Chapter 7. Chemical Bond Concept

Transcription

1 Chapter 7 Covalent Bonds & Molecular Structure Chemical Bond Concept Recall that an atom has core and valence electrons. Core electrons are found close to the nucleus. Valence electrons are found in the most distant s and p (and d for transition metals) energy subshells. It is valence electrons that are responsible for holding two or more atoms together in a chemical bond. Chapter 7 2 1

2 Octet Rule The octet rule states that atoms bond in such a way so that each atom acquires eight electrons in its outer shell. There are two ways in which an atom may achieve an octet. (a) By transfer of electrons from one atom to another (b) By sharing one or more pairs of electrons Chapter 7 3 Types of Bonds Ionic Bonds are formed from the complete transfer of electrons between atoms to form ionic compounds. Covalent Bonds are formed when two atoms share electrons to form molecular compounds. Chapter 7 4 2

3 Covalent Bonds Covalent bonds are formed when two non-metal atoms share electrons The bond is formed by the overlap of the atomic orbitals containing the outer most electrons When two electrons are shared in a bond, they are called bonding electrons Unshared electrons pairs are called non-bonding electrons or lone pairs Chapter 7 5 Covalent Bond Length Every covalent bond has a characteristic length that leads to maximum stability. This is the bond length Chapter 7 6 3

4 Strengths of Covalent Bonds Energy required to break a covalent bond in an isolated gaseous molecule is called the bond dissociation energy (D) Chapter 7 7 Polar Covalent Bonds Often, the two atoms in a covalent bond do not share the electrons equally When one of the atoms holds the shared electrons more tightly, the bond is polarized. A polar covalent bond is one in which the electrons are not shared equally. The unequal sharing of electrons is dependent on the electronegativity of the atoms in the bond Chapter 7 8 4

5 Electronegativity (EN) is the ability of an atom to attract bonding electrons towards itself. Electronegativity We can estimate the polarity of a bond calculating the difference in electronegativity between the two atoms in the bond F Electronegativites: Difference ( EN): 1.9 δ + δ - Delta Notation F Arrow Notation Chapter 7 9 Polar Covalent Bonds % Ionic Character: As a general rule for two atoms in a bond, we can calculate an electronegativity difference ( EN ): EN = EN(Y) EN(X) for X Y bond. If EN < 0.5 the bond is covalent. If 0.5 < EN < 2.0 the bond is polar covalent. If EN > 2.0 the bond is ionic. Chapter

6 Polar Covalent Bonds Using electronegativity values, predict whether the bonds in the following compounds are nonpolar covalent, polar covalent, or ionic: SiCl 4 CsBr FeBr 3 C 4 Cl CCl 4 N 3 2 O Chapter 7 11 Lewis Structures Using electron-dot (Lewis) structures, the valence electrons in an element are represented by dots. Valence electrons are those electrons with the highest principal quantum number (n). For the main group elements, they are the outermost s and p electrons. The number of valence electrons is equal to the group number for the main group elements. Chapter

7 Lewis Structures The Lewis structures provide a simple, but useful, way of representing chemical reactions. Ionic: Covalent: Chapter 7 13 Multiple Bonds More than one pair of electrons can be shared between two atoms. These types of bonds are called multiple bonds One shared pair is a single bond Two shared pairs are a double bond O C O Three shared pairs are a triple bond The electrons shared in these bonds belong to both atoms and can be counted as part of their octet just like in a single bond. Like single bonds, multiple bonds can also be polar or nonpolar. N N Chapter

8 Lewis Structures Single Bonds: C C Double Bonds: C C C C Triple Bonds: C C C C Chapter 7 15 Guidelines for Drawing Lewis Diagrams of Molecules Step 1: Determine the total number of valence electrons for the compound by adding up all of the valence electrons for each atom in the molecule. If the molecule is an anion, you add the charge to the total. If the molecule is a cation, you subtract the Step 2: Determine the arrangement of the atoms. Generally, the least electronegative element in the formula is the central atom with all the other attached to it (ydrogen is NEVER the central atom!) Attach each of the surrounding atoms to the central atom by a bond (a straight line) or with two electrons. Step 3: Determine the number of electrons surrounding the central atom (2 for each bond!) and subtract this value from your total valence electrons. Step 4: Use the remaining electrons to complete the octets around all the other atoms. Remember that hydrogen and helium only need two electrons. Also, review the exceptions to the Octet Rule. Step 5 (only sometimes): If you don t have enough electrons to fill all of the octets (and these atoms are not exceptions to the Octet Rule!), you will need to form a multiple bond. Chapter

9 Guidelines for Drawing Lewis Diagrams of Molecules The group number for the main group elements can also be used to determine the number of bonds that the element usually will form. This does not mean that the element cannot form more or less than this number. Chapter 7 17 Guide for Drawing Lewis Diagrams Let s look at N 3 as an example! STEP 1 Determine the total number of valence electrons Element Group # of Valence Atoms electrons N 5A 1 X 5 e - 1A 3 X 1 e - Total valence electrons for N 3 Total = 5 e - = 3 e - = 8 e - STEP 2 Determine the arrangement of the atoms Generally, the least electronegative element in the formula is the central atom with all the other atoms around it ydrogen is NEVER the central atom! So here, Nitrogen is the central atom, surrounded by the three hydrogens N N Chapter 7 is equal to 18 OR 9

10 Guide for Writing Electron-Dot Formulas, Cont. STEP 4 Determine the number of electrons remaining from your total (Total valence e - ) 2(Number of bonds) = Remaining e - (8 e - ) - 2(3) = 2 e - STEP 5 Place any remaining electrons as lone pairs around the atoms to fill their octets We have two remaining electrons. The Nitrogen still needs two electrons for a full octet so these two electrons are placed on the nitrogen as a lone pair Either structure below would be a correct answer: N OR N STEP 6 (only if needed!) If octets are not complete and you are out of electrons, form a multiple bond. Convert a lone pair into a bonding pair with the central atom. Chapter 7 19 O O S O O C O Lewis Structures Draw Lewis Structures for: C O 2 CO 2 N 2 4 C 5 N C 2 4 C 2 2 Cl 2 CO SF 4 SF 6 XeOF 4 XeF + 5 XeF 4 3 S + CO 3 Chapter

11 Resonance ow is the double bond formed in O 3? Move lone pair from this oxygen? O O O O O O or Or from this oxygen? O O O The correct answer is that both are correct, but neither is correct by itself. These structures are called resonance structures Chapter 7 21 Resonance When multiple structures can be drawn, the actual structure is an average of all possibilities. The average is called a resonance hybrid. A straight doubleheaded arrow indicates resonance. O O O O O O The nitrate ion, NO 3, has three equivalent oxygen atoms, and its electronic structure is a resonance hybrid of three electron-dot structures. Draw them. Chapter

12 Formal Charge Formal Charge is calculated for each element in the compound and used to determine the most energetically favorable resonance structure. Formal Charge = # of Valence e - # of bonding e- # of nonbonding e- 2 The most favorable resonance structure has: Formal charges as close to zero as possible A negative formal charge on the most electronegative atom and a positive formal charge on the least electronegative atom Calculate the FC and determine the most favorable resonance structure for the following: SO 2 CO 3 2- PO 4 3- Chapter 7 23 Shapes of Molecules The 3-D shape of a molecule plays a very important part in determining the chemical and physical properties of a compound. We know that electrons will repel one another As a result, electron pairs surrounding an atom will also repel each other. This phenomena causes molecules to take a 3-D shape that places the electron pairs (bonding and lone!) as far away from one another as possible. This is referred to as Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory allows us to predict the 3-D shape of a molecule There are two shapes that VSEPR gives us: Electron Pair Arrangement (EPA) indicates the arrangement of bonding AND lone pairs of electrons around the central atom. Molecular Geometry (MG) gives the arrangement of atoms (bonding pairs!) around the central atom as a result of electron repulsion. Chapter

13 Rules for Using VSEPR to Predict Shape Step 1: Draw the Lewis Diagram for the molecule Step 2: Determine the number of bonding pairs and lone pairs of electrons around the central atom - You count multiple bonds as ONE pair!! Step 3: Use VSEPR to determine the electron pair arrangement (EPA) for that number of electron pairs Step 4: To determine the molecular geometry (MG), cover up any lone pairs of electrons and rename your shape. - If there are no lone pairs: the MG = the EPA Chapter 7 25 EPA: Two Pairs Molecules with two electron pairs around the central atom form an EPA shape that is linear The bond angle for these molecules is 180 Chapter

14 EPA: Three Pairs Molecules with three electron pairs around the central atom form an EPA shape that is trigonal planar The bond angle for these molecules is 120 Chapter 7 27 EPA: Four Pairs Molecules with four electron pairs around the central atom form an EPA shape that is tetrahedral The bond angle for these molecules is Chapter

15 EPA: Five Pairs Molecules with five electron pairs around the central atom form an EPA shape that is trigonal bipyramidal There are two bond angles for these molecules: 120 and 180 Chapter 7 29 EPA: Six Pairs Molecules with six electron pairs around the central atom form an EPA shape that is octahedral There are two bond angles for these molecules: 90 and 180 Chapter

16 MG: Three Pairs If there are no lone pairs, EPA = MG so Trigonal Planar If there is one lone pair, the molecular geometry is Bent Bent No Lone Pairs One Lone Pair Chapter 7 31 MG: Four Pairs If there are no lone pairs, EPA = MG so Tetrahedral If there is one lone pair, the molecular geometry is now Trigonal Pyramidal If there are two lone pairs, the molecular geometry is now Bent Chapter

17 MG: Four Pairs If there are no lone pairs, EPA = MG so Trigonal Bipyrimidal If there is one lone pair, the molecular geometry is now See-Saw If there are two lone pairs, the molecular geometry is now T- Shaped Chapter 7 33 MG: Four Pairs (Cont.) If there are three lone pairs, the molecular geometry is now Linear Chapter

18 MG: Six Pairs If there are no lone pairs, EPA = MG so Octahedral If there is one lone pair, the molecular geometry is now Square Pyramid If there are two lone pairs, the molecular geometry is now Square Planar Chapter 7 35 Summary of VSEPR Theory Electron Groups Lone Pairs Bonds Geometry Examples Linear BeCl Trigonal planar BF Bent SO Tetrahedral C Trigonal pyramidal N Bent 2 O Trigonal bipyramidal PCl See-saw SF T-Shaped ClF linear - I Octahedral SF Square pyramidal 2- SbCl Square planar XeF 4 Also see Table 7.4 in your book Chapter

19 Molecular Shapes: VSEPR Draw the Lewis electron-dot structure and predict the shapes of the following molecules or ions: O 3 3 O + XeF 2 PF 6 XeOF 4 AlCl 3 BF 4 SiCl 4 ICl 4 Chapter 7 37 Polarity of Molecules A polar molecule contains an unequal distribution of charge resulting in positive and negative ends of the molecule (aka. Poles!) This unequal distribution of charge is the result of an unequal sharing of the electrons in the bonding pairs (polar bonds) Chapter

20 Polarity of Molecules To determine the polarity of an entire molecule you must know two things: The polarity of bonds within the molecule The electron pair arrangement (EPA) For the molecule to be non-polar, all polar bonds and any lone pairs must be cancelled out by the same type of bond pointing in the opposite direction Chapter 7 39 Rules for Determining the Polarity of Molecules Rule 1: If the central atom has an odd number of lone pairs, the molecule is polar One exception is a Linear Molecule (MG) that has a Trigonal Bipyramidal EPA Rule 2: If there are no lone pairs and the central atom is bound to only one type of atom (for example, C 4 ) then the molecule is non-polar. Rule 3: If there are no lone pairs on the central atom, but it is surrounded by more than one type of atom (for example, C 3 Cl), you must look at the shape of the molecule. Linear, Trigonal planar & Tetrahedral = Polar Trigonal Bipyramidal & Octahedral = Look at structure Chapter

21 Valence Bond Theory Lewis Structures describe a bond as a sharing of a pair of electrons. Valence Bond Theory explains how electrons become shared by the overlap of atomic orbitals Covalent bonds are formed by overlapping of atomic orbitals, each of which contains one electron of opposite spin. Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms. The greater the amount of orbital overlap, the stronger the bond. Chapter 7 41 Valence Bond Theory: ybridization ow does the VBT describe the electronic structure and 3-D shape of a polyatomic molecule? Let s look at Methane (C 4 ) C According to the Lewis Structure, we need to make four bonds. Each hydrogen brings one electron to the molecule and the carbon brings four. What is the electron configuration of Carbon? 1s 2 2s 2 2p x1 2p y 1 ow many unpaired electrons do we have? ow many do we need? ow do we get more? Chapter

22 Valence Bond Theory: ybridization So, if the excited state of Carbon uses two types of orbitals (the 2s and 2p), how can it form four equivalent bonds? Also, the 3-D orientation of the p and s orbitals are different. ow do we end up with a tetrahedral shape? To explain the bonding in Methane, VBT uses ybrid Orbitals ybrid Orbitals are atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine in preparation for covalent bond formation. Chapter 7 43 Valence Bond Theory: ybridization sp 3 ybrid Orbital Methane (C 4 ) Chapter

23 Valence Bond Theory: ybridization sp ybrid Orbital What EPA arrangement do you think this is? Chapter 7 45 Valence Bond Theory: ybridization What EPA arrangement do you think this is? sp 2 ybrid Orbital Chapter

24 Valence Bond Theory: ybridization sp 3 d ybrid Orbital What EPA arrangement Chapter 7 do you think this is? 47 Valence Bond Theory: ybridization sp 3 d 2 ybrid Orbital What EPA arrangement Chapter 7 do you think this is? 48 24

25 Rules for Determining ybridization Step 1: Draw the Lewis Diagram for the molecule Step 2: Determine the number of bonding pairs and lone pairs of electrons around the central atom - You count multiple bonds as ONE pair!! Step 3: Use VSEPR to determine the electron pair arrangement (EPA) for that number of electron pairs Step 4: Add s, p and d orbitals until you have a number of orbitals equal to the EPA number. Cl Cl P Cl Cl Cl Let s look at PCl 5 EPA is Trigonal Bipyramidal so we need 5 orbitals s + p + p + p + d = 5 orbitals = sp 3 d Chapter 7 49 ybridization in Molecules with Multiple Bonds Molecules with multiple bonds (double and triple bonds) will have nonhybridized orbitals that will overlap to form the multiple bonds. For a double bond: a hybridized orbital will overlap for the first bond (a sigma (σ) bond) a non-hybridized orbital will overlap for the second bond (a pi (π) bond) For a triple bond: a hybridized orbital will overlap for the first bond (a sigma (σ) bond) two non-hybridized orbitals will overlap for the second and third bonds (two pi (π) bonds) Chapter

26 ybridization in Molecules with Multiple Bonds Let s look at Ethylene (C 2 4 ) C C What is the EPA for each carbon? What is the hybridization of the orbitals for each carbon? Double Bond: 1 Sigma (σ) / 1 Pi (π) Chapter 7 51 ybridization in Molecules with Multiple Bonds Let s look at Acetylene (C 2 2 ) What is the EPA for each carbon? What is the hybridization of the orbitals for each carbon? C C Triple Bond: 1 Sigma (σ) / 2 Pi (π) Chapter

27 Molecular Orbital Theory (MO) The molecular orbital theory (MO) provides a better explanation of chemical and physical properties than the valence bond theory (VB). For Example, let s talk about O 2 : O O Looking at the Lewis Structure, all the electrons appear to be paired making the molecule diamagnetic. Experiments have shown that O 2 is actually paramagnetic with two unpaired electrons ow is this possible? Atomic Orbital: Probability of finding the electron within a given region of space in an atom. Molecular Orbital: Probability of finding the electron within a given region of space in a molecule. Chapter 7 53 Molecular Orbital Theory (MO) Additive combination of orbitals (σ) is lower in energy than two isolated 1s orbitals and is called a bonding molecular orbital. Subtractive combination of orbitals (σ * ) is higher in energy than two isolated 1s orbitals and is called an antibonding molecular orbital. Chapter

28 Molecular Orbital Theory (MO) Molecular Orbital Diagram for 2 Molecule: 1s 1 from each Chapter 7 55 Molecular Orbital Theory (MO) Molecular Orbital Diagrams for 2 - and e 2 Molecules Chapter

29 Molecular Orbital Theory (MO) Additive and subtractive combination of p orbitals leads to the formation of both sigma (σ / σ * ) and pi (π / π * )orbitals. Chapter 7 57 Molecular Orbital Theory (MO) Second-Period MO Energy Level Diagrams Orbital Order has flipped here! Chapter

30 Molecular Orbital Theory (MO) MO Diagrams Can Predict Magnetic Properties: Chapter 7 59 Molecular Orbital Configurations Molecular Orbital Configurations are analogous to Atomic Electron Configurations in that they tell us how the electrons are distributed within the molecule In order to write the electron configuration of the molecule, we first need to arrange the molecular orbitals in order of increasing energy. σ 1s < σ * 1s < σ 2s < σ* 2s < π 2py = π 2pz < σ 2px < π* 2py = π* 2pz < σ* 2px Step 1: Determine the TOTAL number of electrons in the molecule (not just the valence electrons. Step 2: Using the order shown above, start filling the molecular orbitals lowest energy to highest energy. For elements in Groups 6A and 7A, the orbitals in red are switched Step 3: Each atomic orbital can hold two electrons Step 4: When electrons are added to orbitals of the same energy (the pi orbitals), add one to each orbital, then pair. Li 2 C 2 N 2 - F 2 F 2 + Chapter

31 Molecular Orbital Theory (MO) Molecular Orbital Diagrams for heteronuclear molecules have skewed energies for combining the orbitals due to differences in electronegativity of the atoms. The more electronegative elements are lower in energy that the less electronegative elements. 2p 2s 1s σ * 2px π * 2py π * 2pz σ 2px π 2py π 2pz σ * 2s σ 2s σ * 1s 2s CO 2p σ 1s Chapter 7 61 C 1s O Molecular Orbital Theory: Bond Order Bond Order (BO) is the number of electron pairs shared between atoms. # of Bond Order = - Bonding e # of 2 - Antibonding e The BO can give us info about the molecule and/or bonds: The greater the BO, the more stable the molecule or ion. The greater the BO, the shorter the bond length, the stronger the bond. Calculate the BO for the following: O 2 O 2 + O 2 - NO Chapter

32 Molecular Orbital Theory: Delocalization VBT has a second problem in that it cannot describe molecules with resonance structures with a single structure. MOT describes bonds in which the electrons are delocalized (spread out) over the entire molecule. This is a more accurate picture of pi bonds Chapter