8.1 Types of Chemical Bonds List and define three types of bonding. chapter 8 Bonding General Concepts.notebook. September 10, 2015

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1 chapter 8 Bonding General Concepts.notebook Chapter 8: Bonding: General Concepts Mar 13 11:15 AM 8.1 Types of Chemical Bonds List and define three types of bonding. Bonds are forces that hold groups of atoms together and make them functions as a unit. Bonds form because the energy of the system is lower than if bonds did not form. Bond Length the distance where the energy is minimal. Bond Energy the energy required to break the bond. A relationship exists between the number of shred electron pairs and the bond length. As the number of shared electrons increases, the bond length shortens. Mar 3 4:53 PM 1

Mar 3 4:53 PM Covalent Bonding occurs when bonds form between similar kinds of atoms for the same reason as between dissimilar atoms: the energy of the system is lowered as a result of the bond

2 Ionic Bonding is due to electrostatic attraction. It results from the loss of an electron from a metal and its gain by a nonmetal. Ionic Compound are formed when an atom that loses electrons relatively easily reacts with an atom that has high affinity for electrons. Ionic compounds result when a metal reacts with a non metal. Mar 3 4:53 PM Covalent Bonding occurs when bonds form between similar kinds of atoms for the same reason as between dissimilar atoms: the energy of the system is lowered as a result of the bond formation. At a particular distance apart, the combination of repulsive and attractive forces allows the system to have a minimum energy. In this case, we have a covalent bond in which electrons are shared by both nuclei approximately equally. Examples are: S 8, graphite, and diamond. The change in potential energy as a function of H H bond distance Mar 3 5:34 PM 2

Double and Triple Bonds In a double bond, 2 atoms share 2 pairs of electrons (4 electrons). In a triple bond, 2 atoms share 3 pairs of electrons (6 electrons).

Charges are not distributed equally in such a molecule. Positive and negative poles exist. Examples of polar covalent bonds are C Cl, H Cl, and O H.

2 Electronegativity Use the position of atoms in the periodic table and their electronegative to predict relative bond polarities Electronegative is the ability of an atom in a

3 Double and Triple Bonds In a double bond, 2 atoms share 2 pairs of electrons (4 electrons). In a triple bond, 2 atoms share 3 pairs of electrons (6 electrons). Double and triple bonds are stronger than single bonds Mar 3 5:16 PM In the case where there is unequal sharing of electrons, the polar covalent bond exists. Charges are not distributed equally in such a molecule. Positive and negative poles exist. Examples of polar covalent bonds are C Cl, H Cl, and O H. In summary, the nature of the bond will depend upon the ability of each atom in the bond to attract electrons to itself. This is call electronegative. 8.2 Electronegativity Use the position of atoms in the periodic table and their electronegative to predict relative bond polarities Electronegative is the ability of an atom in a molecule to attract shared electrons to itself. The effect of an electric field on hydrogen fluoride molecule. When the field is turned on, the molecules tend to line up with their negative ends toward the positive pole and their positive ends towards the negative pole. Mar 3 5:51 PM 3

For the representative elements, electronegative decreases going down a group and increase going across a period.

If ΔEN= 0, the bond is said to be perfectly covalent. But there are no precise cut offs. All bond have some ionic and some covalent characters. Example 8.

Li H c. H C d. H F e. Rb O Mar 3 5:58 PM An electrostatic potential map of HF.

4 For the representative elements, electronegative decreases going down a group and increase going across a period. Thus francium has the lowest electronegative, and fluorine has the highest. The greater the Δ, the more ionic character the bond has. If ΔEN= 0, the bond is said to be perfectly covalent. But there are no precise cut offs. All bond have some ionic and some covalent characters. Example 8.2 B Bond Polarity Using Figure 8.3 in your textbook, calculate Δ for each of the following bonds, and order the set from most covalent to most ionic character. a. Na Cl b. Li H c. H C d. H F e. Rb O Mar 3 5:58 PM An electrostatic potential map of HF. Red indicates the most electron rich area (the fluorine atom) and the blue indicates the most electron poor region the hydrogen atom) The tree possible type of bonds: (a) a covalent bond formed between identical F atoms (b) the polar bond of HF, with both ionic and covalent components (c) an ionic bond with no electron sharing Mar 14 10:48 AM 4

chapter 8 Bonding General Concepts.notebook 8.3 Bond Polarity of Dipole Moments You will be able to determine whether or not simple molecules have a dipole moment.

The arrow points to the center of negative charge while the tail is at the center of positive charge. It is possible to determine the polarity of a bond by the size of ΔEN.

For instance, we determined that H F was polar, with fluorine being the electronegative atom.

5 chapter 8 Bonding General Concepts.notebook 8.3 Bond Polarity of Dipole Moments You will be able to determine whether or not simple molecules have a dipole moment. A molecule such as HF that has a pole of positive charge and a pole of negative charge is said to be dipolar, or to have a dipole moment. The arrow points to the center of negative charge while the tail is at the center of positive charge. It is possible to determine the polarity of a bond by the size of ΔEN. If a molecule is diatomic (two atoms), there is often only one bond, and that will determine whether the molecule is polar. For instance, we determined that H F was polar, with fluorine being the electronegative atom. A partial negative charge (δ ) resides on the fluorine atom, and a partial positive charge (δ+) resides on the hydrogen atom Mar 3 5:52 PM The situation is clear cut with HF. It becomes more difficult with 3 or more atoms in a molecule because the individual dipoles can cancel each other out. (a) The charge distribution in the water molecule. (b) The water molecule in an electric field. (c) The electrostatic potential diagram of a water molecule. This table shows how individual bond polarities can cancel each other out to yield a molecule with no dipole moment. Mar 3 5:16 PM 5

Example 8.3 A Dipole Moment Does CHCl 3 (a tetrahedral molecule with carbon at the center) have a dipole moment? If so, show the orientation of the dipole moment. Example 8.

6 Example 8.3 A Dipole Moment Does CHCl 3 (a tetrahedral molecule with carbon at the center) have a dipole moment? If so, show the orientation of the dipole moment. Example 8.3 B Practice with Dipole Moments For each of the following determine the orientation of the dipole moment (if any). a. HI b. N 2 c. CCl 2 F 2 (carbon is the central atom) Mar 3 7:32 PM 8.4 Ions: Electron Configuration & Size When you finish this section you will be able to predict: The formulas of simple ionic compounds The relative size of ions Your textbook deals only with nonmetals, representative metals, and ionic bonds in this discussion. It has been observed that atoms that form bonds in stable compounds have a noble gas electron configurations (Each is isoelectronic with a noble gas.) Isoelectronic ions are ions containing the same number of electrons S 2 is isoelectronic with Ca 2+. Because Ca 2+ has two more protons than electrons, while S 2 has more electrons than protons, S 2 is much larger than Ca 2+. Therefore, we can conclude that for an isoelectronic series, the more positive the nuclear charge (Z), the smaller the ion. Mar 3 5:17 PM 6

8.5 Lattice Energies and the Strength of the Ionic Bond The strength of the bond between the ions of opposite charge in an ionic compound depends on the charges on the ions and the distance between

7 8.5 Lattice Energies and the Strength of the Ionic Bond The strength of the bond between the ions of opposite charge in an ionic compound depends on the charges on the ions and the distance between the centers of the ions when they pack to form a crystal. An estimate of the strength of the bonds in an ionic compound can be obtained by measuring the lattice energy of the compound, which is the energy given off when oppositely charged ions in the gas phase come together to form a solid. The bond between ions of opposite charge is strongest when the ions are small. Lattice Energies of Alkali Metals Halides (kj/mol) The lattice energies for the alkali metal halides is therefore largest for LiF and smallest for CsI, as shown in the table. The ionic bond becomes stronger as the charge on the ions becomes larger. Lattice Energies of Salts of the OH and O 2 Ions (kj/mol) The data in the table shows that the lattice energies for salts of the OH and O 2 ions increase rapidly as the charge on the ion becomes larger. Mar 3 4:53 PM 8.6 Partial Ionic Character of Covalent bonds Check your understanding by answering the following questions 1. Why do we say that there are no totally ionic bonds? 2. Why is it ambiguous to say that Na 2 CO 3 or KHSO 3 are ionic compounds? 8.7 The Covalent Chemical Bond: A Model Check your understanding by answering the following questions 1. What is a chemical bond? Why do chemical bonds occur? 2. Knowing the energy of stabilization of CCl 4, discuss how we would determine the bond energy of C H bond in CHCl 3 3. What is a model? Why do we develop models?list the fundamental properties of a model. 4. What are the limitations of models? Is a wrong model useless? Why or why not? 8.9 The Localized Electron Bonding Model Check your understanding by answering the following questions 1. What is the covalent bonding model described in this section? 2. What is the basic assumptions about electron position in this model? 3. What are lone pairs and bonding pairs? Mar 3 7:51 PM 7

8.10 Lewis Structures Draw Lewis structures for a variety of covalent molecules. Lewis structures are often used to depict bonding pairs and lone pairs in molecules.

8 8.10 Lewis Structures Draw Lewis structures for a variety of covalent molecules. Lewis structures are often used to depict bonding pairs and lone pairs in molecules. We are concerned only with valence electrons because these are the ones (for Period 1 and 2 atoms) that are involved in bond making and breaking. Individual atoms are represented with Lewis structures by putting valence electrons (as dots or circles) around the atomic symbol. Example 8.10A Lewis Dot Structures Draw the Lewis dot structures for the following atoms or ions a. N b. N 3 c. I d. Ba e. Ba +2 Mar 3 7:46 PM ELEMENTs in PERIODS 1 & 2 (with the exception of H, He, B, and Be) CAN FORM COMPOUNDS OF LOWEST ENERGY IF THEIR HIGHEST ENERGY (s 2 p 6 ) LEVELS ARE FILLED. THIS IS CALLED THE OCTET RULE Steps for Writing Lewis Structures 1. Sum the number of valence electrons from all the atoms. Do not worry about keeping track of which electrons come from which atoms. It is the total number of electrons that is important. (Add electrons if it is an anion and subtract for cations) 2. Use a pair of electrons to form a bond between each pair of bond atoms. 3. Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the second row elements. Place least electronegative element in center (usually) Give each atom 8 electrons except for hydrogen (only give hydrogen 2 electrons). Hydrogen cannot go between 2 atoms, only around outside. All carbon valence electrons must be bonded. ALWAYS CHECK OCTETS!!!! Example 8.10B More Lewis Structures Write the Lewis structures that obey the octet rule for the following: a. Cl 2 b. CH 2 Cl 2 c. H 2 CO d. NH 3 Mar 3 4:53 PM 8

8.11 Exceptions to the Octet Rule The octet rule for drawing Lewis structures states that all atoms must be surrounded by eight electrons.

Atoms in molecules with odd numbers of electrons, such as NO and NO 2, cannot satisfy the octet rule.

Elements in period 3 or higher have unoccupied d orbitals in the valence shell. These d orbitals can hold up to 10 electrons beyond the 8 usually held in the valence shell.

9 8.11 Exceptions to the Octet Rule The octet rule for drawing Lewis structures states that all atoms must be surrounded by eight electrons. However, boron and beryllium are sometimes stable in compounds with less than an octet. In compounds, Be is generally satisfied with 4 valence electrons and B is generally satisfied with 6. Atoms in molecules with odd numbers of electrons, such as NO and NO 2, cannot satisfy the octet rule. Second period atoms can have only 8 electrons when in molecules because the n = 2 energy level can have only s and p orbitals. Elements in period 3 or higher have unoccupied d orbitals in the valence shell. These d orbitals can hold up to 10 electrons beyond the 8 usually held in the valence shell. In this case, the atom with more than 8 electrons is said to have an expanded octet. 17 electrons 11 electrons Mar 3 4:53 PM Lewis Structures: Comments about the Octet Rule The second row elements C, N, O, and F should always be assumed to obey the octet rule. The second row elements B and Be often have fewer than eight electrons around them in their compounds. These electron deficient compounds are very reactive. The second row elements never exceed the octet rule, since their valence orbitals (2s and 2p) can accommodate only eight electrons. Third row and heavier elements often satisfy the octet rule, but can exceed the rule by using their empty valence d orbitals. When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d orbitals (elements in period 3 or beyond). Example 8.11 Exception to the Octet Rule Write the Lewis structure for the following molecules: a. IF 2 b. SbF 5 c. XeF 2 d. I 3 Mar 3 8:26 PM 9

10 8.12 Resonance For example: CO 3 2. The double bond could have been placed on any of the three oxygens: Measurements in bond lengths suggests that all three C O bond lengths are equivalent. Electrons move around the entire molecule. Therefore, the actual structure is a time average of all these structures. These structures are called resonance structures. The Lewis structure of the molecule can be drawn any of three ways. Resonance Tutorial Mar 3 4:53 PM Formal Charge Formal Charge is the difference between the number of valence electrons on the free atom and the number of electrons assigned to the atom in the molecule. It is based on the localized electron (LE) model. For nitrate ion The sum of the formal charges must always equal the charge on the ion Example 8.12A Formal Charges Expanded Valence Shells Assign formal charges to each atom in the following resonance structures of CO 2. Mar 3 4:54 PM 10

11 Patterns of formal charge Although we can easily calculate formal charge according to the formula above, it is helpful to be able to recognize patterns for selected elements. For example, carbon with four covalent bonds and no lone pairs has a formal charge of zero. Similarly, nitrogen with three covalent bonds and one lone pair and oxygen with two covalent bonds and two lone pairs both have formal charge of zero. Furthermore, for any element, converting a lone pair into a covalent bond changes the formal charge by plus one. Converting a covalent bond into a lone pair changes the formal charge by minus one. Mar 3 9:04 PM Sometimes the formal charges can help one decide between the relative importance of non equivalent resonance structures. The best structure has the fewest formal charges and has the negative charge on the highest electronegative atom, as can be seen in the cyanate anion What is the 'best' structure for the cyanate anion, above? The worst? Example 8.12B Resonance Draw all resonance structures and select the most stable one for SCN Mar 3 9:08 PM 11

12 8.13 Molecular Structure: The VSEPR Model What is VSEPR? The Valence Shell Electron Pair Repulsion (VSEPR) model: based on the number of regions of high electron density around a central atom. can be used to predict structures of molecules or ions that contain only nonmetals by minimizing the electrostatic repulsion between the regions of high electron density. can also be used to predict structures of molecules or ions that contain multiple bonds or unpaired electrons. does fail in some cases. VSEPR Rules 1. Draw the Lewis structure for the molecule or ion. 2. Count the total number of regions of high electron density (bonding and lone electron pairs) around the central atom. Double and triple bonds count as ONE REGION OF HIGH ELECTRON DENSITY. A lone electron pair counts as ONE REGION OF HIGH ELECTRON DENSITY. For molecules or ions that have resonance structures, use the most likely resonance structure(s). Mar 3 9:11 PM Practice VSPER Mar 3 4:54 PM 12

13 VSEPR Rules Continued Identify the most stable arrangement of the regions of high electron density as ONE of the following: linear trigonal planar tetrahedral trigonal bipyramidal octahedral Linear arrangement of 2 regions of high electron density. The central atom and the regions of high electron density are arranged in a straight line. A linear arrangement minimizes electron repulsion for molecules or ions with a total of 2 regions of high electron density. Trigonal planar arrangement of 3 regions of high electron density. The regions of high electron density point at the corners of an equilateral triangle. A trigonal planar arrangement minimizes electron repulsion for molecules or ions with a total of 3 regions of high electron density. Mar 3 4:54 PM Tetrahedral arrangement of 4 regions of high electron density. The regions of high electron density point at the corners of a tetrahedron. A tetrahedral arrangement minimizes electron repulsion for molecules or ions with a total of 4 regions of high electron density. Trigonal bipyramidal arrangement of 5 regions of high electron density. Three regions of high electron density point at the corners of an equilateral triangle. One region of high electron density is directly above the plane of the triangle, and one is directly below the plane. A trigonal bipyramidal arrangement minimizes electron repulsion for molecules or ions with a total of 5 regions of high electron density Mar 3 9:30 PM 13

Determine the positions of the atoms based on the types of electron pairs present (i.e., bonding pairs vs. unshared pairs).

14 Octahedral arrangement of 6 regions of high electron density. The regions of high electron density point at the corners of an octahedron. An octahedral arrangement minimizes electron repulsion for molecules or ions with a total of 6 regions of high electron density. 4. Determine the positions of the atoms based on the types of electron pairs present (i.e., bonding pairs vs. unshared pairs). For trigonal bipyramidal and octahedral arrangements, there can sometimes be more than one possible arrangement of the bonding and unshared pairs: Trigonal bipyramidal place any unshared pairs in the plane of the triangle. Octahedral if you have two unshared pairs, place them on opposite sides of the central atom. 5. Identify the molecular structure based on the positions of the ATOMS (NOT on the regions of high electron density). VSPER Models Mar 3 4:54 PM Mar 3 9:24 PM 14

AP Chemistry A. Allan Chapter 8 Notes - Bonding: General Concepts

AP Chemistry A. Allan Chapter 8 Notes - Bonding: General Concepts 8.1 Types of Chemical Bonds A. Ionic Bonding 1. Electrons are transferred 2. Metals react with nonmetals 3. Ions paired have lower energy