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.

FOR CRIMINOLOGY CHAPTER 1-INTRO TO FORENSIC CHEM-LECTURE

Forensic Chemistry Forensic science is the application of scientifi c principles to matters involving the law. This area of science is generally considered quite fascinating and it continues to experience growing popularity. Many would agree that the current public interest in forensics is a direct result of CSI-related television programming. These weekly shows have brought a once relatively unknown area of science to the forefront of public mainstream. Viewers are captivated and intrigued by well-informed scientists working in spotless labs with ominous lighting and a modern music background. The use of cutting-edge technology provides last-minute revelations culminating in the solution of a complex crime. These programs are entertaining and have certainly increased public awareness to the fi eld of forensics; but alas, television is not reality. Although it is true that forensic science has experienced tremendous growth, few would (or should) believe this to be the result of fi ctional television programming. Media coverage of high-profi le cases has increased over the last decade in both numbers and content. Crime-scene investigation and forensic analysis have been brought out of the lab and into the public’s “scrutinizing eye.” Forensic science, once a broad fi eld, has become segregated into highly specialized disciplines. For example, forensic chemistry, forensic pathology, forensic dentistry, forensic entomology, and forensic DNA analysis have evolved into independent fi elds of forensics. It seems more appropriate – and clearly more realistic – to attribute the unprecedented popularity of forensic investigation to enhanced public awareness and an increase in the availability of career opportunities. Chemistry is the study of the composition of matter and the changes it undergoes. Forensic chemistry is a specialized area of forensic science involving the application of chemical principles and techniques to the fi eld of forensic investigation. The role of forensic chemistry in criminal investigations is vast and ranges from techniques used to collect and preserve evidence, to complex chemical procedures used to identify elements and compounds. Identifi cation procedures are highly reliable and are frequently based on the chemical and physical properties of the substance supported by data obtained from analytical analysis. Most chemical techniques used for isolation, purifi cation, and identifi cation are valid forensic techniques; however, chemical analysis differs from forensic chemical analysis in two ways: regulatory and judiciary. The results of forensic investigation may have a serious impact on lives. Therefore, techniques performed during forensic analysis must be closely regulated to ensure the accuracy and integrity of experimental results. Forensic laboratories must develop two operating manuals designed to meet the specifi c needs of each laboratory. The technical procedures manual outlines the step-by-step details of all procedures used in forensic examinations. The quality- control manual is designed to maintain the highest standards of reliability and integrity of work done by scientists in the lab. Adherence to both the technical procedures manual and lab quality manual is a crucial part of any analysis and is strictly enforced both internally and externally. Internal quality control includes, but is not limited to, periodic instrument calibration, checking reagents for expiration, and performance evaluations on scientists working in the laboratory. In addition, a detailed record is kept of all internal quality procedures performed. Outside regulatory agencies are responsible for external quality control and these agencies may vary from state to state in the US. The American Society of Crime Laboratory Directors (ASCLD) has recently accepted the painstaking task of regulating various fi elds within forensic science worldwide. This includes the forensic chemistry section in the United States. ASCLD is the regulatory organization responsible for supervising, evaluating, and directing all laboratories within its membership. Their designated inspectors evaluate technical staff and conduct periodic site inspections to ensure the highest standards of quality and technical performance. The efforts of ASCLD have helped to streamline and 4 1 Introduction standardize forensic analytical techniques worldwide. In addition, ASCLD provides direction and qualifi ed solutions to potential issues facing member laboratories. Courtroom presentation of scientifi c principles and techniques used during forensic examination is the judiciary responsibility of the forensic chemist. Forensic chemists are often called upon to describe complex chemical procedures to individuals who have a limited understanding of scientifi c principles. This responsibility can present a variety of challenges to the forensic chemist as an expert witness. Courtroom testimony is carefully prepared using common terminology and the presentation must be in a clear, simple manner that avoids confusion and misinterpretation. To achieve this, forensic chemists often use common analogies to describe complex chemical and analytical techniques. For example, a gas chromatograph is an instrument used to separate a gaseous mixture into individual components based on size and/or charge. The description of how a gas chromatograph functions may contain a reference to coin-separating machines frequently found in local grocery stores. A coin machine separates the mixture of coins based on size, and totals each pile based on weights. This analogy would illustrate how a gas chromatograph functions and may help members of a jury be more comfortable with testimony about this complex instrument. Similar analogies will be used in the following chapters to describe complex chemical procedures and analytical techniques frequently used in forensic chemistry. These analogies are designed to promote an understanding of the topic under discussion while adding clarity and continuity to the subject. 1.2 Scientifi c Investigation Imagine yourself in a classroom for an extended period of time without the ability to see outside. When you exit the building, you immediately notice that the ground is wet. Your fi rst thought is that it rained while you were inside. To confi rm this, you look to the sky to identify rain clouds. If the sky is cloudy, you are reasonably sure that it rained. If the sky is clear, you consider another possibility – perhaps sprinklers wet the ground. To confi rm this, you look for sprinklers in the immediate area. If they are found, you are reasonably sure of why the ground is wet. If no sprinklers are found, you consider another possibility and the cycle repeats. Each time you consider a possible cause, you search for supporting evidence to confi rm that cause. You accept or reject a possibility based on the presence or absence of supporting evidence. In the above scenario, you observe a water truck spraying an adjacent construction site. You are now reasonably sure of how the ground became wet. The wet ground was your observation . The possibility of rain was your fi rst hypothesis . Searching the sky for clouds was your experimentation . The absence of clouds in the sky caused you to reject your hypothesis . Other hypotheses were considered and subsequently rejected based on a lack of supporting evidence. Finally, the water truck hypothesis was confi rmed when you saw the truck in the immediate area. Your determination that the water truck wet the ground is your conclusion or theory . This deductive procedure is termed the scientifi c method : the process used to form theories . It begins with an observation : the discovery and recognition of some type of unexplained phenomenon. The observation is followed by the hypothesis : the proposal of a possible cause of the observation. The hypothesis is tested during the experimentation phase using experiments specifi cally designed to prove the hypothesis. If experimental results do not support the hypothesis, another possibility is considered and tested. If the experiments are successful and repeatable, the hypothesis becomes a theory and is presented to the scientifi c community. 1.3 Forensic Investigation Imagine a distant planet, similar to earth, with diverse climates and distinctly different environments across its surface. Now imagine that four space programs on earth send their astronauts to the new planet that, by chance, land in different regions characterized as a desert landscape, a tropical rainforest, a frozen landscape, and mountainous landscape. The astronauts explore their regions collecting samples, data, and video from their distinctly different environments. They return to their respective countries with a description of the planet supported by evidence collected during exploration. Each space program presents their information to the world, but the views are confl icting. Each country defends their position and accuses the others of presenting false or misleading information. Whom do you believe? Intuitively, you trust your astronauts and reject the other three despite the fact that, in reality, each is truthful and correct. It is not uncommon for different forensic scientists to arrive at different conclusions after examining the same piece of evidence. This is acceptable, if not expected, in the fi eld of forensic investigation. The results of forensic examinations must never be accepted or rejected because you know or trust one scientist more than another. You must keep an unbiased, open mind, knowing that two or more scientists may present different perspectives when evaluating the same piece of evidence. 1.5 Physical Properties 5 An unfortunate aspect of forensic investigation is that the results of your examination will always have a negative impact on one party. If the evidence supports the suspect’s innocence, the victim is unhappy; if it supports the suspect’s guilt, the suspect is unhappy. This is both unfortunate and unavoidable; however, it is the duty of the forensic chemist to present the unbiased story of the evidence. 1.4 Properties of Matter Matter is anything that has mass and occupies space. It is diffi cult to imagine something that has mass that does not occupy space, or something that occupies space that does not have mass. Do not spend too much time pondering the previous, I cannot think of anything either (perhaps something on the previously referenced imaginary planet). Despite the apparent redundancy in the defi nition of matter, it must satisfy the two parameters. There is a difference between the mass of an object and its weight. Weight is a force resulting from the pull of gravity on a given mass. Mass is defi ned as a specifi c quantity of matter and is not affected by the pull of gravity. The weight of an object on earth will be different from its weight on the moon because the force of gravity is different. The mass of an object will be constant at these locations despite the differences in gravitational fi eld strength. For this reason, the term “mass” should always be used in any area of science when referring to “weight.” There are three states (or phases) of matter: solid, liquid, and gas. Solids have a defi ned volume and a fi xed shape ; liquids have a defi ned volume and undefi ned shape – they conform to the shape of their container; and gases have an undefi ned volume and undefi ned shape – they take the shape and volume of the container holding the gas. Elements are the fundamental building blocks of all matter . The symbols used to identify all known elements can be found on the periodic table, an arrangement of the elements based on atomic properties. For example, “H” represents the element hydrogen and “O” represents the element oxygen. Compounds are formed through the combination of two or more elements . Chemical formulas are used to represent compounds. They specify the identity and relative number of each atom present using symbols from the periodic table and subscripts attached to each symbol. For example, the chemical formula for water is H 2 O, a compound containing two atoms of the element hydrogen (note subscript 2 attached to H) and one atom of the element oxygen. Elements and compounds may exist as pure substances or as mixtures. Pure substances contain only one component and have the same composition throughout, for example, pure gold, pure sugar, and pure water. Mixtures contain two or more pure substances and may be homogeneous or heterogeneous. Homogeneous mixtures have the same composition and properties throughout . They are not pure substances because they contain more than one component. For example, pure sugar water is a homogeneous mixture containing sugar and water. It has the same sweetness throughout; however, evaporating one component (the water) will produce the other (sugar crystals). Heterogeneous mixtures have distinctly different properties within the mixture ; water and sand would be an example. The sand and water are easily identi- fi ed, regardless of the degree of mixing. There are fundamental properties associated with all forms of matter. These distinguishing characteristics may be physical or chemical in nature and are frequently used to identify and classify a particular substance.

CHAPTER 1. INTRO-FORENSIC CHEMISTRY

CHPTER TEST 1 FOR CRIMINOLOGY The Periodic Table And Atomic Structure Test Multiple Choice Identify the choice that best completes the statement or answers the question. 31) The smallest particle into which an element can be divided and still have the properties of that element is called a(n) A) nucleus. B) electron. C) atom. D) neutron. 32) How would you describe the nucleus? A) dense, positively charged B) mostly empty space, positively charged C) tiny, negatively charged D) dense, negatively charged 33) Where are electrons likely to be found? A) in the nucleus B) in electron clouds C) mixed throughout an atom D) in definite paths 34) Every atom of a given element has the same number of A) protons. B) neutrons. C) electrons. D) isotopes. 35) What is the meaning of the word atom? A) dividable B) invisible C) hard particles D) not able to be divided 36) Which statement is true about isotopes of the same element? A) They have the same number of protons. B) They have the same number of neutrons. C) They have a different atomic number. D) They have the same mass. 37) Which of the following has the least mass in an atom? A) nucleus B) proton C) neutron D) electron 38) If an isotope of uranium, uranium-235, has 92 protons, how many protons does the isotope uranium-238 have? A) 92 B) 95 C) 143 D) 146 39) A carbon atom with 6 protons, 6 electrons, and 6 neutrons would have a mass number of A) 6. B) 18. C) 12. D) 15. 40) Periodic describes A) something that occurs at regular intervals. B) something that occurs very rarely. C) something that occurs frequently. D) something that occurs three or four times a year. 41) The elements in each vertical column on the periodic table usually have similar properties and are called a(n) A) period. B) group. C) element. D) property. 42) The elements to the right of the zigzag line on the periodic table are called A) metalloids. B) conductors. C) metals. D) nonmetals. 43) Most metals are A) solid at room temperature. B) bad conductors of electric current. C) dull. D) not malleable. 44) The properties of elements in a horizontal row of the periodic table follow A) a nonrepeating pattern. B) a periodic pattern. C) no pattern. D) an unpredictable pattern. 45) Elements on the periodic table are arranged in order of A) increasing density. B) decreasing density. C) increasing atomic number. D) decreasing atomic number. 46) Which of the following statements describes most metals? A) They are easily shattered. B) They are gases at room temperature. C) They are dull. D) They are good conductors of electric current. 47) Which of the following is a property of alkali metals? A) They are so hard they cannot be cut. B) They are very reactive. C) They are stored in water. D) They have few uses. 48) Most of the elements in the periodic table are A) metals. B) metalloids. C) gases. D) nonmetals. 49) A horizontal row on the periodic table is called a(n) A) group. B) family. C) period. D) atomic number. 50) Elements lying along the zigzag line on a periodic table are A) metals B) nonmetals C) metalloids D) noble gases 51) How do the physical and chemical properties of the elements change? A) periodically within a group B) periodically across each period C) periodically within a family D) periodically across each group 52) Transition metals are A) good conductors of thermal energy. B) more reactive than alkali metals. C) not good conductors of electric current. D) used to make aluminum. 53) The letter C is carbon’s A) atomic number. B) atomic mass. C) chemical symbol. D) element name. 54) The number at the top is the A) atomic number. B) element name. C) atomic mass. D) chemical symbol. 55) The carbon group has two metalloids, both of which are used to make A) dinnerware. B) foil. C) cans. D) computer chips. 56) Which one of the following tells the physical state of an element at room temperature? A) the atomic number B) the color of the chemical symbol C) the atomic mass D) the element name The Periodic Table And Atomic Structure Test