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Chemistry 101 : Chap. 8 Basic Concepts of Chemical Bonding
Chemical Bonds, Lewis Symbols and the Octet Rule (2) Ionic Bonding (3) Covalent Bonding (4) Bond Polarity and Electronegativity (5) Drawing Lewis Structures (6) Resonance Structures (7) Exceptions to the Octet Rule (8) Strengths of Covalent Bonds
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Chemical Bonds Chemical bond is formed when two atoms or ions are held
together by the attractive force between them. Ionic Bond : a chemical bond formed between cation and anion Covalent Bond : a chemical bond formed between two nonmetallic atoms by sharing one or more pairs of electrons. Metallic Bond : a chemical bond formed when valence electrons of metal atom are attracted by the nuclei of surrounding atoms (electrons are free to move throughout the metal)
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Lewis Symbols Lewis electron dot structure (or Lewis symbol) : Symbol of element surrounded by dots representing the valence electrons in the atom Lewis symbol for sulfur : [Ne]3s23p4 S Maximum 2 electrons on each side This works only for representative elements (main group) Gilbert N. Lewis ( )
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Lewis Symbols Elements Group e- Configuration Lewis Symbol
Hydrogen A s H Helium A s He Lithium A [He]2s Li Berylium A [He]2s Be Boron A [He]2s22p B Carbon A [He]2s22p C
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Lewis Symbols O O = Elements Group e- Configuration Lewis Symbol
Nitrogen A [He]2s22p N Oxygen A [He]2s22p O Fluorine A [He]2s22p F Neon A [He]2s22p Ne Note : All four sides of the symbol are equivalent O = O
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Lewis Symbols Elements in the same group of periodic table have the same Lewis symbols F Cl Br I Elements in the same group have the same valence electron configurations For halogen atoms : ns2np5
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Octet Rule Only the valence electrons are involved in chemical bonding. Octet Rule : When forming chemical bond, atoms tend to gain, loose or share electrons in order to achieve a complete octet of valence electrons (ns2np6) same electron configuration as noble gas atom + K + Cl K+ Cl electron configuration: [Ar] [Ar] Both ions have an octet of electrons !
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Ionic bonding Ionic Bonding : Cations (metals) and anions (non-metal)
combine to form ionic bonds NaCl Alternating positive and negative charges
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Ionic bonding NaCl formation : Na(s) + ½ Cl2(g) NaCl (s) Hof = -490 kJ Metal : small ionization energy Na(g) Na+(g) + e IE = 496 kJ Non-metal : large electron affinity Cl(g) + e- Cl-(g) EA = -349 kJ Removing an electron from Na and transferring it to Cl is NOT exothermic ! Then, why NaCl formation is an exothermic process?
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Ionic bonding The main driving force to form ionic bonds is the electrostatic interaction between positive and negative ions. charges of ions distance between ions Strength of ionic bond depends on Eel the larger Eel, the stronger the bond the greater the charges, the stronger the bond the smaller the distance between the charges, the stronger the bond
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Ionic bonding The stronger the ionic bond the the melting point higher
66, 133 1261oC 66, 140 2852oC SrF2 +2, -2 113, 133 1473oC r1 r2
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Covalent bonding Covalent bond is formed when two atoms share electrons in order to achieve the electron configuration of the nearest noble gas. satisfy octet rule H + H Each hydrogen has the electron configuration of He H H Each fluorine has the electron configuration of Ne F + F F
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Covalent bonding Lewis dot structure for covalent bonds H + H H H
single covalent bond F + F F F F A shared electron pair is drawn as a dash (two bonding electrons) Unshared electrons are drawn as dots (lone-pair electrons)
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Covalent bonding Example : Draw the Lewis dot structures of H2O and NH3
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Covalent bonding Multiple bond F + F F F or F F Single bond O + C O
Double bond N + N N or Triple bond
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Covalent bonding Single and Multiple bond X X X X X X
Distance between atoms (bond length) decreases Bond strength increases X X X X
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Drawing Lewis Structure
Things to know before you start to draw Lewis structure Chemical formulas are often written in the order in which the atoms are connected ex) HCN Hydrogen has only two electrons (shared) and always has only one covalent bond The central atom is usually written first ex) NH3, CCl4, CHCl3, PCl3
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Drawing Lewis Structure
Rules for drawing Lewis structure (1) sum the number of valence electrons from all atoms (2) write the symbols for the atoms and connect them with a single bond (3) complete the "octet rule" for the atoms bonded to central atom (4) place any left over electrons on the central atom (5) If there are not enough electrons to give the central atom 8 electrons, try multiple bonds.
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Drawing Lewis Structure
Lewis Structure of NH3 (1) Total number of valence electrons = 1 = 8 (2) Connect atoms with a single bond H N H H and count the number electrons used for single bond = 6 (3) Complete the octets on the atoms bonded to the central atom : done (4) Place remaining electrons (8-6=2) on the central atom H N H H (5) All atoms are satisfying octet. No need to consider multiple bonds
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Drawing Lewis Structure
Lewis Structure of CO (1) Total number of valence electrons = =10 (2) Connect atoms with a single bond C O and count the number electrons used for single bond = 2 (3) Complete the octets on the atoms bonded to the central atom (6 electrons are used) C O (4) Place remaining electrons ( = 2) on the central atom C O (5) Carbon is NOT satisfying octet rule. Need to have multiple bonds C O C O
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Drawing Lewis Structure
Example : Determine the Lewis structure of HCN (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom. (5) Carbon is NOT satisfying octet rule. Need to have multiple bonds
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Drawing Lewis Structure
Example : Determined the Lewis structure of CH2O (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom. (5) Carbon is NOT satisfying octet rule. Need to have multiple bonds
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Drawing Lewis Structure
Example : Determined the Lewis structure of H2O2 (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom. (5) All atoms are satisfying octet. No need to consider multiple bonds What happens if you choose a different geometry in step (2)?
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Drawing Lewis Structure
Lewis Structure of ClO3- [ion] (1) Total number of valence electrons = 7 + 6 = 26 (2) Connect atoms with a single bond and count the number electrons used for single bond = 6 O Cl O O (3) Complete the octets on the atoms bonded to the central atom (18 electrons are used) O Cl O O (4) Place remaining electrons ( = 2) on the central atom O Cl O O O Cl O O (5) All atoms are satisfying octet. No need to consider multiple bonds
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Drawing Lewis Structure
Example : Determined the Lewis structure of ClO2- (1) Total number of valence electrons (2) Connect atoms with a single bond and count the number electrons used for single bond (3) Complete the octets on the atoms bonded to the central atom (4) Place remaining electrons on the central atom (5) All atoms are satisfying octet. No need to consider multiple bonds
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Drawing Lewis Structure : Exceptions
Atoms having fewer than 8 valence electrons : Group IIA and IIIA (mostly Be, B). Example = BeCl2 (1) Total number of valence electrons = 7 = 16 (2) Connect atoms with a single bond and count the number electrons used for single bond = 4 Cl Be Cl (3) Complete the octets on the atoms bonded to the central atom (12 electrons are used) Cl Be Cl (4) Place remaining electrons ( = 0) on the central atom : None left (5) Be is not satisfying the octet rule, but no electron is available:
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Drawing Lewis Structure : Exceptions
Atoms having more than 8 valence electrons : central atom with n 3, which can use d-orbitals for bonding Example = SF4 (1) Total number of valence electrons = 7 = 34 (2) Connect atoms with a single bond and count the number electrons used for single bond = 8 F F S F F F F (3) Complete the octets on the atoms bonded to the central atom (24 electrons are used) S F F (4) Place remaining electrons ( = 2) on the central atom F F (5) S is not satisfying the octet rule (10 electrons) S F F
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Drawing Lewis Structure : Exceptions
Molecule having an odd number of valence electrons : Example = NO2 (1) Total number of valence electrons = 2 = 17 (2) Connect atoms with a single bond and count the number electrons used for single bond = 4 O N O (3) Complete the octets on the atoms bonded to the central atom (12 electrons are used) O N O (4) Place remaining electrons ( = 1) on the central atom O N O (5) Nitrogen has only 5 electrons. Need to have multiple bonds O N O Free radical
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Drawing Lewis Structure : Exceptions
Example : Determine the Lewis structure of BF3, BrF5 and OH
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Drawing Lewis Structure : Resonance
Lewis Structure of SO3 (1) Total number of valence electrons = 6 = 24 (2) Connect atoms with a single bond and count the number electrons used for single bond = 6 O S O O O (3) Complete the octets on the atoms bonded to the central atom (18 electrons are used) S O O (4) Place remaining electrons on the central atom. No more electron is left ( =0) O O O (5) Sulfur is NOT satisfying octet rule. Need to have multiple bonds S S S O O O O O O All three S-O bonds have the same length Resonance structures
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Drawing Lewis Structure : Resonance
Example : Determine the Lewis structure of O3 and HCO2-
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Properties of Covalent Bond
Bond length : The distance between two bonded atoms bond length Bond length depends on the size of two atoms and the number of covalent bond (single, double or triple) between them.
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Properties of Covalent Bond
Example : Predict which member of each set would have the shortest bond length S S
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Properties of Covalent Bond
Bond Enthalpy: Energy required to completely separate two bonded atoms in gas phase. A short bond is usually harder to break. C (g) H (g) DH = 1660 kJ/mol (g) per C-H bond: C H (g) C (g) H (g) D (C-H) = 1660/4 kJ/mol = 415 kJ/mol
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Properties of Covalent Bond
Bond enthalpy can be used to estimate the enthalpy change of chemical reactions, Hrxn H2(g) + Cl2(g) 2HCl(g) H = ? H1 H2 Hrxn
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Properties of Covalent Bond
H1 = D(H-H) + D(Cl-Cl) = 436kJ/mol + 243kJ/mol = 679 kJ/mol H2 = 2 [ D(H-Cl) ] = 2 kJ/mol = kJ/mol Hrxn = H1 + H2 = 697kJ/mol – 862 kJ/mol = -183 kJ/mol DHorxn = Σ n x Dbroken – Σ m x Dformed moles of bonds
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Properties of Covalent Bond
+ + Hrxn bonds broken H – H Cl – Cl bonds formed H – Cl Bond Enthalpy (kJ/mol) H – H Cl – Cl H – Cl Hrxn = 1 243 – 2 431 = kJ/mol
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Properties of Covalent Bond
Example : Estimate the Hrxn of following reaction CH4 (g) O2 (g) → CO2 (g) H2O (g)
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Electronegativity Electronegativity : A measure of the attraction an atom has for the electron in a bond Metals low electronegativity Nonmetals high electronegativity electronegativity scale: Fluorine = 4 (most electronegative) most strongly attracting electron Cesium = 0.7 (least electronegative) most easily giving up electron Linus Carl Pauling ( )
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Electronegativity Pauling scale of electronegativity Element EN F 4.0
Cl N C H
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non-polar covalent bond:
Bond Polarity Nonpolar covalent bond When two atoms of same element are bonded together, there is equal sharing of the electrons in the bond Cl Cl non-polar covalent bond: equal sharing of electrons
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unequal sharing of electrons
Bond Polarity Polar covalent bond When two different elements are bonded together, there is unequal sharing of the electrons in the bond + - The bonding pair of electrons is pulled toward the chlorine atom (partial charge) H Cl polar covalent bond: unequal sharing of electrons
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Bond Polarity and Electronegativity
+ - Na+ Cl - H Cl ionic bond: electrons are not shared polar covalent bond: unequal sharing of electrons EN = 3.0 – 2.1 = 0.9 EN = 3.0 – 0.9 = 2.1 EN < 0.5 non-polar bond 0.5 EN < 2.0 polar bond EN 2.0 ionic bond
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Bond Polarity and Electronegativity
Example : For each pair of bonds, predict which bond is more polar and the partial charge on the atoms (a) Cl – Br Br – F (c) C – H C – O (b) O – F S – F (d) H – O Na – O
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