Chemical Bonding and Molecular Geometry PDF

6 Chemical Bonding and Molecular Geometry  6.1 Introduction chemical Bonding Atoms with unfilled valence shells are considered unstable. Atoms will try to fill their outer shells by bonding with other atoms. Chemical bond = the attractive force that holds atoms or ions t...

6 Chemical Bonding and Molecular Geometry  6.1 Introduction chemical Bonding Atoms with unfilled valence shells are considered unstable. Atoms will try to fill their outer shells by bonding with other atoms. Chemical bond = the attractive force that holds atoms or ions together in a compound Atoms form atomic bonds to become more stable. – Atoms become more stable by filling their valence shell or at least meeting the octet rule by getting 8 valence electrons. There are three main types of chemical bonds used by atoms to fill their valence shell: – Ionic – Covalent – Metallic  6.1.1 Formation of Ionic compounds Ionic bonds are formed between metals and non-metals. Ionic bonds are formed between positively and negatively charged atoms (ions). Ionic bonds are formed by the transfer of electrons. –One atom loses (gives away) electrons. –One atom gains (receives) electrons. Use the number of valence electrons to determine the # of electrons that are lost or needing to be gained. The transfer of electrons create a positive ion and a negative ion. The opposite charges attract one another, causing a chemical bond to form. Atoms with 4 or less valence electrons want to LOSE (give away) their valence electrons. [Groups 1, 2, 13, 14] Atoms with 4 or more valence electrons want to GAIN (receive) more electrons to satisfy their octet. [Groups 14, 15, 16, 17]  Cont… The normal charge of an ion can be quickly determined using the oxidation number of an element. –The oxidation number of an atom is the charge that atom would have if the compound was composed of ions. To find the oxidation number : Look at Group # Determine # of valence electrons If 4 or less, atom will lose (give away) valence electrons (ion is positive) If 4 or more, atom will gain the needed # to fill valence shell. (ion is negative) 4  Cont… Example: –Beryllium is in Group 2 –Be has 2 e- –Wants to achieve octet –Loses the 2 e- –Oxidation #/Ion charge of +2 Example: –Nitrogen is in Group 15 –N has 5 e- –Needs 3 more for octet –Gains 3 e- –Oxidation #/Ion charge of -3 5  6.1.2 Electronic structure of Cations and Anions Ions are formed when atoms gain or lose electrons. Ions are charged atoms (positive or negative). Positive ions are called cations. –Formed when the atom loses electrons. –Lose negative charge, becomes positive ION –Metals Negative ions are call anions. –Formed when the atom gains electrons. –Gain negative charge, become negative ION –Non-metals  Drawing Ionic Bonds 1 – Draw the Lewis structure for each element. –Ex: Na Cl 2 – Draw arrows to show the TRANSFER (gain/loss) of electrons [draw extra atoms if needed] 3 – Draw ion Lewis diagrams showing the new charge for each ion. Ex: 4- Write the chemical formula for the compound formed represents the ratio of negative ions to positive ions. Ex: NaCl – for every 1 sodium ion, there is also 1 chlorine ion. Chemical Formula = NaCl  6.2 Covalent Bonds 6.2.1 Formation of covalent bonds Covalent bonds form between two non-metals. Groups 14-17 on the Periodic Table Covalent bonds are formed when atoms SHARE electrons. –Both atoms need to gain electrons to become stable, so they share the electrons they have. Atoms can share more than one pair of electrons to create double and triple bonds. Properties of Covalent Compounds Results in a NEUTRAL molecule Weak bonds Physical State usually liquids or gases Low Melting and Boiling Points Poor conductors of electricity (no free electrons to move around)  Cont … Use Lewis structures to draw valence electrons for each atom in the covalent pair. Each chlorine atom wants to gain one electron to achieve an octet. Cl Cl 9  Cont… The octet is achieved by each atom sharing the electron pair in the middle. octet Cl Cl octet Now, each Chlorine atom has 8 valence electrons because it is sharing one pair. 10  Chlorine Molecule It is a single bonding pair so it is called a single covalent bond. The compound is now called a molecule. Cl Cl Cl Cl Cl2 11  Cont… Elements can share up to three pairs of electrons. (6 total electrons). Single Bond (2e) Double Bond (4e) Triple Bond (6e) 12  Cont… Atoms can share their electrons equally or unequally. When atoms share electrons equally, it is called a non-polar covalent bond. –Non-polar covalent bonds form between atoms of the same type. Ex: H2, Cl2, When atoms share electrons unequally it is called a polar covalent bond. –One atom pulls the electrons closer to itself. –The atom that pulls the electrons more gets a slightly negative charge. –The other atom gets a slightly positive charge. Ex: Water molecule Bonding Animation 6.2.2 Polarity of covalent bonds ▪ Molecular dipoles occur due to the unequal sharing of electrons between atoms in a molecule. Those atoms that are more electronegative pull the bonded electrons closer to themselves. ▪ Even though the total charge on a molecule is zero, the positive and negative charges are not completely symmetrical in most molecules. ▪ These molecules are said to be polar because they possess a permanent dipole.  Nonpolar Molecule Nonpolar When atoms bond together to form molecules, they share or give electrons. If the electrons are shared equally by the atoms, then there is no resulting charge and the molecule is nonpolar. Dipoles & Polarity Dipoles maybe symbolized by either Greek letter delta (lowercase), δ , or arrows crossed at the positive end, or both  6.3 Lewis Structures Lewis Dot Diagrams represent the number of valence electrons present in an atom. Lewis structures are often used to indicate the bonds in a covalent molecule. Lines are used to represent bonds – 1 line = single bond, – 2 lines = double bond – 3 lines = triple bond – 2 dots represent lone pairs of electrons  6.3.1. Drawing Lewis Structures 1. Count the total number of valence electrons in the molecule. If you divide this by 2, it will give you the # of bonds needed to draw the structure. 2. Create a “skeleton” structure by connecting surrounding atoms to the central atom. The central atom is the one that there is one of or the least electronegative (generally the first element in the formula is the central). 3. Place electrons between the central atom and surrounding atoms so that it has an octet. (remember hydrogen only needs 2). 4. Complete octets on the outside atoms. 5. If you run out of electrons to complete the octets of the surrounding atoms, then you must move electrons from the central atom to the outside and create double or triple bonds between the central atom and a surrounding atom.  Practice: Drawing Covalent Bonds We can illustrate covalent bonding using Lewis structures. 1 – Draw a Lewis structure for each element. –Ex: C H 2 - Continue adding atoms until all atoms have a full valence H H C H CH4 carbon tetrahydride H  Cont.. Similarly, the formation of CCl4 and CH4 takes place as below: Some other examples of the formation of covalent compounds is shown below: (i) Single bonds: H2 Cl2 (ii) Double bonds: O2 CO2 (iii) Triple bonds: The covalent bonds are of two types: (i) Polar covalent bonds  Practice Draw a Lewis Structure for: – A water molecule: – Ammonia: – Carbon dioxide: – Methane: – Ethane: – Ammonium ion: Answer to practice 22  6.3.2. EXCEPTIONS TO THE OCTET RULE A. Hydrogen Only two electrons are required for H to obtain a noble gas configuration like He No double bonds on hydrogen Never put lone pairs on H in a lewis structure B. Electron Deficient Means the central atom has less than 8 electrons The second row elements B and Be often have fewer than 8 electrons around them in compounds and as a result are highly reactive Boron is satisfied with 6 electrons Berillium is fine with 4 electrons Boron Trifluoride BF3  C. Odd-Electron Molecules Molecules with odd numbers of electrons will result in a Lewis structure with one unpaired electron. Nitrogen is a period 2 element and can be satisfied with less than an octet. Nitrogen dioxide ………. 17 valence electrons D. Expanded Octet Means that the central atom is sharing more than 8 electrons. Third Row and heavier elements can exceed the octet rule by using their empty valence d orbitals. Xenon tetrafluoride XeF4  Resonance Resonance refers to when more than one valid Lewis structure can exist for a molecule. The actual structure lies somewhere in between the two as an average. The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms ignoring electronegativity. You can use formal charges to identify the most reasonable Lewis structure for a given molecule  Easy way to calculate formal charge Formal charge = FC Valence electrons = VE Bonded pairs = sticks Nonbonded electrons= dots FC = VE – (# of STICKS) - (# of DOTS) Guidelines for determining best structure A molecular structure in which all formal charges are zero. The structure with the smallest nonzero formal charges. Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign. When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable. 26 Nitrous oxide, N2 O, commonly known as laughing gas, is used as an anesthetic in minor surgeries, such as the routine extraction of wisdom teeth. Which is the likely structure for nitrous oxide? Answer 27  6.4. Resonance and Formal Charges Sometimes more than one valid Lewis structure (one that obeys the rules outlined) is possible for a given molecule. Actual structure is an average of the resonance structures. Electrons are really delocalized – they can move around the entire molecule. ATOMS do not move! 28  FORMAL CHARGE At times there are several possible Lewis structures for the same molecule. Formal charges are used to evaluate nonequivalent Lewis structures. (move the ATOMS, not electrons) Atoms in molecules try to achieve formal charges as close to zero as possible. Any negative formal charges are expected to reside on the most electronegative atoms. Two possible skeletal structures of formaldehyde (CH2O H C O H H C O H 29  Which one is the correct Lewis structure? To determine the correct structure, the charge on each atom in the various Lewis structures must be estimated and those charges are used to select the most appropriate structure. The formal charge of an atom in a molecule is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule. To determine the formal charge of a given atom in a molecule, we need to know two things. 1.The number of valence electrons on the free neutral atom (which has zero net charge because the number of electrons equals the number of protons). 2.The number of valence electrons “belonging” to the atom in a molecule. If in the molecule the atom has the same number of valence electrons as it does in the free state, the (+) and (-) charges balance and it has a formal charge of zero. 30 If the atom has one more valence electron in a molecule than it has as a free  formal charge total number total number total number ( ) 1 on an atom in a Lewis = of valence electrons in- of nonbonding - 2 of bonding electrons electrons structure the free atom The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. Make the following assumptions: 1. Lone pair electrons belong entirely to the atom in question. 2. Shared electrons are divided equally between the two sharing atoms. 31  -1 +1 C – 4 e- 2 single bonds (2x2) = 4 H C O H O – 6 e- 1 double bond = 4 2H – 2x1 e- 2 lone pairs (2x2) = 4 12 e- Total = 12 formal charge total number total number total number ( ) 1 on an atom in of valence a Lewis = electrons in - of nonbonding - 2 of bonding electrons electrons structure the free atom formal charge =4- 2- ½ x 6 = -1 on C formal charge =6- 2- ½ x 6 = +1 on O 32  0 0 H C – 4 e- 2 single bonds (2x2) = 4 C O O – 6 e- 1 double bond = 4 H 2H – 2x1 e- 2 lone pairs (2x2) = 4 12 e- Total = 12 formal charge total number total number total number ( ) 1 on an atom in of valence a Lewis = electrons in - of nonbonding - 2 of bonding electrons electrons structure the free atom formal charge =4- 0- ½x8=0 on C formal charge =6- 4- ½x4=0 on O 33  Formal Charge and Lewis Structures 1. For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. 2. Lewis structures with large formal charges are less plausible than those with small formal charges. 3. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH2O? -1 +1 H 0 0 H C O H C O H 34  Formal Charge An arithmetic formula for calculating formal charge. Formal charge = Number of number of number of – – valence bonds unshared electrons electrons 35  "Electron Counts" and Formal Charges in NH4+ and BF4-.. 1 H : F: 7 +.. –.. H N H :..F B..F : H :..F: 4 4 36 For OB # valence e- = 6 # nonbonding e- = 2 # bonding e- = 6 X 1/2 = 3 Formal charge = +1 For OC B # valence e- = 6 O # nonbonding e- = 6 O O # bonding e- = 2 X 1/2 = 1 For OA A C Formal charge = -1 # valence e- = 6 # nonbonding e- = 4 # bonding e- = 4 X 1/2 = 2 Formal charge = 0 37 Resonance and Formal Charge EXAMPLE: NCO- has 3 possible resonance forms - N C O N C O N C O A B C 38 Now Determine Formal Charges -2 0 +1 -1 0 0 0 0 -1 N C O N C O N C O Forms B and C have smaller formal charges; this makes them more important than form A. (rule 1) Form C has a negative charge on O which is the more electronegative element, therefore C contributes the most to the resonance hybrid. (rule 2) 39  6.5 strengths of ionic and covalent bonds 6.5.1. Ionic bonding strength and lattice energy Occurs most easily when elements with low ionization energies (metals) react with elements with high electron affinity (non metals). The formula of an ionic crystal (MX), indicates that the compound contains M and X in a 1:1 ratio and the formula does not indicate the ionic nature of the compound, only the ratio of atoms. Example: Na + 1/2F2 → NaF(s) (Na+, F-). Ca + Cl2 → CaF2 (s) (Ca2+, 2Cl-) Facts about ionic bonding; A. A positive ion forms by the loss of one or more electrons. Ions with the following structure can form ionic compounds: i) Ions with inert core electron in the following groups: 1, 2, 3, 4, etc……. Example:Na+, Ca2+, Sc3+, Ti4+, etc… (ii) Ions with 18-electron core-transition metals with configuration ns1-2(n-1)d10. Example: Cu+, Zn2+, Ga3+, etc… (iii) Ions formed from post-transition elements in the following groups: 13, 14, and 15. Example:Ti+, Pb2+, Bi+, etc…  B. A non-metal atom gains sufficient Cont… electrons to form an anion with the same electronic configuration as a noble gas. C. The oppositely charged ions come together to form an ionic compound. D. A formula unit of an ionic compound is the smallest collection of ions that would be electrically neutral. The formula unit is automatically obtained when the Lewis structure of the compound is written. 2.2.2. Structure of Ionic Solids and Lattice energy From X-ray study, components of crystals are ions. In a crystal of an ionic crystal each ion is surrounded by another ion in a regular manner and crystals are classified as MX, MX2 , and MX3 depending on the relative number of positive and negative ions. For example: NaCl is MX type crystal and each sodium ion is surrounded by six chloride ion and each chloride ion is surrounded by six sodium ion. Each ion is surrounded by the greatest possible number of oppositely charged ions and this number is called coordination number. Coordination numbers of 3, 4, 6, and 8 are common are related to the relative size of the ions.  Table 2.1. The relationship between coordination no of cation/anion with type of crystal Ionic crystal Type of crystal Coordination No. of cation Coordination No. of anion ZnS MX Four Four NaCl MX Six Six CsCl MX Eight Eight CaF2 MX2 Eight Four TiO2 MX2 Six Three K2O M2X Four Eight Coordination number can be predicted from simple geometric ratio, radius M+/radiusX- (called limiting radius ratio). Table 2.2. The relationship b/n Coordination number and limiting radius ratio. Coordination number Limiting radius ratio Example 3 0.155-0.255 LiF 4 0.225-0.414 ZnS 6 0.414-0.732 NaCl 8 0.732-1.00 CsCl Remark The formula works for essentially ionic compound.e.g. check ZnS Lattice Energy (U) Cont… Is defined as the energy released (exothermic process) in the formation of a mole of MX (crystal) from gaseous ions separated from each other by infinite distances. M(g)+a + X(g)-b → MbXa (s) Example: Na+(g) + Cl-(g) → NaCl(s) ΔU = -782 KJ/mol Lattice energy cannot be measured directly, but calculated (or theoretical) values from born Lande equation and experimental values obtained from thermodynamic data can be used. Born Lande Equation Born Lande equation assumes the ions as spherical and are not distorted by neighboring ions. He considered mainly simple electrostatic treatment of ions which accounts for about system of atoms in going from one stage to another is independent of the path followed. The formation of one mole of crystalline ionic solid can occur by two alternative processes.  Cont… Path I: The formation of NaCl lattice: The heat of formation is mostly negative for ionic compound formation. Na(s)+1/2Cl2(g) → NaCl(s) ΔHf= -411KJ/mol Path II: The second path has five steps: 1.Sublimation of Na(s) → Na(g) ΔHsub=109KJ/mol 2.Ionization of gaseous sodium atoms: Na(g) → Na(g)+ + e- I.E.=+496KJ/mol 3.Dissociation of gaseous chlorine: 1/2Cl2 (g) → Cl(g) 1/2ΔHdiss= 121KJ/mol 4.Conversion of gaseous chlorine to ions: Cl(g) + e →Cl(g)- E.A. = -356KJ/mol 5.Combination of gaseous ions (lattice energy): Na(g) +Cl(g) → NaCl(s) + - ΔU= ? The two processes are independent of each other and applying Hess’s law of heat summation, Path I = Path II (sum of the five steps) ΔHf(NaCl = ΔHsub + 1/2ΔHDiss + I.E + E.A + ΔU -411 KJ/mol = 109 KJ/mol + 496 KJ/mol + 121 KJ/mol – 356 KJ/mol - ΔU ΔU = -781 KJ/mol Direct measurement of electron affinity is difficult and Born-Haber cycle is  Cotn…. Example: Calculate electron affinity of iodide, based on the following data: ΔHf(NaI) = -270KJ/mol; ΔHsub = 108 KJ/mol 1/2ΔHDiss = 106KJ/mol I. E = 496 KJ/mol E.A = ? ΔU = -691 KJ/mol -ΔHf(NaCl) + ΔHsub + ½ ΔHDiss + I.E + ΔU = E.A E.A = -289 KJ/mol The Born-Haber cycle treatment also permits analysis of the stability of ionic compounds. For example: MgO has negative heat of formation because the lattice energy is highly exothermic. The lattice energy calculated (from Born-Lande equation) for different compounds assuming complete ionic character, show deviations from the values calculated from the thermodynamic data based on experiments (using Born-Haber cycle). Close agreement indicates that the assumption that the bonding is ionic is true. The larger the discrepancy, the compound has some covalency.  Cont… Table 2.4: Lattice energies of some alkali halides Crystal Experimental L.E (-U KJ/mol) Calculated L.E (-U KJ/mol NaF 907.9 907.9 NaCl 769.9 762.5 NaBr 736.4 723.8 NaI 690.4 677.8 KF 803.3 803.3 KCl 702.9 694.5 KBr 674.2 667.3 KI 674.6 665.7 Exercise 3 Based on the following data, calculate the lattice energy of MgCl2. Mg(s) → Mg(g) ΔHo = 147.7 KJ Mg(g) → Mg+(g) ΔHo = 737.7 KJ Mg+(g) → Mg2+(g) ΔHo = 1450.6 KJ Cl2(g) → 2Cl(g) ΔHo = 243.4 KJ Cl(g) + e- → Cl- ΔHo = -348.7 KJ  6.5.2 Covalent Bond Strength H = 242 kJ/mol The strength of a bond is measured by determining how much energy is required to break the bond. This is the bond enthalpy. The bond enthalpy for a Cl—Cl bond, D(Cl—Cl), is 242 kJ/mol. 47  Average Bond Enthalpies Average bond enthalpies are positive, because bond breaking is an endothermic process. 48  Cont… NOTE: These are average bond enthalpies, not absolute bond enthalpies; the C—H bonds in methane, CH4, will be a bit different than the C—H bond in chloroform, CHCl3. 49  Cont… Can use bond enthalpies to estimate H for a reaction Hrxn = (bond enthalpies of bonds broken) (bond enthalpies of bonds formed) This is a fundamental idea in chemical reactions. The heat of a reaction comes from breaking bonds and remaking bonds. 50  Cont… CH4(g) + Cl2(g) CH3Cl(g) HCl(g) In this example, one C—H bond and one Cl—Cl bond are broken; one C—Cl and one H—Cl bond are formed. 51  Cont… CH4(g) + Cl2(g) CH3Cl(g) HCl(g) So, Hrxn = [D(C—H) + D(Cl—Cl) [D(C—Cl) + D(H—Cl) = [(413 kJ) + (242 kJ)] [(328 kJ) + (431 kJ)] = (655 kJ) (759 kJ) = 104 kJ  Enthalpy problem: Calculate the enthalpy of reaction for: CH4 + 3/2O2 ---> CO2 + 2H2O 4(C--H) + 3/2(O==O) - 2(C==O) - 4(OH) 4(413) + 3/2(495) - 2(800) - 4(463) = -563 kJ HC CH + 5/2O2 ------> 2CO2 + H2O 1(CC) + 2CH + 5/2(O=O) - 4(C==O) - 2(OH) 1(834) + 2(413) +5/2(495) - 4(800) - 2(463) = -1229 kJ 53  Quiz Draw the Lewis structure (include resonance and formal charges) for: HOFO NCO - 54  Table 2.6: Comparison of Electrovalent and Covalent Compounds Electrovalent or Ionic Compounds Covalent compound This type of bond is frequently This type of bond is frequently found encountered in inorganic compounds in organic compounds. Electrovalent compounds are soluble Covalent compounds are soluble in in water but insoluble in organic organic solvents and insoluble in solvents water They possess high M.P and B.P They have low M.P. and B.P. Electrovalent compounds are not They are inflammable inflammable They do not possess any They usually posses a characteristic characteristic smell smell. In electrovalent compounds In covalent compounds reactions are reactions slow are rapid  6.6. Molecular structure and polarity The molecule’s structural name is always based on the number of the atoms in the molecule. Basically, you pretend that the lone pairs of electrons are not there, and then you name the molecule by only considering the arrangement of the atoms. It was a simpler way of naming the structures. 2 atoms Linear 3 atoms Bent and Angular 4 atoms Trigonal planar 5 atoms Tetrahedral  6.6.1Molecular Shapes We’ve learned to draw Lewis structures and account for all the valence electrons in a molecule. But: Lewis structures are two dimensional and molecules are 3 dimensional objects. The 3D structure is absolutely critical for understanding molecules. geometry & shape of molecule critical we can easily predict the 3D structure of a molecule just by adding up: bound atoms + lone pairs 57 58  6.6.1.1 What Determines the Shape of a Molecule? atoms and lone pairs take up space and prefer to be as far from each other as possible lone pair shape can be predicted from simple bonds geometry “Things” The central atom has four “things” around it. A “thing” is an atom or a lone pair of electrons. # things = atoms plus lone pairs 59  6.6.2 Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of things is the one that minimizes the repulsions among them.” 60 number of things arrangement geometry bond angles These are the geometries for two through six things around a central atom. You must learn these! 61  Molecular Geometries The geometry is often not the shape of the molecule, however. The “shape” is defined by the positions of only the atoms in the molecules, not the lone pairs. 62  Geometries vs. shape Within each geometry, there might be more than one shape. 63  6.6.3 Polarity of molecule In Chapter 8 we discussed bond dipoles. polar bonds versus polar molecules. We must think about the molecule as a whole. By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule. 64 Polarity “The tractor pull ”