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Ch 9 Covalent Bonding
 

Ch 9 Covalent Bonding

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    Ch 9 Covalent Bonding Ch 9 Covalent Bonding Presentation Transcript

    • Chapter 9 Covalent Bonding
    • 9.1 The Covalent Bond
      • Covalent 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), (O 2 )
    • Covalent Bonds
    • Polar Covalent Bonds: Unevenly matched, but willing to share.
      • 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 O 2 , which tells you that this molecule contains 2 atoms of oxygen.
      • The molecular formula for sucrose (table sugar) is C 12 H 22 O 11 . 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. C 6 H 12 O 6 becomes CH 2 O.
      • 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 structure.
        • 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.
      • 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, then they will share one so that they both will have 8. The sharing of that electron is a single covalent bond .
    • Single Covalent bonding
      • Fluorine has seven valence electrons
      • A second F atom also has seven
      • By sharing electrons
      • Both end with full orbitals
      F F
    • Covalent bonding
      • Fluorine has seven valence electrons
      • A second atom also has seven
      • By sharing electrons
      • Both end with full orbitals
      F F 8 Valence electrons
    • Covalent bonding
      • Fluorine has seven valence electrons
      • A second atom also has seven
      • By sharing electrons
      • Both end with full orbitals
      F F 8 Valence electrons
      • 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.
      • 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 bond
        • A triple dash represents a triple covalent bond.
      • 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.
      • 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 larger
      • Bond length and bond dissociation energy are directly related.
    • 9.2 Naming Molecules
        • Covalent 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. (CO 2 would be called Carbon Di oxide, CCl 4 is called Carbon Tetra chloride)
        • The following prefixes are used:
          • mono-1 di-2
          • tri-3 tetra-4
          • penta-5 hexa-6
          • hepta-7 octa-8
          • nona-9 deca-10
      • Note : When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.
        • Remember to use the “ide” suffix. (Cl is chlor ide not chlorine)
        • Exceptions to the rule :
          • The prefix “mono” is usually not written with the first word of the compounds name.
            • CO 2 is not called mono carbon di oxide, just carbon dioxide.
          • Some compounds have common names. (H 2 O is water or dihydrogen monoxide)
    • Naming Binary Covalent Compounds: Examples N 2 S 4 di nitrogen tetra sulfide NI 3 nitrogen tri iodide XeF 6 xenon hexa fluoride CCl 4 carbon tetra chloride P 2 O 5 di phosphorus pent oxide SO 3 sulfur tri oxide 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 heptaa 8 octa 9 nona 10 deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’
    • Naming Acids
      • An 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.
          • HNO 3 acid will form H+ and NO 3 - ions.
    • Naming Acids
      • The name of an acid usually results from the name of the anion(-).
        • For acids, where the anion ends with the suffix “-ide” (ex.Chloride), mainly binary acids, 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 HNO 3 - derives its name from NO 3 - which is nitrate. HNO 3 is named Nitric Acid.
      • The following are names and formulas of common acids:
    • Naming Acids
      • Anion Corresponding Acid
      • F - , Fluoride HF, hydrofluoric acid
      • Cl - , Chloride HCl, hydrochloric acid
      • Br - , Bromine HBr, hydrobromic acid
      • I - , iodide HI, hydroiodic acid
      • S 2 - , sulfide H 2 S, hydrosulfuric acid
      • NO 3 - , nitrate HNO 3 , nitric acid
      • CO 3 -2 , carbonate H 2 CO 3 , carbonic acid
      • SO 4 -2 , sulfate H 2 SO 4 , sulfuric acid
      • PO 4 -3 , phoshate H 3 PO 4 , Phosphoric acid
      • C 2 H 3 O 2 - , acetate HC 2 H 3 O 2 , acetic acid
    • 9.3 Molecular Structures
      • Structural formulas (Lewis Structures) use letter symbols and bonds to show the position of atoms.
      • Drawing Lewis Structures:
    • Water
      • Each hydrogen has 1 valence electron
      • and wants 1 more
      • The oxygen has 6 valence electrons
      • and wants 2 more
      • They share to make each other “happy”
      H O
    • Water
      • Put the pieces together
      • The first hydrogen is happy
      • The oxygen still wants one more
      H O
    • Water
      • The second hydrogen attaches
      • Every atom has full energy levels
      H H O
    • Multiple Bonds
      • Sometimes 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.
    • Carbon dioxide
      • CO 2 - Carbon is central atom ( I have to tell you)
      • Carbon has 4 valence electrons
      • Wants 4 more
      • Oxygen has 6 valence electrons
      • Wants 2 more
      C O
    • Carbon dioxide
      • Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short
      C O
    • Carbon dioxide
      • Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short
      C O O
    • Carbon dioxide
      • The only solution is to share more
      C O O
    • Carbon dioxide
      • The only solution is to share more
      C O O
    • Carbon dioxide
      • The only solution is to share more
      C O O
    • Carbon dioxide
      • The only solution is to share more
      C O O
    • Carbon dioxide
      • The only solution is to share more
      C O O
    • Carbon dioxide
      • The only solution is to share more
      O C O
    • Carbon dioxide
      • The only solution is to share more
      • Requires two double bonds
      • Each atom gets to count all the atoms in the bond
      O C O
    • Carbon dioxide
      • The only solution is to share more
      • Requires two double bonds
      • Each atom gets to count all the atoms in the bond
      O C O 8 valence electrons
    • Carbon dioxide
      • The only solution is to share more
      • Requires two double bonds
      • Each atom gets to count all the atoms in the bond
      O C O 8 valence electrons
    • Carbon dioxide
      • The only solution is to share more
      • Requires two double bonds
      • Each atom gets to count all the atoms in the bond
      O C O 8 valence electrons
    • 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 2
      • Tells you how many bonds - draw them.
      • Fill in the rest of the valence electrons to fill atoms up.
    • Examples
      • NH 3
      • N - has 5 valence electrons wants 8
      • H - has 1 valence electrons wants 2
      • NH 3 has 5+3(1) = 8
      • NH 3 wants 8+3(2) = 14
      • (14-8)/2= 3 bonds
      • 4 atoms with 3 bonds
      N H
    • Examples
      • Draw in the bonds
      • All 8 electrons are accounted for
      • Everything is full
      N H H H
    • Examples
      • HCN C is central atom
      • N - has 5 valence electrons wants 8
      • C - has 4 valence electrons wants 8
      • H - has 1 valence electrons wants 2
      • HCN has 5+4+1 = 10
      • HCN wants 8+8+2 = 18
      • (18-10)/2= 4 bonds
      • 3 atoms with 4 bonds -will require multiple bonds - not to H
    • HCN
      • Put in single bonds
      • Need 2 more bonds
      • Must go between C and N
      N H C
    • HCN
      • Put in single bonds
      • Need 2 more bonds
      • Must go between C and N
      • Uses 8 electrons - 2 more to add
      N H C
    • HCN
      • Put in single bonds
      • Need 2 more bonds
      • Must go between C and N
      • Uses 8 electrons - 2 more to add
      • Must go on N to fill octet
      N H C
    • Another way of indicating bonds
      • Often use a line to indicate a bond
      • Called a structural formula
      • Each line is 2 valence electrons
      H H O = H H O
    • Structural Examples
      • C has 8 electrons because each line is 2 electrons
      • Ditto for N
      • Ditto for C here
      • Ditto for O
      H C N C O H H
    • Coordinate Covalent Bond
      • A Coordinate Covalent Bond is when one atom donates both electrons in a covalent bond.
      • Carbon monoxide
      • CO
      O C
    • Coordinate Covalent Bond
      • When one atom donates both electrons in a covalent bond.
      • Carbon monoxide
      • CO
      O C
    • Coordinate Covalent Bond
      • When one atom donates both electrons in a covalent bond.
      • Carbon monoxide
      • CO
      O C O C
    • Resonance
      • Resonance 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.
    • 9.4 Molecular Shape (VSEPR)
      • V alence S hell E lectron P air R epulsion.
      • 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
    • VSEPR
      • Based on the number of pairs of valence electrons both bonded and unbonded.
      • Unbonded pair are called lone pair .
      • CH 4 - draw the structural formula
    • VSEPR
      • Single bonds fill all atoms.
      • There are 4 pairs of electrons pushing away.
      • The furthest they can get away is 109.5º.
      C H H H H
    • 4 atoms bonded
      • Basic shape is tetrahedral.
      • A pyramid with a triangular base.
      • Same basic shape for everything with 4 pairs.
      C H H H H 109.5º
    • 3 bonded - 1 lone pair
      • Still basic tetrahedral but you can’t see the electron pair.
      • Shape is called trigonal pyramidal.
      N H H H N H H H <109.5º
    • 2 bonded - 2 lone pair
      • Still basic tetrahedral but you can’t see the 2 lone pair.
      • Shape is called bent.
      O H H O H H <109.5º
    • 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 bond
      C H H O C H H O 120º
    • 2 atoms no lone pair
      • With three atoms the farthest they can get apart is 180º.
      • Shape called linear.
      • Will require 2 double bonds or one triple bond
      C O O 180º
    • Molecular Orbitals
      • The overlap of atomic orbitals from separate atoms makes molecular orbitals
      • Each molecular orbital has room for two electrons
      • Two types of MO
        • Sigma ( σ ) between atoms
        • Pi ( π ) above and below atoms
    • Sigma bonding orbitals
      • From s orbitals on separate atoms
      + + s orbital s orbital + + + + Sigma bonding molecular orbital
    • Sigma bonding orbitals
      • From p orbitals on separate atoms
      p orbital p orbital Sigma bonding molecular orbital      
    • Pi bonding orbitals
      • P orbitals on separate atoms
          Pi bonding molecular orbital    
    • Sigma and pi bonds
      • All single bonds are sigma bonds
      • A double bond is one sigma and one pi bond
      • A triple bond is one sigma and two pi bonds.
    • Hybridization
      • The 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).
    • Hybridization
      • When 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.
    • 9.5 Electronegativity and Bond Type Electro negativity Difference Bond Type <.4 Non-Polar Covalent 0.5 – 1.9 Polar Covalent >2.0 Ionic
    • Polar Bonds
      • When 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 .
    • Electronegativity
      • A measure of how strongly the atoms attract electrons in a bond.
      • The bigger the electronegativity difference the more polar the bond.
      • 0.0 - 0.4 Covalent nonpolar
      • 0.5 - 1.0 Covalent moderately polar
      • 1.0 -2.0 Covalent polar
      • >2.0 Ionic
    • How to show a bond is polar
      • Isn’t a whole charge just a partial charge
      •  means a partially positive
      •  means a partially negative
      • The Cl pulls harder on the electrons
      • The electrons spend more time near the Cl
      H Cl  
    • Polar Molecules
      • Polar molecules have a partially positive end and a partially negative end
      • Requires 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.
    • Polar Molecules
      • Symmetrical shapes are those without lone pair on central atom
        • Tetrahedral
        • Trigonal planar
        • Linear
      • Will be nonpolar if all the atoms are the same
      • Shapes with lone pair on central atom are not symmetrical
      • Can be polar even with the same atom
    • Intermolecular Forces
      • They are what make solid and liquid molecular compounds possible.
      • The weakest are called van der Waal’s forces - there are two kinds
        • Dispersion forces
        • Dipole Interactions
    • Dispersion Force
      • Depends only on the number of electrons in the molecule
      • Bigger molecules more electrons
      • More electrons stronger forces
          • F 2 is a gas
          • Br 2 is a liquid
          • I 2 is a solid
    • Dipole interactions
      • Occur when polar molecules are attracted to each other.
      • Slightly stronger than dispersion forces.
      • Opposites attract but not completely hooked like in ionic solids.
    • Dipole interactions
      • 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 F     H F    
    • Properties of Molecular Compounds
      • Made of nonmetals
      • Poor or nonconducting as solid, liquid or aqueous solution
      • Low melting point
      • Two kinds of crystals
        • Molecular solids held together by IMF
        • Network solids- held together by bonds
        • One big molecule (diamond, graphite)