There are two major types of chemical bonds, ionic and covalent.

2. There are two major types of chemical bonds, ionic and covalent. • Ionic Bond: A strong electrostatic attraction between a positive ion and a negative ion. One or more electrons have been transferred from the valence shell of one atom to the valence shell of the other atom. • Ionic bonds generally occur between metals and nonmetals. • Covalent Bond: Electrons are shared between atoms. • Covalent bonds generally occur between nonmetals.

3. One of the major factors determining whether an ionic bond will form is the ionization energy of the metal atom. • Ionization energy is the energy required to remove a negatively charged electron from an atom, leaving behind a positively charged cation. (Unlike charges attract.) • In general, metals have low ionization energies and nonmetals have high ionization energies.

4. The ionization energy (IE) of the elements shows two general trends: • IE increases as you move from left to right across the periodic table. • IE decreases as you move from top to bottom down the periodic table. • Elements having the largest ionization energy are located at the upper right corner of the periodic table, and those with the smallest ionization energy at the bottom left corner. • The increase in IE can be explained by the fact that the positive nuclear charge is increasing as we move from left to right across the periodic table. The additional valence electrons are all being added at about the same distance from the nucleus and, as a group, see increasing nuclear charge. • The decrease in IE from top to bottom can be explained by the fact that the valence electrons in each row are farther and farther away from the nucleus. The remaining, or “core,” electrons lie between the nucleus and the valence electrons and help to balance the increasing nuclear charge.

6. In Chapter 2 we saw how a single electron was transferred from a sodium atom to a chlorine atom giving each the electronic configuration of an ideal gas. • We will now extend this idea to the group II and group III metals of the periodic table.

7. Lewis Symbols and Formulas of Ionic Compounds Li+ [He]+ 1s2 Mg+2 [Ne]+2 1s22s22p6 Al+3 [Ne]+3 1s22s22p6 F- [Ne]+3 1s22s22p6 • Fluorine accepts only one electron to achieve its octet or ideal gas configuration. • Two fluorine atoms are thus required to accept the two electrons that magnesium must lose to achieve its octet. • Using this information, the chemical formula of magnesium fluoride can be predicted to be: MgF2. Aluminum fluoride must be: AlF3.

8. Formulas of Ionic Compounds • The formulas of all binary ionic compounds between the group I, II, and III metals, and the group V, VI, and VII nonmetals can be predicted using these principles. • Group I metals lose 1 electron to form +1 ions. • Group II metals lose 2 electrons to form +2 ions. • Group III metals lose 3 electrons to form +3 ions. • Group V nonmetals gain 3 electrons to form -3 ions. • Group VI nonmetals gain 2 electrons to form -2 ions. • Group VII nonmetals gain 1 electron to form -1 ions.

9. When cations and anions combine to form a binary salt or acid, they always do so in such a way that the net charge of the substance formed is zero. • The easiest way to apply this principle is the cross over approach, as illustrated for aluminum oxide: • The magnitude of the charge on each ion is used as the subscript for the other ion in the formula for the substance. NaCl Na2O Na3N CaCl2 CaO Ca3N2 AlCl3 Al2O3 AlN

10. In the name of a binary ionic compound, the cation is always given first and the anion second. • The cation is always specified as the name of the metal, unchanged. • The anion is specified by using the first part of the name of the nonmetal and then adding the suffix “ide”. • Sodium chloride (NaCl), aluminum oxide (Al2O3), magnesium sulfide (MgS). • This simple naming convention is sufficient for the group I, II, and III cations.

11. Transition metals in period 4 and higher can often have several stable positive ions (Fe2+ and Fe3+, Cu1+ and Cu2+). • For this reason, names such as iron chloride and copper chloride are ambiguous. • The modern system for naming the salts of these types of metal ions indicates the charge on the cation as a roman numeral within the name of the substance: • Iron (II) chloride FeCl2 Copper (I) chloride CuCl • Iron (III) chloride FeCl3 Copper (II) chloride CuCl2 • An older system, still in common use, uses the suffixes –ous and –ic to indicate the less positive and more positive cations: • Ferrous chloride FeCl2 Cuprous chloride CuCl • Ferric chloride FeCl3 Cupric chloride CuCl2

13. A polyatomic ion is two or more atoms chemically combined into a single charged unit:

14. The method for constructing the formula from the name of a compound containing a polyatomic ion is the same as if it were a simple monatomic ion. • If two or more instances of the polyatomic ion are required in the formula, surround the polyatomic ion with parentheses and then use a subscript as usual: • Aluminum carbonate Al2(CO3)3 Al3+ and CO32- • Calcium phosphate Ca3(PO4)2 Ca2+ and PO43- • Ammonium hydroxide NH4OH NH4+ and OH-

15. Ionic solids do not conduct electricity; neither does pure water. However, • when ionic solids, such as sodium chloride, are dissolved in water or melted, the solutions and the melted solids do conduct electricity.

16. Substances which conduct electricity in this way are called electrolytes. • Substances which do not conduct electricity when dissolved in water, such as glucose, are called nonelectrolytes. • In their solid form, the ions of an ionic compound are locked into place and all their charges are neutralized. Only when the solid melts or is dissolved in water, thereby freeing the individual ions, can the ions carry an electric current by moving to an electrode of opposite charge. Ionic solids consist of three-dimensional arrays of positive ions surrounded by negative ions and negative ions surrounded by positive ions. This arrangement makes it impossible to identify a discrete structural unit that can be represented by the formula of the solid (Figure 3.4). There is no such structural unit as a molecule for an ionic compound.

18. The formulas of compounds such as water (H2O), carbon dioxide (CO2), and oxygen (O2) represent not only the combining ratios of component elements, but also the actual, ultimate structural units that we call molecules. • Molecules consist of atoms connected by covalent bonds.

21. The systematic method for naming simple, binary molecular compounds is based on the fact that some pairs of elements can form more than one covalently bonded compound. • Common examples are: • CO and CO2 • N2O, NO, NO2, N2O3, N2O4, and N2O5

22. The first element in the formula is named first, followed by the second element using the –ide ending, as in the nomenclature for salts. • Greek numeric prefixes indicate the numbers of each type of element (the prefix mono- is never used in naming the first element)

24. Electron dot structures are used to denote covalent bonds between atoms as well as the remaining valence electrons on each atom. • The resultant diagrams are called Lewis structures. The key concept in constructing Lewis structures is the octet rule. • The number of chemical bonds an atom forms to complete its octet is called its combining power. • Group IV elements 4 additional electrons to complete its octet • Group V elements 3 additional electrons to complete its octet • Group VI elements 2 additional electrons to complete its octet • Group VII elements 1 additional electron to complete its octet

25. Each chemical bond involves two electrons, one donated from each atom at the two ends of the bond. The number of additional electrons to complete an atom’s octet is simply the number of bonds the atom normally makes, or its combining power. • Group IV elements 4 additional electrons or 4 bonds • Group V elements 3 additional electrons or 3 bonds • Group VI elements 2 additional electrons or 2 bonds • Group VII elements 1 additional electron or 1 bond

26. Two fluorine atoms each achieve an octet by sharing two electrons between themselves, one from each atom. • Lewis structures are usually simplified by replacing each pair of bonding electrons with a line. • Lewis structures drawn in this way distinguish between bonding electrons (lines) and non-bonding electrons (pairs of dots). • A further simplification omits the non-bonding electrons to form what is called a structural formula.

27. There are several exceptions to the octet rule: • H forms a duet rather than an octet. Hydrogen can never form two bonds, thus it is always on the outer edge of a molecule. • Be only has two electrons to share and can thus only make two bonds to other atoms. The Lewis structures of its compounds only contain 4 electrons about the Be atom. • B only has three electrons to share and can thus only make three bonds to other atoms. The Lewis structures of its compounds only contain 6 electrons about the Be atom. • Molecules formed from atoms whose valence electrons add up to an odd number (i.e. NO or NO2) cannot have eight electrons (an even number) about every atom.

28. To derive a Lewis structure for a compound requires a knowledge of the compound’s molecular formula and a set of simple rules: • Place the symbols for the bonded atoms into an arrangement that will allow you to begin distributing electrons. • Determine the total number of valence electrons in the molecule. • Represent shared pairs of electrons by drawing lines between bonded atoms. • Position any additional electrons as lone pairs to satisfy the octet rule for each atom (duet rule for hydrogen).

29. Construct the Lewis structure for ammonia, NH3.

30. Lewis Structures Containing Multiple Bonds • In some cases it appears that there are too few electrons to satisfy the octet rule for every atom in a structure (C2H4, C2H2, N2, etc.). • In these cases, pairs of atoms are connected by double bonds (sharing 4 electrons, 2 from each atom) or triple bonds (sharing 6 electrons, 3 from each atom).

31. Write the Lewis structure for ethylene, C2H4. Two electrons remain unused.

32. Write the Lewis structure for nitrogen, N2.

33. When dealing with ions, the total number of valence electrons placed on a Lewis structure must be adjusted for the charge on the ion. • For positive ions, subtract the magnitude of the charge from the valence electron count. • For negative ions, add the magnitude of the charge to the valence electron count.

35. Most of the polyatomic ions we will consider arise from reactions taking place in aqueous solutions: • In each of the reactions above, a proton (H+) is transferred from an electron pair (bonding pair) on one molecule to an electron pair (lone pair) on another molecule. • When both electrons in a bond are donated by only one of the atoms (as in the hydronium ion and ammonium ion above), the bond is called a coordinate covalent bond.

36. When both electrons in a bond are attracted equally to either atom, as in H2, F2, O2, N2, or Cl2, the bond is called a nonpolar covalent bond. The electrons lie more or less equally between the two atoms. • When bonds between different types of atoms form, one atom often attracts the electrons more strongly than the other atom. • In this case, the electrons will lie closer to the atom having the stronger attraction. Bonds of this type are called polar covalent bonds.

37. In a nonpolar covalent bond, the charge on either atom will be zero. The two electrons are distributed equally between the two atoms. • In an ionic bond, the charge on one atom is +1 and the charge on the other atom is -1. The two electrons are “owned” by the atom having the -1 charge. • A polar covalent bond has properties somewhere between these two extremes.

38. The ability of an atom within a molecule to draw electrons to itself is called its electronegativity: • The most electronegative elements are at the upper right of the periodic table and the least electronegative at the lower left.

39. The electronegativity difference between two atoms determines whether the bond will be ionic, nonpolar covalent, or polar covalent: • Bonds with differences greater than 2.0 are usually ionic. • Bonds with differences less than 1.5 are usually polar covalent. • Bonds with differences between 1.5 and 2.0 will be ionic if the compound is a metal-nonmetal combination, and polar covalent if it is a nonmetal-nonmetal compound. • Bonds with a difference of 0 will be nonpolar covalent.

40. In polar covalent bonds, the charges on the two atoms involved lie somewhere between 0 and 1. • Partial charges on atoms are often labeled with the symbols: + and -. • The HF molecule contains a polar covalent bond. One end of the molecule has positive character and one end has negative character. The molecule as a whole is electrically neutral, however.

41. A polar molecule is attracted to a second polar molecule, positive end to negative end. The force of the attraction is related to a quantity called the dipole moment of the molecule. • The dipole moment of a molecule depends upon the magnitude of the partial charges on the molecule and their distance of separation: the larger the charges (electronegativity difference) and the greater the distance, the larger the dipole moment. • Dipole moments are customarily represented on a molecule using the symbol pointing from the positive end towards the negative end of the molecule (less-electronegative element to more-electronegative element).

42. If a molecule contains two or more polar bonds, it is possible for the dipole moments of the polar bonds to reinforce each other or cancel each other out. • As an example, consider two molecules having the formula XY2, one linear and one bend: • Each of the X-Y bonds in each molecule is identical and will have identical dipole moments. • In the linear molecule, since the two dipoles point in exactly opposite directions, they will cancel each other out and the molecule will have a 0 net dipole moment. • In the bent molecule, parts of the two dipoles cancel (the horizontal parts), however the vertical parts of the two dipoles reinforce each other. This molecule has a non-zero net dipole moment.

43. An example of a linear molecule containing two dipole moments which exactly cancel each other out is O=C=O. • An example of a linear molecule containing two dipole moments which do not exactly cancel each other out is Cl-Be-F. Although the molecule is linear, the two halogens have different electronegativities and the magnitudes of the two dipoles are not equal. • An example of a bent molecule having a non-zero net dipole moment is water, H2O.

44. Molecules with large dipole moments tend to interact strongly with each other, as well as with other molecules possessing dipole moments. • This results in higher melting and boiling points and lower vapor pressures for polar molecules than for non-polar molecules. • Methane, CH4, has a zero dipole moment and is a gas at room temperature and pressure. • Water, H2O, has a large dipole moment and is a liquid at room temperature and pressure. Because of its polarity, water is an excellent solvent for polar substances.

45. Lewis structures indicate how the atoms within a molecule are connected, whether the bonds are single, double, or triple, and the locations of the extra, or non-bonding electrons. • However, by themselves, Lewis structures do not provide any information about the three-dimensional shape of the molecule.

46. VSEPR Theory • The valence-shell electron-pair repulsiontheory, or VSEPR, enables us to predict the three dimensional shape of a molecule from the completed Lewis structure of the molecule. • This theory is based upon the idea that all of the pair of electrons on a central atom in a molecule will repel each other and try to get as far apart as possible. • This includes both the bonding pairs of electrons and the non-bonding pairs on the atom.

47. We will first consider a series of molecules which possess no non-bonding electrons: • BeCl2 (beryllium chloride) is an electron-deficient compound possessing two bonding pairs of electrons and no non-bonding pairs. • The two pairs of bonding electrons will be as far apart as possible if the bonds point in exactly opposite directions (180o apart). • This molecule is linear. • AlCl3 (aluminum chloride) is an electron-deficient compound possessing three bonding pairs of electrons and no non-bonding pairs. • The three pairs of bonding electrons will be as far apart as possible if the bonds point at angles of 120o to each other. • This molecule is trigonal planar or triangular planar.

48. CH4 (methane) possesses four bonding pairs of electrons and no non-bonding pairs. • Four pairs of electrons can get as far apart in space as possible if the four bonds point an angles in space of 109.5o to each other. • The shape of the methane molecule is tetrahedral.

50. In each of these molecules, BeCl2, AlCl3, and CH4, the individual bonds are polar, however, because of the symmetrical nature of these molecules, the individual dipoles with each molecule cancel each other out. • The VSEPR predictions about the shapes of molecules have been experimentally confirmed using a technique called X-ray diffraction. In this technique, X-rays are passed through a crystal of a substance and the directions and intensities of the emerging X-rays are used to calculate the locations of all of the atoms within the crystal.