How are molecular formulas different from empirical formulas? Can they ever be the same for a particular substance?

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1 Chapter 3 Reading Guide Tro 4 th edition Suggested Problems: balancing equations is very time consuming in MC, so there are no assigned problems. Don't ignore this though, you need to practice. Use worksheet #4 for practice, and 101, 107, 109, 111. Naming compounds: 33, 35, 37, 39, 41, 43, 45, 47, 49, 51, 53, 55, 57. Moles: Try one of each set in these problems, do more if you need to: 59d, 61, 63, 65, 67. Composition and formulas: 71, 81, 85, 87, 89, 93, 95, 97, 99. Need more help with naming compounds? Try this Extra nomenclature handout with key. Section 3.1, Do the parts add up to the whole? Water is a compound made up of two different elements. Does water have similar properties to H 2 and O 2? Does sodium(s) have the same chemical properties as sodium in table salt? Do compounds always have the same ratio of one element to another? Section 3.2, What are the two main types of chemical bonds? We will discuss chemical bonding at length later in the semester. This section gives you a preview of what s ahead. What two types of elements combine to make most ionic compounds? Why do cations and anions stick together to form ionic bonds? Check out Figure 3.2. What happens to the number of electrons in a sodium atom when it reacts with Cl atoms to form NaCl? This is a figure that you really should study. What happens to the number of electrons in a chlorine atom when it reacts with Na atoms to form NaCl? What kind of elements make up most covalently bonded compounds? Are electrons transferred in covalent compounds, or is there a different way that electrons are distributed in these compounds? What is the difference between an ionic bond and a covalent bond? Section 3.3, What are the different types of chemical formulas? How are molecular formulas different from empirical formulas? Can they ever be the same for a particular substance? What s the difference between a molecular formula and an empirical formula? Can they ever be the same? Example 3.1 has a few examples. Section 3.4, Can you classify elements, compounds, and molecules? This section has further descriptions of elements and compounds. You must know the 7 diatomic elements, those that exist as H H or N N at room temperature in their elemental state. There are some mnemonic devices you can use, my favorite is Horses Need Oats For Clear Brown I s. These are highlighted in Figure 3.5. There is an important but subtle distinction drawn between molecules and the formula units of ionic compounds. Is KF a molecular substance with covalent bonds? Chapter 3 reading guide Dr. Baxley p 1

2 Poly-atomic ions are introduced in this section. They are like little charged molecules, and almost always remain intact during chemical reactions. Give Example 3.2 and For Practice 3.2 a going over to ensure that you can properly categorize these substances. The answers are in the back of the text. Conceptual Connection 3.4 is a great way to make sure you can apply concepts in this section. Chapter 3 reading guide Dr. Baxley p 2

3 Section 3.5, How do chemists name and write formulas for ionic compounds Here is a not so subtle reminder to GET TO WORK ON MEMORIZING ELEMENT AND ION NAMES! There is a nice 3 point summary of formulas for ionic compounds at the beginning of this section. Metals that form ions with only one charge: (like Ca 2+ or K + ) Referring to Table 3.2 and back to Figure 2.13 will help you determine charges of elements that make only one type of ion. Other elements make multiple ions and are discussed later. Does the sum of the charges of all of the cations need to be equal to the sum of the charges of the anions? Is it acceptable to write the formula of CaO as Ca 2O 2? The two examples (Examples 3.3 and 3.4) are really good at showing how to write formulas for the simple ionic compounds. Work out For Practice 3.3 and 3.4 too. There is a little flow chart on p 97 and a big one on p 105 (Figure 3.11) for naming ionic compounds, based on the number of different cations that a metal might form. Does sodium form more than one type of ion? What about iron? Work on the Example and For Practice 3.5 items. Are there Greek prefixes (like di or tri) used in names of ionic compounds? Metals that form cations ions with more than one charge: (like Co 2+ and Co 3+ ) Table 3.3 and the assigned elements you need to know highlight some of the metallic elements that form multiple cations, each with different charges. Check out the pretty colored boxes for how to write names of this type of compound. Question A: What is the difference between naming Na 2O and Cu 2O? Question B: Does the Roman numeral indicate the number of anions, number of cations, or the charge of the metal? You don t have to memorize the common charges of these elements, just be aware that they can have cations with variable charges. For Example and For Practice 3.6 will give you more practice. BONUS Question C: My absolute favorite examples illustrating how to name compounds like this are: Write the name for Cu 2O. Write the name for CuO. Ask yourself Question B again. Answers on the next page. Polyatomic ions: (like OH, SO 4 2, and NH 4 +) These are treated like the elemental ions, they are just little charged molecules. Memorize the ones you need to know. Section 3.6, How do you name molecular compounds and acids? Binary compounds: Chapter 3 reading guide Dr. Baxley p 3

4 Part of this section details how to name binary (only 2 elements) molecular compounds. If there are 3 elements, don t use these rules. Answers from previous page Question A: What is the difference between naming Na2O and Cu2O? A: Sodium oxide has no Roman numeral (Na only makes a +1 charge with compounds) Question B: Does the Roman numeral indicate the number of anions, number of cations, or the charge of the metal? A: The Roman numeral indicates the charge of the metal, nothing else. Question C: Write the name for Cu2O. copper(i) oxide, since each copper cation must have a +1 charge to go with 1 O 2 ion Write the name for CuO. copper(ii) oxide, since the copper cation must have a +2 charge to go with 1 O 2 ion Note that these are the only compounds that use Greek prefixes in their formal names. The prefix for seven is a commonly missed prefix. Tetra is easy; it s like Tetris (with 4 boxes in each shape). Practice with Example 3.8. Bonus: not in text!! Molecular compounds are compounds in which electrons are shared between atoms, not transferred. Compared to ionic compounds, molecules often: have low boiling points, and can easily be gases or liquids at room temp are non-crystalline, plastic, or soft and squishy change phase (solid liquid gas) with temperature, can dissolve in solvents, but do not undergo bond breakage when heated gently or when dissolved don t conduct electricity well in liquid state Acids Acids have a H atom that ionizes in water to make H + ions. Follow the schematic on p. 103 for separating acids into two groups. One type is a binary acid. One type is an oxo acid (with a polyatomic ion that has oxygen). Knowing the names of the polyatomic is essential to naming acids. The flow chart is a great way to help summarize naming compounds. Practice with Examples 10 and 11. Section 3.8, How do you calculate the Formula Mass and molar mass for compounds? The formula mass is the mass of the smallest unit (a molecule or set of ions) that make up the substance. The molar mass is the mass of one mole of these units. The value is the same, but the units are different (amu vs g). The first part of the section gives an example of how to calculate the formula mass of a compound. Practice with Example 3.12 and For Practice Note how the only the name of the substance is given, so you must know how to write the correct formula from the name! Is the value of the molar mass (in g/mol) identical to the value of the formula mass (in amu)? Chapter 3 reading guide Dr. Baxley p 4

5 Converting between mass, moles, and actual number of particles is a very important, fundamental task in general chemistry. Practice with Example Section 3.9, What is the mass percent of an element in a compound? This section starts off with an odd definition of mass percent. It s much more general that the text definition. The mass percent of an element in a compound is the mass of one element in a sample divided by the mass of the entire sample. The sample could be 1.0 mole, it could be 2.0 moles, or moles. Doesn t matter. If one is trying to calculate the mass percent of an element in a compound, and you don t have a specific sample size to work with, then use a sample size of one mole, and you can see how to calculate the mass percent in Example Example 3.15 shows how knowing the mass percent can be useful. Know anyone on a low-sodium diet? The sub-section of Conversion Factors from Chemical Formulas on p 112 is super useful and very important. The number of moles of one element in a molecule, or the mole ratio of one element to another element in a compound are shown in the chemical formula. Example 3.16 and For Practice 3.16 are quite useful to work on your own. Section 3.10, How do chemists determine chemical formulas? This section has some of the most difficult calculations of the semester, so get ready. It is vital to keep in mind what the chemical formula of a compound tells you: it s the ratio of the number of one atom to the numbers of each other atom. Numbers. Not mass. Numbers can be moles. Not mass. So, to determine a chemical formula from information about the elements in a sample, one needs to know how many moles of each element are in a particular sample. Not mass; moles. Example 3.17 and 3.18 are great. Example 3.18 even has a video to guide you through this in MC. Look in the Study Area in MC for Chapter 3. Does this type of calculation calculate the empirical formula, or the molecular formula? It s an important distinction. Tips: 1. Use mass, percent by mass, or whatever to calculate moles of each element. 2. Divide each number of moles by the lowest number of moles. This gets at least one element to 1 mole relative to the others. 3. Don t memorize the little table of fractional subscripts. Just multiply by small whole numbers. 4. If a number is +/ 0.1 from a whole number, go ahead and round to the nearest number Molecular formulas vs empirical formulas: The text has a nice little way to calculate the molecular formula from the empirical formula if the molar mass is known. Chapter 3 reading guide Dr. Baxley p 5

6 Combustion Analysis: OK, this is the hard part! For a combustion analysis, one knows the mass of CO 2 and the mass of H 2O that is produced in a combustion reaction from a sample of a compound. Remember, you need to know moles of the elements in the compound. Tips: Moles of C from the compound are equal to moles of CO 2, since any carbon in the sample would turn into CO 2. Moles of H from the compound are NOT equal to moles of H 2O, but you can figure this out with mole ratios. Moles of O can NOT be determined from the masses of CO 2 and H 2O, since extra O 2 is needed for combustion. Mass of O, or of other elements, can be determined by subtraction or some other means. Examples 3.20 and 3.21 are great. Section 3.11 Balancing Equations is pretty straightforward, balancing equations. Make sure you do Examples 3.22, 3.21, and 3.22 but only after working on the Ions to Memorize list. Lots of practice in worksheet #4 in the back of the lab book, with answer keys available. We will skip section Don t forget the self-assessment quiz (#1 15) at the end of the chapter. Chapter 3 reading guide Dr. Baxley p 6