Chapter 8: Covalent Bonding

1. Chapter 8: Covalent Bonding • Matter takes many forms in nature: In this chapter, we are going to learn to distinguish the type of compound that we have already studied, the “ionic compound” (which contains oppositely-charged particles: metal cations and non-metal anions), from a different type of compound – a “molecular compound”. Additionally, we are going to focus on a type of molecular compound known as a binary molecular compound

2. II. Binary compounds: • A “binary” compound contains atoms from two different elements. • A. NaCl”, “CaF2”, and “Al2O3” (3 ionic compounds) are binary ionic compounds. “NH4Cl” is an ionic compound, but because it contains more than 2 elements, it is not a binary ionic compound.

3. B. “N2O5”, “SF6”, and H2O” (3 molecular compounds) are binary molecular compounds. • “C6H12O6” is a molecular compound, but because it contains more than 2 elements it is not a binary molecular compound.

4. Comparison/Contrast between an ionic compound and a molecular substance. • A. Molecular substances are made of molecules. • 1. There is no “molecule” in an ionic compound.

5. B. A “molecule” contains a specific number of atoms, connected in a specific manner, to give a specific shape. • If even one atom is “missing” or “different”, the molecule would be an entirely different substance.

6. 1. Not so with an ionic compound: In an ionic compound there is a specific ratio of atoms.. In salt (NaCl) for example, there is a ratio of 1 Na for every 1 Cl. • If a clump of salt lost 1 Na and 1 Cl, it would still be the same original substance: NaCl

7. C. The formula of a “molecule” should never be simplified. C2H8 is not the same substance as CH4. • 1. The formula of an ionic compound should always be simplified. Ba2O2 is the same substance as BaO. • D. A molecule will not crack apart. • 1. Ionic compounds can crack apart if hammered….. If the cations come close too close together and the anions come too close together, the structure cracks apart.

8. E. Molecular substances may have low melting points. • 1. All ionic compounds have a very high melting point. • F. Molecular substances may, at room temperature, be found as solids, liquids, or gases. • 1. All ionic compounds are solids (at room temperature).

9. G. Molecular substances contain atoms which are held together by covalent bonds. • 1. Ionic substances are held together by ionic bonds.

10. IV. Covalent bonding: • The type of bonding that occurs within a molecular substance, in which atoms share their valence electrons in order to become more stable. • A. Occurs between atoms of nonmetallic elements.

11. B. Not all “molecules” or “molecular substances” are compounds! • In addition to the binary molecular compounds that we will study, there are 7 nonmetallic elements found in nature (in their elemental form) as pairs of atoms. These are the 7 “diatomic” elements: • N2, O2, F2, Cl2, Br2, I2, H2.

12. V. An important review: • A. Metallic elements: Found to the left side of the staircase boundary on the periodic table. • Non-metallic elements: • Elements found to the right side of the staircase boundary on the periodic table. • 2. Hydrogen is a nonmetallic element also.

13. VI. The octet rule: • When a molecule is formed: “Nonmetal atoms share electrons in covalent bonds in order to obtain a full octet of electrons.” An octet = 8 valence electrons. • A. Exception: A hydrogen atom will end up with a total of two electrons by sharing with 1 other atom.

14. B. There are a few other notable exceptions to the octet rule: • 1. A few molecular compounds which contain an odd number of valence electrons are known to exist. • 2. A few molecular compounds have either a boron or an aluminum atom with 6 valence electrons.

15. 2. A few molecular compounds have a central atom with 10 or 12 valence electrons. • (1) One common example is “sulfur hexafluoride”. • In this compound, the central sulfur atom contains 6 x 2 = 12 valence electrons. Be sure to remember that this compound is an example in which thecentral atom does not follow the octet rule.

16. VII. Types of covalent bonds. • A. Single covalent bond – 1 shared pair of valence electrons:2 dots, or a single dash, represent 2 electrons that are simultaneously being attracted by, or “shared” by, the nuclei of two neighboring atoms.

17. 1. The formula in the center is a type of structural formula called a “Lewis dot structural formula”. • The formula on the right is the molecular formula. H – H H H H2

18. B. Double covalent bond • – two pairs of shared valence electrons: 4 dots or 2 parallel dashes. C C C C C2H4 H H H H H H H H

19. C. Triple covalent bond– three pairs of shared valence electrons: 6 dots or 3 parallel dashes. N N or N N N2 • Notice the two “unshared pairs” of electrons (one pair is to the far left and one pair to the far righ)t of the nitrogen structure. You may never use a long dash to represent an unshared pair of electrons.

20. Unshared pairs of electrons don’t bond the atoms together….but, the repulsive forces of unshared pairs of electrons do dramatically influence the shape of a molecule!

21. D. Notice how an ion can react with a molecule to generate a polyatomic ion. In the example below, a hydrogen ion bonds to a molecule of ammonia(NH3) to make the ammonium ion (NH4)+: H + + N H H N H + H H ] [ H H

22. VIII. Drawing a Lewis Dot Structure: • A. Certain elements are known as “central” atoms…. They will be found in the center of a structure. The first element given in a formula is usually the central atom (exception: hydrogen and the halogens). • 1. Position the central atom in the center of your work space.

23. B. Hydrogen and the halogens are known as “peripheral” atoms. They will be found only connected to one other atom. • Position hydrogen and halogen atoms so that they “touch”, or “go around” only 1 other atom.

24. C. Add up all the valence electrons. Position the valence electrons as dots around the atom they belong to - the valence electrons may never leave the original atom. • Position the dots to form a “doorway” with 4 sides, in which the symbol of the element appears centered in the doorway. • Start with no more than 2 dots on each side of the 4 sided doorway.

25. D. If you can’t easily achieve a Lewis dot structure which has each atom (other than hydrogen) surrounded by 8 dots by doing what is described above, then you either need a double bond (2 pairs of shared electrons) or a triple bond (3 pairs of shared electrons). • For CO2, you will need two double bonds.

26. 1. To make a double bond, move one “un-shared electron” simultaneously from each of two neighboring atoms, and place those 2 electrons in between the two neighboring atoms. • 2. To make a triple bond, start with a double bonded pair of atoms, and simultaneously move one more unshared electron from each of the two atoms. Reposition those two electrons in between the atoms.

27. Important points regarding nonmetal atoms and their bonding charcteristics:

31. IX. Lewis Dot structural formulas for polyatomic ions: • A. Covalent bonds occur within a polyatomic ion (notbetween polyatomic ions). • B. When drawing polyatomic ions, place the first element in the center of the structure, and place the second element around the first element (placing 1 atom of the second element along each different side of the first element).

32. C. When the charge of a polyatomic ion is +, you need to subtract the indicated number of electrons from the total of the valence electrons in the molecule. So, for +1 ions: take away 1 electron from the molecular ion’s number of valence electrons. • D. When the charge of a polyatomic ion is –, you need to add the indicated number of electrons to the molecular ion’s number of valence electrons.

33. 1. If the charge is 1-, then add 1 more electron to the molecule’s total number of valence electrons. • 2. If the charge is 2-, then add 2 more electrons; if the charge is 3-, then add 3 more electrons. • E. Last, for a polyatomic ion: Draw a large bracket around the ion; and, place its charge at upper right.

34. # of N valence electrons: 5 5 # of H valence electrons: 4 x 1 = 4 4 Charge of ion = +1, therefore less 1 -1 Therefore, total = 8 Ammonium ion (NH4)+ H N H + H ] [ H

35. Steps for Dot Structures: • Step 1: total # valence electrons. • Step 2. Position central atoms: • carbon atoms form a straight line; • assume only single bonding.

36. Step 3. Position other atoms; remember “special” molecules. • A. Peripheral atoms: • Hydrogen and the halogens- • connect to only 1 other atom, • use only 1 single bond. • B. binary polyatomic ions: first element is central, second element is peripheral. assume all single bonds. • C. hydrocarbons” – molecular formula gives list of atoms (from left to right) connecting to each central atom (usually carbon) • CH3CH2OH means “first carbon touches 3 H atoms, second carbon touches 2 H atoms, then there is an O touching an H. Assume all single bonds. • D. Memorize: CO2 C in the middle; use two double bonds

37. Step 4: Make each atom stable. Work from left to right: Assume all single bonds. Position unshared pairs to provide octets. Exception: Hydrogen atoms = only 2 dots. Step 5: Count the dots you’ve used. Make sure the # you used = the # you were supposed to. Erase extras.

38. Step 6: Make corrections - • If your structure “needs” 2 extra dots, it really needs a double bond…. • Erase 2 unshared dots, and share them (as part of a double bond). • If your structure “needs” 4 extra dots, it really needs a triple bond..... • Erase 4 unshared dots, and share them (as part of a triple bond). • Every atom should now be stable.

39. X. VSEPR Theory – • Valence Shell Electron Pair Repulsion theory. [Remember: Like charges repel!] • A. A theory to predict the 3-dimensional geometry, ie. the“shape” of a molecule

40. 1. The theory is based on “electrostatic repulsion”: Molecules will adjust their shape to keep the negatively-charged pairs of valence electrons as far apart as possible from each other. • B. When NOT to use VSEPR theory: When there are only 2 atoms in a molecule. These molecule’s shapes are calledlinear – it doesn’t matter if there are single bonds, double bonds, triple bonds, or unshared electron pairs.

41. C. Using VSEPR theory: • 1. Draw the Lewis dot structure for the molecule. • 2. Identify its central atom. • 3. Identify the sets of valence electrons as one of two possibilities: • A. Those connecting two atoms. • B. Those that do not connect two atoms. These are called “unshared pairs”.

42. 4. The unshared pairs found on a central atom strongly repel each other; and molecules that would otherwise be linear, will be forced into a bent (or angular) shape. • 5. Unshared pairs also cause a molecule that would be shaped like a flat triangle (trigonal planar), to be forced into a not flat ( trigonal pyramidal) shape.

43. 6.Count the number of connections separately from the number of unshared pairs. • 1 single bond counts as 1 connection. • 1 double bond counts as 1 connection. • 1 triple bond counts as 1 connection. • Each unshared set of 2 dots counts as 1 un-shared pair.

44. D Predicting Shapes Using VSEPR Table • Read horizontally across the table.

45. Linear diatomic E. Shap e s: Trigonal Planar Bent Linear diatomic Linear triatomic Tetrahedral Pyramidal Shapes: , Trigonal planar bent Linear triatomic Trigonal pyramidal tetrahedral

46. “Molecular Polarity” • – A term that is used to distinguish two types of molecules…. Based on the presence or absence of a separation of charge. • Some molecules showcharacteristics indicating that they have oppositely-charged ends (a positive end and a negative end). This is called a separation of the charges (or “separation of charge”).

47. Other molecules show characteristics indicating that their structure doesn’t have a separation of charge, or their structurehides the presence of their oppositely-charged ends.

48. How to determine a molecule’s polarity. • The first part of determining a molecule’s polarity is to calculate each individual bond’s polarity. • Be careful with the vocabulary being used – • An individual bond’s polarity is called the “bond polarity” • The polarity of the entire molecule is called the “molecular polarity”

49. To calculate a bond polarity, first identify the “electronegativity value” of each of the 2 atoms in the bond you are working on. • The electronegativity value is number (from 0 to 4) which informs us of an atom’s ability to attract electrons when in a compound.

50. The electronegativity value is given on your periodic table, side 2, within each element’s square…..upper right corner of the square, in black print. • The closer an element’s electronegativity is to “4, the better that an atom of that element will attract electrons when that atom is found in a compound. • The closer an element’s electronegativity is to “0”, the less likely it is for that atom to be able to attract electrons when in a compound.