What holds an atom together? What are the forces that make a liquid remain a liquid, a solid remain a solid, and a gas a gas? In this unit we will explore the forces that keep atoms together and the properties of these bonds.
The way atoms are bonded can help determine the molecular formula of a compound from the bonds between those elements. However, a chemical analysis of a compound is the only sure way of knowing the formula. Bonding does help explain why certain compounds form, and thereby reinforces the law of multiple proportions (e.g. H2O, H2O2, and CO, CO2).
The shape of molecules can contribute to the type of bonds they form. If a molecule has the wrong shape, if may not bond with other molecule. Similarly, if a key has the wrong shape, it cannot fit into a lock. Therefore, a good knowledge of how a substance bonds can help explain macroscopic observations such as the digestion of food, oil in water, and food production in green plants.
When a bond is formed, or a bond is broken, bond energy is involved. A strong bond may require more energy to break than a weak bond. This can lead to strong molecules as in a solid, and to weak molecules as in a gas. Knowing the nature of the bond, and the bond energy involved, can help you better characterize the nature of a molecule.
What happens when two atoms approach each other? We know that each atom has protons with positive charges and electrons with negative charges. As two atoms approach each other, a proton of one atom attracts an electron of the other, and vice versa, until the atoms are joined. The attractive forces of the atoms play an important role on each other (e.g. a postive charge attracts a negative charge, a negative charge attracts a positive charge). The result is a chemical bond in which two atoms join together by the above processes.
As the two atoms draw near each other, their attractive forces pull them together. In some molecules, the repulsion of the atom's electrons began to repel each other at a certain distance. At this point, the proton of one atom has drawn the electron of the other atom as close as possible before the repulsion forces act. Two atoms of similar properties can form a bond in such a way that they are actually sharing electrons. This type of bond is called a covalent bond. The two atoms would count each electron in the pair as one of their own. For example, in a H2 molecule, each atom would claim to have two electrons at a given point in time, even though each atom had only one electron before the bond formed. In essence, a pair of electrons is being shared between the two atoms involved in the covalent bond.
In the last paragraph, we talked about covalent bonds being formed between two atoms will similar properties. In most cases, atoms in a molecule may not have similar properties. One atom may tend to have more of an attractive pull than other. As in a tug of war, one side would end up with the piece of rope. In a ionic bond, that piece of rope would be electrons. When two atoms with unequal properties form, one atom would tend to grab all the electrons in a molecule, leaving the other with none. The difference in the attractive properties of the atoms is called electronegativity. An atom with the greater electronegativity would tend to grab all of the electrons in an ionic compound. For example, in NaCl, the chlorine (Cl) has a greater electronegativity than sodium (Na), so chlorine would take all the valence electrons (outer electons) of sodium in the compound. In the fact, in a compound with an electronegativity difference of 1.7 or more, the compound is most likely ionic.
NaCl also forms a crystal due to its ionic nature. In an NaCl molecule, each sodium atom loses one electron to become a 1+ ion, and each chlorine atom gains one electron to become a -1 ion. A crystal forms between sodium ions and chloride ions because each sodium ion draws six chloride ions and each chloride ion draws six sodium ions. The crystal continues to form releasing energy as long as equal numbers of sodium and chlorine atoms are present. The three dimesional bonding of negative and postive ions is called a crystal lattice structure.
In a polar molecule, the electrons are shared unequally. One atom has a greater electronegativity than the other, resulting in one end of the molecule having more of the electron pair. The atom with the higher electronegativity pulls on the electrons more, and thus it has a partial negative charge. This also means that the other atom, that has the lower electronegativity, has a slight positive charge. This distrubition of charge among the poles (ends) of the molecule make up a dipole molecule. For example in a HCL molecule, hydrogen has a 2.1 electronegativity and chlorine has a 3.0 electronegativity. The difference in electronegativity is 0.9. Chlorine would take up a partial negative charge (between 0 and 1-) and hydrogen would have a partial postive charge (between 0 and 1+).
Water (H2O), also behaves in the same manner. That is why dipole and polar properties are associated with H2O. Other molecules that exhibit dipole properties include SO2, and NH3. However, the only sure way to find out if a compound is polar, is to carry out a chemical analysis of the compound.
In most cases, you can use a table of electronegativity values to determine the type of bonds that will occur between atoms. A simple of rule of thumb goes as follows:
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Unit 3 - Section 2