Chemical Bonds PDF

1 CHEMICAL BONDS Cocaine 2 Introduction A chemical bond is an attraction between atoms that allows the formation of chemical substances that c...

1 CHEMICAL BONDS Cocaine 2 Introduction A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. The bond is caused by the electrostatic force of attraction between opposite charges, either between electrons and nuclei, or as the result of a dipole attraction. The strength of chemical bonds varies considerably; there are "strong bonds" such as covalent or ionic bonds and "weak bonds" such as dipole–dipole interactions, the London dispersion force and hydrogen bonding.  Definitions 3 Bond: A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. Bond Energy: The energy released when a bond is formed or absorbed when it is broken is called the bond energy. e.g. the C-H bond in methane has a bond energy of 413 KJ/mol The enthalpy change required to break a particular 4 bond in one mole of gaseous molecules is the bond energy. Bond Energy H2 (g) H (g) + H (g) ∆H0 = 436.4 kJ Cl2 (g) Cl (g) + Cl (g) ∆ H0 = 242.7 kJ HCl (g) H (g) + Cl (g) ∆ H0 = 431.9 kJ O2 (g) O (g) + O (g) ∆ H0 = 498.7 kJ O O N2 (g) N (g) + N (g) ∆ H0 = 941.4 kJ N N Bond Energies Single bond < Double bond < Triple bond  5 Average bond energy in polyatomic molecules H2O (g) H (g)+OH (g) DH0 = 502 kJ OH (g)H (g)+O (g) DH0 = 427 kJ Average OH bond energy = 502 + 427 = 464 kJ 2 Electronegativity What is it? Electronegativity is the power of an atom to attract electron density  in a covalent bond  6 The concept of Electronegativity was developed by Linus Pauling. Electronegativity is the ability of an element to attract electrons to itself in a molecule. Electronegativity increases across the periodic table and is at a maximum in the top right hand corner at fluorine, and is at a minimum at the bottom left hand corner at Linus Carl Pauling Fransium. (1901-1994)  Electronegativity 7 Pauling’s electronegativity scale The higher the value, the more electronegative the element Fluorine is the most electronegative element It has an electronegativity value of 4.0 H He 2.1 - Li Be B C N O F Ne 1.0 1.5 2.0 2.5 3.0 3.5 4.0 - Na Mg Al Si P S Cl Ar 0.9 1.2 1.5 1.8 2.1 2.5 3.0 - There are two major bond classifications, each with 8 identifiable sub-groups: The valence electrons are in the outer (highest) energy levels and are the main determiners of chemical activity and bonding. An element will be chemically reactive if it can get to the stable electronic configuration of an inert gas, either by losing one or two electrons to another atom, or by gaining one or two electrons from another atom (at most three), or by sharing three or more electrons Primary Bonds: Chemical (strong) bonding, involves the transfer9 or sharing of electrons: ionic, covalent, or metallic bonds. Ionic—complete transfer of 1 or more electrons from one atom to another. Covalent—some valence electrons shared between atoms. Metallic Bonds—Metallic bonding occurs between the positive atom cores and the "nearly free" electrons. Ionic Bonds: Essentially complete electron transfer from an element of low IE (metal) to an element of high affinity for electrons (nonmetal) 2 Na(s) + Cl2(g) ---> 2NaCl Therefore, ionic compounds. exist primarily between metals at left of periodic table (Grps 1A and 2A and transition metals) and nonmetals at right (O and halogens). Atoms near the left or right sides of the periodic table can loose or gain 1 (or102) electrons to form charged "ions". For example, a Sodium atom (row 3, column IA in the periodic table) can loose one electron to have 8 valence electrons and become a positively charged "cation". A Chlorine atom (row 3, column VIIA) can gain one electron to have 8 valence electrons and become a negatively charged "anion". These two ions then will be attracted to each other by non-directional electrostatic force and form an ionic (or electrovalent) bond. Note: cations are +, anions are -. When large numbers of such ion pairs come together an ionic solid is formed. Common salt (NaCl) is an ionic solid which has the cubic structure shown on the right.  11 This involves the sharing of electrons by 2 or more atoms. The electrons are shared, not taken as with ionic bonds. EXAMPLES INCLUDE THE BONDS BETWEEN: H2 F2 Br2 Cl2 H2O Example: a chlorine molecule, each of the atoms shares a pair of electrons circled in GREEN, making the covalent bond. Since it’s just one pair being shared, this is a SINGLE COVALENT BOND. Each of the chlorine atoms also has three pairs each of UNSHARED ELECTRONS. THEY BOTH GET AN OCTET BY SHARING. In covalent bonding: There are no charge requirements. Each atom does not have to have nearest neighbours of opposite charge. Bonds are only between nearest-neighbour atoms sharing electrons. Unshared Electrons 12 When electrons are written into the Lewis diagrams, they are usually paired. The electrons prefer pairs. When they share themselves with other atoms, one electron from each atom connects together, making a SINGLE BOND. Electrons not involved in bonding are mostly paired away from the bond. They are the “unshared” electrons. If they connect 2 or 3 pairs at a time, they make DOUBLE or TRIPLE BONDS. Water has two hydrogen atoms bonded to one oxygen atom. It has 2 single polar covalent bonds. The oxygen has 2 pairs of unshared electrons as well. Oxygen gets the octet.  Metallic Bond 13 Bond that exists between metal atoms. Alloy – two or more different metal atoms bonded together. Electron Sea Model: Metal atoms give up valence electrons and form +ve ions The released electrons move freely around the +ve metal ions This means that in metallic bonding the atoms pack together as closely as possible. Metallic solids occur when large numbers of atoms bond together in close-packed structures. They can be modelled as ping-pong balls glued together as shown in the diagram. 14 Properties of Metallic Bonds Good conductors of electricity – free electrons Malleable and ductile – not in rigid position so ions can be shaped and drawn into wires. Lusterous – absorb and emit light in regular pattern due to free electrons. Secondary Bonds Secondary or weak bonds are formed when there is effectively a partial and/or momentary charge. They are secondary in terms of strength but not necessarily in terms of importance, as life is only made possible because of them.  Hydrogen bonding 15 Are the attractive force caused by hydrogen bonded to F, O, or N. F, O, and N are very electronegative so it is a very strong dipole. The hydrogen partially share with the lone pair in the molecule next to it. The strongest of the intermolecular forces. d+ d- H O H d+ 16 H O H  Van der Waals Forces 17 Two electrically neutral, closed-shell atoms d- d+ d- d+ Temporary dipole resulting Induced dipole, due to Gives net presence of other dipole from quantum fluctuation attraction Although Van der Waals forces are weak, they are often the only attractive force between molecules. VdW forces are not described by Hartree-Fock theory, because they are due to correlation effects. The dependence on the charge density is nonlocal. Bond Properties 18 bond order, bond length, bond energy, bond polarity Buckyball in HIV-protease  19 Bond Order Number of bonds between a pair of atoms Double bond Single bond Acrylonitrile Triple bond 20 Bond Order Fractional bond orders in resonance structures. Consider NO2- N N O O O O The N—O bond order = 1.5 Total # of e - pairs used for a type of bond Bond order = Total # of bonds of that type 3 e - pairs in NO bonds Bond order = 2 N — O bonds  21 Bond Order Bond order is proportional to two important bond properties: (a) bond strength (b) bond length 414 kJ 110 pm 123 pm 745 kJ 22 Bond Length Bond length is the distance between the nuclei of two bonded atoms.  23 Bond Length Bond length depends on bond order. Bond distances measured using CAChe software. In Angstrom units where 1 A = 10-2 pm. 24 Polarity Boiling point = 100 ˚C Boiling point = -161 ˚C Why do water and methane differ so much in their boiling points? Why do ionic compounds dissolve in water?  25 Polarity of Water In a water molecule two hydrogen atoms form single polar covalent bonds with an oxygen atom. Gives water more structure than other liquids – Because oxygen is more electronegative, the region around oxygen has a partial negative charge. – The region near the two hydrogen atoms has a partial positive charge. A water molecule is a polar molecule with opposite ends of the molecule with opposite charges. 26 Water has a variety of unusual properties because of attractions between these polar molecules. – The slightly negative regions of one molecule are attracted to the slightly positive regions of nearby molecules, forming a hydrogen bond. – Each water molecule can form hydrogen bonds with up to four neighbors. Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings  27 Molecular Polarity Molecules—such as HCl and H2O— can be POLAR (or dipolar). They have a DIPOLE MOMENT. The polar HCl molecule will turn to align with an electric field. Figure 9.15 28 Polar or Nonpolar? Compare CO2 and H2O. Which one is polar?  29 Carbon Dioxide CO2 is NOT polar even though the CO bonds are polar. CO2 is symmetrical. Positive C atom is reason CO2 + H2O gives -0.75 +1.5 -0.75 H2CO3 30 Polar or Nonpolar? Consider AB3 molecules: BF3, Cl2CO, and NH3.  31 Is Methane, CH4, Polar? H C H H H Methane is symmetrical and is NOT polar.