Study guide

CHEM 120 Unit 1, Lab 2 - Laboratory 2; Metric Measurement, Density & Dimensional Analysis

Laboratory 3: Ionic and Covalent Compounds Learning Objectives: • Name ionic and covalent compounds and derive their chemical formulas. • Observe absorption spectra of metal ions using flame test. • Draw Lewis symbols of elements and Lewis formulas of simple covalent compounds. In this laboratory exercise, you will learn how to derive formulas of ionic and covalent compounds and name them. You will also flame test various metal ions and draw models of simple covalent molecules. Notes: • Read the lab before coming to class. The expectation at Chamberlain (CCN/CU) is that you come to class fully prepared for lab. • This lab will involve the use of acids and burners. Be careful and mindful of your surroundings as you handle the chemicals and burners. • Always check with your instructor regarding proper waste disposal. • Listen carefully to the professor’s instructions and work safely. • • Always follow the rules outlined in the safety contract. • If in doubt of how to use a piece of lab equipment, ask your instructor. Improper use of lab equipment can be dangerous or could damage lab materials. • • If you have any safety concerns, see your instructor. • Always dress properly for lab. Be sure to wear closed toed shoes and long pants to lab. In lab, wear your safety equipment including goggles, lab coat, and gloves. • • Use deionized water (DI water) in your experiments. Exploration 1: Explain the following terms Cation: A positively charged ion that are made when a metal loses electrons and when a nonmetal gains those electrons. Anion: An anion is negatively charged ion for example, when Cl gains one electron to form (Cl-) Ionic bond: An Ionic bond is formed between a metal and nonmetal. Polar covalent bond: Covalent bond in which atoms share a pair of electrons but not equally. Nonpolar covalent bond: Covalent bond in which atoms share a set of electrons equally. Hydrogen bond: A chemical bond in which a hydrogen atom of one molecule is attracted to an electronegative atom. Exploration 2: Covalent compounds Covalent compounds are formed by sharing electrons between two nonmetals. Unlike ionic compounds, covalent compounds are soft, waxy, or powdery in nature. Generally, covalent compounds have lower melting points compared to ionic compounds. Covalent compounds dissolve more easily in solvents such as ethanol rather than water. They are weaker in conducting electricity because they do not create ions. The rules for naming ionic compounds are as follows: Rule 1: The name of the less electronegative (treated like a cation) is always written first. The more electronegative ion (treated like an anion) is written second. • For example, in a covalent compound containing phosphorus and chlorine, phosphorus will be written before chlorine as it is less electronegative. Rule 2: The name of the cation does not change. • For example, in a covalent compound containing phosphorus (cation) and chlorine (anion), the name of phosphorus will remain the same. Rule 3: The name of the anion is changed by adding the suffix -ide. • For example, in a compound containing phosphorus (cation) and chlorine (anion), the name of chlorine will be altered by adding the suffix ide and become chloride. Rule 4: The number of each participating atom is indicated by prefixes such as mono- (one), di- (two), tri- (three), tetra-(four), penta- (five), hexa- (six), etc. It is common not to include the prefix mono- for the first element in the compound. • For example, if the compound contains one atom of phosphorus and 5 atoms of chlorine, the name of the compound will be phosphorus pentachloride. Based on the information given here, complete Table 1.   Table 1: Number Chemical Formula Name of the Covalent Compound 1 N2O4 Dinitrogen tetroxide 2 SCl2 Sulfur dichloride 3 P4S3 Phosphorus sesquisulfide 4 N2O Nitrous oxide 5 S2F2 Disulfur difluoride 6 I2O5 Iodine pentoxide 7 BrF3 Bromine trifluoride 8 HF Hydrogen fluoride To derive the chemical formula of the covalent compound, identify the quantities of participating atoms and write those as subscripts next to the chemical symbol of the element. For example, sulfur hexafluoride indicates that one atom of sulfur (S) and six atoms of fluorine (F) are present in one molecule of sulfur hexafluoride. Therefore, its formula will be SF6. Complete the following table (Table 2) by deriving the chemical formulas based upon the names of the covalent compounds. Table 2: Number Name of the Covalent Compound Chemical Formula 1 Dihydrogen monosulfide H2S 2 Tricarbon dioxide C3O2 3 Carbon dioxide C O2 4 Dichlorine penta-oxide Cl2O5 5 Diphosphorus pentasulfide P2S5 6 Nitrogen trichloride NCl3 7 Disulfur decaflouride S2F10 8 Tetraphosphorus trisulfide P4S3 Exploration 3: Flames Test- Identifying the Cation Heating different elements in a flame causes electrons in the element to move to higher energy states. When the electrons return to their ground state, the energy is given off in the form of a photon. In many cases, this photon can be observed as visible light. This phenomenon gives each metal a unique colored flame. In this exploration, you will test the flame color of different cations. You will also test the flame color of an unknown metal. Materials: Solid or 1 M solutions of lithium chloride (LiCl), potassium chloride (KCl), copper sulfate (CuSO4), calcium chloride (CaCl2), sodium chloride (NaCl), 1M hydrochloric acid (HCl), DI water, metal loop, burner Method: Note: Use NaCl last because the bright flame of sodium may interfere with the colors of the other metal ions. 1. Obtain a metal loop and a burner. Turn the burner on. Adjust it to have a mostly blue flame. 2. Clean the metal loop by dipping it into 1M hydrochloric acid (HCl) and then burning it in the flame. 3. Wait 1-2 minutes to allow the metal loop to cool. 4. Over your waste beaker, place a drop of one of the metal ion solutions on your loop and then put the loop in the flame. Record the color of the flame. a. If solid salts are given, touch the loop to the solid salt and burn it in the flame. Record the color of the flame. 5. Clean the loop by dipping in HCl and then burning it in the flame. Make sure that the color specific to a metal is completely cleared before testing another salt. a. After cleaning with HCl, rinse the loop with DI H2O and rest it on your workspace in between flame tests. b. Repeat the analysis for all cations by repeating steps 2-5. Record the observations in Table 3. 6. Table 3: Number Solution of- Color of flame 1 Lithium MAGENTA 2 Potassium PURPLE 3 Copper BLUE/ GREEN 4 Calcium 5 Sodium ORANGE 6 Unknown (STRONIUM) DARK RED Questions: 1. What is the color of the flame of the unknown solution? DARK RED 2. The unknown contains one of the cations in the table above. Based on the color of the flame, what cation is present in the unknown? STRONIUM Exploration 4: Appearance and Solubility of Ionic and Covalent Compounds We will now explore the properties of ionic and covalent compounds. In general, ionic compounds tend to be hard and crystalline in nature, whereas covalent compounds tend to be soft, waxy, or powdery in nature. In terms of solubility, due to their polar nature, ionic compounds tend to dissolve readily in water as the polar bonds of the water molecules allow them to surround the dissociated ions with partial positive or negative charges. On the other hand, ionic compounds dissolve less readily in nonpolar organic solvents, as there are no partial positive or negative charges to balance out the dissociated ions. Covalent compounds tend to dissolve readily in nonpolar organic solvents and less readily in water. Please note that some organic compounds are also soluble in water if they contain polar bonds such as those found in OH groups. Materials: weigh plates, CaCl2, NaCl, sucrose, stearic acid, 8 test tubes, test tube rack, spatula, water, ethanol, and parafilm. Method: 1. Collect a small sample of each of the four compounds on labeled weighing dishes. 2. Observe the appearance of each of the compounds and record your observations into Table 4. 3. Label two test tubes with the name of the compound and fill one with 5 mL of water and the other with 5 mL of ethanol. 4. Transfer small, equal amounts of CaCl2 into the pair of appropriately labeled test tubes. If a sample rapidly dissolves in the solvent, it is soluble in that solvent. 5. If the sample is not dissolving in a given solvent, parafilm the top of the test tube and shake for 30 seconds to see if any additional dissolution is observed. If the sample begins to dissolve, it is partially soluble in that solvent, if it does not dissolve; it is insoluble in that solvent. 6. Record your solubility observations into Table 4. 7. Repeat steps 3-6 for each of your compounds. 8. Retain test tubes containing compounds dissolved in water. Dispose of other materials according to your instructor’s directions. Table 4: Name of Compound Appearance of compound Solubility in Water (soluble, partially soluble, or insoluble) Solubility in Ethanol (soluble, partially soluble, or insoluble) NaCl CRYSTALLINE SOLID SOLUBLE PARTIALLY Sucrose CRYSTALLINE POWDER SOLUBLE PARTIALLY Stearic Acid WHITE SOLID INSOLUBLE SOLUBLE CaCl2 WHITE POWDER SOLUBLE SOLUBLE Questions: 1. Based on the appearance of the compounds, which would you expect to be ionic and which would you expect to be covalent? Explain your decision. NaCl, CaCl2 the compound I would expect to be Ionic since they are a metal and nonmetal compound. Sucrose and Stearic Acid are Covalent because they do not completely dissolve but instead remain whole molecules 2. Based on the results of the solubility study, which compounds would you expect to be ionic and which compounds would you expect to be covalent. Explain your decision. Sucrose and Stearic Acid are Covalent because they do not completely dissolve but instead remain whole molecules. NaCl & CaCl2 completely dissolve in water. Exploration 5: Conductivity of ionic and covalent compounds As we discussed in the last section, ionic compounds tend to be soluble in water and separate into ions. Ionic compounds that dissolve in water are known as electrolytes. The ions are good conductors of electricity and thus solutions made from ionic compounds dissolved in water tend to have good conductivity. On the other hand, covalent compounds are less likely to be soluble in water, and those that do dissolve in water do not produce ions. Therefore, solutions made of covalent compounds and water tend to be poor conductors. Using this information, you will use conductivity measurements to help you determine which of the compounds are ionic and which are covalent. Materials: Water soluble compounds from Exploration 4 and a conductivity meter Method: 1. Check with your instructor for how to use the conductivity meter, 2. Measure the conductivity of each of your water based solutions. 3. Record your observations in Table 5. Table 5: Name of Compound Conductivity Based on your conductivity, is this compound Ionic or Covalent? Questions: 1. Do your identifications of the compounds as ionic or covalent from the conductivity study match the results from Exploration 4? Which ones match and which ones do not? 2. Based on the results of both Explorations 4 and 5, which of your substances are ionic and which are covalent.   Exploration 6: Draw Lewis symbols and formulas of simple ionic and covalent compounds Lewis symbols, also known as electron-dot symbols are used to show the number of valence electrons of the main group elements. Valence electrons are placed around the chemical symbol of the element as dots. In order to write Lewis symbol of an element, we follow these rules: • Write the chemical symbol of the element. • Place valence electrons as single dots on all four sides of the symbol of the element (top, bottom, right and left). • Place a second dot on any one side only after there is one dot on all four sides. • On any given side, there can be no more than two dots. In this exercise, you will write Lewis symbols for simple ionic and covalent compounds and molecules. Part 6A: Lewis Symbols for Ionic Compounds Follow these steps to draw the Lewis symbols for ionic compounds. • Write the Lewis symbols of the participating elements. • Transfer the electrons from metal to nonmetal. • These atoms are now ions and have charges. Show the charges of the ions along with the numbers of electrons. See the example below. Draw the Lewis symbols to show electron transfers for the following compounds: (follow the example above) MgO KBr Li2S NaI Part 6B: Lewis Formulas for Covalent Compounds Follow the steps below to draw the Lewis formulas for covalent compounds. • Write the Lewis symbols of the participating elements. • Arrange the Lewis symbols so that the atom able to make the most bonds (highest number of unpaired electrons) is in the center. • Place the valence electrons around each Lewis symbol following the Lewis symbol structure rules. • Pair off unpaired electrons by making bonds. Start with single bonds. • When two electrons are shared, one covalent bond is formed which is shown as a line. • Each participating element will share electrons to meet the Octet rule (or in the case of hydrogen, the Duet rule). • Double and triple bonds may be formed by pairing off additional unpaired electrons. Draw Lewis structures of the following molecules and compounds. Then, determine the VSPER shape of these molecules. HCN Shape: Linear N2 Shape: Linear CH4 Shape: Tetrahegeal H2O Shape: Bent NH3 Shape: Trigonal Reflection: Refer to the objectives of this lab and reflect on four key concepts that you learned in this lab exercise. Be specific in your answer (this should require 5-10 sentences). Grading Rubric: Activity Deliverable Points Exploration 1 Complete definitions 3 Exploration 2: Complete table 1: 4 points, and table 2: 4 points 8 Exploration 3: Complete table 3: 2 points and the questions: 2 points 4 Exploration 4: Complete table 4: 2 points and the questions: 2 points 4 Exploration 5: Complete table 5: 2 points and the questions: 2 points 4 Exploration 6: Complete each drawing and state shape of the molecule 5 Reflection Write a 5-10 sentence laboratory reflection: 2 All Lab Deliverables Complete ALL explorations and reflection activities 30