Chapter 6 Ionic Bonds and Compounds OK. Now we have a handle on some of the properties of individual atoms. Let s go the next step up and look are the properties of molecules. We will start with the general class of compounds called ionic compounds which are made from a metal combined with a non-metal 6-1 The ionic bond and electrostatic forces Ionic compounds are held together by ionic bonds Ionic bonds have some very interesting, almost contradictory properties Ionic solids are very strong. As seen by very high melting and boiling points o MP H2O = 0 C BP 100 MP NaCl = 800 BP 1413 Yet for all their strength, many ionic solids readily fall apart in water and break up into aqueous ions: NaCl(s) Na (aq) + Cl (aq) When this happens the water which is a non conductor, now starts to conduct electricity In contrast covalent compounds don t so this (DEMONSTRATION) The key lies in the fact that ionic compounds are made of ions (DUH) and it is the ions that carry the current in water. We will talk about the interactions of ions in water in a later chapter. For now let s focus on the ions themselves What is the electron configuration of Na and Cl? [Ne] and [Ar] 2+ 2- Clicker question 1; What are the electron configurations of Mg and O Clicker question 2: predict the ions that will come from elements achieving their noble gas configurations So ions are in stable noble gas configuration Figure 6.3- Note you should memorize the charges of all the ions in this table! But are also charged It is the +/- charge interaction (electrostatic force) that hold molecule together Key Terms: The positively charged ions formed by metals are called cations The negatively charged ions formed by non-metals are called anions
2 Ionic Solids are so strong because aren t really composed of molecules, but are composed of alternating ions, and it is the sum of all these interactions that makes the crystals so strong. (We will talk about this in more depth later) Covalent solids, on the other hand are individual molecules held together by much weaker forces (more later) 6-2 Ionic Charge and Chemical Formulas Knowing the charges of the common ions we just examined, you should be able to come up with the formula of any ionic compound made with these ions (Don t do names yet) Key concept: The net charge in any ionic compound is zero, so the sum of the charge of the cations and the sum of the charge of the anions must cancel each other out Examples: Sodium Chloride Na & Cl Na and Cl 1(+1) + 1(-1) =0 NaCl Sodium oxide Na & O + 2- Na & O 2(+1) + 1(-2) =0 Na O 2 Scandium sulfide Sc & S 3+ 2- Sc & S 2(+3) + 3(-2) = 0 Note that cross multiplying works nicely! Sc S 2 3 Clicker questions: Rubidium chloride, Cesium nitride, Aluminum oxide 6-3 Ionic Charges in Transition Metals Silver is a transition metal that is commonly found in a +1 ionic state, Ag + Why? 2 9 While you might thing the electron configuration of Ag is [Kr]5s 4d Aq is just like Cu above it in the periodic table and has the configuration 1 10 [Kr]5s 4d
3 How would you get to an inert gas configuration? Lose 11 electrons or gain 7 electrons. Neither sounds very appealing 10 But if you lost 1 electron you would at least get to [Kr]4d and that configuration 2 6 10 would fill in the 4 shell (4s 4p 4d ) to achieve some stability This explains the ions found in Figure 6.5 The block of atoms by Indium has another way to go. 2 10 1 The electron configuration of Indium is [Kr]5s 4d 5p By the logic we just looked at it can become a +3 ion by going to the 10 [Kr]4d configuration 2 10 But it can go to the +1 state by going to the [Kr]5s 4d as well So we have another block of common ions in this area Figure 6.6 Note that In, Sc, Tl & Pb overlap with previous figure so there are multiple ionic states for these element This is part of a more general problem. Most of the transition metals can have 2 or more ionic states. How do you properly name and identify ionic compounds when you have metal cations that can have different charged state? 6-4 Ionic Charges in Transition metals with >1 common charge The answer to the above question is to use roman numerals to cearly identfy the charge state you want. Key Concept: Metals with 1 common charge state so you DON T use roman numerals Group 1 (+1) Group 2 (+1) Al (+3) Zn(+2) Ag(+1) All other metals need a roman numeral included in their name to identify the charge on the ion Examples Give the name of CuCl 2 (Copper(II) choride) Give the formula of Cobalt(III) sulfide (Cu S ) Clicker questions: Names and formulas of several ionic compounds with and without roman numerals Note: ths book goes in to some older nomenclatures that I will skip! 2 3
4 6-5 Filling order of Transition Metals I think I will skip this section 6-6 Cations are smaller and Anions larger than parent atoms Key concept: Cations smaller than parent atom Anions larger than parent atom If you think about it this makes sense What do you do to make a cation. You remove and electron from the outermost, largest shell, so this naturally makes the cation smaller. Not only that but you frequently go to a noble gas configuration, but the charge in the nucleus is more positive than that of the noble gas you are mimicing, so this pulls the electrons in even more tightly To make an anion you add an electron to the outermost shell which will make the ion larger than the original atom. Not only that, but all the electrons are repelled from each other, to that makes the anion swell up even larger. This is shown graphically in figure 6.8 Practice problem Sort the following atoms and ions from smallest to largest + 3- + 2- - Li, N, Ne, Na, S, Cl, Ar Clicker question Sorting atoms and ion by size 6-7 Coulomb s Law and Ion Pair energy Early in the chapter I said that ionic bonds were strong bonds, and that is why ionic compounds have high melting points and boiling points Let s actually back that up with a calculation or two and get some real energies to look at. Consider the reaction: Na(g) + Cl(g) Na Cl (g) If this reaction releases energy, then the product NaCl (g) is stable because you have to put that energy back into NaCl(g) to get it to change back into the separate atoms of Na(g) + Cl(g)
We can actually calculate the energy of this reaction by breaking it into individual steps: 1.) Na(g) Na (g) + e - - 2.) Cl(g) + e Cl (g) 3.) Na (g) + Cl (g) NaCl(g) Reaction 1: You should recognize the first reaction as the ionization energy. So we can simply look this energy up. It is 0.824aJ Reaction 2: The second reaction is a new one. It is called the electron affinity. It is a measure of how much energy is gained or lost when an atom grabs an electron Key Concept: - - Electron affinity: Y(g) + e Y (g) 5 For Chlorine - - Cl(g) + e Cl (g) -0.58 0 aj - The negative sign says that this process releases energy, thus Cl (g) is more stable than Cl(g)! st nd rd st nd Just as there were 1, 2, 3... ionization energies, there are 1,2, rd 3...electron affinities. Reaction 3: Early in the chapter I said that ionic compounds were held together by charge-charge - electrostatic interactions. Does anybody know the equation for this kind of interaction? Key Equation: It is called Coulomb s law In this case Q 1 and Q 2 are the charges of the two ions, measured in units of the charge of a proton, so Q for Na = +1, and Q for Cl = -1 1 2 d is the distance between the centers of the ions in pm. In this case the radius of Na is 102 pm and Cl is 181 pm (See table 6.4 page 183) so d = 102+181 = 283 pm
6 and k is a proportionality constant = 231 aj pm So our energy is: 231 aj pm x (+1)(-1)/281 pm = -.816 aj The - sign says that energy is released, so the ions of opposite charges are attracted to each other (DUH!) So let s sum things up RXN 1: +0.824 aj RXN 2: -0.580 aj RXN 3: -0.816 aj Net: -0.572 aj Negative sign, energy released, NaCl(g) is more stable than Na(g) + Cl(g) Key calculation: Given ionization energies, electron affinities and ionic radii be able to calculate the energy X(g) + Y(g) XY(g). And any variation At this point our story isn t complete, because ionic compound are so stable, if is very difficult to find them in the gas phase, they always occur as solids at room temperature and pressure. In an ionic solid we have the + and - ions in an alternating lattice, as shown in figure 6.13 + Look at the Na ion in the middle of this lattice. See how it is actually touching 6 - nearest neighbor Cl ions. That mean we have underestimated our energy in the solid. Key concept: Lattice energy = E gained or lost in reaction X(g) + Y(g) XY(s)
7 The actual lattice energy for the NaCl solid is -1.28 aj, more than twice our simple estimate. Calculating the true lattice energy that properly accounts for all neighbors around each ion, and also includes the fact that the ions are not standing still but are vibrating is beyond what you need in this course. It can be done, and the calculation and the experimental number are in very good agreement. (See table 6.6 in your text if you are interested)