CHM012 Chemistry for Engineers PDF

CHM012 Chemistry for Engineers TOPIC 2 Chemical Bonding  Types of Chemical Bond ⮚ Chemical bond – force of attraction that binds atom together in a given molecule ⮚ Types: 1. Covalent – between two non-metal atoms (ex. CH4) 2. Ionic – between...

CHM012 Chemistry for Engineers TOPIC 2 Chemical Bonding  Types of Chemical Bond ⮚ Chemical bond – force of attraction that binds atom together in a given molecule ⮚ Types: 1. Covalent – between two non-metal atoms (ex. CH4) 2. Ionic – between a metal and a non-metal atoms (ex. KCl) 3. Metallic – between metal atoms ✔ Space-filling model: ⮚ Molecules can be represented using: Indicates the relative sizes of the atoms as well as their relative orientation in ✔ Chemical Formula CH4 the molecule ✔ Ball and stick method: A three-dimensional model ✔ Structural Formula using spheres and rods  Ways of Gaining Stability ”Atoms interact to gain stability by changing the outermost configuration to become isoelectronic to a noble gas.” - G.N. Lewis (1916) - 1. Sharing electrons (Covalent bonding) 2. Transferring electrons (Ionic Bonding) ✔ Atoms that form covalent bonds are non-metals or elements that has high ionization energy or ionization potential. ✔ Ionic bonds occurs between an atom that easily loses electrons (low ionization energy) and an atom that has a high affinity for electrons  Lewis Symbol ⮚ G. N. Lewis developed a method to denote potential bonding electrons by using one dot for every valence electron around the element symbol. ⮚ When forming compounds, atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (the octet rule).  Ionic Bond Formation ✔ Atoms tend to lose (metals) or gain (nonmetals) electrons to make them isoelectronic to the noble gases.  Energetics of Ionic Bonding – Born-Haber Cycle ✔ Many factors affect the energy of ionic bonding. ✔ Start with the metal and nonmetal elements: Na(s) and Cl2(g). ✔ Make gaseous atoms: Na(g) and Cl(g). ✔ Make ions: Na+(g) and Cl–(g). ✔ Combine the ions: NaCl(s). We already discussed making ions (ionization energy and electron affinity). ✔ It takes energy to convert the elements to atoms. (endothermic) ✔ It takes energy to create a cation (endothermic). ✔ Energy is released by making the anion (exothermic). ✔ The formation of the solid releases a huge amount of energy (exothermic). ✔ This makes the formation of salts from the elements exothermic.  Lattice Energy ⮚ That huge, exothermic transition is the reverse of the lattice energy, - the energy required to completely separate a mole of a solid ionic compound into its gaseous ions. ⮚ The energy associated with electrostatic interactions is governed by Coulomb’s law ⮚ Lattice energy increases with: – increasing charge on the ions – decreasing size of ions  Covalent Bond Formation ⮚ In covalent bonds, atoms share electrons. ⮚ There are several electrostatic interactions in these bonds: ✔ attractions between electrons and nuclei, ✔ repulsions between electrons, and ✔ repulsions between nuclei. ⮚ For a bond to form, the attractions must be greater than the repulsions.  Types of Covalent Bonds 1. Based on donor atom a. Normal covalent bond – a bond formed when each atom donates an electron. b. Dative or coordinate covalent bond – only one (1) atom donates the electron pair to form a bond.  Types of Covalent Bonds 2. Based on electronegativity of the bonded atoms. a. Non-polar covalent bond – when atoms bonded has almost the same electronegativities. H – H (H2) , O = O (O2) , C – H (CH4) b. Polar covalent bond – having different electronegativities.  Electronegativity ⮚ Electronegativity is the ability of an atom in a molecule to attract electrons to itself. ⮚ On the periodic table, electronegativity generally increases as you go ✔ from left to right across a period. ✔ from the bottom to the top of a group.  Polar Covalent Bonds ✔ The electrons in a covalent bond are not always shared equally. ✔ Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. ✔ Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.  Polar Covalent Bonds ⮚ When two atoms share electrons unequally, a polar covalent bond results. ⮚ Electrons tend to spend more time around the more electronegative atom. The result is a partial negative charge (not a complete transfer of charge). It is represented by δ–. The other atom is “more positive,” or δ+. ⮚ The greater the difference in electronegativity, the more polar is the bond. Points to Note  Lewis Structure Basic concepts ⮚ Duet, octet, expanded octet ⮚ Covalency – the number of bonds formed by an atom in forming a molecule. normal covalency – the number of electrons needed by an atom to have an octet of electrons.  Lewis Structure Simple guidelines in Lewis Structure writing: 1. Determine the total number of valence electrons e.g. SO2 : 6 x 3 = 18 e- SO32- : (6 x 4) + 2 = 26 e- NH4+ : (5 + 4) -1 = 8 e- 2. Determine the central atom. ⮚ the central atom must have the highest normal covalency. ⮚ if the same normal covalency, the less electronegative atom is the central atom. 3. Write the skeletal structure using single bonds to join two (2) atoms. 4. Count the number of electrons for each atoms: ⮚ 2nd period elements must have octet. ⮚ 3rd period and above can have expanded octet. ⮚ All outer atoms obey the octet rule.  Lewis Structure 5. Introduce multiple bond whenever the octet rule is not satisfied. 6. If the substance has a charge, enclose the Lewis Structure with a bracket [ ] and indicate the charge. e.g. 1+, 2+, 3+, 3-, etc. 7. For an AXX molecule, where all atoms obey the octet rule, use Langmuir formula to determine the number of covalent bonds formed. example: NO2- Total No. of e- : 5 + 2 (6) + 1 = 18 e-  Exceptions to the Octet Rule ✔ Boron tends to form compounds in which the boron atom has fewer than eight electrons around it. ✔ It does not have a complete octet ✔ The Lewis structure of boron fluoride ✔ Boron has only 6 electrons around it ✔ Adding a double bond satisfies the octet rule ❖ In stable compounds such as H3NBF4, boron has an octet of electrons  Exceptions: Expanded Octet ✔ Observed only in elements in Period 3 of the periodic table and beyond ✔ Consider the Lewis structure of sulfur hexafluoride: 6 + 6 (7) = 48 electrons ✔ When it is necessary to exceed the octet rule for one of several third-row (or higher) elements, assume that the extra electrons should be placed on the central atom ✔ Calculating the valence electrons present in the triiodide ion (I3−) 3(7) + 1 = 22 valence electrons  Importance of Lewis Structure 1. Determination of formal charges of atoms. 2. Determination of geometry of molecules or ions.  Formal Charge ⮚ Formal Charge - it is an indication of the charge an atom would carry if it shared the electrons in a covalent bond equally. Assumptions: 1. There is equal sharing of electron pair in a given bond. 2. 100% covalent character 3. Polarity of the bond is NOT considered. 4. It does not represent the actual charge distribution. 5. It is dependent on the Lewis Structure. G = Group no. U = No. of unshared e- C = No. of covalent bonds  Formal Charge ⮚ Guidelines in determining FC: 1. For atoms obeying the octet rule a. In a given Lewis Structure, if an atom exhibits its normal covalency, then FC = 0 b. If the atom exceeded its normal covalency then its FC is positive (+) c. If the atom has less no. of bonds compared to its normal covalency, then its FC is negative (-) 2. The sum of all the FC must be equal to the net charge of the ion. If there is no charge, then the sum is equal to zero (molecule).  Example ⮚ Example: SO3 No. of e- = 24 e- - [ ] ⮚ Example: ClO4- No. of e- = 32 e- Importance of Formal Charge  Resonance ⮚ Some molecules or ions can be represented by several Lewis structures/valence bond formulas. Note: Localized bonds – refers to single bonds Delocalized bonds – refers to multiple bonds e.g. CO32-  Resonance Guidelines on determining important Contributing Structures, C.S. 1. C.S. must have the same atomic positions, they may only differ on the positions of e- (when writing C.S. do not change the relative position of the atomic nuclei.) Resonance  Resonance 4. Position of formal charges distribution ✔ Adjacent like charge is unfavorable (less important) ✔ Adjacent opposite charge is favorable (more important) ✔ Important ✔ Important ✔ Reasonable ✔ Adjacent opposite distribution of FC. ✔ Adjacent charge opposite charge ✔ Less important due to adjacent like charge  Resonance 5. Electroneutrality principle ✔ A C.S. where all atoms have zero formal charge will be the most stable, therefore more important. ✔ Important ✔ Most important ✔ Less important ✔ Two atoms has zero ✔ Reasonable charge ✔ One atoms has zero formal charge distribution formal charge ✔ Two atoms has zero formal charge  Resonance 6. The greater the number of covalent bonds, the greater the importance of the C.S. increasing importance of contributing structure