2 9.1 The Covalent BondCovalent Bond – A chemical bond that results from the sharing of valence electrons.Between nonmetallic elements of similar electronegativity.A group of atoms united by a covalent bond is called a molecule.A substance made up of molecules is called a molecular substance.Molecules may have as few as two atoms. (CO), (O2)
4 Polar Covalent Bonds: Unevenly matched, but willing to share.
5 To describe the composition of a molecular compound we use a molecular formula. The molecular formula tells you how many atoms are in a single molecule of the compound.The molecular formula for oxygen is O2, which tells you that this molecule contains 2 atoms of oxygen.
6 The molecular formula for sucrose (table sugar) is C12H22O11 The molecular formula for sucrose (table sugar) is C12H22O11. This indicates that one sucrose molecule contains 12 carbon atoms, 22 hydrogen atoms and 11 oxygen atoms.The empirical formula shows the ratio of atoms in the molecule. You get it by dividing all subscripts by the smallest subscript and rounding. C6H12O6 becomes CH2O.
7 Molecules are often shown using a structural formula. A structural formula specifies which atoms are bonded to each other in a molecule.One type is the Lewis structures.The Lewis structure uses dots as before and has a circle drawn around the dots that represent valence electrons that are joined in a covalent bond.
8 The principle that describes covalent bonding is the Octet rule. Atoms within a molecule will share electrons so as to acquire a full set (8) of valence electrons.If each atom has 7 valence electrons, than they will share one so that they both will have 8. The sharing of that electron is a Single covalent bond.
9 Single Covalent bonding Fluorine has seven valence electronsA second atom also has sevenBy sharing electronsBoth end with full orbitalsFF
10 F F Covalent bonding Fluorine has seven valence electrons A second atom also has sevenBy sharing electronsBoth end with full orbitalsFF8 Valence electrons
11 F F Covalent bonding Fluorine has seven valence electrons A second atom also has sevenBy sharing electronsBoth end with full orbitalsFF8 Valence electrons
12 If more than two atoms join in a covalent bond the atoms with the fewest valence electrons will pair up with the atom that has the most up to a total of 8.If a pair of electrons is not shared they are called an unshared pair.If two atoms share two pairs of electrons it is called a double covalent bond.If two atoms share three pairs of electrons it is called a triple covalent bond.
13 The Lewis structures that you have seen so far use a pair of dots to represent covalent bonds. Another kind of Lewis structure uses dashes to represent covalent bonds.A single dash represents a single covalent bond.A double dash represents a double covalent bondA triple dash represents a triple covalent bond.
14 Covalent compounds (molecules) tend to have low melting and boiling points When an electron pair is shared by the direct overlap of bonding orbitals, a sigma bond results. Single bonds are sigma bonds.The overlap of parallel orbitals forms a pi bond. Multiple bonds involve both sigma and pi bonds.
15 Bond length depends on the sizes of the bonded atoms and the number of electron pairs they share. Bond dissociation energy is the energy needed to break a covalent bond.Double bonds have larger bond dissociation energies than single. Triple even largerBond length and bond dissociation energy are directly related.
16 9.2 Naming MoleculesCovalent compounds (Molecules) are named by adding prefixes to the element names.Covalent’ (in this context) means both elements are nonmetals.The names of molecular compounds will include prefixes that indicate the number of atoms in the molecule. (CO2 would be called Carbon Dioxide, CCl4 is called Carbon Tetrachloride)The following prefixes are used:mono-1 di-2tri-3 tetra-4penta-5 hexa-6hepta-7 octa-8nona-9 deca-10Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.
17 Remember to use the “ide” suffix. (Cl is chloride not chlorine) Exceptions to the rule:The prefix “mono” is usually not written with the first word of the compounds name.CO2 is not called monocarbon dioxide, just carbon dioxide.Some compounds have common names. (H2O is water or dihydrogen monoxide)
18 Naming Binary Covalent Compounds: Examples 1mono2di3tri4tetra5penta6hexa7heptaa8octa9nona10deca* Second element in ‘ide’ from* Drop –a & -o before ‘oxide’N2S4dinitrogen tetrasulfideNI3nitrogen triiodideXeF6xenon hexafluorideCCl4carbon tetrachlorideP2O5diphosphorus pentoxideSO3sulfur trioxide
19 Naming AcidsAn acid is a molecular substance that dissolves in water to produce hydrogen ions (H+).Acids behave like an ionic compound because when they dissolve in water, they create cations and anions.The cation is always H+ and the anion depends on the particular acid.Example: When the acid HCL dissolves in water it forms H+ cations and CL- anions.HNO3 acid will form H+ and NO3- ions.
20 Naming AcidsThe name of an acid usually results from the name of the anion(-).For acids, where the anion ends with the suffix “ide”(Chloride) the name of the acid begins with the prefix hydro.The acids name also includes the root name of the anion(-) and the word acid.In addition you need to change the suffix of the anion from “ide” to “ic”.By these rules HCL is named hydrochloric acid.Other acids are named without the prefix “hydro”. The acid HNO3- derives its name from NO3- which is nitrate. HNO3 is named Nitric Acid.The following are names and formulas of common acids:
22 9.3 Molecular StructuresStructural formulas (Lewis Structures) use letter symbols and bonds to show the position of atoms.Drawing Lewis Structures:
23 H O Water Each hydrogen has 1 valence electron and wants 1 more The oxygen has 6 valence electronsand wants 2 moreThey share to make each other “happy”HO
24 H O Water Put the pieces together The first hydrogen is happy The oxygen still wants one moreHO
25 H O H Water The second hydrogen attaches Every atom has full energy levelsHOH
26 Multiple BondsSometimes atoms share more than one pair of valence electrons.A double bond is when atoms share two pair (4) of electrons.A triple bond is when atoms share three pair (6) of electrons.
27 C O Carbon dioxide CO2 - Carbon is central atom ( I have to tell you) Carbon has 4 valence electronsWants 4 moreOxygen has 6 valence electronsWants 2 moreCO
28 Carbon dioxideAttaching 1 oxygen leaves the oxygen 1 short and the carbon 3 shortCO
29 Carbon dioxideAttaching the second oxygen leaves both oxygen 1 short and the carbon 2 shortOCO
30 Carbon dioxideThe only solution is to share moreOCO
31 Carbon dioxideThe only solution is to share moreOCO
32 Carbon dioxideThe only solution is to share moreOCO
33 Carbon dioxideThe only solution is to share moreOCO
34 Carbon dioxideThe only solution is to share moreOCO
35 Carbon dioxideThe only solution is to share moreOCO
36 O C O Carbon dioxide The only solution is to share more Requires two double bondsEach atom gets to count all the atoms in the bondOCO
37 O C O Carbon dioxide The only solution is to share more Requires two double bondsEach atom gets to count all the atoms in the bond8 valence electronsOCO
38 O C O Carbon dioxide The only solution is to share more Requires two double bondsEach atom gets to count all the atoms in the bond8 valence electronsOCO
39 O C O Carbon dioxide The only solution is to share more Requires two double bondsEach atom gets to count all the atoms in the bond8 valence electronsOCO
40 How to draw them To figure out if you need multiple bonds Add up all the valence electrons.Count up the total number of electrons to make all atoms happy.Subtract.Divide by 2Tells you how many bonds - draw them.Fill in the rest of the valence electrons to fill atoms up.
41 N H Examples NH3 N - has 5 valence electrons wants 8 H - has 1 valence electrons wants 2NH3 has 5+3(1) = 8NH3 wants 8+3(2) = 14(14-8)/2= 3 bonds4 atoms with 3 bondsNH
42 H H N H Examples Draw in the bonds All 8 electrons are accounted for Everything is fullHHNH
43 Examples HCN C is central atom N - has 5 valence electrons wants 8 C - has 4 valence electrons wants 8H - has 1 valence electrons wants 2HCN has = 10HCN wants = 18(18-10)/2= 4 bonds3 atoms with 4 bonds -will require multiple bonds - not to H
44 H C N HCN Put in single bonds Need 2 more bonds Must go between C and NHCN
45 H C N HCN Put in single bonds Need 2 more bonds Must go between C and NUses 8 electrons - 2 more to addHCN
46 H C N HCN Put in single bonds Need 2 more bonds Must go between C and NUses 8 electrons - 2 more to addMust go on N to fill octetHCN
47 Another way of indicating bonds Often use a line to indicate a bondCalled a structural formulaEach line is 2 valence electronsHOHHOH=
48 H C N H C O H Structural Examples C has 8 electrons because each line is 2 electronsDitto for NDitto for C hereDitto for OH C NHC OH
49 Coordinate Covalent Bond A Coordinate Covalent Bond is when one atom donates both electrons in a covalent bond.Carbon monoxideCOOC
50 Coordinate Covalent Bond When one atom donates both electrons in a covalent bond.Carbon monoxideCOCO
51 Coordinate Covalent Bond When one atom donates both electrons in a covalent bond.Carbon monoxideCOCOOC
52 ResonanceResonance is a condition when more than one dot diagram with the same connections is possible.Resonance structures differ only in the position of the electron pairs, never the atom positions.
53 9.4 Molecular Shape (VSEPR) Valence Shell Electron Pair Repulsion.Predicts three dimensional geometry of molecules.The name tells you the theory.Valence shell - outside electrons.Electron Pair repulsion - electron pairs try to get as far away as possible.Can determine the angles of bonds.And the shape of molecules
54 VSEPRBased on the number of pairs of valence electrons both bonded and unbonded.Unbonded pair are called lone pair.CH4 - draw the structural formula
55 H H C H H VSEPR Single bonds fill all atoms. There are 4 pairs of electrons pushing away.The furthest they can get away is 109.5º.HHCHH
56 H C H H H 4 atoms bonded Basic shape is tetrahedral. A pyramid with a triangular base.Same basic shape for everything with 4 pairs.H109.5ºCHHH
57 N H N H H H H H 3 bonded - 1 lone pair Still basic tetrahedral but you can’t see the electron pair.Shape is called trigonal pyramidal.NHNHHH<109.5ºHH
58 O H O H H H 2 bonded - 2 lone pair Still basic tetrahedral but you can’t see the 2 lone pair.Shape is called bent.OHOH<109.5ºHH
59 C H O H C O H 3 atoms no lone pair The farthest you can the electron pair apart is 120º.Shape is flat and called trigonal planar.Will require 1 double bondCHO120ºHCOH
60 2 atoms no lone pairWith three atoms the farthest they can get apart is 180º.Shape called linear.Will require 2 double bonds or one triple bondCO180º
61 Molecular OrbitalsThe overlap of atomic orbitals from separate atoms makes molecular orbitalsEach molecular orbital has room for two electronsTwo types of MOSigma ( σ ) between atomsPi ( π ) above and below atoms
62 Sigma bonding orbitals From s orbitals on separate atoms++ ++++Sigma bonding molecular orbitals orbitals orbital
63 Sigma bonding orbitals From p orbitals on separate atomsp orbitalp orbitalSigma bonding molecular orbital
64 Pi bonding molecular orbital Pi bonding orbitalsP orbitals on separate atomsPi bonding molecular orbital
65 Sigma and pi bonds All single bonds are sigma bonds A double bond is one sigma and one pi bondA triple bond is one sigma and two pi bonds.
66 HybridizationThe mixing of several atomic orbitals to form the same number of hybrid orbitals.When an atom approaches another atom to form a bond, the orbitals of its electrons may be changed.The changing of the orbitals forms what is called hybrid orbitals. This is a cross between all of the orbitals involved (s,p,d,f).
67 HybridizationWhen an “s” and “p” orbital come together they form a linear shaped molecule.When an “s” and “p2” orbital come together they form a trigonal planer molecule.When an “s” and “p3” orbital come together they form tetrahedral, pyramidal or bent molecules.Atoms with higher orbital numbers form longer bonds. Multiple bonds are shorter than single bonds because bonds with more electrons attract the nuclei of the bonding atoms more strongly.
68 9.5 Electronegativity and Bond Type DifferenceBond Type<.4Non-Polar Covalent0.5 – 1.9Polar Covalent>2.0Ionic
69 Polar BondsWhen the atoms in a bond are the same, the electrons are shared equally.This is a nonpolar covalent bond.When two different atoms are connected, the electrons may not be shared equally.This is a polar covalent bond.
70 ElectronegativityA measure of how strongly the atoms attract electrons in a bond.The bigger the electronegativity difference the more polar the bond.Covalent nonpolarCovalent moderately polarCovalent polar>2.0 Ionic
71 How to show a bond is polar Isn’t a whole charge just a partial charged+ means a partially positived- means a partially negativeThe Cl pulls harder on the electronsThe electrons spend more time near the Cld+d-HCl
72 Polar MoleculesPolar molecules have a partially positive end and a partially negative endRequires two things to be true:The molecule must contain polar bonds.This can be determined from differences in electronegativity.Symmetry can not cancel out the effects of the polar bonds.Must determine geometry first.
73 Polar MoleculesSymmetrical shapes are those without lone pair on central atomTetrahedralTrigonal planarLinearWill be nonpolar if all the atoms are the sameShapes with lone pair on central atom are not symmetricalCan be polar even with the same atom
74 Intermolecular Forces They are what make solid and liquid molecular compounds possible.The weakest are called van der Waal’s forces - there are two kindsDispersion forcesDipole Interactions
75 Dispersion Force F2 is a gas Br2 is a liquid I2 is a solid Depends only on the number of electrons in the moleculeBigger molecules more electronsMore electrons stronger forcesF2 is a gasBr2 is a liquidI2 is a solid
76 Dipole interactionsOccur when polar molecules are attracted to each other.Slightly stronger than dispersion forces.Opposites attract but not completely hooked like in ionic solids.
77 Dipole interactions H F d+ d- H F d+ d- Occur when polar molecules are attracted to each other.Slightly stronger than dispersion forces.Opposites attract but not completely hooked like in ionic solids.H Fd+ d-H Fd+ d-
78 Properties of Molecular Compounds Made of nonmetalsPoor or nonconducting as solid, liquid or aqueous solutionLow melting pointTwo kinds of crystalsMolecular solids held together by IMFNetwork solids- held together by bondsOne big molecule (diamond, graphite)