Unit 03B: Covalent Bonds and Molecule Naming

Transcription

1 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 1 NAME: Date: Period: Unit 03B: Covalent Bonds and Molecule Naming Driving Questions Why do covalent molecules form? Why do some covalent compounds react differently than others? How do scientists name covalent compounds? Connections to Past/Future Units Understand how the difference between ionic and molecular compounds Identify valence electrons and draw Lewis Dot structures Use the naming rules with chemical compounds and chemical reactions Objectives: o SWBAT: Differentiate between covalent and ionic compounds, and name each type. o SWBAT: Diagram covalent bonding using Lewis Dot Structures o SWBAT: Apply VSEPR theory to identify molecular shapes. o SWBAT: Explain how differences in electronegativity lead to polarity in molecular bonds. o SWBAT: Identify polarity in a molecule based on molecular geometry. 3 Essential Vocabulary Covalent bond Single bond Double bond Triple bond Sigma bond Pi bond Linear Trigonal Planar tetrahedral trigonal bipyramidal octahedral electronegativity polarity REFLECTION: What is one thing that you learned of importance this unit? What further questions do you have about this unit? What can you improve upon about your class or study work ethic? What did you do well this unit?

2 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 2 SWBAT: Differentiate between covalent and ionic compounds, and name each type. IONIC BONDS COVALENT BONDS Types of Elements Electron Behavior Cause of Bonding Bond Strength Compound/ Molecule Structure Consider a molecule of hydrogen gas, which has a formula of H 2. Hydrogen atoms have one valence electron. By gaining an electron, an atom of hydrogen can achieve the configuration of helium, a very stable noble gas. Both hydrogen atoms in a hydrogen gas molecule can achieve this configuration simultaneously if they share, rather than transfer their electrons. Hydrogen only needs one electron to emulate a noble gas, and so the two atoms will share one pair of valence electrons, forming a single covalent bond. Likewise, two atoms of fluorine each need to gain one electron to become similar to the nearest noble gas, neon. They will each share one electron in the formation of fluorine gas, F 2, making another single covalent bond. In the formation of carbon dioxide, CO 2, carbon has four valence electrons, and oxygen has six. Carbon needs four more electrons, and each oxygen atom needs two, if they are to achieve stable, noble gas configurations. Therefore, the formation of single covalent bonds will not be enough. However, if each oxygen shares two electrons with carbon, the oxygens will gain the two electrons they need to get stable configurations. The carbon atom will also receive the four electrons it needs to become like neon. These bonds, formed through the sharing of two pairs of valence electrons, are called double covalent bonds.

3 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 3 For molecules, we need a new system of naming since a ratio of elements is not important like in ionic bonding, but rather the exact number of atoms in a molecule. The NAME of a molecular compound uses prefixes to tell how many atoms of each element are present in a molecule. The prefix mono- indicates the presence of a single atom of a given type. The prefix di- indicates the presence of two such atoms. A chart showing the most commonly used prefixes can be found at the right. The steps for writing out the name of a molecular compound are as follows: EXAMPLES Compound NI 3 P 4 O 10 I 2 O Name The FORMULA of a compound is easy to write, given the name of the compound. Use the prefixes to tell you the subscript of each element in the formula. Then write the correct symbols for the two elements, with the appropriate subscripts. Use this method to verify the following molecular compound formulas. Name Tetraiodine nonoxide Sulfur trioxide Phosphorus pentafluoride Compound

4 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 4 It is IONIC Yes Does it begin with a metal or NH 4? No It is COVALENT Does it begin with a metal from the first two columns of the periodic table, or Al, Cd, Zn, Ag, or NH 4? Write the name of the 1 st prefix. If the prefix is mono-, do not write it down. Yes No Write the name of the 1 st element. Write the name of the cation. Write the name of the cation, followed by its Roman numeral in parentheses. Write the name of the 2 nd prefix. Write the name of the 2 nd element. Write the name of the anion. Change the ending to ide. Write the names of the following molecular compounds. 1. P 4 O BBr 3 3. N 2 O 5 4. P 2 O 5 5. SCl 6 6. SCl 2 7. CCl 4 8. As 2 S 5 9. PCl SeCl SI OF 2

5 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 5 Write the formulas of the following molecular compounds. 1. silicon dioxide 2. dinitrogen pentasulfide 3. dinitrogen heptoxide 4. sulfur trioxide 5. triarsenic monoxide 6. iodine monochloride 7. tellurium difluoride 8. selenium dibromide On the left, identify the following compounds as Ionic or Molecular. On the right, name the compounds: a. FeCrO 4 b. (NH 4 ) 2 SO 4 c. N 5 O 7 d. C 4 Cl 2 e. Mn(C 2 H 3 O 2 ) 3 f. Li 2 CO 3 g. Na 2 SO 3 h. P 2 O 5 i. C 2 Cl 6 j. Mn(C 2 O 4 ) 2 k. CsHCO 3 l. O 2 F 3 m. Au(OH) 3

6 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 6 SWBAT: Diagram covalent bonding using Lewis Dot Structures REVIEW: What was the octet rule? Why is it important? Consider your answer to that question as you complete the following activity: In the boxes on the left, draw the Lewis Dot Structure of the elements specified (remember, valence electrons only!). Then, using what you know about valence electrons and covalent bonds, draw and name the compound formed when the two elements are combined. Remember these two helpful hints: (1) two electrons are needed to make a single covalent bond (2) the atom that needs the most electrons goes in the middle, when applicable. Element #1 Element #2 Compound fluorine carbon carbon tetrafluoride fluorine sulfur Name: nitrogen bromine Name:

7 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 7 Multiple Bonding: A covalent bond is caused by the sharing of a pair of valence electrons. If two pairs of valence electrons are shared, a double covalent bond is formed. If three pairs of valence electrons are shared, a triple bond is formed. The strength of a bond increases with its multiplicity, or bond order that is, a triple bond is stronger than a double bond, which is stronger than a single bond. Bond Bond Order Enthalpy (Strength) Length C C kj/mol 1.54 Å C=C kj/mol 1.34 Å C C kj/mol 1.20 Å Finding the Center Atom: If there are more than two atoms in a molecule, the molecule usually ends up being one center atom bonded to other atoms that are projecting out of the center. The center atom is usually the atom that needs the most valence electrons (and therefore the most bonds). Lines Instead of Dots: It is a common and time- saving practice to draw shared electron pairs as a line, rather than two dots. A double bond appears as two lines, and a triple bond appears as three lines. This is acceptable, but you must always draw the unshared electrons as dots, as in the Lewis Dot structure of water (H 2 O) at the right. Only the shared electrons can be drawn as lines. Keep in mind that a line is equal to two electrons. Exceptions to the Octet Rule: Some elements are so small that they cannot support a full compliment of eight electrons. Beryllium atoms only hold 4 electrons due to their small size. Boron tends to end up with only 6 electrons. Some elements are so big that they can support more than eight valence electrons. Phosphorus can support either 8 or 10 electrons, depending on the situation. Likewise, Sulfur can hold either 8 or 12 valence electrons, and Xenon can have 12. How to get started:

8 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 8 Draw Lewis Dot Structures for the following molecules, using what you have learned. 1. OCl 2 2. NCl 3 3. CS 2 4. SCl 2 5. SCl 6 6. O 2 7. H 2 O 8. PF 3 9. PF 5

9 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 9 Draw Lewis Dot Structures for the following molecules, using what you have learned. 1. CCl 4 2. BF 3 3. NF 3 4. PCl 3 5. Br 2 6. XeCl 4 7. OH 1 8. O 3 9. SO 3

10 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 10 SWBAT: Apply VSEPR theory to identify molecular shapes. A photograph or sketch may fail to do justice to your appearance. Similarly, electron dot structures fail to reflect the three- dimensional shapes of molecules. The electron dot structure and structural formula of methane (CH 4 ), for example, show the molecule as if it were flat and merely two dimensional. In reality, methane molecules are three dimensional. As seen in the diagrams below, the hydrogens in a methane molecule are at the four corners of a geometric solid called a regular tetrahedron. In this arrangement, all of the H C H angles are 109.5, and not 90. electron dot structure structural formula actual three- dimensional shape The valence- shell electron- pair repulsion theory, or VSEPR theory, explains the three dimensional shape of molecules. Because electron pairs repel each other, molecular shape adjusts so the valence- electron pairs are as far apart as possible. To determine the shape of a molecule, you must first count the number of ways electron pairs are attached to the central atom: 2 ways: If pairs of electrons are leaving the center atom in two ways, the only possible structure is linear. The electron pairs position themselves on either side of the central atom, causing an angle of 180. Examples include BeH 2, CO 2, and HCN. The molecule is said to be sp- hybridized 3 ways: If electron pairs are leaving the center in three ways, the most likely structure is trigonal planar. The bond angle formed is 120. Examples include BCl 3 and AlCl 3. The molecule is said to be sp 2 - hybridized. 4 ways: When electron pairs are leaving the center in four different ways, it gets a little bit more complex. If all the electron pairs leaving the center are bonds, then a tetrahedral shape results, as shown in the earlier example of methane. If one of the bonds is instead substituted for a lone pair of electrons, as in ammonia, NH 3, the lone pair still repels the bonds into a three- dimensional shape called a trigonal pyramid. However, the lone pair repels the other electrons a little bit more than the bonds do, causing the bond to narrow to 107. If electrons are leaving the center in four ways, but two of the ways are lone pairs, as in water molecules, the bond angle tightens further to and the structure is called bent. All of these molecules are said to be sp 3 - hybridized. tetrahedral trigonal pyramid 107 bent 104.5

11 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 11 5 ways: If electrons are leaving the center in five ways, the molecule is called trigonal bipyramidal, and is said to be sp 3 d- hybridized. The bond angles are 90 and ways: If electron pairs are leaving the center in six ways, the molecule is called octahedral, and is said to be sp 3 d 2 - hybridized. The bond angles are 90 and 180. trigonal bipyramid octahedral In each box below, draw the Lewis Dot Structure for the molecule. Then, identify the VSEPR shape, the Bond Angle, and the hybridization. 1. CF 4 2. H 2 S 3. SF 6 Shape: Bond Angle: Hybridization: Shape: Bond Angle: Hybridization: Shape: Bond Angle: Hybridization: 4. PF 3 5. PF 5 6. NF 3 Shape: Bond Angle: Hybridization: Shape: Bond Angle: Hybridization: Shape: Bond Angle: Hybridization:

12 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page BeCl 2 8. OCl 2 9. BI 3 Shape: Bond Angle: Hybridization: Shape: Bond Angle: Hybridization: Shape: Bond Angle: Hybridization: Use VSEPR theory to predict the shapes of the following species. Draw the electron dot structures at the right. a. CO 2 b. SiCl 4 c. SCl 2

13 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 13 Laboratory Activity Marshmallow Molecule Building Students will be required to use marshmallows (atoms) and toothpicks (bonds) to create molecule models. The table provided below should be expanded to include all of the molecules used in this activity. Students will complete a table for eleven of the following: CCl 4 NF 3 CO 2 H 2 S GeCl 4 1+ PH 4 SO 3 PCl 5 SCl 6 SCl 2 AsCl 3 BBr 3 SiO 2 SeCl 6 BF 3 Molecular Formula Lewis Dot Structure 3- Dimensional Drawing VSEPR Shape σ Bonds π Bonds Hybridization

14 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 14 Molecular Formula Lewis Dot Structure 3- Dimensional Drawing VSEPR Shape σ Bonds π Bonds Hybridization

15 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 15 SWBAT: Explain how differences in electronegativity lead to polarity in molecular bonds.

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19 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 19 The Bare Essentials of Polarity After reading the cartoon, answer the following questions. 1. How does the comic book define a polar molecule? 2. Define electronegativity as you understand it, after reading the first two pages of the comic book. 3. Interpret the picture at the bottom of page 14. Explain how the iceberg, penguins, and polar bears represent trends in electronegativity. 4. What is the artist trying to represent when there are two polar bears arm wrestling together, or two penguins arm wrestling together? 5. What three types of bonds are represented on page 16 of the comic book? What happens to the bonding electrons in each type of bond? 6. Explain why there are four scoops of ice cream in the illustration of O 2 on page What do the six scoops of ice cream represent in the illustration of N 2 on page 17? 8. Describe what you think is happening to the penguin in the CO 2 molecule in the picture on page Name three things that the picture of CO 2 on page 17 illustrates about the molecule. 10. Describe what you think is happening to the penguins in the illustration of H 2 O on page Explain what you think the crossed arrow represents in the comic book.

20 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 20 READING ASSIGNMENT: Read and take notes on pages in your textbook. SWBAT: Identify polarity in a molecule based on molecular geometry. The most important statistic when determining the character of a covalent bond is electronegativity. As you might recall, electronegativity is the degree to which an atom holds on to a shared pair of electrons. The electronegativity of each atom in a bond can be found in an electronegativity table, like the one below. The difference in electronegativities of the atoms in the bond can be measured with simple subtraction. H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 K 0.8 Ca 1.0 Sc 1.3 Ti 1.5 V 1.6 Cr 1.6 Mn 1.5 Fe 1.8 Co 1.9 Ni 1.9 Cu 1.9 Zn 1.6 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8 Rb 0.8 Sr 1.0 Y 1.2 Zr 1.4 Nb 1.6 Mo 1.8 Tc 1.9 Ru 2.2 Rh 2.2 Pd 2.2 Ag 1.9 Cd 1.7 In 1.7 Sn 1.8 Sb 1.9 Te 2.1 I 2.5 Cs 0.7 Ba 0.9 La 1.0 Hf 1.3 Ta 1.5 W 1.7 Re 1.9 Os 2.2 Ir 2.2 Pt 2.2 Au 2.4 Hg 1.9 Tl 1.8 Pb 1.9 Bi 1.9 Po 2.0 At 2.2 nonpolar: = 0.0 polar: = 0.9 ionic: = 2.1 Nonpolar bonds occur when the electrons are evenly shared between the two atoms in the bond. The difference of electronegativities is 0.0 or 0.1. Polar bonds occur when the electron pair making up the bond is shared, but not evenly. The difference of electronegativities is between 0.1 and 2.0. Ionic bonds occur when one atom completely takes an electron away from the other atom. The difference in electronegativities is greater than 2.0 The bonds between the following pairs of elements are covalent. Arrange them according to polarity, naming the most polar bond first. a. H Cl b. H C c. H F d. H O e. H H f. S Cl

21 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 21 The presence of a polar bond in a molecule often makes the entire molecule polar. In a polar molecule, one end of the molecule is slightly negative and the other end is slightly positive. This kind of molecule is also referred to as a dipole. The easiest way to determine whether a molecule is polar is to imagine that the atom in the center is a wagon, and that the attached atoms are horses, pulling in opposite directions. Consider the electronegativity of each atom to be relative to the size and strength of each horse. (Yes, this is silly, but it works every time.) If the wagon ends up moving, the molecule is polar. If the horses cancel each other out, the molecule is nonpolar. Consider the molecule BeHCl. Beryllium needs the most electrons, and will go in the center. The Lewis Dot structure of BeHCl would therefore look like this: Now, consulting the electronegativity table, we see that hydrogen only has an electronegativity of 2.1, while chlorine has the much higher electronegativity of 3.0. Therefore, using the horse and wagon analogy, the molecule can be reimagined as a horse and wagon system that looks like this: Obviously, the wagon will be moved to the right, and so we can say that the molecule is polar. Of course, just because a molecule contains polar bonds does not mean that it is polar. Consider a molecule of carbon dioxide, which has a Lewis Dot structure like this: Using the horse and wagon analogy, the molecule would be reimagined to look like this: Since the atoms attached to the central atom are the same, the horses are the same size, and they would cancel each other out. The wagon would not move, and the molecule is therefore nonpolar. Not every molecule with polar bonds is polar. Explain this statement, using CCl 4 as an example.

22 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 22 In each box, draw the Lewis Dot Structure for the molecule. Then, write out the name of the VSEPR shape, the bond angle, and the hybridization of each molecule. Finally, state whether or not the molecule is polar. If there exists a molecular polarity, draw a crossed arrow in the direction of the partial negative charge. 1. NF 3 2. BCl 3 VSEPR Shape: Polar? Bond Angle: Hybridization: VSEPR Shape: Polar? Bond Angle: Hybridization: 3. OCl 2 4. SiO 2 VSEPR Shape: Polar? Bond Angle: Hybridization: 5. H 2 S VSEPR Shape: Polar? Bond Angle: Hybridization: 6. SCl 2 VSEPR Shape: Polar? Bond Angle: Hybridization: VSEPR Shape: Polar? Bond Angle: Hybridization:

23 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 23 Unit 07 Bonding Take- Away Sheet NAME: 1. DRAW the Lewis Dot Diagrams for the following. IDENTIFY the shape of the molecule just under each drawing. a) SCl 2 b) OF 2 c) AsCl 3 d) CF 4 e) SiO 2 f) PI 5 2. Which of the above compounds are polar molecules? 3. HCN, has a few bonds. Draw the structure of HCN and identify the number of Sigma and Pi bonds it has. Sigma: Pi: 4. What will form stronger bonds, NH 3 or NF 3? 5. What is stronger, Ionic Bond or Hydrogen Bond?

24 Sophomore Chemistry - Deak Unit 03B Covalent Bonding Page 24 Fill in the blanks. 6. Compound: Fe 2 (CO 3 ) 3 7. Compound: N 2 O 7 Bond Type (Ionic or Covalent): Bond Type (Ionic or Covalent): Structure (Molecule or Lattice): Structure (Lattice or Molecule): Name of compound: Name of compound: 8. Compound: Manganese (II) Sulfate 9. Compound: Oxygen Monochloride Bond Type (Covalent or Ionic): Bond Type (Ionic or Covalent): Structure (Molecule or Lattice): Structure (Molecule or Lattice): Formula of compound: Formula of compound: 10. Compound: Al 2 S Compound: Tricarbon Heptachloride Bond Type (Ionic or Covalent): Bond Type (Ionic or Covalent): Electrons (Sharing or Transfer): Electrons (Sharing or Transfer): Name of compound: Formula of compound: