Forensic Chemistry STudents : Experiment No. 4 and 5

Laboratory Manual Experiment # 4 Name_______________________________ Mixtures and Compounds Reference: Chapters 1, 2, and 3 Objective: Students will observe common properties of mixtures and compounds. Materials: methanol, DI water, sugar, salt, analytical balance, watch glass, and oven. Introduction: Elements and compounds may exist as pure substances or as mixtures. Pure substances contain only one component and have the same composition throughout, i.e., pure gold, pure sugar, pure water, etc. Mixtures contain two or more pure substances and may be homogeneous or heterogeneous. Homogeneous mixtures have the same composition and properties throughout. However, they are not pure substances because they contain more than one component. Heterogeneous mixtures have distinctly different properties within the mixture; water and sand would be an example. In any binary solution (a solution that contains only two components), the solvent is the component present in greatest amount and the solute is the component present in least amount. The following mixtures will be provided. Classify each mixture by circling homogeneous or heterogeneous. Solution #1 Sugar in water (sat.) homogeneous or heterogeneous Solution #2 Salt in water (sat.) homogeneous or heterogeneous Solution #3 Sugar in methanol (sat.) homogeneous or heterogeneous Solution #4 Salt in methanol (sat.) homogeneous or heterogeneous Part A: Clean and dry four watch glasses and label each 1, 2, 3, or 4. Weigh each empty watch glass on an analytical balance and record the mass in the table below under “watch glass.” Be sure to weigh the watch glasses after they are labeled! Place 1.0 ml of the corresponding solutions above on each of the labeled watch glasses, i.e., place 1.0 ml of solution #1 on watch glass labeled 1, etc. Using the same analytical balance that was used to weigh the empty watch glasses, carefully weigh the watch glasses containing each solution. Determine the mixture mass for each solution by subtracting the mass of the empty watch glass from the mass of the watch glass containing solution. Record the mass of each mixture in the table below under “mixture mass.” Save the watch glasses containing each solution for Part B. Clean and dry two small test tubes and place one test tube into a small beaker. Place the test tube/beaker on the analytical balance and tare the balance (zero the balance with the test tube/beaker on the pan). Place 1.0 ml of water into the test tube and record the mass in the table below under Solvent Mass for water. Repeat the procedure using the other test tube and 1.0 ml of methanol. Record the mass below under Solvent Mass for methanol (water is the solvent in solutions 1 and 2, methanol is the solvent in solutions 3 and 4). Subtract the solvent mass from the mixture mass and record the difference in the table below under “solute mass.” Watch Glass Mixture Mass Solvent Mass Solute mass #1- Water- #2- Water- #3- Methanol- #4- Methanol- What is the mass of sugar in mixture #1? What is the mass of salt in mixture #2? What is the mass of sugar in mixture #3? What is the mass of salt in mixture #4? Laboratory Manual 291 Did the mixture mass exceed the solvent mass in any solution? If so, explain. Did the solvent mass exceed the mixture mass in any solution? If so, explain. Did the solvent mass equal the mixture mass in any solution? If so, explain. Explain your observed results using your knowledge of homogeneous and heterogeneous mixtures. Part B: Carefully place the watch glasses containing each solution in the oven and evaporate the solvent to dryness. When evaporation is complete, weigh each watch glass and record the mass in the table below under “watch glass/residue.” Determine the mass of the residue by subtracting the mass of the empty watch glass (measured in Part A) from the mass of the watch glass/ residue. The residue is the actual mass of solute contained in each solution. Record the residue mass in the table below under “solute mass (actual).” Watch Glass/Residue Solute Mass (actual) What is the actual mass of sugar in mixture #1? What is the actual mass of salt in mixture #2? What is the actual mass of sugar in mixture #3? What is the actual mass of salt in mixture #4? Conclusion: (Hint: Defi nition of solution; did any of the 1.0 ml solutions actually contain 1.0 ml of solvent?) 292 Laboratory Manual Chemical Formulas and Nomenclature Reference: Chapters 2 and 3 Objective: Students will gain experience writing chemical formulas for ionic compounds. Students will learn formal procedures used to name ionic and covalent compounds. Introduction: Substances are either elements or compounds. A compound is a substance that consists of two or more elements bonded together in a specifi c way. The forces that hold atoms together in a compound are called chemical bonds. An ionic bond involves the transfer of electrons from a metal to a nonmetal . A covalent bond consists of a pair of electrons shared between two nonmetals . Ions and the Octet Rule Atoms are electrically neutral because they have an equal number of electrons and protons. An atom can be converted into a charged particle called an ion by losing or gaining one or more electrons. The loss of electron(s) from a neutral atom produces a positively charged ion called a cation (pronounced cat-ion). The gain of electron(s) by a neutral atom produces a negatively charged ion called an anion (pronounced an-ion). Generally, the charge on an ion can be predicted from the position of the element on the periodic table. The metals (on the left-hand side of the table) lose electrons to form cations. The Group IA elements lose ONE electron to achieve an octet and take a charge of 1 positive. This is correctly written using a superscript “+” attached to the upper right side of the elemental symbol, i.e., Na + . Notice that the number “1” is not written when the cation carries a positive one charge. The Group IIA elements lose TWO electrons and take a charge of 2 positives. This is correctly written using a superscript “2+” attached to the upper right side of the elemental symbol, i.e., Mg 2+ . When cations carry a charge greater than one, the number is written fi rst, followed by the sign. The Group IIIA elements lose THREE electrons and take a charge of 3 positives which is written as a superscript “3+”, i.e., Al 3+ . The nonmetals (on the right-hand side of the table) gain electrons to form anions. The Group VIIA elements gain ONE electron to achieve an octet and take a charge of 1 negative. This is written using a superscript “-“ attached to the elemental symbol, i.e., Cl - . Once again, the “1” is not written. The Group VIA elements gain TWO electrons and take a charge of 2 negatives which is written as a superscript “2-“, i.e., O 2- . The Group VA elements gain THREE electrons and take a charge of 3 negatives which is written as superscript “3-“, i.e., N 3- . Some transition metals and metals in Group IVA have variable charges (more than one positive ion is possible). See table below. Some common ions and their location on the periodic table. IA IIA IIIA IVA VA VIA VIIA H + Li + Be 2+ N 3- O 2- F - Na + Mg 2+ Al 3+ P 3- S 2- Cl - K + Ca 2+ Fe 2+ Fe 3+ Co 2+ Co 3+ Ni 2+ Cu + Cu 2+ Zn 2+ Br - Rb + Sr 2+ Ag + Sn 2+ Sn 4+ I - Cs + Ba 2+ Hg 2 2+ Hg 2+ Pb 2+ Pb 4+ Experiment # 5 Name __________________________________ Laboratory Manual 293 Writing Formulas for Ionic Compounds Ionic compounds are electrically neutral . Therefore, when writing formulas, the cations (positive) and anions (negative) must combine to produce a net charge of zero. In the formula, the cation (metal) is always written fi rst, followed by the anion (nonmetal). The number and types of each element must be clearly shown in the formula; the type of element is indicated using the elemental symbol, and the number of each element is indicated using a subscript attached at the lower right side of the symbol. The number “1” is not written in cases requiring only a single element. Formulas for ionic compounds are called formula units . The correct ratio required to produce a net charge of zero when Na + ions combine with Cl - ions is one to one because one Na + cancels one Cl - . Therefore, the formula is NaCl. Notice this is not written Na 1 Cl 1 . The correct ratio when Na + ions combine with O 2- ions is two to one because two Na + are required to cancel one O 2- . The 2 atoms of Na are indicated in the formula using a subscript “2” directly attached to Na. The formula is Na 2 O. The correct combining ratio when Na + ions and P 3- ions combine is: Na 3 P (three to one). Practice Examples: Write the formula for the ionic compound that is formed when each of the following pairs of ions interact: a) K + and S 2- b) Mg 2+ and O 2- c) Ca 2+ and I - d) Li + and N 3- e) Al 3+ and S 2- Solution (a) The cation has a charge of 1+ because K is a member of Group IA. The anion has a charge of 2- because S is member of Group VIA. Thus, two positive ions (2+) are required for each negative ion (2-) to produce an electrically neutra l formula unit. The formula is K2 S . (b) The cation has a charge of 2+ and anion has a charge of 2-. The ratio is 1:1. The formula is MgO . (c) The cation has a charge of 2+ and anion has a charge of 1-. Two negative ions are required for each positive ion. The formula is CaI2 . (d) The cation has a charge of 1+ and anion has a charge of 3-. Three positive ions are required for each negative ion. The formula is Li3 N . (e) The cation has a charge of 3+ and anion has a charge of 2-. Two positive ions are required for three negative ions. Here, the lowest common factor of 3+ and 2- is 6 (without sign). The formula is Al2 S3 . Naming Ions The names of cations and anions are determined by a system developed by the International Union of Pure and Applied Chemistry (IUPAC). Metals That Form Only One Type of Positive Ion Elements in Groups IA, IIA, IIIA, and some transition elements form only one type of cation. For these ions, the name of the cation is the elemental name of the metal followed by the word “ion.” Cations can now be differentiated from their corresponding neutral forms using specifi c names. For example, K is potassium (neutral form) and K + is potassium ion (cationic form). Na + sodium ion K + potassium ion Mg 2+ magnesium ion Al 3+ aluminum ion Ag + silver ion Zn 2+ zinc ion Metals That Form Two Different Positive ions Metals in Group IVA, and most transition metals, form more than one type of cation and the charge must be included in the name. For these ions, the name of the cation is the elemental name of the metal followed by a Roman numeral in parentheses, with no space after the name. The Roman numeral indicates the positive charge on the ion. Technically, the names do not end with the word “ion,” although some still prefer to include it. 294 Laboratory Manual Sn 2+ tin(II) Sn 4+ tin(IV) Pb 2+ lead(II) Pb 4+ lead(IV) Cu + copper(I) Cu 2+ copper(II) Fe 2+ iron(II) Fe 3+ iron(III) Co 2+ cobalt(II) Co 3+ cobalt(III) Hg 2 2+ mercury(I) Hg 2+ mercury(II) Naming Anions Anions are named by replacing the last part of the elemental name with the suffi x –ide , and adding the word “ion”. Anions can now be differentiated from their corresponding neutral forms using specifi c names. For example, S is sulfur (neutral form) and S 2- is sulf ide (anionic form). F - fl uoride ion Cl - chloride ion Br - bromide ion I - iodide ion O 2- oxide ion S 2- sulfi de ion N 3- nitride ion P 3- phosphide ion Polyatomic Ions A polyatomic ion is an ion that contains two or more elements. It is recommended that you memorize the names and formulas of the following polyatomic ions: NH 4 + ammonium SO 3 2- sulfi te CN - cyanide SO 4 2- sulfate OH - hydroxide HSO 3 - hydrogen sulfi te C 2 H 3 O 2 - acetate HSO 4 - hydrogen sulfate CrO 4 2- chromate PO 3 3- phosphite Cr 2 O 7 2- dichromate PO 4 3- phosphate MnO 4 - permanganate HPO 4 2- hydrogen phosphate NO 2 - nitrite ClO - hypochlorite NO 3 - nitrate ClO 2 - chlorite CO 3 2- carbonate ClO 3 - chlorate HCO 3 - hydrogen carbonate ClO 4 - perchlorate The common name for HCO 3 - , HSO 3 - , and HSO 4 - are bicarbonate, bisulfi te, and bisulfate respectively. Naming Ionic Compounds Binary ionic compounds containing metals that form only one type of positive ion These compounds contain only two types of elements; a metal ion and a nonmetal ion. Note: binary refers to element types, not total number of atoms. For example; MgBr 2 contains 3 total atoms, 1 Mg and 2 Br’s, but it contains only two types, Mg and Br. Also, recall all elemental symbols begin with capital letters, if you have a binary compound, your formula contains only two capital letters. The metal is always named fi rst using the elemental name of the metal. The nonmetal is named second using the anionic name of the nonmetal (elemental name modifi ed with the suffi x –ide ). Practice Examples: Name the following binary ionic compounds: NaCl MgBr 2 AlP K 2 S SrF 2 ZnI 2 Solution NaCl sodium chloride K 2 S potassium sulfi de MgBr 2 magnesium bromide SrF 2 strontium fl uoride AlP aluminum phosphide ZnI 2 zinc iodide Binary ionic compounds containing metals that form two different positive ions These compounds are essentially named using the same procedure developed for metals forming only one type of cation. The distinction is that the charge on the cation must be written as a Roman numeral in parentheses immediately after (with no space) the metal name. Laboratory Manual 295 Practice Examples: Name the following binary ionic compounds: FeBr 3 CoF 2 SnO PbI 4 HgS Cu 3 P Solution FeBr 3 iron(III) bromide PbI 4 lead(IV) iodide CoF 2 cobalt(II) fl ouride HgS mercury(II) sulfi de SnO tin(II) oxide Cu 3 P copper(I) phosphide Ionic compounds containing polyatomic ions Identifying compounds containing polyatomic ions is somewhat simplifi ed by the fact that all elemental symbols begin with capital letters. If you identify more than two capital letters in the formula, your compound contains a polyatomic ion that you must immediately recognize (the value of memorizing!). Naming these compounds is simply based on your familiarity with the polyatomic ions. The cation is named fi rst using its elemental name, followed by the name of the polyatomic ion. Practice Examples: Name the following polyatomic ionic compounds: Ca(NO 3 ) 2 ZnSO 4 NH 4 CN Li 3 PO 4 Na 2 CO 3 Mg(HCO 3 ) 2 Solution Ca(NO 3 ) 2 calcium nitrate ZnSO 4 zinc sulfate NH 4 CN ammonium cyanide Li 3 PO 4 lithium phosphate Na 2 CO 3 sodium carbonate Mg(HCO 3 ) 2 magnesium hydrogen carbonate Compounds containing only nonmetals (molecular compounds) This type of compound contains covalent bonds, so the concept of cations and anions is somewhat obscure and does not necessarily apply. Regardless, they are essentially named using the previously developed procedure for naming binary compounds containing a metal. The fi rst element is named using its elemental name and the second is named using its anionic name. The difference is the use of Greek prefi xes attached to each name, which indicate the number of each element present in the formula. There is one important exception; the prefi x “mono” is never attached to the name of the fi rst element in the formula. Let us look at SF 6 as an example. You should immediately notice that S and F are nonmetals because both are located in the nonmetal section of the periodic table (right side). This observation should immediately trigger “prefi xes” in your mind. SF 6 is sulfur hexafl uoride because the formula indicates one sulfur (note the absence of the subscript “1” attached to S) and 6 fl uorides (subscript 6 attached to F). This is not monosulfur hexafl uoride because “mono” is never used with the fi rst element. It is possible, however, to attach all other prefi xes to the name of the fi rst element, i.e., N 2 O 5 is dinitrogen pentoxide. The Greek prefi xes are listed below, note that all end in a vowel. When the prefi x ends with an “a” or “o”, and the elemental name begins with an “a” or “o”, the vowel of the prefi x is usually dropped to simplify pronunciation. Notice in our N 2 O 5 example above, the name of the anion is pentoxide, not pentaoxide. Greek prefi xes 1 (mono-) 2 (di-) 3 (tri-) 4 (tetra-) 5 (penta-) 6 (hexa-) 7 (hepta-) 8 (octa-) 9 (nona-) 10 (deca-) Practice Examples: Name the following binary molecular compounds: N 2 O 5 CO 2 P 4 S 3 XeF 6 ICl NH 3 I 4 O 9 CO H 2 O H 2 O 2296 Laboratory Manual Solution N 2 O 5 dinitrogen pentoxide NH 3 nitrogen trihydride (ammonia) CO 2 carbon dioxide I 4 O 9 tetraiodine nonoxide P 4 S 3 tetraphosphorous trisulfi de CO carbon monoxide XeF 6 xenon hexafl uoride H 2 O dihydrogen monoxide (water) ICl iodine monochloride H 2 O 2 dihydrogen dioxide Exercises Inspect the periodic table and list ALL elements with a ONE letter abbreviation. Symbol Name Period number Group number metal, nonmetal, or metalloid Formulas of Ionic Compounds Complete the following table with the formula of the compound. Br - O 2- NO 3 - PO 4 3- Na + Mg 2+ Al 3+ Pb 4+ NH 4 + Fe 2+ Fe 3+ Ionic Compounds Write the correct name for each of the following ionic compounds. Formula Name Pb(HCO 3 ) 4 Al 2 S 3 LiHSO 4 Zn 3 (PO 4 ) 2 CoF 3 Ca(CN) 2 SnO 2 Na 2 CrO 4 K 2 CO 3 Cu 3 P Sr(OH) 2 (NH 4 ) 2 HPO 4 Hg 2 Cl 2Laboratory Manual 297 Formula Name BaSO 4 Sn(NO 3 ) 2 AgClO 3 Cu(HSO 3 ) 2 Molecular Compounds Name each of the following compounds. Formula Name BBr 3 Br 3 O 8 CI 4 C 3 O 2 Cl 2 O 7 IF 5 I 2 O 5 NCl 3 N 2 O 5 OF 2 P 4 S 3 P 4 S 9 P 4 O 10 SF 6 S 2 Cl 2 SiS 2 SiBr 4 XeF 2 XeO 4 XeF 6298 Laboratory Manual Experiment # 6 Name ________________________________ Solubility Reference: Chapter 3 Objective: Students will test the solvent properties of various liquids to observe and understand the chemical nature of solubility and miscibility. Materials: acetone, chloroform, ammonia, methanol, water, test tubes, and pipettes. Introduction: Water is a common solvent in many solutions and substances like salt and sugar readily dissolve in water. Any substance that dissolves appreciatively in a specifi ed solvent is said to be soluble in that solvent. Technically, the term solubility refers to a quantitative maximum amount of substance that can dissolve in a given volume of solvent at a specifi c temperature. The ability of a substance to dissolve in a particular solvent depends on the identity of both the solvent and the substance; the general rule is “likes dissolve likes.” Water is a polar covalent molecule and, as a solvent, can dissolve similar molecules (polar covalent). The polarity of water is responsible for its remarkable solvent properties and explains why ionic compounds (i.e., NaCl) and polar covalent compounds (i.e., sucrose, ammonia) are soluble, whereas nonpolar molecules (i.e., organic, gasoline, oils) are not. Terms such as soluble, slightly soluble, insoluble, and solubility are used to describe the ability of a substance to dissolve in a solvent. Intuitively, we associate the term “dissolving” with solids and liquids; however, a liquid may also be soluble in another liquid. For example, when 25.0 ml of ether is added to 25.0 ml of water, the resulting total volume is not 50.0 ml, in fact, it is slightly less (approx. 48.5 ml). This is the result of solubility; ether and water are slightly soluble in one another and consequently, the volumes are not additive. The solubility of one liquid in another is diffi cult to determine and is usually not readily observed upon mixing. For this reason, liquids are often characterized using their ability to mix with other liquids rather than their solubility in other liquids. The degree of mixing between two liquids is described using the terms miscible and immiscible . Two liquids are miscible (soluble) if a uniform solution results after mixing (i.e., water and ammonia). Two liquids are immiscible (insoluble) if two distinct layers form after mixing (i.e., oil and water). Water is miscible in polar liquids and immiscible in nonpolar (organic) liquids. Part A: Add 1.0 ml of each of the following reagents to four separate, clean, and dried test tubes: chloroform, ammonia, methanol, water. Add 1.0 ml of acetone to each test tube and mix. Observe the results and characterize the liquids as miscible ( M ) or immiscible ( I ). Record your results in the table below under the corresponding reagents. Clean and dry three test tubes and add 1.0 ml of each of the following reagents to separate test tubes: ammonia, methanol, water. Add 1.0 ml of chloroform to each test tube and mix. Record your results in the table below under the corresponding reagents. Clean and dry two test tubes and add 1.0 ml of each of the following reagents to separate test tubes: methanol, water. Add 1.0 ml of ammonia to each test tube and mix. Record your results in the table below under the corresponding reagents. Clean and dry one test tube and add 1.0 ml of methanol and 1.0 ml of water to the test tube and mix. Record your results in the table below under the corresponding reagents. Reagents Chloroform Ammonia Methanol Water Acetone Chloroform -------------- Ammonia -------------- ------------ Methanol -------------- ------------ ------------ Laboratory Manual 299 Identify the organic liquids: Identify the inorganic liquids: What liquids are polar? What liquids appear to be non-polar? List the liquid-pairs that are miscible: List the liquid-pairs that are immiscible: Are immiscible liquids soluble in one another? Briefl y explain. Are miscible liquids soluble in one another? Briefl y explain.