Mega Molecules, LLC!!!!! Name: Hands-On Science with Molecular Models!! Date:!!!!!!!! Hour: Introduction Covalent Molecules and Lewis Structures Time required: two 50-minute periods To study covalent molecules, chemists find the use of models helpful. Colored plastic spheres represent atoms. The spheres have poles which represent the number of covalent bonds they can form. The poles are at specific angles that approximate the accepted bond angles. The links are used to represent bonds. Single bonds are represented by one link between two atoms, while double and tripe bonds are shown with two or three links between two atoms. In this activity, you will determine the type of bonds in a substance and draw and construct models of molecules. Materials Covalent Bonding and Lewis Structures Model Set Part I: Types of Bonds There are two types of bonds - ionic bonds and covalent bonds. They differ in their structure and properties. Covalent bonds consist of pairs of electrons shared by two atoms. Whether two atoms can form a covalent bond depends upon their electronegativity. Electronegativity is a scale used to determine an atom s attraction for an electron in the bonding process. 1
Differences in electronegativities are used to predict whether the bond is nonpolar covalent, polar covalent, or ionic. Bonds with an electronegativity difference of 0-0.4 are considered to be nonpolar covalent. Bonds that exhibit an electronegativity difference of more than 0.4 but less than 1.7 are classified as polar covalent. If two atoms have an electronegativity difference of 1.7 or more, such as sodium and chloride, then one of the atoms will lose its electron to the other atom. This results in a positively charged ion (cation) and negatively charged ion (anion). The bond between these two ions is called an ionic bond. Table 1: Electronegativity and Bond Type Electronegativity Difference Type of Bond Percent Ionic Character 0-0.4 nonpolar covalent bond 0-4% Greater than 0.4 to 1.7 polar covalent bond 5-49% Greater than or equal to 1.7 ionic bond 50-100% Use the electronegativity values on the periodic table of the elements and the information in Table 1 to determine the electronegativity difference, percent ionic character, and type of bond for each formula in Table 2. Data Table 2: Type of Bond Formula Electronegativity Difference Type of Bond Percent Ionic Character KCl 3.0-0.8 = 2.2 ionic 50-100% K2O Br2 MgI2 CaCl2 MgS Al2S3 NaCl F2 HCl 2
Part II: Lewis Structures Valence electrons are the electrons in the last shell or energy level of an atom. Valence electrons participate in the formation of chemical bonds. A Lewis Symbol consists of the element symbol surrounded by "dots" to represent the number of valence electrons. The octet rule states that atoms of low (<20) atomic number tend to combine in such a way that they each have eight electrons in their highest energy level giving them the same electron configuration as a noble gas. Oxygen has six valence electrons and will form two covalent bonds in order to attain the noble gas configuration of neon. For each of the elements in Table 3, use a periodic table to determine the number of valence electrons, write the Lewis Symbol, and determine number of bonds needed for the element to have eight electrons in their highest energy level. Record you observations in Table 3. Data Table 3: Atoms Color of Atom Model Identity of Element Number of Valence Electrons Lewis Symbol Number of Bonds Needed for an Octet red oxygen orange bromine white hydrogen Hydrogen forms 1 bond, not an octet. green chlorine blue nitrogen black carbon 3
Rules for Lewis Structures 1. Use the electronegativity values of the elements and the information in Table 1 to determine the electronegativity difference and the type of bond for each formula. 2. For covalent substances, determine the total number of valence electrons (E) available to the molecule or ion by: a. summing the valence electrons of all the atoms and b. adding one electron for each net negative charge or subtracting one electron for each net positive charge. Then divide the total number of available electrons by 2 to obtain the number of electron pairs available. 3. Calculate the number of electrons forming a multiple bond: M = 6n + 2 E where n is the number of atoms in the molecule excluding the hydrogen atoms and E is the total number of valence electrons in the molecule. 4. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central atom. 5. Determine a provisional electron distribution by arranging the electron pairs in the following manner until all available pairs have been distributed: a. one pair between the central atom and each ligand atom. A covalent bond is represented by either one line or two dots between two atoms. b. three more pairs on each outer atom (except hydrogen, which has no addition electron pairs) yielding an octet around each ligand atom when the bonding pair is included in the count. c. remaining electron pairs (if any) on the central atom. These electrons pairs are unshared by another atom and are called lone pairs of electrons. Example 1: Methane, CH4 1. This compound is covalent. EN = 2.5-2.1 = 0.4 2. The total number of valence electrons, E = 8. One carbon has 4 valence electrons. Each hydrogen atom has one valence electron. 3. M = 6(n) + 2 - E There is 1 atom excluding H, so n = 1. M = 6(1) + 2-8 = 0 (no multiple bonds) 4. Carbon is the central atom. 5. One pair of electrons is placed between the central atom and each hydrogen atom, forming 4 single bonds. The shape of this molecule is tetrahedral. Record the information in Data Table 4. 4
Example 2: Oxygen Gas, O2 1. This compound is covalent. EN = 3.5-3.5 = 0 2. The total number of valence electrons = 12. Each oxygen atom has six valence electrons. 3. M = 6(2) + 2-12 = 2 A value of 2 indicates there are 2 extra electrons between the oxygen atoms, forming a double bond. 4. Oxygen is the central atom. 5. Two pairs of electrons are placed between the oxygen atoms. The shape of this molecule is linear. Record the information in Data Table 4. Example 3: Carbon Dioxide, CO2 1. This compound is covalent. EN = 3.5-2.5 = 1.0 2. The total number of valence electrons = 16. One carbon has 4 valence electrons. Each oxygen atom has six valence electrons. 3. M = 6(3) + 2-16 = 4 A value of 4 indicates there are 4 extra electrons, forming two double bonds. 4. Carbon is the central atom. 5. Two pairs of electrons are placed between the central atom and each oxygen atom. The shape of this molecule is linear. Record the information in Data Table 4. Example 4: Nitrogen Gas, N2 1. This compound is covalent. EN = 3.0-3.0 = 0 2. The total number of valence electrons = 10. Each nitrogen atom has five valence electrons. 3. M = 6(2) + 2-10 = 4 A value of 4 indicates there are 4 extra electrons. Six electrons are shared between the nitrogen atoms forming a triple bond. 4. Nitrogen is the central atom. 5. Three pairs of electrons are placed between the nitrogen atoms. The shape of this molecule is linear. Record the information in Data Table 4. Look at the atom models and bonds in the Covalent Bonding and Lewis Structures Model Set. Identify the different pieces that represent atoms, single bonds, and double/triple bonds. Use Table 3 to find the colors used to represent each element. When building molecular models, 25mm links are used to attach hydrogen atoms to the molecule. For all other single covalent bonds, use the 40mm gray links. Use two 51mm gray links to form a double bond and three to form a triple bond. Use the atom models and bonds to build, methane, oxygen gas, carbon dioxide, and nitrogen gas. Sketch a diagram of each of the molecules in Data Table 4. Continue to construct the Lewis Structures and build the models for each of the molecules as you complete Data Table 4. 5
Data Table 4: Molecules with Single Covalent Bonds Molecule Total Valence Electrons Electrons for Multiple Bonds Central Atom Lewis Structure Sketch of Model CH4 8 M=6(1)+2-8=0 C O2 CO2 N2 HBr NH3 H2O CCl4 N2H2 CH2O CH3OH C2H3N 6
Part III: Analyze and Conclude 1. How many covalent bonds form around each of the following atoms: a. hydrogen!!! d. nitrogen b. carbon!!! e. chlorine c. oxygen!!! f. bromine 2. In a molecular model, a single gray link (single bond) between two elements represents electrons. 3. Identify the type of bond described for each of the following as ionic, polar covalent, or nonpolar covalent.! a. The C O bonds in CO2!!! b. The bonds in F2!! c. The C C bonds in C3H8!!! d. The bonds in K2O 4. How many unshared pairs of electrons are found on the nitrogen atom in NH3? 5. How many electrons are shared in the triple bond in N2? 6. Water and carbon dioxide molecules are both made up of 3 atoms. Referring to Data Table 4, state several differences between water molecules and carbon dioxide molecules. 7. How many double bonds does carbon dioxide contain? 8. Does the carbon atom in carbon dioxide have an octet of electrons? 9. Which of the following elements (C, H, O, or N) will NOT be surrounded by an octet of electrons in a correctly drawn Lewis structure? 10. Which of the diatomic elements (N2, O2, F2, or Br2) has a double bond between its atoms? 7
Teacher s Key Data Table 2: Type of Bond Formula Electronegativity Difference Type of Bond Percent Ionic Character KCl 3.0-0.8 = 2.2 ionic 50-100% K2O 3.5-0.8 = 2.7 ionic 50-100% Br2 2.8-2.8 = 0 nonpolar covalent 0-4% MgI2 2.5-1.2 = 1.3 polar covalent 5-49% CaCl2 3.0-1.0 = 2.0 ionic 50-100% MgS 2.5-1.2 = 1.3 polar covalent 5-49% Al2S3 2.5-1.5 = 1.0 polar covalent 5-49% NaCl 3.0-0.9 = 2.1 ionic 50-100% F2 4.0-4.0 = 0 nonpolar covalent 0-4% HCl 3.0-2.1 = 0.9 polar covalent 5-49% Data Table 3: Atoms Color of Atom Model Identity of Element Number of Valence Electrons Lewis Symbol Number of Bonds Needed for an Octet red oxygen 6 2 orange bromine 7 1 white hydrogen 1 Hydrogen forms 1 bond, not an octet. green chlorine 7 1 blue nitrogen 5 3 black carbon 4 4 8
Data Table 4: Molecules with Single Covalent Bonds Molecule Valence Electrons π Bonding Electrons Central Atom Lewis Structure Model CH4 8 M=6(1)+2-8=0 C O2 12 M=6(2)+2-12=2 O CO2 16 M=6(3)+2-16=4 C N2 10 M=6(2)+2-10=4 N HBr 8 M=6(1)+2-8=0 Br NH3 8 M=6(1)+2-8=0 N H2O 8 M=6(1)+2-8=0 O CCl4 32 M=6(5)+2-32=0 C N2H2 12 M=6(2)+2-12=2 N CH2O 12 M=6(2)+2-12=2 C CH3OH 14 M=6(2)+2-14=0 C C2H3N 16 M=6(3)+2-16=4 C 9
Part III: Analyze and Conclude 1. How many covalent bonds form around each of the following atoms: a. hydrogen! 1!!! d. nitrogen! 3 b. carbon! 4!!! e. chlorine! 1 c. oxygen! 2!!! f. bromine! 1 2. In a molecular model, a single gray link (single bond) between two elements represents 2 electrons. 3. Identify the type of bond described for each of the following as ionic, polar covalent, or nonpolar covalent.! polar covalent a. The C O bonds in CO2 nonpolar covalent c. The C C bonds in C3H8 nonpolar covalent b. The bonds in F2 ionic d. The bonds in K2O 4. How many unshared pairs of electrons are found on the nitrogen atom in NH3? 1 5. How many electrons are shared in the triple bond in N2? 6 6. Water and carbon dioxide molecules are both made up of 3 atoms. Referring to Data Table 4, state several differences between water molecules and carbon dioxide molecules. Water is composed of 1 oxygen atom and 2 hydrogen atoms. Water has only single bonds. Water has two pairs of unshared electrons. Water is a bent molecule. Carbon dioxide is composed of one carbon atom and 2 oxygen atoms. Carbon dioxide has 2 double bonds. The carbon atom does not have any unshared electrons. Carbon dioxide is a linear molecule. 7. How many double bonds does carbon dioxide contain? 2 8. Does the carbon atom in carbon dioxide have an octet of electrons? yes 9. Which of the following elements (C, H, O, or N) will NOT be surrounded by an octet of electrons in a correctly drawn Lewis structure? H, hydrogen 10. Which of the diatomic elements (N2, O2, F2, or Br2) has a double bond between its atoms? O2, oxygen 10