AP Chemistry - Chapter 8 Basic Concepts of Bonding At the end of this unit I will be able to . . . Target # 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. Target Define ionic bond, covalent bond and metallic bond. Draw the Lewis symbols for elements. Define lattice energy and its relationship to ionic bonds. I will also be able to describe the factors which affect lattice energy and determine the lattice energy using the Born-Haber cycle. Describe a covalent compound using Lewis symbols. Describe a single, double, and a triple covalent bond in terms of number of shared electrons, length and strength. Define electronegativity and be able to list its periodic trends. Predict the relative polarities of bonds using either the periodic table or electronegativity values. I can also desrcibe the factors which affect bond polarity and dipole moment. Write the Lewis structures for molecules and ions containing covalent bonds using the periodic table. Explain the concept of resonance and draw resonance structures for molecules or polyatomic ions. Use the concept of formal charge to predict the most stable resonance structure. Describe the 3 common exceptions to the octet rule and provide examples of each. Relate bond enthalpies to bond strengths and and use bond enthalpies to estimate ∆H for reactions. Book Section 8.1 8.1 8.2 8.3 8.3 8.4 8.4 8.5 8.6 8.5 8.7 8.8 Target 1: I can define ionic bond, covalent bond and metallic bond. Ionic Bond: Ions transfer electrons from one atom to another. Formed by Electrostatic forces that exist between ions of opposite charge. These compounds have a rigid 3-D structure. Covalent Bond: Results from sharing electrons between two nonmetallic atoms. They are found in molecular compounds and polyatomic ions. Metallic Bond: They are found in metals. Each metal atom is bonded by several neighboring atoms. Bonding electrons move through 3-D structure allowing for electrical and heat conductivity. 1. What type of bond would each of the following substances have? CO2 Cu BaCl2 NH3 AgBr Zn Target 2: I can draw the Lewis symbols for elements. Valence(Latin word means to be strong) electrons are the electrons that take part in a chemical reaction. e- dot structure: a simple and convenient way of illustrating valence electrons in an atom. Let’s do the dot structures for the Period 2 elements Li 2. Be B C N O F Ne What rule is being followed in the dot structures above? 3. What is Lewis Symbol (dot structure) for each of the following atoms or ions: Br Mg Ca2+ F- Octet Rule: atoms tend to gain, lose or share electrons until they are surrounded by 8 valence electrons. III. IONIC BONDING (8.2): + Brittle, high melting points, usually crystalline + Generally, forms from the interaction of metals from the far left(1A) on the Periodic Table with nonmentals on the far right(7A) + Ionic compounds have high melting points due to the attraction between ions of unlike charge. This attraction brings ions together, releasing energy, and the ions form a LATTICE. 3. Using Lewis Symbols, diagram the reaction that occurs between magnesium and chlorine. 4. Do these compounds exist? Rb2O LiO Na3N BeF3 Target 3: I can define lattice energy and its relationship to ionic bonds. I will also be able to describe the factors which affect lattice energy. Lattice Energy = energy needed to separate 1 mole of an ionic solid compound into gaseous ions; lattice energy is a measure of stability. (Table 8.2) + Therefore, forming ionic compounds is exothermic. + The energy released by the attraction between ions of unlike charges more than makes up for the ionization energy required to form ions and ionic compounds. Equation representing the lattice energy of NaCl(s) . . . Lattice energy depends upon . . . abE = k Q1Q2 d E: Lattice Energy k: is a constant: 8.99 x 109 J . m/C2 Q1 and Q2 : Charges on the particles d: distance between their centers 5. Arrange the following in increasing Lattice Energy. KBr, MgO, KF, AlN, MgS. Justify your answers. Importance of Size of Ions O2-, F-, Na+, Mg2+, and Al3+ comprise an isoelectronic series since each has the same total number of electrons, 10. Yet the sizes of the ions varies. As the radius decreases, d decreases, lattice energy increases. In later chapters ionic size will also impact how ions pack in a solid and determines some properties of ions in solution. O2- F- Na+ Mg2+ Al3+ The Born-Haber Cycle: The Indirect Calculation of Lattice Energies example to try . . . 8.29 on page 324: Use the data from Appendix C, Figure 7.11, 7.12, and Table 7.4 to calculate the lattice energy of RbCl. Is this value greater than or less than the lattice energy of NaCl? Explain. Homework: Calculate the lattice energy of calcium fluoride. Write the equations representing each step in the Born-Haber process. Use Hess’s Law and values from your text to help you determine the lattice enrgy. (Note: The second ionization energy of calcium is 1145 kJ/mol). Target 4: I can describe a covalent compound using Lewis symbols. 1. COVALENT BONDING (8.3) These compounds can to be gas, liquid or solid at room temperature Many are flexible or pliable like plastic and rubber Acquire noble gas configuration by sharing electrons which is called a covalent bond “The shared pair of electrons in a covalent bond acts as a “glue” to bind atoms together.” May share one pair of electrons(single bond); two pairs of electrons (double bond); or three pairs of electrons(triple bond) “The distance between bonded atoms decreases as the the number of shared electron pairs increases.” 2. BOND POLARITY AND THE ELECTRONEGATIVITY TREND(8.4) Bond polarity is the sharing of electrons between atoms A Nonpolar Covalent Bond is where e-‘s are shared equally In a Polar Covalent Bond one atom has a greater attraction for the bonding electrons that the other atom. We say that this atom has a greater electronegativity(the ability of an atom in a molecule to attract electrons). If the Electronegativity Difference(E.D.) is between 0 – 0.4, then the bond is Nonpolar. If the E.D. is 0.5 – 2.0 then the bond is Polar and if the E.D. is greater than 2.0, then the bond is Ionic. Most bonds show some charge separation and some electron sharing. Examples: Electronegativity Difference Type of Bond O2 H2O NaCl Target 5: I can describe a single, double, and a triple covalent bond in terms of number of shared electrons, length and strength. Single, Double and Triple Covalent Bonds Summary # of shared electrons Length Strength example Single covalent Bond Double Covalent Bond Triple Covalent Bond Target 6: I can define electronegativity and be able to list its periodic trends. Electronegativity = Electronegativity periodic trends in a family: Electronegativity periodic trends in a period: Target 7: I can predict the relative polarities of bonds using either the periodic table or electronegativity values. I can also desrcibe the factors which affect bond polarity and dipole moment. Nonpolar Bond = Polar Bond = The Relationship Between Electronegativity and Bond Type Electronegativity difference between bonded atoms Bond Type Covalent Character Ionic Character Polar molecules are called dipoles! The greater the difference in electronegativity between two atoms, the . . . 1. 2. Dipole Moment = a quantitative measure of the magnitude of a dipole. “Whenever two electrical charges of equal magnitude but opposite sign are separated by a distance, a dipole is established. The quantitative measure of the magnitude of a dipole is called its dipole moment.” = Qr where Q represents two equal, but opposite charges and r is the distance between those 2 charges The unit for dipole moment is D, debye, which is equal to 3.34 x 10 -30 Cm. A dipole moment will increase when . . . ABTarget 8: I can write the Lewis structures for molecules and ions containing covalent bonds using the periodic table. Procedure for Drawing Lewis Structures A. Find the total number of valence electrons B. Use symbols for the atoms and a dash represents two electrons. C. The element that is fewest in number is the central atom. D. The central atom is often less electronegative that the surrounding atoms. E. C, N, O, P and S make great central atoms. F. H is never a central atom and can only form single bonds. G. Complete the octets on all atoms bonded to the central atom. H. Place any extra electrons on the central atom. I. Not enough elecrons around the central atom? Try a double or triple bond. J. A helpful formula to determine number of covalent bonds is b = t - v Where b = # of bond pairs 2 t = total # of e- when attaining noble gas configuration (t is 8 for most elements except H is 2 and Boron is 6) v is the number of valence electrons You try these! PCl3 CH2Cl2 HCN N2 NO+ BrO3- PO43- N2H4 Target 9: I can explain the concept of resonance and draw resonance structures for molecules or polyatomic ions. Resonance Structures = is one of two or more Lewis structures for a single molecule that cannot be represented accurately by one Lewis structure. A. Resonance means the use of two or more Lewis structures to represent a particular molecule. B. A double headed arrow represents resonance structures. C. Resonance is a human invention used to deal with limitations in bonding models. D. The same atoms must be bonded with each other and only electrons may be rearranged. E. Each resonance structure is equivalent, only a different electron placement. F. Resonance in benzene results in “stability” and a higher mpt and bpt. PRACTICE: Draw the resonance structures for the following ions. Ozone: Nitrate: Benzene Chapter 8 Review Targets #1-9 Name __________________________ 1. The Lewis symbol for nitrogen contains ______ pair(s) of electrons and ______ unpaired electrons. 2. For a given salt, lattice energy increases as ionic radius __________________ and as ionic charge __________________. 3. Match the salt on the left with the lattice energy on the right: LiCl CsCl MgCl2 2,326 kJ/mol 834 kJ/mol 657 kJ/mol 4. What are the periodic trends for electronegativity? 5. The the following bonds in order of increasing polarity: O-O, O-C, O-N 6. List 2 covalent bonds which are more polar than the F-O bond. 7. Consider the following bonds: C-O, C=O, CO a) Which is the longest bond? b) Which is the strongest bond? c) Which has the most shared electrons? 8. Draw the dot structures for . . . a) AsF3 b) C2H2 c) SF4 9. Which of the following exhibit resonance? a) NH4+ b) CO32- c) SO32- Target 10: I can use the concept of formal charge to predict the most stable resonance structure. Analysis of Resonance Structures . . . Formal Charge If more than one structure is possible (resonance structures) you can make a reasonable guess as to the most likely structure busing the formal charges of the atoms. Rules for determining formal charges (FC): * The formal charge on an atom is determined by taking the valence electrons minus the electrons assigned to that atom. * For neutral atoms, the formal charges add up to zero and for ions the formal charges add up to equal the charge on the ion. * Lewis structures with large formal charges are unlikely to exist. * A lewis structure is valid if a negative formal charge is on the most electronegative element. * The most probable or ‘stable’ resonance structure will have FC’s around zero. Determine the formal charge on the SCN- resonance structures! Determine the formal charge on the CO2 resonance structures! Target 11: I can describe the 3 common exceptions to the octet rule and provide examples of each. Exceptions to the Octet Rule: A. 1. Molecules with odd numbers of electrons such as: ClO2, NO and NO2. Note: Molecules with odd # of electrons end to form “dimers” 2. Molecules in which an atom has less than an octet (B, Be, Al) 3. Molecules in which an atom has more than an octet. Expanded valence shell octets only occur for elements in period 3 and beyond (never in period 2!). This is because, starting with period 3, d-orbitals can be used in bonding. In addition, the larger the central atom, the larger the number of atoms can surround it. It should also make sense to you that smaller, highly electronegative elements (F, Cl, O) are the most likely candidates to be the elements that are attached to the central atom in an expanded octet. Target 12: I can relate bond enthalpies to bond strengths and and use bond enthalpies to estimate ∆H for reactions. A. B. Bond enthalpy (bond dissociation energy) is the energy needed to break a bond. Bond enthalpy (bond dissociation energy) energy is always +; so energy is always required to break chemical bonds(endothermic). C. The stability of a molecule is related to the strength of the covalent bond. The greater the bond enthalpy … the stronger the covalent bond. D. 90% of the earth’s crust is composed of SiO2(silicates) which have very strong covalent network bonds. E. The enthalpy of a reaction, ∆Hrxn = (sum of the bond enthalpies of bonds broken) – (sum of the bond enthalpies of bonds formed). Let’s work 8.70 b on P.326 Example #2: Use the table below of bond energies in order to calculate the enthalpy of combustion of dimethyl ether, CH3OCH3. (Note: the structure of dimethyl ether is given!) Bond Bond energy (kJ/mol) H-H 436 H-C 414 H-O C-C C=C C-O C=O O-O O=O 464 347 611 360 736 142 498 Example #3: Consider the following reaction: 2 H2O2(g) 2 H2O(g) + O2(g) ; ∆H = -203 kJ If it takes 495 kJ to break the O=O bond, find the bond enthalpy for O-O.
