CHM101 Lab Chemical Compounds Grading Rubric

Transcription

1 Name Team Name CHM101 Lab Chemical Compounds Grading Rubric To participate in this lab you must have splashproof goggles, proper shoes and attire. Criteria Points possible Points earned Lab Performance Printed lab handout and rubric was brought to lab 3 Safety and proper waste disposal procedures observed 2 Followed procedure correctly without depending too much on instructor or lab partner 3 Work space and glassware was cleaned up 1 Lab Report Part A (tables correct and complete) 4 Part B (tables correct and complete) 3 Part C (table correct and complete, observations accurately recorded, detailed explanation provided) 4 Total 20 Subject to additional penalties at the discretion of the instructor.

2 Chemical Compounds When two or more elements come together, a chemical compound is formed. In this lab you will practice naming compounds and predicting their structure and properties. Naming Compounds Many everyday and compounds have common names. For example, water is the common name for H 2 O, baking soda is the common name for NaHCO 3. However, there are too many compounds (80 million +) for this memorizing common names to be practical so chemists have developed rules for naming. In this exercise you will practice rules for naming simple ionic and covalent compounds and learn to predict the ratios that chemicals combine in based on their ionic charge. Ionic Compounds with Metal Charges that Do Not Vary Simple ionic compounds are formed from a metal and nonmetal ions. The metal is the cation, or positively charged ion and the nonmetal is the anion, or negatively charged ion. In the chemical formula, the cation is always written first and the anion is written second. Cations with charges that don t vary include Group 1A alkali metals, which always form 1+ cations and Group 2A alkaline earth metals which always form 2+ cations, and Group 3A metals, which always form 3+ ions. To name this type of ionic compound, the cation gets written first using its full element name. The anion is written second using the stem of its element name and adding an ide ending. For instance, fluorine becomes fluoride and oxygen becomes oxide. Examples: NaCl sodium chloride MgBr 2 magnesium bromide Ionic compounds are electronically neutral, meaning the charges on the anions and cations balance to zero. To determine a formula from a chemical name, you must first determine the charges on the ions in the compound and then figure out how many of each are needed to balance the charges. For instance calcium chloride is made of Ca 2+ and Cl ions. For the charges to balance there needs to be two Cl ions, therefore the formula is CaCl 2. Ionic Compounds with Metal Charges that Vary Some metals can form ions with different charges. For instance, iron in ionic compounds is found as both Fe 2+ and Fe 3+. Metals with charges that vary include most transition metals (zinc, cadmium and silver are exceptions) plus tin and lead. Since more than one charge is possible, more than one chemical formula is possible. Iron can form two compounds with chlorine: Fe 2+ forms FeCl 2, Fe 3+ forms FeCl 3. These two different compounds have different physical and chemical properties. For example, FeCl 2 has a melting point of 306 C, while FeCl 3 melts at 677 C. Therefore, we cannot simply call both compounds iron chloride. When naming compounds with metal that have variable charge, a roman numeral in parentheses is written after the metal name and indicated the charge on the metal. For instance, Fe 2+ is called iron (II) and Fe 3+ is called iron (III). The anion is written as the ide form of the element.

3 When starting with the formula, the charge of the metal has to be deduced from the anion charge. The formula CuO indicates one atom of oxygen carrying a 2 charge. In order for the charges to sum to zero, the copper atom must have a 2+. Therefore the name is copper (II) oxide. The formula SnCl 4 has 4 chloride ions, each with a 1 charge, so the total negative charge is 4. The charge on the tin must be 4+, so the compound is named tin (IV) chloride. Ionic Compounds with Polyatomic Ions Polyatomic ions are ions composed of two or more covalently bonded atoms that can be thought of as a single unit carrying a charge. Ionic compounds that have three or more elements are generally composed of a metal cation and polyatomic anion. Ammonium (NH 4 + ) is the only common polyatomic cation. The names, formulas and charges of polyatomic ions must be memorized (see table of common polyatomic ions below). Ionic compounds with polyatomic ions get named in the same way as Type I and II binary ionic compounds with the name of the polyatomic ion in the place of the anion name. If more than one polyatomic ion is needed to balance the charges, the whole ion gets put in parentheses. Examples: Li 2 CO 3 lithium carbonate Pb(NO 3 ) 2 lead (II) nitrate NH 4 Cl ammonium chloride Ions with 1 charge Ions with 2 charge NO 2 Nitrite 2 SO 3 Sulfite NO 3 Nitrate 2 SO 4 Sulfate ClO 2 Chlorite 2 CO 3 Carbonate ClO 3 Chlorate CN Cyanide OH Hydroxide Ions with 3 charge HCO 3 Hydrogen carbonate 3 PO 3 Phosphite 3 PO 4 Phosphate Ions with 1+ charge + NH 4 Ammonium Covalent (Molecular) Compounds Covalent compounds are made of two or more nonmetals, as opposed to a metal and a nonmetal. The system for naming covalent compounds is different from the one for ionic compounds and it is important to keep the two systems separate. To name a covalent compound, the first element is written as its element name, the second element is written as its ide form, and both elements get a prefix that indicates the number of each element in a compound. If the prefix of the first element is mono, the prefix is dropped. Prefixes 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca Examples: P 2 S 5 diphosphorous pentasulfide N 2 O dinitrogen monoxide NO 2 nitrogen dioxide

4 Lewis Structures Lewis structures show how valence electrons are distributed and shared in covalent molecules. We represent the elements by their symbols. The shared electron pair is shown as a line/bond between the two atoms. All the other valence electrons are shown as dots or lines around the symbol of the element. Steps for drawing Lewis Structures Example with SO 2 1. Sum the valence electrons of all the atoms. The total number of valence electrons is 18 (6 from sulfur and 6 from each oxygen). 2. Use single bonds (lines) to connect the central atom to the surrounding atoms. The central atom is usually the element there is one of, such as C in CBr 4. Each line represents two electrons. Sulfur is central. After making connections to the outer atoms, we have used 4 electrons. 3. Complete the octet for all outer atoms (other than H) by drawing lone pairs. If you have not reached the total number of valence electrons, give remaining electrons to the central atom 4. If all atoms (except H) at this point have an octet the structure is finished. If, and only if, the central atom does not have an octet, turn lone pairs on outer atoms into double and triple bonds until it does have an octet. Each oxygen was given three lone pairs, with the bond they are surrounded by 8 electrons. That brought the count to 16. The two remaining electrons were put on the central sulfur for a to 18 In the last step sulfur only had 6 electrons. One lone pair on an oxygen is shared as a double bond. All atoms now have octets and the Lewis structure is complete. Molecular Shape the Valence Shell Electron Pair Repulsion (VSEPR) Theory The electron groups around the central atom repel each other and therefore prefer to be as far apart from each other as possible. This is the main idea of the VSPER theory. We can apply the VSEPR theory to predict the molecular shape/geometry of a molecule. 1. Draw the Lewis structure for the molecule in question. 2. Count the total number of electron groups on the central atom = the number of atoms bonded to the central atom plus the number of lone pairs on the central atom. Note that multiple bonds to one outer atom still count as one electron group. 3. The shape of the molecules is determined by minimizing the repulsion between the electron groups. (Spacing them out as much as possible) The table to the right summarizes the shapes based on electron groups and lone pairs:

5 Polarity of Molecules A covalent bond is polar if there is a difference in electronegativity between the bonded atoms. An entire molecule will be polar if the bond dipoles do not cancel. Polar molecules have a positive and a negative end and behave like tiny magnets. The shape of the molecule determines if dipoles cancel or not. Polar molecules bond dipoles do not cancel. Have a lone pair on the central atom OR different outer atoms. Nonpolar molecules bond dipoles cancel. Have no lone pairs on the central atom AND have all outer atoms the same. Attractive Forces Attractive forces hold molecules together into solids and liquids. The stronger the attractive forces, the higher the boiling and melting point of a compound. Ionic compounds have the strongest attractive forces and therefore have very high melting points. All the forces between covalent compounds are significantly weak, so covalent molecules tend to have low boiling and melting points. Attractive forces are summarized in the table below. found in ionic compounds found in covalent compounds strongest # weakest Attractive Force Found in Description Examples ionic forces ionic compounds electrostatic attraction between NaCl, CaBr 2 + and hydrogen bonds molecules with HF, HO, extreme case of dipoledipole. H 2 O, NH 3 or HN bonds Occurs with small, highly electronegative atoms dipoledipole forces polar molecules attraction between CH 2 Br 2, PH 3 δ + and δ dispersion forces all molecules (only force in nonpolar molecules) attraction between temporary dipoles CH 4, CO 2 Materials: buret with water buret with hexane balloon model kits acetone dropper ethanol dropper Procedure A. Naming Compounds 1. Fill in the tables with the missing name, formula or ion. Refer to naming rules in the lab background or your textbook. B. Covalent structures 2. Complete the table and use molecular models to make the shapes of each one. C. Properties of Covalent Compounds. 1. Turn on burets of water and hexane to allow a fine stream of liquid. Rub a balloon on your hair or shirt, bring it near each liquid and record your observations. 2. Put two drops of acetone and two drops of ethanol on either end of a paper towel and observe their evaporation. 3. Explain your observations based on attractive forces.

6 Chemical Compounds Name A. Complete the following table for ionic compounds Formula cation anion Name Al 2 O 3 Al 3+ O 2 aluminum oxide Cu(NO 3 ) 2 Cu 2+ NO 3 copper (II) nitrate iron (III) phosphide Pb 2+ Br Li 2 S potassium carbonate Ni(OH) 3 Mg 2+ N 3 cesium nitrite CoF 2 tin (IV) oxide (NH 4 ) 2 SO 4 calcium cyanide Complete the following table of covalent compounds Formula NO 2 Name nitrogen dioxide SO 3 diphosphorous trisulfide CBr 4 oxygen difluoride IBr Report Page 1 of 3

7 Chemical Compounds Name B. Complete the following table Formula Structure Shape Angle Polar? Y/ N Attractive Forces CS 2 linear 180 o N dispersion SBr 2 HOF (O is central) SO 3 COBr 2 (C is central) CF 4 NBr 3 Report Page 2 of 3

8 C. Fill in the table for the balloon experiment water (H 2 O) hexane (CH 3 CH 2 CH 2 CH 2 CH 2 CH 3 ) Lewis structure (all carbons connected in a chain) Polar? Y/N Attractive forces Observations Explain the difference in behavior between the two liquids: Fill in the table for the evaporation experiment acetone (CH 3 COCH 3 ) Lewis structure (connectivity is ) ethanol (CH 3 CH 2 OH) (connectivity is C C O H ) Polar? Y/N Attractive forces Observations Explain the difference in behavior between the two liquids: Report Page 3 of 3