2/14/13CHM 101 Ch 4: Covalent CompoundsCovalent Compounds Covalent bonds form when atoms share electrons to complete octets. Covalent bonds are typically between two nonmetal atoms.
Naming Covalent CompoundsBinary Covalent Compounds – molecule that contains atoms of only 2 elements CS2, SO3, PCl5, etc. General Rules for Naming 1. Element with lower group number is named first. 2. If elements are in the same group, element with higher period number is named first. 3. Second element named as root + -ide. 4. Greek prefixes are used to designate the number of atoms. You must know these!
Naming Covalent CompoundsWhat is the name of SO3?1. The first nonmetal is S sulfur.2. The second nonmetal is O named oxide.3. The subscript 3 of O is shown as the prefix tri. SO3 sulfur trioxide The subscript 1 (for S) or mono is understood.
Naming Covalent Compounds Give names for the following binary compounds: 1. PCl5 phosphorous pentachloride 2. CS2 carbon disulfide 3. P4S3 tetraphosphorous trisulfideGive molecular formulas for the following compounds:1. Dinitrogen monoxide N2 O2. Selenium hexafluoride SeF63. Dichlorine heptaoxide Cl2O7
Naming Covalent CompoundsUse the table below to quiz yourself. Use the formula, write the name. Usethe name, write the formula.
Covalent BondingMain group elements Usually make the number of bonds necessary to have noble gasconfigurationNoble gases (except He) have 8 valence electrons Octet Rule: Main group elements gain, lose, or share e- to achieve a stable e- configuration with 8 valence e- (except H and He – only need 2 valence e- for stability)
Covalent BondingLewis Dot Symbol - Representation of the number of valence electronsin an atom. Usually only used for main group elements. Maximumnumber is 8. X H He Li Be B C N O F Ne
Covalent Bonding Cl2 molecule Cl: 7 valence electrons, needs to make 1 bond to have octet Cl Lewis Dot SymbolLewis Dot Structure Lone pair of e- (unshared) Cl Cl Cl Cl Bonding pair of e- (shared)Lewis Dot Structure – showshow the atoms of a moleculeare connected. Shows lonepairs of electrons and bonding Cl Clpairs of electrons. Count the “shared” e- for both atoms
Covalent BondingMultiple Bonds – Double & Triple In carbon dioxide, CO2, the C atom shares 4 electrons with each O atom in a double bond.
Covalent BondingMultiple Bonds – Double & Triple In nitrogen molecule, N2, each N atom shares 6 electrons with the other N atom in a triple bond.
Covalent BondingWe can often predict the number of covalent bonds an atom will formbased on the number of valence electrons. # Valence e- # Bonds for Octet ExampleGroup 7(A) 7 1 HClGroup 6(A) 6 2 H2 OGroup 5(A) 5 3 NH3Group 4(A) 4 4 CH4 Cl H Cl O H O H
Covalent BondingRules for Drawing Lewis Dot Structures1. Count the total number of valence electrons in the molecule NF3 N: 5 e- F: 7e- x 3 = 21 e- 26 val e- total for the molecule2. Arrange atoms next to each other and connect with bonds Central atom is usually written first in the formula and has lower group number (If atoms are in same group, central atom is from higher period) 26 val e- F N F - 6 bonding e- F 20 remaining e-
Covalent Bonding Rules for Drawing Lewis Dot Structures 3. Place lone pairs around each atom to satisfy octet rule, starting with terminal atoms 20 remaining e- F N F -20 lone e- F 0 remaining e- Make sure each atom has an octet. If it doesn’t, check your work.Practice: Draw Lewis Structures for CH3Br and H2Se.
Covalent BondingRules for Drawing Lewis Dot StructuresIF there are electrons remaining…4. Place leftover e- on the central atom. It is ok to exceed the octet (have more than 8 electrons on an atom) if the central atom is in Period 3 or higher. SF4If the central atom has < 8 e-…5. Change single bonds to multiple bonds (double or triple) using lone pairs from terminal atoms. H2CO, CH3COOH, HCN
Covalent BondingDraw the Lewis structure for formaldehyde: H2CO1. Count the total number of valence electrons in the molecule H2CO H: 1 e- x 2 = 2 e- C: 4e- x 1 = 4 e- O: 6e- x 1 = 6 e- 12 val e- total for the molecule2. Place atoms relative to each other and connect with bonds H can only make 1 bond, so it has to be a terminal atom. C is a very common central atom (lower group # than O). H C H 12 val e- - 6 bonding e- O 6 remaining e-
Covalent BondingDraw the Lewis structure for formaldehyde: H2CO3. Place lone pairs around each atom to satisfy octet rule, starting with terminal atoms H can only accommodate 2 electrons, which it has with the bonding electrons. Never put lone pairs on H. H C H 6 val e- - 6 bonding e- O 0 remaining e-The central atom (carbon) has < 8 e-4. Change single bonds to multiple bonds (double or triple) using lone pairs from terminal atoms. H C H Carbon and oxygen now both have “octets”. O
Lewis Structures of IonsFor anions – add electrons to total valence equal to charge NO3– : N: 5e- x 1 = 5 O: 6e- x 3 = 18 + 1e- (b/c of ion charge) 24 e- total - 6 e- bonding O 18 e- - 18 e- lone pairs O N O 0 e- – Does each atom have an octet? O How can we give N an octet? O N O For ions, always put the Lewis structure in brackets and write the charge as superscript.
Lewis Structures of IonsFor cations – subtract electrons from total valence equal to charge NH4+ : N: 5e- x 1 = 5 H: 1e- x 4 = 4 - 1e- 8 e- total - 8 e- bonding + H 0 e- H N H H Examples for you to practice: ClO3–, ClF4+
Covalent BondingExceptions to the Octet Rule 1. Fewer than 8 electrons: Molecules with Be or B as central atom are often electron deficient. Be usually only needs 4 electrons for stability and B only needs 6. BeCl2 BF3 2. Odd # of valence electrons: NO2: Free radicals: atoms or molecules with unpaired electrons, highly reactive
Covalent BondingExceptions to the Octet Rule 3. More than 8 valence electrons: “expanded valence shells” - only allowed for atoms in Period 3 and beyond H2SO4 O H O S O H O Example for you to practice: PCl5
Covalent BondingResonance Structures Draw the Lewis Structure for ozone, O3: + + O O - - O O O O There are 2 possible structures only differing in the location of electrons and the double bond. Experimental data says that both bonds are identical. The actual structure of O3 is neither of them, but a composite of the two, called a resonance hybrid. Each Lewis structure is a resonance structure of O3.
Covalent Bonding Resonance Structures + + O O - - O O O O Use a double-headed arrow for resonance structures • Resonance structures differ only in the assignment of e- pair positions, not in atom positions.Draw resonance structures for NO3-:
Covalent BondingMolecular Shape – VSEPR Theory The properties of a molecule are heavily influenced by molecular shape.Valence-Shell Electron-Pair Repulsion (VSEPR) Theory Bonding and lone pair electrons surrounding a central atom repel each other. To minimize the repulsions, electron groups are oriented as far apart as possible. How many electron “groups” are on a central atom? Each of the following counts as one electron group: • Lone pair of e- • Single bond • Double bond • Triple bond
Covalent BondingMolecular Shape – VSEPR Theory
Covalent Bonding Ideal bond anglesElectron pair geometry – Arrangement of e- groups (bonds and lone pairs) around the central atom.Molecular geometry – Arrangement of atoms in space, shape of molecule. This is different than e- group geometry if lone pairs are present
Covalent BondingMolecular Shape – VSEPR TheoryIf there are 2 electron “groups” on the central atom… Linear electron pair geometry Linear molecular geometry 2 atoms attached to a central atom 180 bond angle Can’t have any lone pairs, so electron pair geometry is same as molecular geometry Examples: BeCl2, CO2, CS2, HCN
Covalent BondingMolecular Shape – VSEPR TheoryIf there are 3 electron “groups” on the central atom and 0 lone pairs… Trigonal Planar electron pair geometry Trigonal Planar molecular geometry 3 atoms attached to a central atom 120 bond angle If 0 lone pairs, electron pair geometry is same as molecular geometry Examples: BF3, SO3, NO3–, CO32–
Covalent Bonding Molecular Shape – VSEPR Theory If there are 4 electron “groups” on the central atom & 0 lone pairs… Tetrahedral electron pair geometry Tetrahedral molecular geometry 4 atoms attached to a central atom 109.5 bond angleElectron pair geometry If 0 lone pairs, electron pair geometry isMolecular geometry same as molecular geometryExample Examples: CH4, SiCl4, SO42–, ClO4–
Covalent Bonding Molecular Shape – VSEPR Theory If there are 4 electron “groups” on the central atom & 1 lone pair… Tetrahedral electron pair geometry Trigonal pyramidal molecular geometry 3 atoms & 1 lone pair attached to a central atom 109.5 bond angleElectron pair geometryMolecular geometryExample Examples: NH3, PF3, ClO3–, H3O+
Covalent Bonding Molecular Shape – VSEPR Theory If there are 4 electron “groups” on the central atom & 2 lone pairs… Tetrahedral electron pair geometry Bent (Angular) molecular geometry 2 atoms & 2 lone pairs attached to a central atom 109.5 bond angleElectron pair geometryMolecular geometryExample Examples: H2O, SCl2
Covalent BondingCovalent Bonds & Electronegativity Electrons are shared in covalent bonds, but they usually are not shared equally. One atom usually pulls the electrons more strongly than the other.The bond between H and Clin HCl is covalent.Cl pulls the shared electronsmore strongly than H.This creates a “partiallynegative” region around Cland a “partially positive”region around H. The bond between H and Cl is a polar covalent bond.
Covalent Bonding Covalent Bonds & Electronegativity The electronegativity value for an atom indicates the attraction of that atom for shared electrons (in covalent bonds). Electronegativity strengthusually increases as the size of an atom decreases. WHY???
Covalent BondingNonpolar Covalent Bonds A nonpolar covalent bond is an equal or almost equal sharing of electrons. The atoms involved have almost no electronegativity difference (0.0 to 0.4). Examples: Electronegativity Atoms Difference Type of Bond N-N 3.0 - 3.0 = 0.0 Nonpolar covalent Cl-Br 3.0 - 2.8 = 0.2 Nonpolar covalent H-Si 2.1 - 1.8 = 0.3 Nonpolar covalent H-C ??? ???
Covalent BondingNonpolar Covalent BondsDiatomic molecules are molecules that contain 2 atoms of the same element.They are the natural state for elements H, O, N, Cl, Br, I, and F.These are the only truly nonpolar covalent bonds.
Covalent BondingPolar Covalent BondsA polar covalent bond is an unequal sharing of electrons.The atoms involved have a moderate electronegativity difference (0.5 to 1.7). Examples: Electronegativity Atoms Difference Type of Bond O-Cl 3.5 - 3.0 = 0.5 Polar covalent Cl-C 3.0 - 2.5 = 0.5 Polar covalent O-S 3.5 - 2.5 = 1.0 Polar covalent
Covalent BondingIonic Bonds An ionic bond occurs between metal and nonmetal ions, and is a result of electron transfer from the metal to the nonmetal. There is a large electronegativity difference (1.8 or more) between the atoms. Examples: Electronegativity Atoms Difference Type of Bond Cl-K 3.0 – 0.8 = 2.2 Ionic N-Na 3.0 – 0.9 = 2.1 Ionic S-Cs 2.5 – 0.7 = 1.8 Ionic
Covalent BondingMolecular PolarityA covalent bond is polar if it connects atoms with different electronegativityvalues, i.e. H-Cl, C-F, C-Cl, etc. How do you know if a molecule is polar? - Bond Polarity - Molecular Shape
Covalent BondingMolecular Polarity Determine the molecular polarity of CO2: O C O Each C=O bond is polar, but the molecule is linear and the two dipoles cancel each other. Therefore, the CO2 molecule is nonpolar. COS (carbonyl sulfide): O C S C and S have the same EN. There is only 1 dipole, pointing towards O. Therefore, the COS molecule is polar.
Covalent BondingMolecular Polarity What about molecular shape? H2O: H O H Is the molecule nonpolar? NO! The H2O molecule is not linear, it is tetrahedral! O H H
Covalent BondingMolecular Polarity Indicate the dipole moment (if any) for CF4 and CHF3.
Covalent BondingMolecular Polarity Indicate the dipole moment (if any) for CH2Cl2.