# Semester 1 Bonding Lecture1

Information about Semester 1 Bonding Lecture1

Published on October 16, 2007

Author: Shariyar

Source: authorstream.com

Bonding, Molecular Shape & Structure:  Bonding, Molecular Shape & Structure By Dr. Fawaz Aldabbagh http://www.nuigalway.ie/chem/Fawaz/fawaz.htm Slide2:  The Periodic Table Slide3:  Lewis Symbols Represent the number of valence electrons as dots Valence number is the same as the Periodic Table Group Number For example, Groups 1 2 3 4 5 6 7 8 n = 1 n = 2 Slide4:  Elements want to achieve the stable electron configuration of the nearest noble gas Atoms tend to gain, lose or share electrons until they are surrounded by 8 electrons Octet Rule n = 2 n = 3 Slide5:  Nobel Gas Has a Stable Electron Configuration Electronic configuration of Neon achieved in both cases Example of Ionic Bonding 10 11 9 Slide6:  Ionic Bonding refers to electrostatic forces between ions, usually a metal cation and a non-metal anion Covalent Bonding results from the sharing of two electrons between two atoms (usually non-metals) resulting in molecules There are two types of bonding; Octet Rule applies Each Covalent Bond contains two electrons Triple bond Slide7:  Covalent Bonding – Atoms Share Electrons Slide8:  Hydrogen molecule, H2 Concentration of negative charge between two nuclei occurs in a covalent bond 7A elements (e.g. F) have one valence electron for covalent bonding, so to achieve octet 6A elements (e.g. O) use two valence electrons for covalent bonding, so to achieve octet 5A elements (e.g. N) use three valence electrons for covalent bonding, so to achieve octet 4A elements (e.g. C) use four valence electrons for covalent bonding, so to achieve octet Slide9:  Carbon dioxide, CO2 Double bonds Rules for Drawing Lewis Structures First sum the number of valence electrons from each atom The central atom is usually written first in the formula Complete the octets of atoms bonded to the central atom (remember that H can only have two electrons) Place any left over electrons on the central atom, even if doing so it results in more than an octet If there are not enough electrons to give the central atom an octet , try multiple bonds E.g. 1. PCl3 Total Number of valence electrons = 5 + (3 x 7) = 26 Total Number of valence electrons = 4 + (2 x 6) = 16 Slide10:  E.g. 2; CHBr3 Total Number of valence electrons = 4 + 1 + (3 x 7) = 26 Exceptions to the Octet Rule in Covalent Bonding Molecules with an odd number of electrons Other Natural Radicals, which do not obey Lewis Structures (e.g. O2) Molecules in which an atom has less than an octet 3. Molecules in which an atom has more than an octet Slide11:  1. Odd Number of Electrons NO Number of valence electrons = 11 NO2 Number of valence electrons = 17 O2 Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons) Molecules and atoms which are neutral (contain no formal charge) and with an unpaired electron are called Radicals Slide12:  2. Less than an Octet Includes Lewis acids such as halides of B, Al and compounds of Be BCl3 Group 3A atom only has six electrons around it However, Lewis acids “accept” a pair of electrons readily from Lewis bases to establish a stable octet Slide13:  AlX3 Aluminium chloride is an ionic solid in which Al3+ is surrounded by six Cl-. However, it sublimes at 192 °C to vapour Al2Cl6 molecules B2H6 A Lewis structure cannot be written for diborane. This is explained by a three-centre bond – single electron is delocalized over a B-H-B Slide14:  Octet Rule Always Applies to the Second Period = n2 ; number of orbitals 2s, 2px, 2py, 2pz ---orbitals cannot hold more than two electrons Ne [He]; 2s2, 2px2, 2py2, 2pz2 n = 2 n = 3 Third Period ; n2 = 32 = 9 orbitals:  Third Period ; n2 = 32 = 9 orbitals Ar [Ne]; 3s2, 3px2, 3py2, 3pz2 3d0 3d0 3d0 3d0 3d0 n = 3 Slide16:  3. More than an Octet PCl5 Elements from the third Period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding P : (Ne) 3s2 3p3 3d0 Number of valence electrons = 5 + (5 x 7) = 40 10 electrons around the phosphorus SF4 S : (Ne) 3s2 3p4 3d0 Number of valence electrons = 6 + (4 x 7) = 34 The Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency. Slide17:  Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself Prof. Linus Pauling Nobel Prize for Chemistry 1954 Nobel Prize for Peace 1962 Electronegativity is a function of two properties of isolated atoms; The atom’s ionization energy (how strongly an atom holds onto its own electrons) The atom’s electron affinity (how strongly the atom attracts other electrons) For example, an element which has: A large (negative) electron affinity A high ionization (always endothermic, or positive for neutral atoms) Will: Attract electrons from other atoms and Resist having electrons attracted away Such atoms will be highly electronegative Slide18:  Pauling scale of electronegativity; Fluorine is the most electronegative element followed by O and N, Cl are equal third. Cs is least. Electronegativity increases from left to right along the Periodic Table. For the representative elements (s & p block), the electronegativity decreases as you go down a group. No trend in the transition metals. Slide21:  Electronegativity is dictated by The number of protons in the nucleus across a period you are increasing the number of protons, but filling electrons in the same Bohr quantized energy level. You are only filling sub-shells, so electronegativity increases from left to right The distance from the nucleus down groups, you are placing electrons into new quantized energy levels, so moving further away from the attractive power of the nucleus. Outer shell becomes further away from the nucleus. The amount of screening by the inner electrons level of screening upon bonding electrons increases down groups, and adds to the reduction in electronegativity. Screening is caused by repulsion of electrons for each other. In hydrogen atom, energy of orbital depends on the principle quantum number, n. But in many electron atoms, electron-repulsions cause different sub-shells to have different energies, Sub-shell energy increases (with increasing l) s < p < d Slide22:  The three major types of intramolecular bond can be described by the electronegativity difference: Non-Polar Covalent – Bonds which occur between atoms with little or no electronegativity difference (less than 0.5). Polar Covalent – Bonds which occur between atoms with a definite electronegativity difference (between 0.5 and 2.0). Ionic – Bonds which occur between atoms with a large electronegativity difference (2.0 or greater), where electron transfer can occur. E.g. F-F (4.0 – 4.0 = 0) is non-polar covalent H-F (4.0 – 2.1 = 1.9) is polar covalent LiF (4.0 – 1.0 = 3.0) is ionic + - Slide24:  Dipole Moment occurs in any polar covalent bond, because of an unequal sharing of the electron pair between two atoms E.g. Which of the following bonds is most polar: S-Cl, S-Br, Se-Cl or Se-Br? S-Cl (3.0 – 2.5) = 0.5 S-Br (2.8-2.5) = 0.3 Se-Cl (3.0-2.4) = 0.6 Se-Br (2.8-2.4) = 0.4 Therefore, Se-Cl is the most polar! We should be able to reach the same conclusion using the Periodic Table, Cl is furthest to the right and to the top of the Periodic Table, so is the most electronegative. Se is furthest to the left (‘metallic like’) and towards the bottom. Therefore, difference in electronegativity should be the greatest! Slide25:  Electronegativity difference decreases as bond length increases Dipole Moment: µ = Qr Dipole moment is defined as the magnitude of charge (Q) multiplied by the distance between the charges; units are D (Debye) = 3.36 x 1030 C.m Prof. Peter Debye Noble Prize 1936 Slide26:  When proton & electron 100 pm apart, the dipole moment is 4.80 D 4.8 D is a key reference value! It represents a pure charge of +1 and -1, which are 100 pm (100pm = 1Å) apart. The bond is said to be 100% ionic! H-F; µ = 1.82 D (measured) bond length = 0.92 Å If 100% ionic, µ = 92/100 (4.8 D) = 4.42 D % ionic = 1.82/4.42 x 100 = 41 % ionic H-Cl; µ = 1.08 D (measured) bond length = 1.27 Å If 100% ionic, µ = 127/100 (4.8 D) = 6.10 D % ionic = 1.08/6.10 x 100 = 18 % ionic H-Br; µ = 0.82 D (measured) bond length = 1.41 Å If 100% ionic, µ = 141/100 (4.8 D) = 6.77 D % ionic = 0.82/6.77 x 100 = 12 % ionic Slide27:  Polar Molecules = Molecules with permanent dipole moments HCl has only one covalent bond (which is polar). Therefore, its dipole moment = H-Cl bond dipole In a molecule with two or more polar bonds, each bond has a dipole moment contribution = bond dipole Net dipole moment = vector sum of its bond dipoles Linear Molecules: CO2 is Non-polar Because CO2 dipoles are orientated in opposite directions. The dipoles have equal magnitudes; they cancel Net dipole = 0 Slide28:  Symmetrical molecules (e.g. CCl4, CH4) are non-polar. The four dipoles are of equal magnitude and neutralize one another at the center of a tetrahedron Non-symmetrical molecules (e.g. CHCl3, CO(CH3)2, H2O) are Polar. The dipoles are not all equal or in opposite directions (partial charges and bond lengths are all different in C-Cl, C-H, C=O, C-H) (H2O is a bent molecule not linear, see later notes) Slide29:  Formal Charges: the number of valence electrons in the isolated atom minus the number of electrons assigned to the atom in the Lewis structure. These are not real charges, but help with keeping count of electrons in Lewis structures. E.g. CN- Question: Draw the Lewis structures of NO+ and determine the formal charges of the atoms. Which Lewis structure is the preferred one? Number of valence electrons = 9 + 1 =10 Number of valence electrons = 11 - 1 = 10 Structure 1 is preferred because the positive charge is on the least electronegative atom. 1 Slide30:  Lewis structures of Charged Molecules: Predict the most likely structure! E.g. NCS- Number of valence electrons = 15 + 1 =16 Structure 1 is preferred because the negative charge is on the most electronegative atom with the lowest formal charge. 1 Tutorial Questions: Use the electronegativities of C (2.5) and Cl (3.0) to describe the character of the C-Cl bond in CCl4, and explain why CCl4 is a non-polar molecule. CHCl3 has a C-Cl bond of 178 pm, and measurements reveal 1.87 D. Calculate the percentage ionic character. Is this a polar molecule? Draw the most plausible Lewis Structure for NO2+, H2SO4 and SO42- Describe the molecule (ClO2)- using three possible Lewis structures, which is the most important? Slide31:  Shapes of Molecules We use Lewis structures to account for formula of covalent compounds. Lewis structures also account for the number of covalent bonds. Lewis structures however do not account for the shapes of molecules. Molecules of ABn have shapes dependent on the value of n AB2 must be either linear or bent: Examples of Linear molecules Linear - No non-bonding electrons Slide32:  Linear Molecules have a bond angle = 180° Bent molecules have a bond angle ≠ 180° AB3 most common shapes place the B atoms at the corners of an equilateral triangle: bent Trigonal Planar The A atom lies in the same plane as the B atoms (Flat) Bond angle = 120° No non-bonding electrons Slide33:  The A atom lies above the plane of the B atom. Pyramid with an equilateral triangle as the base. Trigonal Pyramidal Slide34:  The ideal tetrahedron has a bond angle = 109.5° The lone electron pair exerts a little extra repulsion on the three bonding hydrogen atoms to create a slight compression to a 107° bond angle. VSEPR model explains distortions of molecules Less repulsion is exerted by a bonding pair of electrons because they feel attraction from two nuclei, while a non-bonding pair feels attraction from only one nucleus. Non-bonding pairs spread out more! Slide35:  AB4 is Tetrahedral The carbon has 4 valence electrons and thus needs 4 more electrons from four hydrogen atoms to complete its octet. The hydrogen atoms are as far apart as possible at 109° bond angle. This is tetrahedral geometry. The molecule is three dimensional. Slide36:  Valence-Shell Electron-Pair Repulsion Theory (VSEPR) In molecules there are 2 types of electron 1. Bonding Pairs 2. Non-bonding or lone pairs   The combinations of these determine the shape of the molecule   Single bonds have a big impact on shape, double bonds have little effect The outer pairs of electrons around a covalently bonded atom minimize repulsions between them by moving as far apart as possible Slide37:  Water is a bent molecule with bond angles of 104.5° Notice – the bond angle decreases as the number of non-bonding pairs increases AB2 - classification H2O Slide38:  Ozone O3 ; number of valence electrons = 18 electrons Resonance structures AB3 - classification Slide39:  Valence Shell Electron-Pair Repulsion Theory (VSEPR) Procedure Sum the total Number of Valence Electrons Drawing the Lewis Structure 2. The atom usually written first in the chemical formula is the Central atom in the Lewis structure Complete the octet bonded to the Central atom. However, elements in the third row have empty d-orbitals which can be used for bonding. If there are not enough electrons to give the central atom an octet try multiple bonds. Predicting the Shape of the Molecule Sum the Number of Electron Domains around the Central Atom in the Lewis Structure; Single = Double = Triple Bonds = Non-Bonding Lone Pair of Electrons = One Electron Domain From the Total Number of Electron Domains, Predict the Geometry and Bond Angle(s); 2 (Linear = 180º); 3 (Trigonal Planar = 120º); 4 (Tetrahedral = 109.5º); 5 (Trigonal Bipyramidal = 120º and 90º); 6 (Octahedral = 90º) Lone Pair Electron Domains exert a greater repulsive force than Bonding Domains. Electron Domains of Multiple Bonds exert a greater repulsive force than Single Bonds. Thus they tend to compress the bond angle. Slide40:  Further Examples: Tutorial Questions : Draw Lewis structures and the molecular geometry of the following molecules: H3O+, NH4+, CS2, SCl2 Slide42:  Molecules with Expanded Valence Shells When the central atom of a molecule is from the third period of the Periodic Table and beyond, that atom may have more than four pairs of electrons around it Five pairs of electrons around the central atom are based on the Trigonal Bipyramidal structure. Three pairs define an Equatorial Triangle (Equatorial electrons) Two pairs lie above and below the triangle plane (Axial electrons) AB5: e.g. PCl5 The repulsion between pairs located 90° apart are much greater than for those 120° apart: Slide43:  Because repulsion is greater for non-bonding than for bonding electron pairs, then non-bonding pairs occupy equatorial positions on the Trigonal Bipyramidal structure SF4 : The non-bonding pair occupies an equatorial position. The axial and equatorial S-F bonds are slightly bent back because of the larger repulsive effect of the lone pair. BrF3 : T-shaped 116° and 186º 90° Slide44:  Third Period ; n2 = 32 = 9 orbitals Ar [Ne]; 3s2, 3px2, 3py2, 3pz2 3d0 3d0 3d0 3d0 3d0 n = 3 Slide45:  Six pairs of electrons around the central atom are based on the Octahedron structure. AB6 : e.g. SF6 The central atom can be visualized as being at the centre of an octahedron, with the six electrons pointing to the six vertices – all bond angles are 90° Octahedral Square Pyramidal E.g. BrF5 Square Planar E.g. XeF4 Should be less than 90º 90° Slide46:  Intermolecular Forces: are generally much weaker than covalent or ionic bonds. Less energy is thus required to vaporize a liquid or melt a solid. Boiling points can be used to reflect the strengths of intermolecular forces (the higher the Bpt, the stronger the forces) Hydrogen Bonding : the attractive force between hydrogen in a polar bond (particularly H-F, H-O, H-N bond) and an unshared electron pair on a nearby small electronegative atom or ion Very polar bond in H-F. The other hydrogen halides don’t form hydrogen bonds, since H-X bond is less polar. As well as that, their lone pairs are at higher energy levels. That makes the lone pairs bigger, and so they don't carry such an intensely concentrated negative charge for the hydrogens to be attracted to. Slide47:  Hydrogen Bonding & Water Slide48:  One of the most remarkable consequences of H-bonding is found in the lower density of ice in comparison to liquid water, so ice floats on water. In most substances the molecules in the solid are more densely packed than in the liquid. A given mass of ice occupies a greater volume than that of liquid water. This is because of an ordered open H-bonding arrangement in the solid (ice) in comparison to continual forming & breaking H-bonds as a liquid. Slide49:  Weaker Intermolecular Forces Ion-Dipole Forces An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. A positive ion (cation) attracts the partially negative end of a neutral polar molecule. A negative ion (anion) attracts the partially positive end of a neutral polar molecule. Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases. Slide50:  Dipole-dipole Attractive Forces A dipole-dipole force exists between neutral polar molecules Polar molecules attract one another when the partial positive charge on one molecule is near the partial negative charge on the other molecule The polar molecules must be in close proximity for the dipole-dipole forces to be significant Dipole-dipole forces are characteristically weaker than ion-dipole forces Dipole-dipole forces increase with an increase in the polarity of the molecule Slide51:  Boiling points increase for polar molecules of similar mass, but increasing dipole: Slide52:  London Dispersion Forces – significant only when molecules are close to each other Prof. Fritz London Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom Slide53:  The ease with which an external electric field can induce a dipole (alter the electron distribution) with a molecule is referred to as the "polarizability" of that molecule The greater the polarizability of a molecule the easier it is to induce a momentary dipole and the stronger the dispersion forces Larger molecules tend to have greater polarizability Their electrons are further away from the nucleus (any asymmetric distribution produces a larger dipole due to larger charge separation) The number of electrons is greater (higher probability of asymmetric distribution) thus, dispersion forces tend to increase with increasing molecular mass Dispersion forces are also present between polar/non-polar and polar/polar molecules (i.e. between all molecules) Slide54:  Group 4A hydrides Groups 4, 5, 6A hydrides Van der Waals forces are made of dipole-dipole and London dispersion forces

26. 03. 2008
0 views

19. 10. 2007
0 views

01. 11. 2007
0 views

08. 10. 2007
0 views

12. 10. 2007
0 views

15. 10. 2007
0 views

15. 10. 2007
0 views

17. 10. 2007
0 views

24. 10. 2007
0 views

16. 10. 2007
0 views

11. 12. 2007
0 views

02. 11. 2007
0 views

22. 08. 2007
0 views

22. 10. 2007
0 views

07. 11. 2007
0 views

14. 11. 2007
0 views

19. 11. 2007
0 views

20. 11. 2007
0 views

10. 10. 2007
0 views

25. 09. 2007
0 views

22. 08. 2007
0 views

22. 08. 2007
0 views

22. 08. 2007
0 views

22. 08. 2007
0 views

02. 08. 2007
0 views

02. 08. 2007
0 views

02. 08. 2007
0 views

02. 08. 2007
0 views

25. 09. 2007
0 views

24. 10. 2007
0 views

25. 09. 2007
0 views

02. 08. 2007
0 views

16. 10. 2007
0 views

04. 01. 2008
0 views

07. 12. 2007
0 views

17. 10. 2007
0 views

25. 09. 2007
0 views

19. 02. 2008
0 views

26. 02. 2008
0 views

22. 08. 2007
0 views

29. 02. 2008
0 views

16. 11. 2007
0 views

04. 03. 2008
0 views

22. 08. 2007
0 views

18. 03. 2008
0 views

25. 03. 2008
0 views

30. 03. 2008
0 views

09. 04. 2008
0 views

10. 04. 2008
0 views

13. 04. 2008
0 views

04. 10. 2007
0 views

14. 04. 2008
0 views

17. 04. 2008
0 views

17. 04. 2008
0 views

02. 08. 2007
0 views

19. 06. 2007
0 views

19. 06. 2007
0 views

19. 06. 2007
0 views

19. 06. 2007
0 views

19. 06. 2007
0 views

22. 08. 2007
0 views

05. 01. 2008
0 views

18. 06. 2007
0 views

18. 06. 2007
0 views

18. 06. 2007
0 views

22. 08. 2007
0 views

03. 10. 2007
0 views

02. 08. 2007
0 views

30. 12. 2007
0 views

26. 10. 2007
0 views

02. 08. 2007
0 views

15. 06. 2007
0 views

15. 06. 2007
0 views

15. 06. 2007
0 views

15. 06. 2007
0 views

15. 06. 2007
0 views

15. 06. 2007
0 views

15. 06. 2007
0 views

22. 08. 2007
0 views

26. 11. 2007
0 views

19. 06. 2007
0 views

24. 11. 2007
0 views

18. 06. 2007
0 views

27. 11. 2007
0 views

19. 06. 2007
0 views

19. 06. 2007
0 views

23. 10. 2007
0 views

11. 10. 2007
0 views

15. 06. 2007
0 views

14. 03. 2008
0 views

24. 02. 2008
0 views

22. 10. 2007
0 views

17. 11. 2007
0 views

15. 11. 2007
0 views

25. 09. 2007
0 views

22. 08. 2007
0 views