Experiment 3

The Language of Chemistry (Part I): Chemical Formulas, Nomenclature &

Mass-Mole Calculations

Learning Goals & Objectives

This activity provides an overview of the chemistry language and introduces mass-mole calculations. All too often because of time constraints, simple core concepts like nomenclature and mass-mole calculations receive, at most, a quick superficial treatment in chemistry lecture. Gaining proficiency in using the chemistry language along with understanding the nature of mass-mole calculations is vital to your success. Students who do not grasp these essential principles quickly will fall behind and most likely not pass either the lecture or laboratory courses. Now is the time to receive extra help from your instructor so as to clear up any points of confusion.

Upon completing the exercises in this worksheet, students will be able to: 1) identify numerous compounds, ions and molecules by name and chemical formula; 2) write the formulas for compounds formed from selected ions; 3) predict how ionic compounds dissociate in water; and 4) perform calculations involving mass and mole conversions.

Introduction

Have you ever enrolled in a foreign language class? The first thing one learns is the alphabet (letters) for that language, they form the building blocks. Combinations of letters produce words and words stringed together lead to sentences, and so forth. Before people can communicate effectively, they must all understand what information individual words convey. Chemistry has its own language; elements found in the periodic table serve the same function as letters. One must understand how elements (letters) are linked together to form compounds and molecules (words) and how when combined together compounds and molecules lead to new substances being formed (chemical reactions or sentences). Toward this end, the following pages contain a tabulation of names for the most common ionic building blocks to form compounds.

The goal of this activity is not to make you a nomenclature expert, rather we want you to become familiar with (remember) unit names so that you can write the chemical formula for a

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compound when given its name. Likewise, you are expected to be able to convert a chemical formula into a chemical name.

Naming Compounds: An Overview

We find two kinds of names for some compounds, the trivial name and chemical name. The chemical name is much more informative because it tells one the composition of the compound. Table 3.1 shows examples of trivial and chemical names for common substances used in the home and in industry. How many do you immediately recognize?

Table 3.1. Common ChemicalsTrivial Name Chemical Formula Chemical NameGrain Alcohol (drinkable) C2H5OH Ethyl Alcohol, EthanolWood Alcohol (poison) CH3OH Methyl Alcohol, MethanolBaking Soda NaHCO3 Sodium hydrogen carbonateBleach NaOCl Sodium hypochloriteBorax Na2B4O7 · 10 H2O Sodium tetraborate decahydrateBrimstone S SulfurCane or Beet Sugar C12H22O11 SucroseEpsom Salts MgSO4 · 7 H2O Magnesium sulfate heptahydrateGypsum CaSO4 · 2 H2O Calcium sulfate dihydrateLaughing Gas N2O Dinitrogen oxideLime CaO Calcium oxideLimestone or Marble CaCO3 Calcium carbonateLye, Caustic Soda NaOH Sodium hydroxideMilk of Magnesia Mg(OH)2 Magnesium hydroxideMuriatic acid HCl Hydrochloric acidOil of Vitriol H2SO4 Sulfuric acidPotash K2CO3 Potassium carbonateQuicksilver Hg MercurySaltpeter NaNO3 Sodium nitrateTable Salt NaCl Sodium chlorideWater H2O Water

As you convert chemical formulas into a name, try to visualize the individual chemical building blocks (formulas and charges). Inorganic compounds are usually composed of two parts, one has a positive charge and the other carries a negative charge. Usually, a metal combining with a second metal, forms an alloy. Two non-metal elements regularly combine to form a covalent substance. A metal and a non-metal usually combine to form an ionic compound.

To name a compound, the positive part (cation) is named and written first. The positive part will be a metal, a hydrogen ion, or a positively charged group made up of several elements. The negative

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part (anion) is named and written second. The negative part will usually be a non-metal or a negatively charged group composed of several non-metal atoms (polyatomic anion) linked together by strong covalent bonds.

By examining Table 3.2, one sees it is easy to name a cation. In most cases, the name consists of the metal element’s name or stem name with an added –ium ending. Elements which exhibit two or more oxidation states, such as Fe2+ and Fe3+ are named iron (II) and iron (III) respectively. The oxidation number, written in roman numerals, is enclosed in parenthesis immediately following the elemental name. Older nomenclature rules make use of –ous (ferrous) and –ic (ferric) endings to denote the metal’s lower and higher oxidation states respectively. Although these classic names are being phased out, we include them here for completeness. Note that the charge associated with any ion is always written as a superscript, with the number preceding the plus or minus charge. A cation is written as Cu 2+, not Cu+2.

Table 3.2. Names for Common Cationic GroupsCharge of 1+ Charge of 2+ Charge of 3+

H+ hydrogen Be2+ Beryllium Al3+ aluminumLi+ Lithium Mg2+ Magnesium Cr3+ chromium (III)

chromicNa+ Sodium Ca2+ Calcium Co3+ cobalt (III)

cobalticK+ potassium Sr2+ Strontium Fe3+ iron (III)

ferricRb+ Rubidium Ba2+ Barium Mn3+ manganese (III)

manganicCs+ Cesium Cd2+ Cadmium Sc3+ scandiumNH4

+ ammonium Cr2+ chromium (II)chromous

Ti3+ titanium

Ag+ Silver Co2+ cobalt (II)cobaltous

Cu+ copper (I)cuprous

Cu2+ copper (II)cupric

Hg+, Hg22+ mercury (I)

mercurousFe2+ iron (II)

ferrousMn2+ manganese (II)

manganousHg2+ mercury (II)

mercuricZn2+ Zinc

The vocabulary problem is greatest for the anions. Table 3.3 lists by charge, the majority of common anions. As you learn the name of each cation and anion, also remember its charge and individual formula unit. This information is required in order to write the formula for any compound. When writing

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the formula for a compound, the combined oxidation numbers (i.e., charge balance) of the cation and anion add up to zero. For example, in the compound Na2CO3, the charge balance is 2(+1) + 1(-2) = 0. Examine the trends in table 3.3 and note that the names for anions end in –ide, -ate, and –ite. Except for cyanide and hydroxide, the –ide ending is used to denote a monoatomic anion such as chloride, oxide, and nitride. The –ate and –ite endings suggest that a non-metal (or metal) is covalently bonded to one or more oxygen atoms. For two anions having the same central non-metal anion combined with oxygen, the –ate ending denotes which anion has the greatest number of combined oxygen atoms. Sulfate has 4 oxygen atoms whereas sulfite has only 3 oxygen atoms.

Table 3.3. Names for Nonmetallic Anions and Polyatomic AnionsCharge of 1- Charge of 2- Charge of 3-

CH3COO- acetate CO32- Carbonate AsO4

3- arsenateHCO3

- hydrogen carbonate (bicarbonate)

CrO42- Chromate AsO3

3- arsenite

Br- bromide Cr2O72- Dichromate BO3

3- borateClO3

- chlorate HPO42- hydrogen

phosphateN3- nitride

Cl- chloride O2- Oxide PO43- phosphate

ClO2- chlorite C2O4

2- Oxalate P3- phosphideCN- cyanide O2

2- Peroxide PO33- phosphite

H2PO4- dihydrogen

phosphateSiO3

2- Silicate

F- fluoride SO42- Sulfate

H- hydride SO32- Sulfite

OH- hydroxide S2- SulfideClO- hypochlorite S2O3

2- ThiosulfateI- iodideNO3

- nitrateNO2

- nitriteClO4

- perchlorateMnO4

- permanganate

Although we have listed a number of common cations and anions in tables 3.2 and 3.3, please do not get the impression that a chemist can go to the reagent shelf and obtain a bottle that contains only calcium ions or only hydroxide ions. Solid ionic compounds are an electrically neutral combination of both a cation and an anion. It is not possible to have one without the other.

As one encounters structurally more complex compounds; nomenclature rules become more specific. Let us consider briefly the naming of four other classes of compounds; small covalent molecules, acids, salts of acids, and bases.

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Naming Covalent Molecules

To name a covalent molecular compound formed from two non-metal elements one uses Greek prefixes preceding without hyphen the name of the element to which they refer. The prefixes for the numbers 1-12 are as follows: mono, di, tri, tetra, penta, hexa, hepta, octa, ennea (or Latin nona), deca, hendeca (or Latin undeca), and dodeca. Occasionally one may have to use prefixes to denote fractions like hemi (1/2) and sequi (3/2) to name hydrates.

The first element in the chemical formula is named (with Greek prefix) followed by the second element (with Greek prefix) with its element name ending in –ide. If the first member has one atom, the prefix mono is dropped from the name. Some examples include the following: SF4 sulfur tetrafluoride; CO carbon monoxide; PBr5 phosphorus pentabromide.

Naming Acids

Compounds that produce hydrogen ions when dissolved in water are called acids. An acid has one or more ionizable hydrogen atoms bonded to an electronegative non-metal atom. Naming acids in a single systematic way is almost impossible because of the structural variety that exits for acids. The common or trivial names survive year after year in spite of nomenclature rules. True, there are systematic reasons behind the trivial names, but it is easier to learn the names than it is to learn the rules. Most of the common acids you are likely to encounter are listed in table 3.4. Examine the name prefixes and endings and where appropriate, correlate them with the rules that follow.

Acids giving rise to the –ide anions should be named as hydrogen -ide. Examples include: HCl, hydrogen chloride; HI, hydrogen iodide; and HCN, hydrogen cyanide. In writing the formula for an acid, hydrogen precedes the anion. Names such as hydrochloric acid refer to aqueous solutions, and percentages such as 36% HCl denote the weight/volume of hydrogen chloride in the solution. The acids listed in table 3.4 retain their trivial name due to long established usage. Anions may be formed from these trivial names by changing –ous acid to –ite and –ic acid to –ate. The prefix hypo- is used to denote a lower oxidation state; the prefix per- designates a higher oxidation state.

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Table 3.4. Names for Common AcidsMonoprotic Acids Diprotic Acids Triprotic Acids

HF Hydrofluoric acid H2SO4 Sulfuric acid H3BO3 Boric acidHCl Hydrochloric acid H2SO3 Sulfurous acid H3PO4 Phosphoric acidHBr Hydrobromic acid H3PO3 Phosphorous acidHI Hydroiodic acid H2CO3 Carbonic acid

H2C2O4 Oxalic acidHNO3 Nitric acidHNO2 Nitrous acid

HClO4 Perchloric acidHClO3 Chloric acidHClO2 Chlorous acidHClO Hypochlorous acid

HC2H3O2 Acetic acidHBrO3 Bromic acidHIO3 Iodic acid

Naming Acid Salts

Salts are formed when acids having one or more hydrogen atoms are neutralized. Replacing the hydrogen ion(s) with one or more metal ions produces an acid salt. Salts containing acid hydrogen are named by adding the word hydrogen without a space before the name of the anion. The most electropositive group is named first, followed by others, then the correct salt ending is added. The prefix bi- is commonly used to indicate a compound in which only one of two hydrogens has been replaced by a metal ion. Examples: NaHCO3 sodium hydrogencarbonate (sodium bicarbonate); NaH2PO4 sodium dihydrogenphosphate; NaKSO4 sodium potassiumsulfate.

Naming Bases

Bases are compounds that produce hydroxide ions when dissolved in water. A base has one or more ionizable hydroxide ions associated with an electropositive metal cation. Recognizing a base by name is very easy; the word hydroxide is present in nearly every base name. Since the word hydroxide is found in the names of most bases, it is not surprising that the hydroxide group will be a part of the chemical formula for a base. Examples of common bases encountered in chemistry and biology include lithium hydroxide, LiOH; sodium hydroxide, NaOH; potassium hydroxide, KOH; and ammonium hydroxide, NH4OH.

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Introduction to Mass-Mole Calculations

If there is one concept that gives students great difficulty in chemistry, it is the mole and its applications. As you learned in chapter one of the lecture course, the mole is the SI unit for the amount of a substance. In simplest terms, a mole is a counting unit that chemists use to track items and their interaction with other substances.

There are two scenarios that are typically encountered with the mole. First, a mole of any sample corresponds to the formula mass of that substance in grams . Recall that formula masses are calculated from knowing the chemical formula of the substance and the individual atomic mass values of each element found in the periodic table. The second relationship you need to be aware of is that a mole is equivalent to a specific fixed number of “entities” which can represent atoms, molecules, ions, etc. The aforementioned number of entities is represented by Avogadro’s number (NA), which is equal to 6.022 x 1023 entities per mole.

Concept mapping is quite useful in helping students solve problems. Figure 3.1 depicts the mole concept roadmap and how one can transition from mass to moles, individual particles (atoms, molecules, ions) to moles, and even from molar volumes of gases to moles. In each instance, the mole is the focal point. To quote a favorite axiom, “Moles carry us everywhere we want to go in chemistry.”

Figure 3.1. Mole Concept Roadmap.

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Experimental

The work that follows is meant to supplement lecture material and allow you to practice naming compounds in the presence of an experienced instructor. It will also help you learn the formulas and ionic charges of many basic building blocks that make up chemical compounds. In these exercises, you will practice the everyday conversational vocabulary of chemistry. Your instructor will be happy to help you with pronouncing compound and ion names.

Your instructor will have you complete either the even- or odd-numbered exercises on the report pages. When you have finished the exercises, your instructor will quiz you using several of the unassigned exercises. The quiz will be turned in to the instructor for evaluation; the remaining exercises are to be completed outside of lab and turned in the following week for a grade.

You are to complete the exercises on the data report pages. Part one of the lab activities asks you to write the chemical formulas for compounds formed from cations and anions with various charges. This exercise will teach you how to construct compound formulas, given the ions and their charges. Part two asks you to name a series of compounds. To do this, you need to visualize the individual ions that make up the formula and then name them appropriately. In part three, you will be given the names of a series of compounds and you will have to deduce the correct chemical formulas. Part four asks you to predict how compounds dissociate (“break apart”) into their individual ions; this is a necessary skill you need to master in order to construct “chemical sentences” (chemical reactions). Although we will examine chemical reactions in a later laboratory activity; predicting how substances dissociate will provide valuable insights into how starting materials transform into final products.

In a chemistry laboratory, the quantities of matter are recorded as a number of grams and as a calculated number of moles. The exercises in part five will teach you the relationship between moles and grams. Do not be afraid to call on your instructor for help if you find these last two pages difficult. Develop these basic skills now in laboratory and your lecture work will become much more enjoyable.

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Experiment 3 Language of Chemistry (Part I)

Name: Section:

Instructor: Date:

Data Report Sheet

Part I: Write the chemical formula for the compound that would be formed for the ions given.

F- OH- S2- CrO42- BO3

3-

K+ KF

Sr2+ Sr(OH)2

Zn2+ ZnS

Cr3+ Cr2(CrO4)3

Al3+ AlBO3

From this exercise, describe how the formula for an ionic compound is written.

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Experiment 3 Language of Chemistry (Part I)

Name: Section:

Instructor: Date:

Data Report Sheet

Part II: Write the chemical formula for each of the following compounds.

1. Sodium cyanide

2. Lithium sulfite

3. Tin (II) chloride

4. Chromium (III) oxide

5. Carbon disulfide

6. Ammonia

7. Barium nitrate

8. Aluminum sulfate

9. Magnesium bromide

10. Nitric acid

11. Sodium peroxide

12. Calcium bicarbonate

13. Potassium permanganate

14. Zinc phosphate

15. Sodium carbonate

16. Ammonium perchlorate

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Experiment 3 Language of Chemistry (Part I)

Name: Section:

Instructor: Date:

Data Report Sheet

Part III: Write the chemical name for each of the following compounds.

1. BaI2

2. ZnSO4

3. SrCrO4

4. P2O5

5. Ag2CO3

6. HClO4

7. Hg(NO2)2

8. Al(HSO4)3

9. LiF

10. N2Cl4

11. Co2O3

12. (NH4)2C2O4

13. CaCr2O7

14. Mg(OH)2

15. Ca(H2PO4)2

16. Cu3(AsO4)2

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Experiment 3 Language of Chemistry (Part I)

Name: Section:

Instructor: Date:

Data Report Sheet

Part IV: Write the product ions that form when the following compounds are dissolved in water.

1. NaCl yields→

2. NH4CN yields→

3. K3PO4 yields→

4. Zn(C2H3O2)2 yields→

5. Lithium perchlorate yields→

6. Copper (II) bromide yields→

7. Potassium chlorate yields→

8. Ammonium borate yields→

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Experiment 3 Language of Chemistry (Part I)

Name: Section:

Instructor: Date:

Data Report Sheet

Part V: Complete the following exercises. Show all work with proper units and significant figures.

Assume you have exactly 2.50 pounds of each compound listed below. How many moles of each compound would you have? (Note: number of moles = mass /formula mass)

Calcium oxide

Sodium nitrate

Borax

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Experiment 3 Language of Chemistry (Part I)

Name: Section:

Instructor: Date:

Data Report Sheet

Assume you have 1.25 moles of each compound listed below. How many grams of each compound would you have?

H2SO4

MgSO4 * 7 H2O

Potash

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