DOCSLIB.ORG

2 the Other Sp Orbitals Hold the Oxygen Lone Pairs

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
2 the Other Sp Orbitals Hold the Oxygen Lone Pairs Worksheet 14 - Hybridization The hybridized molecular orbitals have different shapes and energy levels than the atomic orbitals. The number of molecular orbitals created by hybridization depends on the number of atomic orbitals that are mixed to form them. In forming sp3 hybridized orbitals, four atomic orbitals are mixed, one s and three p. The energy diagram for this process is shown below. The hybridized orbitals are higher in energy than the s orbital, but lower in energy than the p orbitals. p energy sp3 s hybridization atomic orbitals hybridized orbitals Carbon has 4 valence electrons. Add these electrons to the atomic and molecular orbitals. This hybridization gives tetrahedral geometry. With this hybridization, C will form four equivalent σ bonds. Draw a similar energy diagram for sp3 hybridized oxygen. p energy sp3 s hybridization How many σ bonds will be formed? 2 How are the other sp3 orbitals used? The other sp3 orbitals hold the oxygen lone pairs Do the same for sp3 hybridized nitrogen. p energy sp3 s hybridization In some Lewis structures, there are only three equivalent bonds formed. To create three equivalent hybridized orbitals, mix three atomic orbitals. Draw and name the orbitals formed in this hybridization, then add the electrons for sulfur. Since the hybridized orbitals are close in energy, every orbital is filled with one electron before electrons are paired. p p energy sp2 s hybridization The hybridized orbitals will form __1___ σ bond(s). The unhybridized orbital will form __1___ π bond(s). There will be __2___ lone pair(s). This hybridization gives trigonal planar geometry. In linear molecules, like CO2, the central atom has only two equivalent bonding orbitals. Draw the energy levels and name the orbitals formed in this hybridization. p p energy sp s hybridization Fill in the electrons for carbon and determine the number and type of bonds formed. There will be 2 sigma bonds (with sp) and 2 pi bonds (with p) In CO2, determine the hybridization of the oxygen atoms. Complete the energy diagram for the oxygens. Draw the structure of CO2. p p energy sp2 : : s hybridization O: C O: Each oxygen atom forms 1 sigma bond (with the sp2) and 1 pi bond (with the p) with the sp and p orbitals on the carbon atom In atoms with n=3 or larger, the d orbitals can also be hybridized. In molecules with five molecular orbitals, five atomic orbitals are mixed: _d_ + _s_ + _p_ + _p_ + _p_ This will give trigonal bipyramidal geometry and is called dsp3 hybridization. Finally, molecules with octahedral geometry, will have __6 __ molecular orbitals. This hybridization is called __d2sp3______ . Shown below is a portion of the chart from Worksheet 13. Fill in the hybridization for each of the compounds. compound bonds lone geometry shape hybridization pairs bonds + lone pairs = 6 2 3 SF6 6 0 octahedral octahedral d sp bonds + lone pairs = 4 3 NH3 3 1 tetrahedral trigonal sp pyramidal bonds + lone pairs = 6 - 2 3 ICl4 4 2 octahedral square d sp planar bonds + lone pairs = 4 3 CF4 4 0 tetrahedral tetrahedral sp bonds + lone pairs = 3 2 SO3 3 0 trigonal trigonal sp planar planar bonds + lone pairs = 5 3 SF4 4 1 trigonal seesaw dsp bipyramidal bonds + lone pairs = 2 CO2 2 0 linear linear sp bonds + lone pairs = 4 3 H2O 2 2 tetrahedral V-shaped sp bonds + lone pairs = 3 - 2 NO2 2 1 trigonal V-shaped sp planar Fill in the chart below and then complete the Lewis structures for the molecules shown below and fill in those charts. element Lewis symbol # bonds # lone pairs C 4 0 C N 3 1 N H 1 0 H O 2 2 O Halogen F 1 3 1 H O σ bonds __7__ atom # bond angle hybridization 1 120 sp2 H C C OH π bonds __ 1__ 3 2 2 109 sp3 2 H 3 120 sp acetic acid H H 1 atom # bond angle hybridization C C C N σ bonds _6___ 1 180 sp H 3 2 2 180 sp π bonds _3___ 3 120 sp2 acrylonitrile H 2 O H H CH σ bonds 25__ atom # bond angle hybridization 2 H C C N 1 120 sp N C π bonds 4__ 2 120 sp2 H C H 3 C C 1 3 109 sp O N N H CH 3 H caffeine The molecule shown to the left is riboflavin H (vitamin B2). Answer the following questions HHO C about its structure. a) how many carbons are sp3 hybridized? 8 H C O H H O C H sp2 hybridized? 10 H C O H sp hybridized? 0 H C H H H HHC b) How many nitrogens are sp3 hybridized? 2 H C CNN O C C C C sp2 hybridized? 2 H H C C C N C C NCH sp hybridized? 0 H H H O c) How many oxygens are sp3 hybridized? 4 sp2 hybridized? 2 sp hybridized? 0 d) How many σ bonds are there in total? 52 e) How many π bonds are there in total? 7 f) Which of the three rings are planar? The ring on the left, in which all carbon atoms are sp2 hybridized - The acetate ion, C2H3O2 , has both oxygens bonded to the same carbon. a) Draw the Lewis structure and all resonance forms. H sp2 H sp3 O C C H O C C H sp3 O H O H sp2 b) Label the hybridization around each carbon. c) Pick one resonance structure and label the hybridization of each oxygen. d) How many σ and π bonds are present? 6 sigma bonds, 1 pi bond e) Which atom carries the formal negative charge? the oxygen with 3 lone pairs .
Load more

Recommended publications
  • No Rabbit Ears on Water. the Structure of the Water Molecule
    No Rabbit Ears on Water The Structure of the Water Molecule: What Should We Tell the Students? Michael Laing University of Natal. Durban, South Africa In the vast majority of high school (I)and freshman uni- He also considers the possibility of s character in the bond- versity (2) chemistry textbooks, water is presented as a bent ing orbitals of OF2 and OH2 to improve the "strength" of the molecule whose bonding can be described bv the Gilles~ie- orbital (p 111) from 1.732 for purep to 2.00 for the ideal sp3. Nyholm approach with the oxygen atomsp3 Gyhridized. The Inclusion of s character in the hybrid will be "resisted in the lone uairs are said to he dominant causine the H-O-H anele case of atoms with an nnshared pair.. in the s orbital, to beieduced from the ideal 1091120to 104' and are regul&ly which is more stable than the p orbitals" (p 120). Estimates shown as equivalent in shape and size even being described of the H-E-H angle range from 91.2 for AsH3 to 93.5O for as "rabbit ears" (3). Chemists have even been advised to Hz0. He once again ascribes the anomalies for Hz0 to "re- wear Mickey Mouse ears when teaching the structure and uulsion of the charzes of the hvdroaeu atoms" (D 122). bonding of the H20molecule, presumably to emphasize the - In his famous hook (6)~itac~oro&kii describks the bond- shape of the molecule (and possibly the two equivalent lone ing in water and HzS (p 10).
    [Show full text]
  • Valence Bond Theory
    UMass Boston, Chem 103, Spring 2006 CHEM 103 Molecular Geometry and Valence Bond Theory Lecture Notes May 2, 2006 Prof. Sevian Announcements z The final exam is scheduled for Monday, May 15, 8:00- 11:00am It will NOT be in our regularly scheduled lecture hall (S- 1-006). The final exam location has been changed to Snowden Auditorium (W-1-088). © 2006 H. Sevian 1 UMass Boston, Chem 103, Spring 2006 More announcements Information you need for registering for the second semester of general chemistry z If you will take it in the summer: z Look for chem 104 in the summer schedule (includes lecture and lab) z If you will take it in the fall: z Look for chem 116 (lecture) and chem 118 (lab). These courses are co-requisites. z If you plan to re-take chem 103, in the summer it will be listed as chem 103 (lecture + lab). In the fall it will be listed as chem 115 (lecture) + chem 117 (lab), which are co-requisites. z Note: you are only eligible for a lab exemption if you previously passed the course. Agenda z Results of Exam 3 z Molecular geometries observed z How Lewis structure theory predicts them z Valence shell electron pair repulsion (VSEPR) theory z Valence bond theory z Bonds are formed by overlap of atomic orbitals z Before atoms bond, their atomic orbitals can hybridize to prepare for bonding z Molecular geometry arises from hybridization of atomic orbitals z σ and π bonding orbitals © 2006 H. Sevian 2 UMass Boston, Chem 103, Spring 2006 Molecular Geometries Observed Tetrahedral See-saw Square planar Square pyramid Lewis Structure Theory
    [Show full text]
  • Lewis Structure Handout
    Lewis Structure Handout for Chemistry Students An electron dot diagram, also known as a Lewis Structure, is a representation of valence electrons in a single atom, and can be further utilized to depict the bonds that form based on the available electrons for covalent bonding between multiple atoms. -Each atom has a characteristic dot diagram based on its position in the periodic table and this reflects its potential for fulfillment of the octet rule. As a general rule, the noble gases have a filled octet (column VIII with eight valence electrons) and each preceding column has successively one fewer. • For example, Carbon is in column IV and has four valence electrons. It can therefore be represented as: • Each of these "lone" electrons can form a bond by sharing the orbital with another unbonded electron. Hydrogen, being in column I, has a single valence electron, and will therefore satisfy carbon's octet thusly: • When a compound has bonded electrons (as with CH4) it is most accurate to diagram it with lines representing each of the single bonds. • Using carbon again for a slightly more complicated example, it can also form double bonds. This is the case when carbon bonds to two oxygen atoms, resulting in CO2. The oxygen atoms (with six valence electrons in column VI) are each represented as: • Because carbon is the least electronegative of the atoms involved, it is placed centrally, and its valence electrons will migrate toward the more electronegative oxygens to form two double bonds and a linear structure. This can be written as: where oxygen's remaining lone pairs are still shown.
    [Show full text]
  • Polar Covalent Bonds: Electronegativity
    Polar Covalent Bonds: Electronegativity Covalent bonds can have ionic character These are polar covalent bonds Bonding electrons attracted more strongly by one atom than by the other Electron distribution between atoms is not symmetrical Bond Polarity and Electronegativity Symmetrical Covalent Bonds Polar Covalent Bonds C – C + - C – H C – O (non-polar) (polar) Electronegativity (EN): intrinsic ability of an atom to attract the shared electrons in a covalent bond Inductive Effect: shifting of sigma bonded electrons in resppygonse to nearby electronegative atom The Periodic Table and Electronegativity C – H C - Br and C - I (non-pol)lar) (po l)lar) Bond Polarity and Inductive Effect Nonpolar Covalent Bonds: atoms with similar EN Polar Covalent Bonds: Difference in EN of atoms < 2 Ionic Bonds: Difference in EN > 2 C–H bonds, relatively nonpolar C-O, C-X bonds (more electronegative elements) are polar Bonding electrons shift toward electronegative atom C acqqppuires partial positive char g,ge, + Electronegative atom acquires partial negative charge, - Inductive effect: shifting of electrons in a bond in response to EN of nearby atoms Electrostatic Potential Maps Electrostatic potential maps show calculated charge distributions Colors indicate electron- rich (red) and electron- poor (blue ) reg ions Arrows indicate direction of bond polarity Polar Covalent Bonds: Net Dipole Moments Molecules as a whole are often polar from vector summation of individual bond polarities and lone-pair contributions Strongly polar substances soluble in polar solvents like water; nonpolar substances are insoluble in water. Dipole moment ( ) - Net molecular polarity, due to difference in summed charges - magnitude of charge Q at end of molecular dipole times distance r between charges = Q r, in debyes (D), 1 D = 3.336 1030 coulomb meter length of an average covalent bond, the dipole moment would be 1.60 1029 Cm, or 4.80 D.
    [Show full text]
  • Introduction to Molecular Orbital Theory
    Chapter 2: Molecular Structure and Bonding Bonding Theories 1. VSEPR Theory 2. Valence Bond theory (with hybridization) 3. Molecular Orbital Theory ( with molecualr orbitals) To date, we have looked at three different theories of molecular boning. They are the VSEPR Theory (with Lewis Dot Structures), the Valence Bond theory (with hybridization) and Molecular Orbital Theory. A good theory should predict physical and chemical properties of the molecule such as shape, bond energy, bond length, and bond angles.Because arguments based on atomic orbitals focus on the bonds formed between valence electrons on an atom, they are often said to involve a valence-bond theory. The valence-bond model can't adequately explain the fact that some molecules contains two equivalent bonds with a bond order between that of a single bond and a double bond. The best it can do is suggest that these molecules are mixtures, or hybrids, of the two Lewis structures that can be written for these molecules. This problem, and many others, can be overcome by using a more sophisticated model of bonding based on molecular orbitals. Molecular orbital theory is more powerful than valence-bond theory because the orbitals reflect the geometry of the molecule to which they are applied. But this power carries a significant cost in terms of the ease with which the model can be visualized. One model does not describe all the properties of molecular bonds. Each model desribes a set of properties better than the others. The final test for any theory is experimental data. Introduction to Molecular Orbital Theory The Molecular Orbital Theory does a good job of predicting elctronic spectra and paramagnetism, when VSEPR and the V-B Theories don't.
    [Show full text]
  • Heterocycles 2 Daniel Palleros
    Heterocycles 2 Daniel Palleros Heterocycles 1. Structures 2. Aromaticity and Basicity 2.1 Pyrrole 2.2 Imidazole 2.3 Pyridine 2.4 Pyrimidine 2.5 Purine 3. Π-excessive and Π-deficient Heterocycles 4. Electrophilic Aromatic Substitution 5. Oxidation-Reduction 6. DNA and RNA Bases 7. Tautomers 8. H-bond Formation 9. Absorption of UV Radiation 10. Reactions and Mutations Heterocycles 3 Daniel Palleros Heterocycles Heterocycles are cyclic compounds in which one or more atoms of the ring are heteroatoms: O, N, S, P, etc. They are present in many biologically important molecules such as amino acids, nucleic acids and hormones. They are also indispensable components of pharmaceuticals and therapeutic drugs. Caffeine, sildenafil (the active ingredient in Viagra), acyclovir (an antiviral agent), clopidogrel (an antiplatelet agent) and nicotine, they all have heterocyclic systems. O CH3 N HN O O N O CH 3 N H3C N N HN N OH O S O H N N N 2 N O N N O CH3 N CH3 caffeine sildenafil acyclovir Cl S N CH3 N N H COOCH3 nicotine (S)-clopidogrel Here we will discuss the chemistry of this important group of compounds beginning with the simplest rings and continuing to more complex systems such as those present in nucleic acids. Heterocycles 4 Daniel Palleros 1. Structures Some of the most important heterocycles are shown below. Note that they have five or six-membered rings such as pyrrole and pyridine or polycyclic ring systems such as quinoline and purine. Imidazole, pyrimidine and purine play a very important role in the chemistry of nucleic acids and are highlighted.
    [Show full text]
  • Lewis Structures
    Lewis Structures Valence electrons for Elements Recall that the valence electrons for the elements can be determined based on the elements position on the periodic table. Lewis Dot Symbol Valence electrons and number of bonds Number of bonds elements prefers depending on the number of valence electrons. In general - F a m i l y → # C o v a l e n t B o n d s* H a l o g e n s X F , B r , C l , I → 1 bond often C a l c o g e n s O 2 bond often O , S → N i t r o g e n N 3 bond often N , P → C a r b o n C → 4 bond always C , S i The above chart is a guide on the number of bonds formed by these atoms. Lewis Structure, Octet Rule Guidelines When compounds are formed they tend to follow the Octet Rule. Octet Rule: Atoms will share electrons (e-) until it is surrounded by eight valence electrons. 4 unpaired 3unpaired 2unpaired 1unpaired up = unpaired e- 4 bonds 3 bonds 2 bonds 1 bond O=C=O N≡ N O = O F - F Atomic Connectivity The atomic arrangement for a molecule is usually given. CH2ClF HNO3 CH3COOH H2SO4 Cl N H O O O O H C F O H C C H O S O H H H H O H O In general when there is a single central atom in the 3- molecule, CH2ClF, SeCl2, O3 (CO2, NH3, PO4 ), the central atom is the first atom in the chemical formula.
    [Show full text]
  • Draw Three Resonance Structures for the Chlorate Ion, Clo3
    University Chemistry Quiz 3 2014/11/6 + 1. (10%) Explain why the bond order of N2 is greater than that of N2 , but the bond + order of O2 is less than that of O2 . Sol. + In forming the N2 from N2, an electron is removed from the sigma bonding molecular orbital. Consequently, the bond order decreases to 2.5 from 3.0. In + forming the O2 ion from O2, an electron is removed from the pi antibonding molecular orbital. Consequently, the bond order increases to 2.5 from 2.0. - 2. (10%) Draw three resonance structures for the chlorate ion, ClO3 . Show formal charges. Sol. Strategy: We follow the procedure for drawing Lewis structures outlined in Section 3.4 of the text. After we complete the Lewis structure, we draw the resonance structures. Solution: Following the procedure in Section 3.4 of the text, we come up with − the following Lewis structure for ClO3 . O − + − O Cl O We can draw two more equivalent Lewis structures with the double bond between Cl and a different oxygen atom. The resonance structures with formal charges are as follows: − − O O O + − − + − − + O Cl O O Cl O O Cl O 3. (5%) Use the molecular orbital energy-level diagram for O2 to show that the following Lewis structure corresponds to an excited state: Sol. The Lewis structure shows 4 pairs of electrons on the two oxygen atoms. From Table 3.4 of the text, we see that these 8 valence electrons are placed in the σ p p p p 2p , 2p , 2p , 2p , and 2p orbitals.
    [Show full text]
  • Molecular Orbital Theory
    Molecular Orbital Theory The Lewis Structure approach provides an extremely simple method for determining the electronic structure of many molecules. It is a bit simplistic, however, and does have trouble predicting structures for a few molecules. Nevertheless, it gives a reasonable structure for many molecules and its simplicity to use makes it a very useful tool for chemists. A more general, but slightly more complicated approach is the Molecular Orbital Theory. This theory builds on the electron wave functions of Quantum Mechanics to describe chemical bonding. To understand MO Theory let's first review constructive and destructive interference of standing waves starting with the full constructive and destructive interference that occurs when standing waves overlap completely. When standing waves only partially overlap we get partial constructive and destructive interference. To see how we use these concepts in Molecular Orbital Theory, let's start with H2, the simplest of all molecules. The 1s orbitals of the H-atom are standing waves of the electron wavefunction. In Molecular Orbital Theory we view the bonding of the two H-atoms as partial constructive interference between standing wavefunctions of the 1s orbitals. The energy of the H2 molecule with the two electrons in the bonding orbital is lower by 435 kJ/mole than the combined energy of the two separate H-atoms. On the other hand, the energy of the H2 molecule with two electrons in the antibonding orbital is higher than two separate H-atoms. To show the relative energies we use diagrams like this: In the H2 molecule, the bonding and anti-bonding orbitals are called sigma orbitals (σ).
    [Show full text]
  • Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Review Questions
    Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Review Questions 10.1 J The properties of molecules are directly related to their shape. The sensation of taste, immune response, the sense of smell, and many types of drug action all depend on shape-specific interactions between molecules and proteins. According to VSEPR theory, the repulsion between electron groups on interior atoms of a molecule determines the geometry of the molecule. The five basic electron geometries are (1) Linear, which has two electron groups. (2) Trigonal planar, which has three electron groups. (3) Tetrahedral, which has four electron groups. (4) Trigonal bipyramid, which has five electron groups. (5) Octahedral, which has six electron groups. An electron group is defined as a lone pair of electrons, a single bond, a multiple bond, or even a single electron. H—C—H 109.5= ijj^^jl (a) Linear geometry \ \ (b) Trigonal planar geometry I Tetrahedral geometry I Equatorial chlorine Axial chlorine "P—Cl: \ Trigonal bipyramidal geometry 1 I Octahedral geometry I 369 370 Chapter 10 Chemical Bonding II The electron geometry is the geometrical arrangement of the electron groups around the central atom. The molecular geometry is the geometrical arrangement of the atoms around the central atom. The electron geometry and the molecular geometry are the same when every electron group bonds two atoms together. The presence of unbonded lone-pair electrons gives a different molecular geometry and electron geometry. (a) Four electron groups give tetrahedral electron geometry, while three bonding groups and one lone pair give a trigonal pyramidal molecular geometry.
    [Show full text]
  • PS 6 Answers, Spring 2003, Prof
    Chem 332, PS 6 answers, Spring 2003, Prof. Fox 1. Which of the following are aromatic. Draw "Frost Circles" that support your answers 2– 2+ 2+ + all bonding MO's although all electrons all bonding MO's this is an interesting case: this is are in bonding orbitals, are filled are filled a 4n+2 molecule, and it is flat-- two bonding MO's are aromatic aromatic characteristic of aromatic only half-filled molecules. However, it has four anti-aromatic electrons in non-bonding orbitals, so we might have considered this to be non-aromatic. So it comes down to how you define aromaticity. For now we will call it aromatic with a question mark. 2. Which of the following are aromatic. Use the Huckel rule to support your answer 2– P H B 14 electrons, aromatic H 6 electrons 10 electrons 2 6 electrons sp3 center (filled sp on aromatic (empty sp2 on not conjugated phosphorus) boron) and not aromatic aromatic aromatic 3. Compound 1 has an unusually large dipole moment for an organic hydrocarbon. Explain why (consider resonance structures for the central double bond) We need to remember that for any double bond compound, we can draw two polar resonance structure. Normally, these resonance structures are unimportant unimportant unimportant However, for 1 the story is different because there is a polar resonance form that has two aromatic components. In other words, 'breaking' the double bond gives rise to aromaticity, and therefore the polar resonance structure is important and leads to the observed dipole. dipole 1 8 e– 4 e– 6 e– 6 e– anti- anti- aromatic aromatic aromatic aromatic important resonance unimportant resonance structure structure 5.
    [Show full text]
  • Structure of Benzene, Orbital Picture, Resonance in Benzene, Aromatic Characters, Huckel’S Rule B
    Dr.Mrityunjay Banerjee, IPT, Salipur B PHARM 3rd SEMESTER BP301T PHARMACEUTICAL ORGANIC CHEMISTRY –II UNIT I UNIT I Benzene and its derivatives 10 Hours A. Analytical, synthetic and other evidences in the derivation of structure of benzene, Orbital picture, resonance in benzene, aromatic characters, Huckel’s rule B. Reactions of benzene - nitration, sulphonation, halogenationreactivity, Friedelcrafts alkylation- reactivity, limitations, Friedelcrafts acylation. C. Substituents, effect of substituents on reactivity and orientation of mono substituted benzene compounds towards electrophilic substitution reaction D. Structure and uses of DDT, Saccharin, BHC and Chloramine Benzene and its Derivatives Chemists have found it useful to divide all organic compounds into two broad classes: aliphatic compounds and aromatic compounds. The original meanings of the words "aliphatic" (fatty) and "aromatic" (fragrant/ pleasant smell). Aromatic compounds are benzene and compounds that resemble benzene in chemical behavior. Aromatic properties are those properties of benzene that distinguish it from aliphatic hydrocarbons. Benzene: A liquid that smells like gasoline Boils at 80°C & Freezes at 5.5°C It was formerly used to decaffeinate coffee and component of many consumer products, such as paint strippers, rubber cements, and home dry-cleaning spot removers. A precursor in the production of plastics (such as Styrofoam and nylon), drugs, detergents, synthetic rubber, pesticides, and dyes. It is used as a solvent in cleaning and maintaining printing equipment and for adhesives such as those used to attach soles to shoes. Benzene is a natural constituent of petroleum products, but because it is a known carcinogen, its use as an additive in gasoline is now limited. In 1970s it was associated with leukemia deaths.
    [Show full text]