonds, with noble gas configurations as they will remain incomplete. Covalent compounds are bad conductors of electricity and are generally liquids; properties opposed to metal formations. Thus, metals have a different model of bonding.

Electron sea model for metallic bonding

To account for the bonding in metals, Lorentz proposed a model known as electron gas model or electron sea model. This model is based on the following characteristic properties of metals:

Low ionization energies

Metals generally have low ionization energies. This implies that the valence electrons of metal atoms are not strongly held by the nucleus. Valence electrons can move freely out of the influence of their kernels (atomic orbit/structure minus valence electrons). Thus, metals have free mobile electrons.

Large number of empty orbitals

It has been observed that in metals a number of valence orbitals remain empty as the number of valence electrons in metals is generally less than the number of valence orbitals.
For example, lithium {(Li, Z = 3) 1s²2s¹} has 2p-orbitals vacant; Sodium {(Na, Z = 11) 1s²2s²2p⁶ 3s¹} has 3p-and 5d-orbitals vacant;
Magnesium {(Mg, Z = 12) 1s²2s²2p⁶ 3s²} has 3p-and 3d-orbitals vacant The important features of electron sea model are:

• The positively charged kernels of metal atoms are arranged in a regular fashion in a metallic lattice.

• Loosely held valence electrons, surround each kernel in metallic lattice. Being loosely held to its kernel, the valence electrons enjoy complete freedom in the metallic lattice and are regarded as mobile electrons.

In short, the metal may be regarded as 'a sea of electrons (common pool of electrons) in which there is a three dimensional ordered arrangement of positively charged kernels, surrounded throughout by mobile valence electrons'. This explanation is also responsible for its name electron sea model. Thus, the simultaneous force of attraction between the mobile electrons and the positive kernels that binds the metal atoms together, is known as metallic bond.

Fig: 6.19 - Electron sea model

Comparison of Ionic bond Covalent bond and Metallic bond

Ionic Bond	Covalent Bond	Metallic Bond
The transfer of electrons between two atoms having different electro negativities forms this bond.	This bond is formed by the mutual sharing of electrons between same or different elements .	This bond is formed due to the attraction between kernels and the mobile electrons in a metal lattice.
This is a strong bond due to electrostatic force of attraction.	This is also a fairly strong bond because the electron pair is strongly attracted by two nuclei.	This is a weak bond due to the simultaneous attraction of the electrons by a large number of kernels
This is a non-directional bond.	This is a directional bond.	This is a non-directional bond.
This bond makes substances hard and brittle.	This bond makes substances hard and incompressible.	This bond make substances malleable and ductile.

Explanation of physical properties of metals

All the characteristic metallic properties can be explained on the basis of the electron.

Metallic lustre

The bright lustre of metals is due to presence of delocalised mobile electrons.
When light falls on the surface of the metal, the loosely held electrons absorb photons of lights. They get promoted to higher energy levels (excited state), oscillating at a frequency equal to that of the incident light. These oscillating electrons readily return from the higher to the lower levels of energy by releasing energy, thus becoming a source of light radiations. Light appears to be reflected from metal surface and the surface acquires a shining appearance, which is known as metallic lustre.

Electrical conductivity

The presence of mobile electrons causes electrical conductivity of a metal. When a potential difference is applied across the metal sheet, the free mobile electrons in the metallic crystal start moving towards the positive electrode. The electrons coming from the negative electrode simultaneously replace these electrons. Thus, the metallic sheet maintains the flow of electrons from negative electrode to positive electrode. This constitutes electrical conductivity.

Thermal conductivity

When a part of the metal is heated, the kinetic energy of the electrons in that region increases. Since the electrons are free and mobile, these energetic electrons move rapidly to the cooler parts and transfer their kinetic energy by means of collisions with other electrons. Therefore, the heat travels from hotter to cooler parts of the metals.

Malleability and ductility

Metals can be beaten into sheets (malleability) and drawn into wires (ductility). Metallic bonds are non-directional in nature. Whenever any stress is applied on metals, the position of adjacent layers of metallic kernels is altered without destroying the crystal. The metallic lattice gets deformed but the environment of kernels does not change and remains the same as before. The deforming forces simply move the kernels from one lattice site to another.

Fig: 6.20 - Displacement of metal kernels in a metallic lattice

High tensile strength

Metals have high tensile strength. Metals can resist stretching without breaking. A strong electrostatic attraction between the positively charged kernels and the mobile electrons surrounding them is the reason for tensile strength.

Hardness of metals

The hardness of metals is due to the strength of the metallic bond. In general, the strength of a metallic bond depends upon:

• The greater the number of valence electrons for delocalisation the stronger is the metallic bond.

• Smaller the size of the kernel of metal atom, greater is the attraction for the delocalised electrons. Consequently, stronger is the metallic bond.

For example, alkali metals have only one valence electron and larger atomic kernels, which makes the metallic bonds weak. Consequently these metals are soft metals.

Opaqueness

The light that falls on metals is either reflected or completely absorbed by the delocalised electrons. Because of this, no light is able to pass through metals and they are termed as opaque.

Melting and boiling points

Metals have metallic bond strengths, which is intermediate to that of covalent and ionic bonds. Therefore in general, metals have boiling and melting points in between to that of covalent and ionic compounds.

Question Chemical Bonding

(1) What does Gilbert Lewis's theory of chemical bonding state?

(a) A chemical bond is formed when an atom's electrons interact with each other (b) All of the choices (c) Atoms bond together to fill their valence shells (d) The most stable configuration for many elements is one that contains eight valence electrons

(2) What are the two main types of chemical bonds?

(a) ionic bonding and covalent bonding (b) nonpolar covalent bonding and covalent bonding (c) polar bonding and covalent bonding (d) simple bonding and ionic bonding

(3) Which particles play the most active role in chemical bonding?

(a) electrons (b) neutrons (c) protons (d) valence electrons

(4) An ionic bond is formed when electrons are:

(a) completely destroyed (b) completely transferred (c) divided (d) equally shared

(5) Due to the fact that Ionic compounds have strong intermolecular forces they are ____________ at room temperature.

(a) bonded covalently (b) gases (c) liquids (d) solids

(6) Ionic bonds form between which elements in the periodic table?

(a) metals and metalloids (b) metals and nonmetals (c) noble gases and metalloids (d) noble gases and nonmetals

(7) A covalent bond is formed when electrons are:

(a) destroyed (b) shared (c) split (d) transferred

(8) Covalent bonds form between which elements in the periodic table?

(a) a metal and metalloid (b) two metalloids (c) two metals (d) two nonmetals

(9) In a polar covalent bonding, electrons are:

(a) completely transferred (b) destroyed (c) equally shared (d) unequally shared

(10) In a nonpolar covalent bonding, electrons are:

(a) completely transferred (b) destroyed (c) equally shared (d) unequally shared

(11) What kind of bonding occurs in the water molecule?

(a) covalent bond (b) ionic bond (c) nonpolar covalent bond (d) polar covalent bond

(12) What type of bonding exists in the N2 molecule?

(a) covalent-ionic (b) ionic bond (c) nonpolar covalent bond (d) polar covalent bond

(13) When a partial electrical charge exists across a molecule, it is called a(n):

	(a) dipole
	(b) ion
	(c) ionic compound
	(d) nonpolar bond

Chemical Bonding

by Anthony Carpi, Ph.D.

Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because atoms can react with one another to form new substances called compounds (see our Chemical Reactions module). Formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.
Let's look at an example.  The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet.  The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I.  When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!

enlarge image

In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.
While some of Lewis' predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding. We now know that there are two main types of chemical bonding; ionic bonding and covalent bonding.

Ionic bonding

In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond.
For example, during the reaction of sodium with chlorine:

sodium (on the left) loses its one valence electron to chlorine (on the right),

resulting in

a positively charged sodium ion (left) and a negatively charged chlorine ion (right).

The reaction of sodium with chlorine

Concept simulation - Reenacts the reaction of sodium with chlorine.

(Flash required)

Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After the reaction takes place, the charged Na⁺ and Cl^- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common:

• Ionic bonds form between metals and nonmetals.
• In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride).
• Ionic compounds dissolve easily in water and other polar solvents.
• In solution, ionic compounds easily conduct electricity.
• Ionic compounds tend to form crystalline solids with high melting temperatures.

This last feature, the fact that ionic compounds are solids, results from the intermolecular forces (forces between molecules) in ionic solids. If we consider a solid crystal of sodium chloride, the solid is made up of many positively charged sodium ions (pictured below as small gray spheres) and an equal number of negatively charged chlorine ions (green spheres). Due to the interaction of the charged ions, the sodium and chlorine ions are arranged in an alternating fashion as demonstrated in the schematic. Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single molecule does not apply to ionic crystals because the solid exists as one continuous system. Ionic solids form crystals with high melting points because of the strong forces between neighboring ions.

	Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1
Sodium Chloride Crystal	NaCl Crystal Schematic

Covalent bonding

The second major type of atomic bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons. Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H₂. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.
Covalent bonding between hydrogen atoms

Concept simulation - Recreates covalent bonding between hydrogen atoms.

(Flash required)

Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed.  Thus covalent molecules are not  strongly attracted to one another.  As a result, covalent molecules move about freely and tend to exist as liquids or gases at room temperature. 
Multiple Bonds: For every pair of electrons shared between two atoms, a single covalent bond is formed.  Some atoms can share multiple pairs of electrons, forming multiple covalent bonds.  For example, oxygen (which has six valence electrons) needs two electrons to complete its valence shell.  When two oxygen atoms form the compound O₂, they share two pairs of electrons, forming two covalent bonds. 
Lewis Dot Structures: Lewis dot structures are a shorthand to represent the valence electrons of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons. The Lewis structures for the elements in the first two periods of the periodic table are shown below.

	Lewis Dot Structures

Lewis structures can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H₂ and O₂ are shown below.

H2	H:H	or	H-H
O2

Polar and nonpolar covalent bonding

There are, in fact, two subtypes of covalent bonds. The H₂ molecule is a good example of the first type of covalent bond, the nonpolar bond. Because both atoms in the H₂ molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed.
A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.

H2O: a water molecule

Water molecules contain two hydrogen atoms (pictured in red) bonded to one oxygen atom (blue). Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells.
Polar covalent bonding simulated in water
The primary difference between the H-O bond in water and the H-H bond is the degree of electron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond.  

The dipole

Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it called a dipole. The water dipole is represented by the arrow in the pop-up animation (above) in which the head of the arrow points toward the electron dense (negative) end of the dipole and the cross resides near the electron poor (positive) end of the molecule.