UNIT 1 1

HONORS CHEMISTRY

HARVARD-WESTLAKE

UNIT 1

Introduction to Chemistry and

Compound Nomenclature,

Mole Concept,

Basic Calculations

A chemistry teacher was berating the students for not learning the Periodic Table of the

Elements. She said "Why when I was your age I knew both their names and weights." One kid popped up, "Yeah, but teach, there were so few of them back then."

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Reference Sheet

Avogadro’s Number (N): 1 mole= 6.02x1023 particles

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Naming Examples NaCl FeCl3

CaCl2 ZnSO4

K2O AgBr

AlBr3 Co2O3

S2Cl2 Al(NO3)3

MgS NiF2

PCl3 KI

Li2SO4 SO3

NaNO3 SrCO3

NH4Cl NH4CH3COO

Ca(OH)2 K2CrO4

Na2Cr2O7 IF7 N2O5 FeSO4

Mg3(PO4)2 SnS

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Naming Examples (continued)

potassium nitride lithium bromide sodium nitrate tin(II) oxide dinitrogen tetroxide barium hydroxide ammonium carbonate copper(II) sulfate calcium phosphate silver acetate zinc cyanide manganese(II) nitrate mercury(I) chloride iodine pentafluoride sodium nitrite cadmium iodide potassium permanganate copper(I) oxide strontium fluoride

potassium sulfite aluminum sulfate lead(II) oxide silver phosphate copper(II) acetate magnesium nitride cesium chloride barium carbonate mercury(II) hydroxide

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Writing Formulas for Inorganic Compounds

yes* no yes

no

* There are a few minor exceptions in which a prefix is part of the historical name of the ion. The most common example is the dichromate ion, Cr2O7

2-

Are there prefixes?

Write the formula: sulfur trioxide SO3 carbon tetrachloride CCl4 dinitrogen tetroxide N2O4

Are there Roman numerals?

The numeral tells the chargeof the metal. You must know the charge on the negative ion in order to make the compound neutral. Iron(III) chloride is FeCl3 because chloride is -1.

Write the formula, paying attention to the charges on the ions. Use enough of each ion to make the compound neutral. Magnesium fluoride is MgF2 because the charge on magnesium is +2 and the charge on fluoride is -1.

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Naming Inorganic Compounds

yes* no

no yes * There is an important exception, the ammonium ion: NH4

+. This ion is treated as if it were a metal cation (its chemical properties are similar to alkali metal ions).

Is a metal involved?

Can the metal have more than one charge?

Use prefix system: CO carbon monoxide CO2 carbon dioxide N2O5 dinitrogen pentoxide (mono- is optional)

Use only the names of the ions: NaCl sodium chloride AlPO4 aluminum phosphate BaBr2 barium bromide CaS calcium sulfide

Use Roman numerals to indicate the charge of the metal ion: FeO iron(II) oxide Fe2O3 iron(III) oxide CuS copper(II) sulfide Cu2S copper(I) sulfide

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Mole Concept

How Big is Avogadro’s Number? 602,000,000,000,000,000,000,000

…..if you had Avogadro’s number of tiny grains of sand you could spread them out over the entire state of California to make layers, 10 stories high….. …..if you could divide Avogadro’s number of sheets of paper into a million stacks, the piles would be tall enough to reach from here to beyond the sun….. …..Avogadro’s number of marshmallows spread uniformly over all 50 states would yield a blanket of marshmallows more than 600 miles deep….. …..if Avogadro’s number of pennies were distributed equally among all the people on the earth, each person would have enough money to spend a million dollars every hours, day and night, throughout a lifetime, and would still not spend half of it (of course, that many pennies would be more than 50 miles deep…)….. …..Avogadro's number of soft drink cans would cover the surface of the earth to a depth of over 200 miles….. …..if you spread Avogadro's number of unpopped popcorn kernels across the USA, the entire country would be covered in popcorn to a depth of over 10 miles….. …..if we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole….. 1. How many moles are in 111.8 g of iron? 2. How many grams of water would contain 2.5 moles of molecules? 3. How many moles of atoms are in 1 mole of methane molecules?

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4. How many atoms of gold are in 10.0 g? 5. How many water molecules are in 36.0 g?

6. How many moles of aluminum hydroxide are contained in 39.0 g? 7. What is the mass of 0.75 mole of calcium sulfate? 8. How many grams of gold would contain 0.25 moles of atoms? 9. How many moles of sodium nitrate are in 43.6 g?

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Percent Composition and Empirical Formulas 1. What is the percent composition of each element in calcium sulfate? 2. The formula for rust is sometimes represented as Fe2O3. How many moles of Fe are present in 24.6 g of the compound? 3. During physical activity, lactic acid (MM = 90.1 g/mol) forms in muscle tissue and is responsible for muscle soreness. The compound contains 40.0 mass % C, 6.71 mass % H and 53.3 mass % O. What are the empirical and molecular formulas?

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4. 2.45 g of silicon reacts completely with 12.4 g of chlorine to form a compound. What is the empirical formula of the compound? 5. A compound contains 81.8% by mass C and 18.2% by mass H. The molar mass of the compound is between 40 and 50 g/mol. What are the empirical and molecular formulas of the substance? 6. A sample of 0.370 mole of a metal oxide, M2O3, has a mass of 55.4 g. What element is best represented by the letter “M”?

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Name:___________________________ Per.:____ Date:_______________

What are Compounds Like? Purpose: to observe some properties of compounds and use the observations to categorize the compounds as “ionic” or “molecular” Method: I will first try to dissolve a very small amount of each compound in distilled water. I will then test the electrical conductivity of each mixture (whether the solid appears to dissolve or not). Finally I will compare the conductivity of distilled water and water from the faucet. I will use my visual observations to separate the compounds by solubility and compare this information to groupings based on conductivity in distilled water (the comparison between tap water and distilled water is for a control). I can use these observations to decide whether the compounds are ionic or molecular with the general properties given in the lab handout for these types of compounds.  

Observations:

compound

appearance of

compound

dissolves?

conducts?

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compound

appearance

dissolves?

conducts?

distilled water

tap water

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LAB: What are compounds like? An element may be described as a substance in which all of the atoms are alike. Elements are the "stuff" of chemistry, but not quite the only stuff. If that were true, chemistry would be a lot simpler since there are currently known only 113 elements. Elements are interesting in their own right and you will have the opportunity to learn a lot about them this year. But when elements combine in a special way that we describe as chemical, then they can form compounds. A compound is thus a substance in which not all of the atoms are alike. To be more precise: in a compound elements are combined in some fixed ratio (by mass) and in such a manner that they cannot be separated by physical means (e.g., using a microscope and tweezers...), but will only yield to more chemical action if we want to separate them again. Compounds are responsible for most of the material and biological diversity that surrounds us. Some compounds are good news (aspirin comes to mind). Others we could better do without (sulfur dioxide and nitrogen dioxide which contribute to acid rain, for example). So the study of compounds is really pretty important in chemistry. And therefore the question which is the title of this activity: what are compounds like? It would be more interesting for you to answer that question yourself. But before you begin, here are a few things to keep in mind: 1. The properties of compounds are often not similar to the elements that compose them; e.g., sodium metal reacts violently with water and chlorine is a poisonous gas, but sodium chloride is—to most people—very harmless. 2. There are two very large categories of compounds which are based on properties that you can observe 3. There are a lot of compounds and it is impossible to fit every one into neat generalizations Preparing to experiment You will be provided with the following materials: 1. six compounds 2. conductivity device

3. 24-well plate 4. plastic stirrer

Design an experiment that will help you classify compounds as either ionic or molecular by testing whether the compounds are soluble in water and if their solutions conduct electricity. [see Technique] Technique 1. The 24-well plates are convenient for making a battery of tests on SMALL amounts of compounds. You cannot judge solubility properly if you fill up the well with solid!!! This seems so obvious and yet….. Using your scoop, take a little of each solid (the size of a match head) and place it in one of the wells in the plastic plate provided. Fill the well up about half-way with distilled water from your wash bottle. Jiggle the plate a little. You can use a plastic stirrer if needed. DO NOT use your glass stirring rod as it will scratch the bottom of the well. 2. The little conductivity device you have is powered by a 9 volt battery and it measures the conductivity of a substance by passing electricity through the material. The relative amount of electric current that gets through is indicated by the LEDs (small colored lights).

 

 

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All you need to do is pick up the device and place the two metal electrodes sticking out from one end into the solutions you prepared earlier in the wells. You should rinse the wires off between solutions with a squirt of distilled water. Record your conductivities in relative terms (no light = no conductivity, red = fair , red and green = good). You should also test tap water and distilled water just to have some kind of comparison.

If you have a compound which is not to be opened, it will have two wires inserted in the stopper. Touch the electrodes to these to check the conductivity. Questions 1-3 could easily be answered in a table. Feel free to use one if you like. 1. One way to divide compounds into categories is helpful for naming them. Compounds containing only two different elements are called binary. Carbon dioxide (CO2) is a familiar example. Compounds containing more than two different elements are known as ternary. Sodium hypochlorite (NaClO), which is in bleach, is one such compound. Divide your compounds into these categories. 2. Look in your text book (inside front cover) and determine which of your compounds contained metals (or the ammonium ion*, NH4

+) with non-metals and which contained non-metals only (excluding the ammonium ion). (for the purposes of naming compounds, metalloids are considered non-metals) *the ammonium ion is a troublesome exception--look out for it 3. To name your compounds, you need to know a few rules. These are covered in your text in detail, so here are just a few basic reminders: 1. metals always come first in mixed compounds 2. all binary compounds end in -ide 3. ternary compounds have various endings MORAL: look them up until you learn them 4. Greek numerical prefixes are only used with binary non-metal compounds (very few) Name your compounds (in most cases you can name that compound in 2 words...). Some of your compounds may have already had names on them. Obviously you don't need to name them again! These are ternary compounds composed of only non-metals and are typically carbon compounds of some sort. Their nomenclature is beyond the scope of this course. 4. There are, in fact, two major groups of chemical compounds, as you will discover in your reading. One group is ionic, the other is molecular. Ionic compounds consist of ions (charged atoms) which often separate in water and allow electricity to flow through the solutions. These ions (and others) are on the list of cations and anions you are learning. Molecular compounds consist of neutral molecules. Summarize the general behavior and constituents of ionic and molecular compounds based on your observations. 5. During the experiment you checked the conductivity of both distilled water and tap water. Based on your observations and the behavior of ionic compounds when dissolved in distilled water, what is a possible explanation for the difference in conductivity of tap and distilled water?

Analysis 

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LAB: The Synthesis and Characterization of Zinc Iodide It may be a little difficult to believe today, but in the late 1700's and early on into the 19th century the raging controversy among chemists was the idea that matter was particulate and had fixed chemical properties, i.e., that it was composed of what we today call atoms. The evidence for the earliest atomic hypotheses is fairly simple:

1. There are elements and compounds; unlike mixtures, compounds do not retain the properties of the elements in them 2. A given compound always contains the same proportion of its elements by mass (Proust)

3. When two elements form several different compounds, their relative masses are ratios of small whole numbers (Dalton)

4. Mass is conserved in chemical processes (Lavoisier) such as the formation of compounds from elements Because we take the concept of atoms for granted today, it is sometimes a stretch to imagine that from this rather simple experimental evidence people like John Dalton were able to postulate the existence of atoms (1808) and use the word in much the same way that we do today. Therefore in this experiment you will see if you can reproduce some of this evidence by making zinc iodide from the elements zinc (Zn) and iodine (I), investigating the proportions in which they combine (to determine the chemical formula), comparing the properties of the compound with those of the original elements, and decomposing the compound to (hopefully) obtain the original elements. The composition of a compound is often expressed in percentages by mass. The entry at the end of this experiment for zinc iodide has this information deleted but it is standard data which can be found in any good handbook. A simple example will show how the values are calculated:

The formula mass of water, H2O , is [2(1.0) + 16.0] or 18.0 Therefore the %, by mass of water that is hydrogen is:

2(1.0)

x 100 = 11%18.0

In a similar manner the % that is oxygen can be obtained:

16.0

x 100 = 89%18.0

Note that these are merely simple fractions which represent the mass of a particular element in the formula divided by the total formula mass. Notice also that in the case of hydrogen, the subscript 2 is used in calculating the % H since there are two hydrogens in the formula for water. No matter how simple or complex the formula for a substance may be, its composition may be expressed in this manner. Proust's statement (number 2 on the facing page) simply says that the numbers calculated for a compound are always the same for that compound, regardless of how it is prepared or from where it is obtained. This is equivalent to saying that water is always H2O, never H4O, HO2, or some other arbitrary combination. Fundamental atomic theory also tells us that atoms cannot be divided into smaller parts without losing their chemical identities. During chemical processes, atoms remain intact. Thus atoms which combine to form molecules must do so in whole-number ratios (Dalton). The smallest whole-number ratio of atoms in a molecule (compound) is known as the empirical formula. If discrete molecular units exist which have this ratio, then the formula is also the molecular formula. For example, the empirical and molecular formulas for water are both H2O. In contrast, some compounds exist only in a multiple of their empirical formula. Hydrogen peroxide is an example. The actual molecule contains two atoms of hydrogen and two atoms of oxygen. Thus the molecular formula is H2O2. The empirical formula is simply HO. In general once the empirical formula of a substance has been

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established in the laboratory, other tests are required to determine if there is a different molecular formula. To find the empirical formula, a ratio of numbers of atoms (or moles) is needed. What is measured in the lab is typically mass. A simple example illustrates the process of conversion:

1.50 g of Mg metal is placed in a crucible and heated until it ignites. After cooling, the sample is found to have a mass of 2.49 g. Assuming that the product contains only Mg and O, the mass of O would be:

2.49 g - 1.50 g = 0.99 g

Moles of each element are then calculated:

1.50 g

moles Mg = = 0.0617 mol24.3 g/mol

0.99 g

moles O = = 0.0619 mol16.0 g/mol

Within experimental error this is a 1:1 ratio so the empirical formula is MgO In order to know if this is the actual molecular formula we would need information such as the measured molar mass--but that's another experiment! Sometimes it is not so obvious from a simple calculation just what the integer ratio should be. Consider the next example:

7.3 g of finely divided Al is heated in oxygen. The final mass after complete reaction is 13.8 g. What is the empirical formula of aluminum oxide? The mass of oxygen would be: 13.8 g - 7.3 g = 6.5 g O

The moles of each element then are: 1 mole

7.3 g Al x = 0.27 mol Al27.0 g

1 mole

6.5 g O x = 0.41 mol O16.0 g

It is not exactly clear what the ratio is here but you can get a better idea by dividing the smallest number of moles into each value:

0.27 0.41 = 1.0 = 1.5

0.27 0.27

The corresponding integer ratio is thus 2:3 and so the empirical formula is Al2 O3 What about zinc and iodine? By doing a little detective work in your book you can probably figure out what the formula for the compound of these two elements should be. Then you could calculate the % composition. But what if you didn't know the formula? The challenge in this experiment is to put yourself in the position of an early chemist: you are still trying to figure out what things are made of (and that's not far from the truth!). One other important aid in this endeavor is your own senses. Compounds and their elements seldom have the same physical properties or appearance (certainly not the same chemical properties). Your careful observations about physical appearance, water solubility, electrical conductivity, and so on can be very useful. Many compounds like zinc iodide are held together by electrical forces and so it is sometimes possible to break them up into their elements again using electricity. At the end of the experiment you can also try this to see if you recognize the original starting materials, zinc and iodine.

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Preparing to experiment You will be provided with the following materials: 1. granular zinc (use about 0.5 g) 2. iodine [at balances] (use about 0.5 g) 3. slightly acidic water (use about 3 mL) 4. two 18x150 test tubes 5. test tube clamp 6. 24-well plate 7. conductivity tester 8. 9-volt battery and wire clip Design an experiment to make zinc iodide by reacting zinc and iodine in acidic water, recovering and measuring the amount of compound formed and the leftover zinc remaining. Test the compound for water solubility and electrical conductivity in solution. Technique Zinc can react with iodine pretty vigorously--the combination can be dangerous. So a little respect is due to these elements, especially iodine. Be sure to read the information at the end of this handout on the substances in this experiment. Because the two elements are solids at room temperature, reaction between them is very slow unless they are ground together vigorously. THIS IS VERY DANGEROUS! So we won't do it. 1. Quantitative transfers are very difficult to make. Stuff gets left behind. In an experiment of this kind, with small masses involved, loss of even a tiny amount can lead to significant experimental error. The best way to avoid this problem is not to transfer unless necessary. Measure the zinc directly into a pre-massed test tube. You can stand the test tube up in a beaker which has been “tared-out” on the balance. It takes a little more time to place solid into the narrow test tube, but it’s all in there when you’re done. To add iodine, just tare the whole mess and keep adding----but be careful: if you add too much it all has to be discarded since the materials are contaminated with each other. 2. Recovery from solution requires different techniques depending on whether the material is soluble or not. Fortunately iodine is slightly soluble in water and bringing at least one element into solution speeds up this reaction significantly, but constant swirling is needed to get the job done in the class time. However, zinc will slowly react with water to form zinc hydroxide. This will interfere with your experiment, so rather than use the standard distilled water from your water bottles, be sure to use the slightly acidic water provided for this experiment. It will prevent the formation of zinc hydroxide during the reaction between zinc and iodine. At the end of the reaction, zinc is left over at the bottom of the test tube but the product is invisible in the solution. The solution can be decanted into another clean, dry pre-massed test tube, i.e., the liquid on top (supernatant) is carefully poured off, either into another container (in this lab, another clean, dry pre-massed test tube) or the sink. Generally, to ensure that all of the substance in solution is removed/recovered, several small rinses of the solid and walls of the original container are also made. These are combined with the original decanted solution. In this experiment, any rinsings will add to the total volume of solution which must eventually be boiled away. It is therefore wise to keep the rinsings small, perhaps 3 rinses of 1 mL each---be sure to use the acidic water. Solids in test tubes can be troublesome to dry other than slowly in an oven, but it is possible to heat the tube gently at first, then more strongly, and drive off any liquid present, holding the test tube in a spring clamp. Keeping the test tube moving through the flame will help prevent little steam explosions that can eject the contents. Also, the clamp itself will become too hot to hold if it is held in position over the flame constantly. Once the initial heating is done, deliberate heating, moving from the bottom of the test tube to the mouth--driving out any moisture--should be done. Any water remaining will make the final massing inaccurate.

 

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Liquids in test tubes boil very rapidly and have a tendency to "foam" up and --occasionally--out. To minimize this, hold the test tube at an angle above the flame and keep it moving all the time. Be sure not to point the mouth of the test tube at anyone. Because solutions sometimes bump and spurt out when being heated (especially as they become more concentrated), a "boiling stone" is often added to aid in the formation of smaller bubbles during boiling. These stones do add mass, however, and we will not use them in this experiment. Zinc iodide absorbs water rapidly from the air and should be heated until it is slightly yellow. Overheating will decompose the compound, liberating iodine gas and ruining your results. It may be that your class will do the experiment on two consecutive days. In that case much or all of the drying may be done in the oven overnight. 3. Placing hot objects on the balance is a no-no. The heat they give off creates air currents which often cause the sensitive balance pans to move up and down. Also there is a possibility that very hot objects will damage the balance. The general rule is that you must be able to pick up the object with your hand and walk to the balance with it. If you can't, it's still too hot. ALWAYS USE THE SAME BALANCE FOR AN ENTIRE EXPERIMENT. ____________________________________________________ Equipment

Bunsen burner test tube test tube clamp ____________________________________________________ The chemicals Iodine can be obtained from many natural sources including brines (concentrated salt solutions) associated with oil wells and seaweed. In rocky minerals it is present only to the extent of 3 x 10-5 %, in seawater, 5 x 10-8 %. Once the compounds of iodine have been isolated they can be treated with chlorine which displaces the elemental iodine. At room temperature iodine is a shiny, gray-black, non-metallic solid. In the gas phase and in some solvents, it is violet. In water and alcohols (like ethanol) it has a reddish-brown color. Tincture of iodine is a solution of iodine in alcohol and is used as an antiseptic. Iodized table salt contains potassium iodide (KI) as a nutritional supplement to help prevent iodine deficiency diseases such as goiter. In pure form and in solution iodine is very reactive and can cause staining on skin and clothing. Its vapors are irritating to the eyes, nose and throat. Iodine shares an unusual property with another common substance, carbon dioxide: it does not pass through a liquid phase at ordinary pressures but instead goes directly from solid to gas upon heating. This process is known as sublimation. Zinc is generally obtained from ores of zinc containing sulfur. Its abundance in the earth's crust is about 0.02%. Zinc is a fairly reactive metal which combines readily with oxygen, sulfur and the halogens. Pure zinc, when exposed to air gradually becomes coated with white zinc carbonate (ZnCO3). Most zinc compounds are colorless in solution (or white as solids). Zinc is readily attacked by dilute acids, releasing hydrogen as it dissolves. It is used in corrosion protection (galvanizing) and its compounds are employed as paint pigments and disinfectants. Zinc iodide is a white (hydrated), or pale yellow (dry), odorless powder when pure. It becomes brown on exposure to air and light, slowly releasing iodine. One gram dissolves in 0.3 mL of water. It is sometimes used as a topical antiseptic or astringent. ____________________________________________________

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Analysis 1. Use your data to determine how many grams of zinc reacted and how many grams were left over. 2. Assuming that all of the iodine reacted, what is the percent composition (by mass) of each element in the compound? (be sure to use your data for this, not some theoretical calculations) [hint: this determination should be based on the zinc reacted--from #1-- the iodine reacted and the recovered mass of zinc iodide] 3. Based on your answer to question 2, what is the empirical formula of zinc iodide? 4. Use cation and anion charges to determine the formula for a compound between zinc and iodine. 5. Does the theoretical % composition calculation [show this!] from the expected formula match your result in question 2? If not, can you suggest where errors might have been made? (be specific--for example, if your mass percent of zinc is too large, what does this suggest? or if your mass percent of iodine is too large, what measuring error might you have made?) 6. Compare the mass of zinc iodide recovered (value from lab data) with the expected total mass of product based on the masses of the elements which reacted. Should these masses be the same or different? In addition to the answers in question #5, what errors might have contributed to any inequality between the expected and experimental masses of zinc iodide? 7. What evidence did you see that both zinc and iodine were present in the final recovered solid? [hint: what happened when electric current was run through a solution of the product? Be specific for each of the two wires.]

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Unit 1 Sample Test The test will consist of 5 multiple choice questions, 4 required problems, and one essay question. The following are representative of typical multiple choice questions but do not necessarily indicate topics to be addressed on your actual test. 1. The majority of elements on the periodic table fall into the category a. metalloids b. noble gases c. metals d. non-metals 2. In the atom 35

17 Cl, there are___ protons and ___ neutrons.

a. 17, 35 b. 35, 17 c. 18, 17 d. 17, 18 3. The name of the compound S2Cl2 is a. sulfur dichloride b. disulfur dichloride c. sulfur chloride d. disulfur dichlorine 4. The empirical formula for a substance is CH3. Which of the following cannot be a possible molecular

formula for this substance? a. C2H6 b. C4H12

c. C5H10

d. C3H9 5. The existence of compounds such as N2O3 and N2O4 is explained by a. the Law of Definite Proportions b. the Law of Multiple Proportions c. Avogadro's number d. none of these

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The next section consists of representative problems which might be found in the problem section. 6. A compound upon analysis gave 38.67% K, 13.85% N, and 47.48% O. Find the empirical formula of this compound. If the molar mass is known to be approximately 100 g/mol, what is the most likely molecular formula? What is the name of this compound? 7. Small pieces of copper metal are placed in a crucible and covered with powdered sulfur. The crucible is heated strongly until all evidence of reaction ceases. Assuming that the amount of sulfur used was in excess, what is the empirical formula for the compound formed if the original mass of the copper metal was 2.47 g and the mass of the final compound was 3.72 g? What is the name of the compound? 8. Name the following: ____________________________________________ a. Cu3 PO4

____________________________________________ b. Br2O5

____________________________________________ c. NH4Cl ____________________________________________ d. Al(OH)3

____________________________________________ e. HgS

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9. Write formulas for the following: ______________________ a. sodium sulfide ______________________ b. carbon disulfide ______________________ c. cobalt(III) oxide ______________________ d. lithium oxide ______________________ e. potassium phosphate 10. The formula for the compound sodium oxalate may be written as Na2C2O4 . a. Find the % composition by mass for each element in the compound. b. How many moles of sodium oxalate would be in 34.5 g of the compound? c. What is the empirical formula for sodium oxalate? 11. All of the substances listed below are fertilizers that contribute nitrogen to the soil. Which of these is the richest source of nitrogen on a mass percentage basis? SHOW WORK! a. urea, (NH2)2CO b. ammonium nitrate, NH4NO3

c. guanidine, HNC(NH2)2

d. ammonia, NH3

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Definitions from unit 1, that are important to understand, not just memorize. ______12. ion A. referring to numerical information ______13. binary compound B. the smallest unit of a compound ______14. qualitative C. a charged atom or group of atoms ______15. molecule D. a negatively charged particle found outside the nucleus ______16. electron E. a vertical column in the periodic table ______17. group F. a compound consisting of only two different kinds atoms G. referring to non-numerical information H. a diatomic molecule The essay questions for this first unit will cover the topics listed below. Only one of these topics will appear on your particular test so you should be prepared to write intelligently about any of them. -- laboratory safety equipment -- early chemical laws -- bunsen burner -- determining the empirical formula for an iodide of aluminum (aluminum has some chemical properties similar to zinc)