science by boysemperor

Ionic Bonds

What is an ionic bond?

An ionic bond is the force of attraction between the opposite charges of an ion. One element in an ionic bond loses electrons, and another element must gain the electrons. Some atoms lose electrons to make the outside energy levels become more stable. Atoms become more stable when their outer most energy level has 8 electrons. Pure ionic compounds usually are crystalline solids, liquids, or gases. Many ionic compounds are binary compounds.

This shows an atom losing an electron.

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Ionic compounds usually have much higher melting and boiling points than covalent compounds.

Ionic Compounds:

Many ionic compounds dissolve easily in water. The human body must keep a precise amount of ions in order to function properly, these ions are called electrolytes. Without the right concentration of electrolytes your nerve impulses can't travel to your brain.

*When you sweat you lose electrolytes. Athletes drink certain drinks to keep the electrolytes balanced*

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FORMULAS

A good way of showing information about the atoms that constitutes a particular chemical compound is in a table. Formulas identify types of chemical elements by its chemical symbol. A chemical formula tells what elements a compound contains and the exact number of atoms of each element in a unit of that compound.

Oxidation Numbers

The oxidation number tells you how many electrons an atom has gained, lost, or shared to become stable. The sum of oxidation numbers of all the atoms in the compound must equal zero. Oxidation numbers may be used to determine the most likely formulas for binary compounds. Oxidation number may be criss crossed to create subscripts in the formula.

example: Aluminum has an oxidation number of +3 and oxygen has an oxidation number of -2. The formula for Aluminum Oxide is Al₂O₃.

The oxidation numbers on Aluminum (Al) and Oxygen (O) are switched, and written as subscripts on the formula for Aluminum Oxide.

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QUIZ-_

1. One that is composed of two elements.

   1.   Binary compound
   2.   Chemical bond
   3.   Hydrate
   4.   Ion

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!

	+
sodium metal		chlorine gas		table salt

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).

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 becomes blurred in ionic crystals because the solid exists as one continuous system. Forces between molecules are comparable to the forces within the molecule, and ionic compounds tend to form crystal solids with high melting points as a result.

	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.

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.

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 hydr Next post: Clostridium Previous post: Amoeba (genus) « Back to posts

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