MARIA JONALYN E.VALENCIA MAED SCIENCE are formed by the joining of two or more atoms. A stable compound occurs when the total energy of the combination has lower energy than the separated atoms. The bound state implies a net attractive force between the atoms ... a chemical bond. Covalent bond: bond in which one or more pairs of electrons are shared by two atoms. Ionic bond: bond in which one or more electrons from one atom are removed and attached to another atom, resulting in positive and negative ions which attract each other. Covalent chemical bonds involve the sharing of a pair of valence electrons by two atoms, in contrast to the transfer of electrons in ionic bonds. Such bonds lead to stable molecules if they share electrons in such a way as to create a noble gas configuration for each atom. Hydrogen gas forms the simplest covalent bond in the diatomic hydrogen molecule. The halogens such as chlorine also exist as diatomic gases by forming covalent bonds. The nitrogen and oxygen which makes up the bulk of the atmosphere also exhibits covalent bonding in forming diatomic molecules. Covalent bonds in which the sharing of the electron pair is unequal, with the electrons spending more time around the more nonmetallic atom, are called polar covalent bonds. Atoms can either transfer or share their valence electrons. In the extreme case where one or more atoms lose electrons and other atoms gain them in order to produce a noble gas electron configuration of ionic bonds are those in the alkali halides such as sodium chloride, NaCl. Typical The properties of metals suggest that their atoms possess strong bonds, yet the ease of conduction of heat and electricity suggest that electrons can move freely in all directions in a metal. The general observations give rise to a picture of "positive ions in a sea of electrons" to describe metallic bonding. NaCl This is the formation of an ionic bond. Na + Cl - electron transfer and the formation of ions Cl2 This is the formation of a covalent bond. Cl sharing of a pair of electrons and the formation of molecules Cl Step 1 Determine the total number of valence electrons for the molecule or polyatomic ion. (Remember that the number of valence electrons for a representative element is equal to its group number, using the A-group convention for numbering groups. For example, chlorine, Cl, is in group 7A, so it has seven valence electrons. Hydrogen has one valence electron.) For uncharged molecules, the total number of valence electrons is the sum of the valence electrons of each atom. For polyatomic cations, the total number of valence electrons is the sum of the valence electrons for each atom minus the charge. For polyatomic anions, the total number of valence electrons is the sum of the valence electrons for each atom plus the charge. Step 2 Draw a reasonable skeletal structure, using single bonds to join all the atoms. One or more of the following guidelines might help with this step. Try to arrange the atoms to yield the most typical number of bonds for each atom. Apply the following guidelines in deciding what element belongs in the center of your structure. Hydrogen and fluorine atoms are never in the center. Oxygen atoms are rarely in the center. The element with the fewest atoms in the formula is often in the center. The atom that is capable of making the most bonds is often in the center. Oxygen atoms rarely bond to other oxygen atoms. The molecular formula often reflects the molecular structure. Carbon atoms commonly bond to other carbon atoms. o Step 3 Subtract two electrons from the total for each of the single bonds (lines) described in Step 2 above. This tells us the number of electrons that still need to be distributed. Step 4 Try to distribute the remaining electrons as lone pairs to obtain a total of eight electrons around each atom except hydrogen and boron. We saw in Chapter 3 that the atoms in reasonable Lewis structures are often surrounded by an octet of electrons. The following are some helpful observations pertaining to octets. In a reasonable Lewis structure, carbon, nitrogen, oxygen, and fluorine always have eight electrons around them. Hydrogen will always have a total of two electrons from its one bond. Boron can have fewer than eight electrons but never more than eight. The nonmetallic elements in periods beyond the second period (P, S, Cl, Se, Br, and I) usually have eight electrons around them, but they can have more. The bonding properties of the metalloids arsenic, As, and tellurium, Te, are similar to those of phosphorus, P, and sulfur, S, so they usually have eight electrons around them but can have more. Step 5 Do one of the following. If in Step 4 you were able to obtain an octet of electrons around each atom other than hydrogen and boron, and if you used all of the remaining valence electrons, go to step 6. If you have electrons remaining after each of the atoms other than hydrogen and boron have their octet, you can put more than eight electrons around elements in periods beyond the second period. If you do not have enough electrons to obtain octets of electrons around each atom (other than hydrogen and boron), convert one lone pair into a multiple bond for each two electrons that you are short. If you would need two more electrons to get octets, convert one lone pair in your structure to a second bond between two atoms. If you would need four more electrons to get octets, convert two lone pairs into bonds. This could mean creating a second bond in two different places or creating a triple bond in one place. If you would need six more electrons to get octets, convert three lone pairs into bonds. Etc. Step 6 Check your structure to see if all of the atoms have their most common bonding pattern If each atom has its most common bonding pattern, your structure is a reasonable structure. Skip Step 7. If one or more atoms are without their most common bonding pattern, continue to Step 7. Step 7 If necessary, try to rearrange your structure to give each atom its most common bonding pattern. One way to do this is to return to Step 2 and try another skeleton. (This step is unnecessary if all of the atoms in your structure have their most common bonding pattern.) Recall that the valence electrons for the elements can be determined based on the elements position on the periodic table. Lewis Dot Symbol Lewis dot structures show the valence electrons around at atom and for most molecules and compounds a complete octet for the elements N most Al monatomic ions have an electron configuration of noble gases 1s22s22p5 F + e- F 1s22s22p6 Ne Electronegativity The electronegativity difference - DEN = ENhigher – EN lower Chlorides of Period 2 compound LiCl BeCl2 BCl3 CCl4 NCl OCl2 DEN Compound 2.2 1.6 1.1 0.6 0 3 Cl2 0.6 0 Chlorides of Period 3 NaC MgCl AlCl SiCl PCl3 SCl6 2 3 4 l 2.2 1.9 1.6 1.3 1.0 0.6 Cl2 0 DEN large difference small difference Using electronegativities to determine bond type DEN > 1.7 ionic bond - transfer DEN < 1.7 covalent bond - sharing So we have a range of electronegativity difference of 0 to 1.7 for sharing an electron pair. Is the sharing of electrons in molecules always equal? non-polar bond Which element is more electronegative? X X X Y increasing polarity of bond DEN = 0 Y Y DEN = 0.3 ENY > ENX DEN = 0.6 polar bond 0 < EN < 1.7 X X Y Y DEN = 0.9 DEN = 1.2 Direction of electron migration More sharing examples O2 O O Share until octet is complete. O O O O double bond (2 pairs) N2 N N NN N N octet complete N N triple bond (3 pairs) Is breaking a bond an endothermic or exothermic process? X2 + energy F2 O2 X + X increasing bond strength single bond BE = 142 kJ/mole double bond BE = 494 N2 triple bond BE = 942 http://wulfenite.fandm.edu/Data%20/Table_6.html Some more sharing examples NH3 H N H H normal covalent bond (each atom supplies an electron) NH4+ NH3 + H+ NH4+ H+ H N H H coordinate covalent bond (the pair of electrons from the same atom) Using the EN trends to predict bond type Increasing EN 105 107 Db Bh NO RbF FeS H2S Modified from http://www.cem.msu.edu/~djm/cem384/ptable.html Increasing EN Comparison of Bonding Types ionic ions molten salts conductive transfer of electrons high mp DEN > 1.7 covalent molecules nonconductive valence electrons sharing of electrons low mp DEN < 1.7 100% covalent 100% ionic Bonding spectrum A B A B A + B - Increasing DEN Increasing polarity Transfer