T2 Chemical Bonding Part 2 PDF

T2- Chemical Bonding_Part2 PART 2 – COVALENT BONDING Learning objectives Covalent bond Describe the concept of chemical bonding as the electrostatic attraction between protons and electrons. Describe covalent bonding as sharing of electrons between atoms using Lewis’ diagram...

T2- Chemical Bonding_Part2 PART 2 – COVALENT BONDING Learning objectives Covalent bond Describe the concept of chemical bonding as the electrostatic attraction between protons and electrons. Describe covalent bonding as sharing of electrons between atoms using Lewis’ diagrams. Name covalent bonds as single, double, or triple covalent bonds when two, four, or six electrons are, respectively, shared. Explain that bond length decreases and bond strength increases as the number of shared electrons increases. Distinguish between lone pairs and bonding pairs. Dative or co-ordinate covalent bond Illustrate and define dative bond formation in molecules such as H3O+, NH4+. Bond polarity and electronegativity Define electronegativity as the measure of the force of attraction of an atom towards the electrons engaged in a bond. Define non-polar and polar covalent bond. Shapes of molecules – VSEPR theory Use VSEPR theory to predict shape of simple molecules and ions Construct the following molecules using molecular models (e.g.Molymod®): H 2O, CH4, NH3, CO2, C2H4, HCHO. Construct molecular models of H3O+, NH4+. Polarity of molecules Predict molecule polarity from bond polarity and molecular geometry. Lewis structures Apply Lewis’ model to represent molecules and polyatomic ions. Draw resonance structures where more than one structure can be drawn. Calculate formal charge of simple polyatomic molecules where more structures can be drawn. Elucidate that the most likely structure(s) is(are) the one(s) which minimise formal charges, and where negative charge(s) is(are) on most electronegative atom(s). Valence bond theory Describe covalent bonding in terms of atomic orbital overlap. Explain sigma and pi bonding orbitals as head-to-head or side-to-side combination of atomic orbitals. Explain the concept of hybridization (sp, sp2, sp3 only) by examining the shapes of simple molecules. Describe the structure of benzene as cyclic, planar, and with an uninterrupted ring of electrons (delocalised electrons), with the help of two resonance forms. Predict the shapes of, and bond angles in, molecules and ions such as, or similar to: BF3 (trigonal planar), CO2 (linear), CH4 (tetrahedral), NH3 (pyramidal), H2O (bent), SF6 (octahedral), PF5 (trigonal bipyramidal) and define dative bond formation in molecules such as H3O+, NH4+. Construct molecular models of H3O+, NH4+. 1/43  T2- Chemical Bonding_Part2 1. Formation of covalent bonds When two non-metal atoms react together, both need to gain electrons to achieve stable outer shells. Hence an exchange (simultaneous gain and loss) is not possible in this case. In covalent bonds, atoms gain stable outer shells by sharing electrons with each other. Example 1: chlorine Cl2 Isolated chlorine atoms Chlorine molecule Explanation: The isolated Cl atoms contain 7 outer electrons and are hence not stable. By sharing their unpaired electrons, the electron pair forming the bond (=bonding electrons) belong now to both atoms connected by this bond, i.e. each Cl atom has now 8 electrons making the atoms stable (octet rule) within the structure formed. The formation of a covalent bond stabilizes the atoms and energy is released as the bond forms. The forces of attraction between the nuclei of the 2 atoms and the shared electrons are in balance with the forces of repulsion between the 2 nuclei. Example 2: formation of hydrogen H2 This figure shows for the formation of hydrogen H2 that the energy between 2 H atoms is at its lowest and hence at maximum stability, when the distance between 2 H atoms corresponds to the bond length in an H2 molecule. 2/43  T2- Chemical Bonding_Part2 2. Recall – simple Lewis structures Lewis symbols In Chemistry, valence electrons are very important as they are responsible for the chemical behaviour of the element, whereas inner electrons don’t intervene in chemicals reactions. Symbols of atoms showing only the outer electrons of atoms are called LEWIS SYMBOLS. The first four valence electrons are represented by a cross (or bullet) and are called unpaired or single electrons. The remaining electrons will form an electron pair with the single electrons to create a pair represented by a line. Lewis symbols of the 20 first chemical elements Lewis structures Lewis structures are used to illustrate the covalent bonds in molecules. Simple Lewis structures are set up by connecting unpaired electrons of Lewis symbols of the atoms in a molecule so that bonding pairs are formed. Example: water Structural formulae vs. Lewis structures Lewis structures are structured formulae that also show the non-bonding electron pairs also called lone electron pairs. Example: Water H2O Structured formula Lewis structure 3/43  T2- Chemical Bonding_Part2 More examples: (a) NOF (b) C2H3Br (c) CSCl2 Remember An atom can have unpaired electrons (single electrons) or electron pairs in its valence shell. The electron pair forming the bond is called bonding pair. The electron pairs that do not take part in the covalent bonds, are called non- bonding or lone pairs Covalent bonds can be single bonds (2 shared electrons) or double bonds (4 shared electrons) or triple bonds (4 shared electrons). Example: In a DOUBLE bond, 4 electrons are shared (2 bonding pairs) In a TRIPLE bond, 4 electrons are shared (2 bonding pairs) 4/43  T2- Chemical Bonding_Part2 3. Bond length and bond strength Remember Bond length: a measure of the distance between the 2 bonded nuclei. Bond strength: (also called bond enthalpy) a measure of the energy required to break a bond Example: As we go down the group of the halogens, the bond length increases at the atomic radius becomes bigger. At the same time, the shared electron pair is further from the pull of the nuclei in the larger molecules and the bond becomes weaker, the bond strength hence decreases down the group What about the strength of multiple covalent bonds? In multiple covalent bonds, the number of shared electrons and so the electron density and the related electronic charge is higher than in simple covalent bonds, resulting in stronger electrostatic attraction between the bonding electrons and the nuclei of the bonded atoms. ➔ The MORE ELECTRONS ARE SHARED, the STRONGER THE BOND, the closer the nuclei of the bonded atoms are, resulting in SHORTER BONDS 5/43  T2- Chemical Bonding_Part2 The relationship between bond type (single/double/triple), bond length and bond strength is shown in the above table. General trend: Check your knowledge https://quizizz.com/admin/quiz/60b98a075175ac001d98c44d/bond-length-strength-energy?source=HeroSearchBar&page=FeaturedPage&searchLocale=&fromSearch=true 6/43  T2- Chemical Bonding_Part2 4. Dative bonds or coordinate bonds Remember A coordinate bond or dative bond is a covalent bond in which BOTH bonding electrons are provided by ONE of the atoms. WHEN does a coordinate bond form? When one atom has a lone pair of electrons and is able to donate it to another atom has a free orbital so that it can accept the lone pair donated. HOW is a coordinate bond represented ? An arrow on the head of the bond is used to show a coordinate bond, with the direction indicating the origin of the electrons. Examples Practice Check your knowledge https://quizizz.com/admin/quiz/5ebd5448591c53001b197587/dative-bond Answer: C 7/43  T2- Chemical Bonding_Part2 5. Recall – naming of covalent substances The naming of molecules uses a prefix system: it adds numerical prefixes to identify the number of atoms in the molecule. The prefix ‘mono’ is left away if the first atom in the name is only a single atom. Examples: CO2 - carbon dioxide CO - carbon monoxide N2O3 - dinitrogen trioxide SO3 - sulfur trioxide PCl5 - phosphorus pentachloride I2O3 - Diiodine trioxide Remark Names of ionic substances never contain prefixes! CaCl2 – calcium chloride AlCl3 – aluminium chloride CuCl2 – copper(II) chloride 8/43  T2- Chemical Bonding_Part2 6. Bond polarity and electronegativity COVALENT or IONIC BOND? Rather than saying that ionic and covalent are two distinct types of bonding, it is more accurate to say that they are at the two extremes of a same scale. Let’s consider a covalent bonding, where electrons are shared between 2 atoms. Both atoms have an electronegativity which represents the ability to attract the bonding electrons when bot h atoms make a bond. Depending on the difference in electronegativity (ΔEN) between the bonding atoms, different cases are considered. Less polar bonds have more covalent character. More polar bonds have more ionic character. The more electronegative atom attracts the electrons in the bond enough to ionize the other atom. Electronegativity trends according to PAULING scale 9/43  T2- Chemical Bonding_Part2 The following Table shows the different types of bondings formed when 2 atoms bond together. Identical atoms, ΔEN=0 Example: H2 - Equal sharing of electrons: bonding electrons are distributed evenly. - Equal overlapping of orbitals - No charges on atoms - Name of bond: NON-POLAR COVALENT BOND Different atoms, ΔEN=low Example: HCl - Uneven sharing of electrons: bonding electrons are closer to the more electronegative atom. Increasing polarization - the bigger ΔEN, the more uneven the electron distribution - Partial overlapping of orbitals - Partial charges on atoms - Name of bond: POLAR COVALENT BOND Different atoms, ΔEN=high Example NaCl - Transfer of electrons to the more electronegative atom - No overlapping of orbitals - Full charges on the atoms, - Name of bond: IONIC BOND Bond polarity trends in hydrogen halides: 10/43  T2- Chemical Bonding_Part2 General relationship between electronegativity difference and bond type: Practice Remark: 1. 1.The C-H bond, omnipresent in organic molecules, is considered to be largely NON-POLAR, although there is in fact a ΔEN=0.4 (Carbon has an electronegativity of 2.5, while the value for hydrogen is 2.1. The very low polarity of the C-H bond is an important factor in determining the properties of organic molecules. 2. Don’t mix up BOND polarity and MOLECULE polarity 11/43  T2- Chemical Bonding_Part2 7. COMPLEX Lewis structures Before explaining how to set up more complex LEWIS STRUCTURES, a new concept has to be introduced: FORMAL CHARGES in molecules or ions Remember In molecules, a formal charge (FC) is the hypothetical charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. How to check for formal charges? Remember Formal charge (FC) is the difference between the number of valence electrons (VE) of an atom in isolated state and the number of electrons (bonding (B) and non-bonding (NB) )assigned to that atom in a Lewis structure. FC = VE – NB – B VE - number of valence electrons (VE) the neutral atom NB - number of non-bonding/lone electrons (NB) on the atom in the molecule. (Each lone pair counts as 2, and each unpaired electron counts as 1.) B - number of bonding electrons (B) to the atom in the molecule. This is the number of bonding electron pairs divided by 2. Examples Calculate the formal charge of the central atom of all the following molecules: Answer 12/43  T2- Chemical Bonding_Part2 Setting up complex LEWIS STRUCTURES RULE 1 Symmetry in the atom arrangement RULE 2 Number of electrons in a Lewis structure (bonding and non-bonding/lone electrons) is the number of valence electrons plus the charge of the ion B + NB = VE - charge RULE 3 Octet rule must be met for all atoms (with higher atomic number than Carbon): 8 valence electrons for every atom RULE 4 Option with lowest amount of formal charges is valid (in case of more options) (this rule outweighs rule 3 so that exceptions to octet rule are possible!!) RULE 5 Charges are shown outside of square brackets in final Lewis structures Example 1: ammonium ion NH4+ RULE 1 There is only one possible atom arrangement: RULE 2 VE(N) 5 VE(H) 4x1= 4 Charge -(+1) 8 electrons → 4 electron pairs (B and NB) RULE 3 Octet rule is fulfilled for all the atoms. RULE 4: does not apply here RULE 5: Formal charges for central nitrogen atom (hydrogen has no formal charge) FC = VE – NB - B VE(N)=5 / NB(N)=0 / B(N)=4 ➔ FC=+1 Structure with formal charge: Final LEWIS structure shows the charge outside of square brackets: 13/43  T2- Chemical Bonding_Part2 Example 2: carbonate ion CO32- 3 Options for atom arrangement: RULE 1 The following option is preferred: RULE 2 VE(C) 4 VE(O) 3x6=18 Charge -(-2) 24 electrons → 12 electron pairs (B and NB) RULE 3 For each atom, the octet rule is met. BUT we have now: B=6 and NB=20 Resulting in 13 electron pairs instead of 12 electron pairs (see above) The following structure fullfills the rules 1, 2 and 3 Bonding electrons: B=8 Non-bonding/lone electrons: NB=16 RULE 4 does not apply as there is only 1 option RULE 5 Formal charges: Final LEWIS structure shows the charge outside of square brackets: 14/43  T2- Chemical Bonding_Part2 Expert question*** Answer 15/43  T2- Chemical Bonding_Part2 8. Resonance structures (sb p.164-167) Let’s consider again the structure of a carbonate polyatomic ion. When setting up the final structure, the question arises which of the 3 oxygen atoms will finally form the double bond with the central carbon atom. Three different possibilities result from this question. In fact, all 3 Lewis structures are possible and co-exist. These multiple Lewis structures for one chemical formula are called resonance structures. These different resonance structures are possible as there are some electrons in the molecule that are not associated with a single atom or one covalent bond: they are delocalized electrons. The true structure is an intermediate form between all the resonance structures and is called resonance hybrid. Resonance structures and resonance hybrid are represented in the following way. Resonance structures for a carbonate ion: Double arrow between every resonance structure Resonance hybrid for a carbonate ion: Dotted lines are used to show the position of the delocalised electrons 16/43  T2- Chemical Bonding_Part2 SUMMARY Context Some species don’t seem to fit with a single Lewis structure, but several Lewis structures seem possible. What does it mean when a molecule is said to have a resonance structure? Resonance structures mean that there are several possible Lewis Structures, but the true structure is an intermediate form, known as the resonance hybrid. When do resonance structures occur? when there is more than one possible position for a double bond in a molecule. The arrangement of atoms does not change, but the arrangement of electrons does. In fact, electrons from a double bond are capable of moving from one part of a molecule to another What are resonance structures due to? The delocalization of electrons can explain the structures. Delocalized electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or one covalent bond What evidence is there that there is not one sinle Lewis structure? Example: polyatomic ion nitrate NO3- However, studies of the electron density and bond length in the nitrate ion indicate all the bonds are equal in length and the electron density is spread evenly between the three oxygen atoms. The study shows that the bond length is intermediate between a single and a double bond. The actual structure is something in between the resonance structures and is known as a resonance hybrid. What Is new in drawing resonance structures? Several Lewis structures are drawn and the true, intermediate structure: the resonance hybrid. Example: polyatomic ion nitrate NO3- 17/43  T2- Chemical Bonding_Part2 Three possible resonance structures are possible, consisting of one double and two single bonds: The resonance hybrid (intermediate, true structure) is the following. Dotted lines are used to show the position of the delocalised electrons Remark: the most likely structure is the one which minimises formal charges, and where negative charge(s) is(are) on most electronegative atom(s) 18/43  T2- Chemical Bonding_Part2 More examples: SPECIES Resonance structures Resonance hybrid Carbonate ion Nitrite ion Benzene Ozone Methanoate ion 19/43  T2- Chemical Bonding_Part2 Exceptions to the OCTET RULE The Octet Rule is violated in three scenarios: (a) INCOMPLETE OCTET: When there are too few valence electrons (b) EXPANDED OCTET: When there are too many valence electrons (c) When there is an odd number of valence electrons Reminder: Always use the Octet Rule when drawing Lewis Structures, these exceptions will only occur when exceptionally and this will be announced in a test. Examples (a) INCOMPLETE OCTET: One example for an BH3 B has an incomplete octet; it only has six electrons around it. Hydrogen atoms can naturally only have only 2 electrons in their outermost shell, and as such there are no spare electrons to form a double bond with boron. (b) EXPANDED OCTET More common than incomplete octets are expanded octets where the central atom in a Lewis structure has more than eight electrons in its valence shell. In expanded octets, the central atom can have ten electrons, or even twelve. Molecules with expanded octets involve highly electronegative terminal atoms, and a non-metal central atom found in the third period or below, which those terminal atoms bond to. 20/43  T2- Chemical Bonding_Part2 For example, phosphorus pentachloride PCl5 is a legitimate compound: P has an expanded octet: it has 10 electrons around it Another example is the sulfate ion SO42- S has an expanded octet: it has 12 electrons around it (c) These structures contain one unpaired electron. There are very few stable molecules with odd numbers of electrons that exist, since that unpaired electron is willing to react with other unpaired electrons. This scenario is not relevant for S6 curriculum. Check your knowledge https://quizizz.com/admin/quiz/58d2d7116ea9076e03259c2e/lewis- structures?source=MainHeader&page=QuizPage&searchLocale=&fromSearch=true https://quizizz.com/admin/quiz/5fa222e7235a92001b033674/resonance- structures?fromSearch=true&source= https://quizizz.com/admin/quiz/6020117d3b4640001b1c12f9/exceptions-to-the-octet-rule 21/43  T2- Chemical Bonding_Part2 9. Shapes of molecules – VSEPR theory Valence Shell Electron-Pair Repulsion THEORY (VSEPR) This MODEL is used to predict the SHAPE of individual molecules or ions from the number of electron pairs surrounding their central atoms. The theory is based on the following considerations: ▪ The shape of an ion or a molecule is defined by the repulsion of valence electron pairs of the central atom (bonding or lone pairs) ▪ Repulsion between lone electron pairs is greater than between bonding electron pairs ▪ Valence electron pairs arrange to positions of maximum separation. This produces different shapes for different number of electron pairs. STEPS TO FOLLOW to determine the shapes of an ion/molecule: Step 1: Draw the LEWIS STRUCTURE for the molecule or ion Step 2: Count the NUMBER OF ELECTRON DOMAINS of the central atom. o Multiple bonds (double or triple bonds) count as 1 electron domain o Lone electron pairs count as 1 electron domain o (For molecules or ions that have resonance structures, you may use any one of the resonance structures.) Step 3: Determine the ELECTRON DOMAIN GEOMETRY of the ion/molecule according to the following rules: o 2 electron pairs → LINEAR SHAPE o 3 electron pairs → TRIGONAL PLANAR SHAPE o 4 electron pairs → TETRAHEDRAL SHAPE o 5 electron pairs → TRIGONAL BIPYRAMIDAL SHAPE o 6 electron pairs → OCTAHEDRAL SHAPE Step 4: Determine the SHAPE of the ion/molecule considering the effect of lone electron pairs: o biggest angles between lone-pair/lone-pair bonds o smallest angles between bonding-pair/bonding-pair bonds REMARK: Step 4 only applies if the ion/molecule contains lone electron pairs! Otherwise the SHAPE is equivalent to the ELECTRON DOMAIN GEOMETRY! 22/43  T2- Chemical Bonding_Part2 MOLECULAR SHAPES - OVERVIEW TABLE Number of Geometry Number of Number of electron Around central BONDING LONE SHAPE Examples domains atom electron electron domains domains 2 linear 2 0 linear BeCl2, 180° (AX2) CO2 trigonal trigonal – BF3, NO3 3 planar 3 0 planar (AX3) 120° 2 1 Non- SnCL2 linear/(bent) (AX2E) 4 tetrahedral 4 0 tetrahedral CH4, NH4+ (AX4) (109,5°) 109.5° 3 1 pyramidal NH3 (AX3E) (107°) 2 2 Non- H 2O linear/(bent) (104,5°) (AX2E2) trigonal trigonal 5 bipyramidal 5 0 bipyramidal PF5 120°(in plane) (AX5) 90° (above & 4 1 Seesaw* SF4 below) (AX4E) 3 2 T-shaped* ClF3 (AX3E2) 2 3 Linear* XeF2 (AX2E3) octahedral Octahedral* 6 90° 6 0 (AX6) SF6(g) 5 1 Square BrF5 Pyramidal* (AX5E) 4 2 Square XeF4 planar* (AX4E2) 3 3 T-shaped* (AX3E3) 2 4 Linear* (AX2E4) * you don’t need to memorize these geometries 23/43  T2- Chemical Bonding_Part2 10. Polarity of molecules In a simple molecule like HCl, there is a is a polar bond, which results in HCl being a POLAR MOLECULE (also called sometimes a DIPOLE). Polar molecules have got a DIPOLE MOMENT: it is a measure of the polarity of a chemical bond or a molecule. The dipole arrow always points to the more electronegative atom (in other words from the POSITIVE to the NEGATIVE pole) What about more complicated molecules? In CCl4 Cl δ- 4δ+ δ- Clδ- C Cl δ- Cl All 4 bonds are polar covalent bonds, but the molecule as a whole is non-polar! Why? Due to the tetrahedral shape of the molecule, the individual dipole moments cancel each other out. 24/43  T2- Chemical Bonding_Part2 What about CH3Cl? There are 4 polar covalent bonds in this molecule that do NOT cancel each other out causing CH3Cl to be a polar molecule: the overall dipole moment is towards the electronegative chlorine atom. A molecule is polar if the following conditions are both fulfilled: 1. Molecule must contain one or more polar bonds 2. The individual dipole moments of the bonds do not cancel each other out by symmetry. Steps to follow: Step 1 Polar bonds? Check difference in electronegativity between the bonded atoms Step 2 (only if step 1 shows polar bonds!) Determine the shape of the molecule Step 3 Check if individual dipole moments cancel out or not due to the shape of the molecule PRACTICE – fill out the table below 25/43  T2- Chemical Bonding_Part2 Molecule Lewis structure Displayed formula # of # of Electron domain SHAPE Bond Dipole bonding lone geometry (name & sketch) angle YES/NO? pairs pairs BeH2 2 0 linear linear nonpolar BH3 3 0 trigonal planar trigonal planar nonpolar SO2 2 1 trigonal planar bent polar CH4 4 0 tetrahedral tetrahedral nonpolar NH3 3 1 tetrahedral trigonal polar pyramidal H 2O 2 2 tetrahedral bent or V- polar shaped PCl5 5 0 trigonal trigonal nonpolar bipyramidal bipyramidal SF6 6 0 octahedral octahedral nonpolar 26/43  T2- Chemical Bonding_Part2 11. Valence bond theory (see also studentbook p. 185 / 198-205) The VSEPR theory predicts the shape of molecules and ions, but does not give any details about orbitals present in molecules. Likewise, no information is provided on the nature of covalent bonds. The valence bond theory gives answers to these questions. Two types of covalent bonds - Sigma bonds (σ) - Pi bonds (π) A. SIGMA BONDS The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond. The electrostatic attraction between the electrons and nuclei bonds the atoms to each other. 27/43  T2- Chemical Bonding_Part2 B. Pi BONDS Remark: Pi bonds are weaker than sigma bonds!! Explanation: lower electrostatic attraction between the bonding electrons and the nuclei of the atom: In pi bonds, the electron density in Pi bonds is further away (above and below the line connecting the nuclei) from the positive charge of the nuclei and so they break more easily. In sigma bonds, the electron density is concentrated and close (between the nuclei) to the positive charge of the nuclei and so they break less easily. Consequence: organic molecules with carbon-carbon double bonds are MORE reactive than organic molecules with only carbon-carbon single bonds. 28/43  T2- Chemical Bonding_Part2 29/43  T2- Chemical Bonding_Part2 Q11 B Q5 B Answers: 30/43  T2- Chemical Bonding_Part2 Hybrid orbitals sp3 hybridisation Explanation? Bond hybridization – a two-step process Step 1: Promotion from ground state to excited state: s electrons occupy empty p orbitals Step 2: Hybridization: mixing atomic orbitals to form new hybrid atomic orbitals 31/43  T2- Chemical Bonding_Part2 The 2s and 2p subshells blend together and form four new hybrid orbitals (called sp3 orbitals, after the merger of one s and 3p orbitals) This would give four unpaired electrons of equal energy, capable of forming four identical covalent bonds. More information: sp3 hybridization takes place in: Organic molecules with only single carbon-carbon bonds Molecules with 4 electron domains and a TETRAHERDRAL molecular geometry Examples: CH4 H2O NH3 32/43  T2- Chemical Bonding_Part2 sp2 hybridization More information: Example: ethene C2H4 sp2 hybridization takes place in: Organic molecules with carbon forming carbon-carbon double bonds Molecules with 3 electron domains and a PLANAR molecular geometry Examples: C 2H4 CH2O CO32- 33/43  T2- Chemical Bonding_Part2 sp hybridization sp hybridization takes place in: Organic molecules with carbon forming carbon-carbon triple bonds Molecules with 2 electron domains and with a LINEAR molecular geometry Examples: C 2H2 HCN Remark The overlap of a hybrid orbital with any other atomic orbital (s, p, or another hybrid orbital) always forms a sigma bond! Q15 A Answer: 34/43  T2- Chemical Bonding_Part2 Identifying hybridization 35/43  T2- Chemical Bonding_Part2 Table: LINK between number of electron domains shape of molecules hybridization 36/43  T2- Chemical Bonding_Part2 Revision – QUICKCHECK Quickcheck 1 Molecule to consider: HCONH2 (a) Shortest bond (b) Strongest bond (c) How many Pi bonds? (d) How many Sigma bonds? (e) Dative bond? (f) Bond with sideways overlap? (g) How many bonds with end-on overlap? (h) Between ___ bond and ___ bond, there is a bond angle of 120° (i) Between ___ bond and ___ bond, there is a bond angle of 109.5° (j) How many bonding e- ? (k) How many lone pairs of e- ? (l) How many valence electrons ? (m) Strongest polar bond (n) Dipole, yes or no ? 37/43  T2- Chemical Bonding_Part2 Quickcheck 2 ozone (O3) (a) How many valence electrons? (b) How many bonding pairs and how many lone pairs? (c) Draw 1 resonance structure for ozone. (d) Draw the hybrid structure for ozone. Bond length, bond strength, bond polarity 38/43  T2- Chemical Bonding_Part2 Lewis structures Molecular shape, bond angles, electron domains 39/43  T2- Chemical Bonding_Part2 Polarity of molecules Lewis structures and Bond length 40/43  T2- Chemical Bonding_Part2 Valence bond theory 41/43  T2- Chemical Bonding_Part2 COMPARISON VSEPR VALENCE BOND THEORY 42/43  T2- Chemical Bonding_Part2 43/43