Chemical Bonding AP Chemistry Ms. Grobsky

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1 Chemical Bonding AP Chemistry Ms. Grobsky

2 What Determines the Type of Bonding in Any Substance? Why do Atoms Bond? The key to answering the first question are found in the electronic structure of the atoms involved In general, bonding is the interplay between interactions between atoms There are two types of interactions between atoms: Energetically favored Electrons on one atom interacting with protons of another atom Energetically unfavorable Electrons on one atom interacting with electrons of another atom Protons on one atom interacting with protons of another atom A bond will form if the system can LOWER its total energy in the process

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4 Energy Changes & Chemical Bonding As you know, most chemical reactions can be explained in terms of the rearrangements of bonds Energy must be added to break existing bonds in the reacting substances Energy is released when new bonds are formed in the products

5 Types of Bonds Ionic Bond Defined as the energy released when gaseous ions react to form one MOLE of a solid ionic compound For the formation of ionic compounds, the H rxn is based upon all of the energy changes involved in the transfer of electrons between the metal and nonmetal But, how does an electron transfer occur? Based upon two atomic properties discussed earlier in the year ionization energy and electron affinity

6 The Energetic Process of Ionic Bond Formation Ionization energy must be added to completely remove electrons from the metal in the gaseous state Since most metals are solids at room temperature, energy must be added to vaporize the metal before this ionization occurs This energy value is called the heat of formation, Hº f Then, energy is released when the nonmetal gains electrons, as measured by the electron affinity value Energy is also released when the ions bond to form a crystal lattice structure, called the lattice energy

7 A Specific Example of Ionic Bond Formation 2Na(s) + Cl 2 (g) ---> 2NaCl(s) 2Na(s) ---> 2Na(g) Add Hº f = (108 kj/mole Na x 2 moles Na) = 216 kj 2Na(g) ---> 2Na 1+ (g) + 2e - Add ionization energy = (496 kj/mole Na x 2 moles Na) = 992 kj Cl 2 (g) ---> 2Cl(g) Add bond energy = 243 kj/mole Cl 2 x 1 mole Cl 2 ) = 243 kj 2Cl(g) + 2e > 2Cl 1- (g) Release electron affinity = 349 kj/mole Cl 1- x 2 moles Cl 1- = 698 kj 2Na 1+ (g) + 2Cl 1- (g) ---> 2NaCl(s) Release lattice energy = 788 kj/mole NaCl x 2 moles NaCl = 1576kJ H rxn = (216 kj kj kj) (698 kj kj) = -823 kj per mole of reaction

8 Energetics of Ionic Bond Formation and Coulomb s Law The principal reason ionic compounds are stable is the attraction between ions of opposite charge Electrostatic in nature This attraction draws the ions together, releasing energy and causing the ions to form a crystal lattice A measure of how much stabilization results from arranging oppositely charged ions in an ionic solid is given by the lattice energy Defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous ions Mathematically, this relationship is represented by Coulomb s Law

9 Coulomb s Law and Lattice Energy Lattice Energy = k( Q`Q 2 ) r k = proportionality constant that depends on structure of solid E = joules r = distance between ion centers in nanometers Q1 and Q2 = numerical ion charges Negative E indicates an attractive force - ion pair has lower energy than separated ions Positive E indicates repulsive energy - two like-charged ions are brought together For a given arrangement of ions, the lattice energy increases as the charges on the ions increase and as their radii decrease

10 Ionic Bonds and Lattice Energy Lattice Energy = k( Q`Q 2 ) r Most negative lattice energy occurs between large charges and small ionic radii Highly favorable Least negative lattice energy occurs between small charges and large ionic radii Less favorable Looking at the above equation, it is clear that small, highly charged ions form strong, favorable ionic compounds The lattice energies of the following ionic compounds can be arranged as such: NaCl < NaF < MgS < MgO Strong interactions between ions have a profound effect on melting points and solubilities Large melting points Solids at room temperature

11 Electron Configurations of Ions of s- and p- Block Elements The energetics of bond formation helps explain why many ions tend to have noble-gas electron configurations For example, sodium readily loses one electron to form Na +, which has the same electron configuration as Ne Even though lattice energy increases with increasing ionic charge, we never find ionic compounds that contain Na 2+ ions This is because the second electron removed would have to come from an inner shell of the sodium atom, and removing electrons from an inner shell require a very large amount of energy This increase in lattice energy is not enough to compensate for the energy needed to remove an inner-shell electron

12 Another Type of Bond - Covalent Bonds Defined as a type of bond in which valence electrons are shared between nuclei of bonding atoms As a result, physical properties vary wildly Bond order is the number of bonding electron pairs shared by two atoms in a molecule Bonds can be single, double, or triple as shown by Lewis Dot structures Single bond (bond order = 1) One pair of electrons shared Called a sigma bond Double bond (bond order = 2) Two pairs of electrons shared One sigma and one pi bond Triple bond (bond order = 3) Three pairs of electrons shared One sigma and two pi bonds

13 Wait What s a Sigma Bond? A sigma (σ) bond is the bonding pair held directly between both nuclei in a molecule In other words, the electron pair occupies the space between the nuclei This is a result of an overlap of the orbitals Example: The H-H bond forms as a result of the overlap of the 1s orbitals from each atom 74 pm

14 What s a Pi Bond? Multiple bonds are formed from sharing an electron pair in the space above and below the sigma bond located between the two nuclei Called a pi (π) bond Therefore, a multiple bond always consists of: One σ bond where the electron pair is located directly between the atoms At least one π bond where the shared pair occupies the space above and below the σ bond Example - C 2 H 4 π bond H C σ bond C H H H π overlap

15 Bond Energies of Sigma and Pi Bonds Multiple bonds increase electron density between two nuclei Decreases nuclear repulsions while enhancing the nucleus to electron density attractions As a result, nuclei move closer together Bond lengths from shortest to longest are as follows: Triple bond < Double bond < Single bond So, the greater the number of electron pairs between a pair of atoms, the shorter the bond This implies that atoms are held together more tightly when there are multiple bonds So, it s obvious that combinations of sigma and pi bonds are stronger than sigma alone Pi bonds are weaker than sigma but NEVER exist alone Sure enough, there IS a relation between bond order and the energy required to separate them

16 Energetics of Covalent Bonding For the reactions of covalent compounds, the net energy of the reaction, H rxn = (sum of energy added to break bonds) - (sum of energy released when bonds form) H = D bonds broken D (bonds formed) D represents the bond energy per mole of bonds (ALWAYS POSITIVE) To calculate the enthalpy change for a reaction involving covalent molecules, write the Lewis dot structures for the reactants and products Then, determine which bonds must be broken and which bonds must be formed Use the average bond energy values to calculate the energy needed to break the bonds in the reactants and the energy released when the bonds form in the products

17 An Example of a Formation of a Covalent Bond Consider the reaction for the formation of water from elemental hydrogen and oxygen: 2H 2 (g) + O 2 (g) ---> 2H 2 O(g) To break the H-H bond requires 436 kj/mole H 2 (g) Since there are two moles of H 2 (g) in the reaction, this value is doubled to 872 kj The oxygen atoms in the O 2 (g) are held together by a double bond which requires 495 kj/mole to break For each water molecule produced, there are two O-H bonds formed For the 2H 2 O(g), the total energy released when four moles of O- H bonds form is: 4 moles O-H x 463 kj/mole = 1852 kj H = kj mol kj mol kj mol = 485 kj

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19 Using Localized Electron Model to Describe Covalent Bonding The Localized Electron Model bonding theory assumes that electrons are localized on an atom or the space between atoms in a covalent bond Lone pair electrons Bonding pair electrons Has 3 parts: Lewis Dot structures describe valence electron arrangement VSEPR theory predicts molecular geometry Hybridization describes the type of atomic orbitals blended by the atoms to share electrons or hold lone pairs

20 How to Illustrate Bonding -Lewis Dot Structures Illustrate valence electrons and subsequent bonding A line shows each shared electron pair Dots represent unpaired electrons called lone pairs

21 How to Draw Lewis Dot Structure Determine total number of valence electrons Add for anions, subtract for cations Predict # of bonds by counting the number of unpaired electrons in Lewis structure Least electronegative atom is the center atom Remember the trend! Draw a single bond, -, (2 electrons) to each atom Subtract from total # of electrons Add lone pair electrons, :, to terminal atoms to satisfy octet rule Extras go to central atom If central atom is not octet, use terminal electrons to make double bond Carbon bonded to N, O, P, S tend to form double bonds Hydrogen is ALWAYS a terminal atom Only makes 1 bond

22 Bond Polarity and Electronegativity (E n ) Bond polarity is a measure of how equally or unequally the electrons in any covalent bond are shared A non-polar covalent bond is one in which the electrons are shared equally A polar covalent bond is one in which one of the atoms exerts a greater attraction for the bonding electrons than the other Bond polarity is based on electronegativty Defined as the ability of an atom IN A MOLECULE (meaning it s participating in a bond) to attract shared electrons to itself F is most electronegative Highest Zeff and smallest radius so that the nucleus is closed to the action Fr is least electronegative Lowest Zeff and largest radius so that the nucleus is farthest from the action Can use the difference in electronegativities to determine type of bond formed Ionic electronegativity difference greater than 1.67 Covalent electronegativity difference less than 1.67 Non-polar covalent electronegativity differences less than 0.4

23 Using the Idea of Bond Polarity to Draw Lewis Dot Structures Write the electron dot diagrams for each element in the compound Check the electronegativity difference between the elements to determine if electrons are transferred or shared If the electronegativity difference > 1.67, the reaction forms ions Remove the electrons from the metal and add them to the nonmetal

24 Lewis Dot Structure of Ionic Compounds Write the charges of the ions formed and use coefficients to show how many of each ion are needed to balance the overall charge + 2-2Na,[ O ] Ionic sodium oxide, Na 2 O

25 Lewis Dot Structure of Covalent Compounds If the electronegativity difference < 1.67, then the atoms will share electrons Position shared electron pairs between the two atoms, and connect them with a single line to represent a covalent bond Place the extra pairs of electrons around atoms until each has eight

26 Lewis Dot Structure of Covalent Compounds If an atom other than hydrogen or a metal has less than eight electrons, move unshared pairs to form multiple bonds Add extra atoms, if needed, to obtain the octets Atoms with positive oxidation numbers should be bonded to those with negative oxidation numbers If extra electrons still remain, add them to the central atom All oxidation numbers should add up to zero for a compound

27 Exceptions to the Octet Rule Not all compounds obey the octet rule Three types of exceptions: Species with more than eight electrons around an atom Species with fewer than eight electrons around an atom Species with an odd total number of electrons

28 : : Atoms with Fewer than Eight Electrons Referred to as electron deficient Examples - Beryllium and boron will both form compounds where they have less than 8 electrons around them.. :Cl Be Cl: :F.. B F:.. :F:..

29 Atoms with More than Eight Electrons Except for species that contain hydrogen, this is the most common type of exception For elements in the third period and beyond, the d orbitals can become involved in bonding Examples 5 electron pairs around P in PF 5 5 electron pairs around S in SF 4 6 electron pairs around S in SF 6

30 Species with an Odd Total number of Electrons A very few species exist where the total number of valence electrons is an odd number This must mean that there is an unpaired electron which is usually very reactive A radical is a species that has one or more unpaired electrons They are believed to play significant roles in aging and cancer

31 : Species with an Odd Total Number of Electrons Nitrogen monoxide (aka nitric oxide) is an example of a compound with an odd number of electrons It has a total of 11 valence electrons - six from oxygen and 5 from nitrogen The best Lewis structure for NO is:. :N O:

32 Single Covalent Bonds Cl H Be H C H F F Cl H Do atoms (except H or metals) have octets?

33 Lewis Dot Structures with Multiple Bonds Example CO 2 Step 1 Draw any possible structures C-O-O O-C-O You may want to use lines for bonds. Each line represents 2 electrons.

34 Lewis Dot Structures with Multiple Bonds Step 2 Determine the total number of valence electrons. CO 2 1 carbon x 4 electrons = 4 2 oxygen x 6 electrons = 12 Total electrons = 16

35 Lewis Dot Structures with Multiple Bonds Step 3 Try to satisfy the octet rule for each atom - all electrons must be in pairs - make multiple bonds as required Try the C-O-O structure C O O No matter what you try, there is no way satisfy the octet for all of the atoms.

36 Lewis Dot Structures with Multiple Bonds O C O This arrangement needs too many electrons. How about making some double bonds? O=C=O That works! = is a double bond, the same as 4 electrons

37 Multiple Bonds So how do we know that multiple bonds really exist? The bond energies and lengths differ! Bond Bond Length Bond energy type order pm kj/mol C C C C C C

38 Practice! # on page

39 Alternative Lewis Dot Structures - Resonance Structures Defined as a compound that has multiple equivalent structures A compound with resonance is best described as the average of all the equivalent structures Intermediate bond lengths Clarification of common misconceptions: Structures with resonance do not flip between equivalent structures Structures differ by placement of electrons, not atoms Octet rules still apply Not all compounds have resonance

40 Resonance Structures Sometimes we can have two or more equivalent Lewis structures for a molecule: O - S = O O = S - O They both: Satisfy the octet rule Have the same number of bonds Have the same types of bonds So, which is right?

41 Resonance Structures They both are! O -- S = O O = S -- O O S O This results in an average of 1.5 bonds between each S and O To quickly figure out the correct resonance structure, check formal charges

42 Formal Charges Formal charges is a bookkeeping system for electrons that is used to predict which possible Lewis Dot structure is more likely Defined as the charge assigned to an atom in a molecule assuming that electrons in a bond are shared equally between atoms Note oxidation state assumes NO sharing Important for determining which resonance structure is most significant contributor All atoms have formal charge of 0 (most important) Charges are consistent with electronegativity Atoms that are higher in electronegativity have negative formal charges

43 Determining Formal Charges Atom FC = # valence electrons [# of lone electrons # of bonding groups ]

44 Formal Charges O=C=O O C=O Structure 1 Structure 2 For each oxygen (4 electrons which assigned have: from unshared e e - from the bonds) = 6 total Formal charge = 6-6 = 0 For carbon 4 e - assigned from the bonds = 4 total Formal charge = 4-4 = 0 For the single-bond oxygen (6 e - from unshared e - + 1e - from bond) = 7 total Formal charge = 6-7 = -1 For the triple-bond oxygen (2 e - from unshared e - + 3e - from bonds) = 5 total Formal charge = 6-5 = +1 For carbon 4 e - from the bonds = 4 total Formal charge = 4-4 = 0 The most likely Lewis structures are those all atoms obeying the octet rule, all atoms with a formal charge of zero, or the most electronegative element with the negative formal charge.

45 Another Example of Formal Charges C=O C=O Structure 1 Structure 2 For oxygen (4 electrons assigned from unshared e e - from the bonds) = 6 total Formal charge = 6-6 = 0 For carbon (2 electrons assigned from unshared e e - from the bonds) = 4 total Formal charge = 4-4 = 0 Although Structure 1 has all atoms with a formal charge of zero, the carbon atom does not obtain an octet. Therefore, Structure 2 is the most likely Lewis structure since all atoms obey the octet rule. For oxygen (2 e - from unshared e - + 3e - from bonds) = 5 total Formal charge = 6-5 = +1 For carbon (2 electrons assigned from unshared e e - from the bonds) = 5 total Formal charge = 4-5 = -1

46 Practice! # on page

47 Putting This All Together- Molecular 3-D Geometry

48 Determining Molecular Geometries Lewis dot structures only show the number and types of bonds they do not provide any information about the shape of the molecule The shape of a molecule is determined by its bond angles and bond lengths These are affected by the energetically favorable and energetically unfavorable interactions of electrons and protons In order to predict molecular shape, we use the Valence Shell Electron Pair Repulsion (VSEPR) theory

49 The VSEPR Model This theory proposes that the geometric arrangement of groups of atoms about a central atom in a covalent compound is determined solely by the repulsions between electron pairs present in the valence shell of the central atom The molecule adopts whichever 3-D geometry minimizes the repulsion between valence electrons To determine the shape of a molecule, we must consider all electron domains in a molecule Defined as the regions where electron pairs may be found Therefore, we must distinguish between: Lone pairs (non-bonding pairs) Bonding pairs (those found between two atoms) Multiple bonds are considered as ONE bonding pair even though in reality, they have multiple pairs of electrons In general, each non-bonding pair, single bond, or multiple bond produces a single electron domain around the central atom in a molecule

50 Why Do We Have Different Electron Domains? This is because each electron domain affects the amount of electron repulsion around the atom differently Electron pairs of bonding atoms are shared by two atoms, whereas the nonbonding electron pairs (lone pairs) are attracted to a single nucleus As a result, lone pairs can be thought of as having a somewhat larger electron cloud near the parent atom This crowds the bonding pairs and the geometry is distorted! Multiple bonds exert a greater repulsive force on adjacent electron pairs than do single bonds as a result of higher electron density All electrons are considered when determining 3-D shape

51 Factors Affecting Bond Angles and Molecular Geometries Multiple Bonds Non-Bonding (Lone) Pairs

52 The Five Electron-Domain Geometries 2 e- domains 3 e- domains 4 e- domains 5 e- domains 6 e- domains

53 Determining a Molecule s Geometry Note the arrangement of all electron domains about the central atom in an AX n molecule or ion is called its electron-domain geometry The molecular geometry is the arrangement of only the atoms in a molecule or ion Any non-bonding pairs in the molecule are not part of the description of the molecular geometry

54 Steps in Determining a Molecular Shape Draw the Lewis dot structure Count number of electron domains (bonding pairs and lone pairs) around the central atom For molecules with MORE THAN ONE CENTRAL ATOM, work with one central atom at a time Determine the electron-domain geometry by arranging the electron domains about the central atom so that the repulsions among them are minimized Determine the molecular geometry with the help of the following formula : AX m E n A - central atom X surrounding atom E non-bonding valence electron group m and n - integers

55 The Single Molecular Shape of Linear Electron-Group Arrangement AX2 X A X

56 Example - CO2

57 The 2 Molecular Shapes of Trigonal Planar Electron-Domain Arrangement Trigonal Planar AX 3 Bent AX 2 E Example NO2 X X A X X E A X

58 Example of Trigonal Planar - BCl3

59 The 3 Molecular Shapes of the Tetrahedral Electron-Group Arrangement Tetrahedral AX 4 Trigonal Pyramidal AX 3 E Bent AX 2 E 2 X E E X A X X X A X X X A X E

60 Example of Tetrahedral - CH4

61 Example of Trigonal Pyramidal - NH3

62 Example of Bent H 2 O O H H

63 The 4 Molecular Shapes of the Trigonal Bipyramidal Electron-Group Arrangement Trigonal Bipyramidal AX 5 Examples PCl 5, PF 5, AsF 5, SOF 4 See-Saw T-Shaped Linear AX 4 E Examples SF 4, XeO 2 F 2, IF 4 +, IO 2 F 2 - AX 3 E 2 Examples ClF 3, BrF 3 AX 2 E 3 Examples XeF 2, I 3 -, IF 2 -

64 The 3 Molecular Shapes of the Octahedral Electron-Group Arrangement Octahedral AX 6 Examples SF 6, IOF 5 Square Pyramidal AX 5 E Examples BrF 5, XeOF 4, TeF 5 - Square Planar AX 4 E 2 Examples XeF 4, ICl 4 -

65 Time to Explore the Different Molecule Arrangements! VSEPR Simulation

66 VSEPR Theory and Polar/Nonpolar Molecules Most bonds between atoms of different elements in a molecule are polar This does not mean that the molecule will be polar O = C = O The electronegativity values show that the C-O bond would be polar with electrons being pulled towards the oxygens The vectors represent a dipole moment, µ, which is a separation of the charge in a molecule (slightly positive/slightly negative poles) The dipole moment increases as the magnitude of the charge that is separated increases and as the distance between the charges increases However due to its linear geometry, the pull happens in equal and opposite directions Electronegativities: Oxygen = 3.5 Carbon = 2.5 Difference 1.0 (Polar Bond)

67 VSEPR Theory and Polar/Nonpolar Molecules As we just saw, in order for a molecule to be polar, the effects of bond polarity must not cancel out One way is to have a geometry that is not symmetrical like the bent geometry in water H O H Electronegativity Difference = 1.3 Here, the effects of the polar bonds do not cancel, so the molecule is polar

68 Polar vs. Nonpolar Molecules A molecule is nonpolar if the central atom is symmetrically substituted by identical atoms CO 2, CH 4, CCl 4 A molecule will be polar if the geometry is not symmetrical H 2 O, NH 3, CH 2 Cl 2 The degree of polarity is a function of the number and type of polar bonds as well as the geometry

69 Geometry and Polar Molecules Again, in order for a molecule to be polar: Must have polar bonds Must have the proper geometry Examples: CH 4 Non-polar CH 3 Cl Polar CH 2 Cl 2 Polar CHCl 3 Polar CCl 4 Non-polar WHY?

70 Why is Polarity So Important? It affects physical properties such as melting point, boiling point and solubility Chemical properties also depend on polarity

71 VSEPR Symmetry and Molecular Polarity

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