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Chapter 8 “Covalent Bonding”. Covalent Bonds. covalent co- (Latin - “together”) valere - “to be strong” 2 e- shared - strength holding 2 atoms together Smallest particle - “ molecule ”. Molecules. Some elements in nature are molecules : neutral group of atoms covalently bonded

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Chapter 8 covalent bonding

Chapter 8“Covalent Bonding”


Covalent bonds
Covalent Bonds

  • covalent

    • co- (Latin - “together”)

    • valere - “to be strong”

  • 2 e- shared - strength holding 2 atoms together

    • Smallest particle - “molecule”


  • Molecules
    Molecules

    • Some elements in nature are molecules:

      • neutral group of atomscovalently bonded

        • Ex. - air contains O molecules, 2 O atoms joined covalently

          • Called “diatomic molecule” (O2)


    How does h 2 form

    +

    +

    +

    +

    How does H2 form?

    (diatomic hydrogen molecule)

    • The nuclei repel each other, (both have + charge)


    How does h 2 form1

    +

    +

    How does H2 form?

    • nuclei attraction to e-’s stronger than repulsion of nuclei

    • e-’s shared

      • covalent bond

        • Only NONMETALS!


    Covalent bonds1
    Covalent bonds

    • Nonmetals

    • don’t lose e-’s

      • still want NGC

      • share valence e- w/ each other = covalent bonding

    • both atomscount shared e- for NGC

    Covalent bonding w/ Fluorine atoms


    Covalent bonding

    • Fluorine has 7 valence e-

    • A second atom also has seven

    • By sharing electrons…

    • …both end with full orbitals

    F

    F

    8 Valence electrons


    Covalent bonding

    • Fluorine has 7 valence e-

    • A second atom also has 7

    • By sharing electrons…

    • …both full orbitals

    F

    F

    8 Valence electrons

    single covalent bond between 2 H atoms


    Molecular formulas
    Molecular Formulas

    • molecular compounds - Compounds bonded covalently

    • Molecular compounds have:

      • lower melting/boiling pts

      • Generally stronger bond than ionic

      • gases or liquids @ RT

      • molecular formula:

        • Shows # atoms of each element in molecule


    Reminder from ch 7
    Reminder from Ch. 7

    • No “molecule” of sodium chloride

      • Ionic cmpds exist as collection of + & - charged ions arranged in repeating 3D patterns.

        • Formula unit


    Molecular formulas1
    Molecular Formulas

    • water H2O

      • Subscript “2” means 2 atoms of H

      • subscript 1 omitted

    • Molecular formulas do not give info about structure (arrangement of atoms)


    - Page 215

    3. The ball and stick model is BEST, because it shows 3D arrangement.

    These are some of the different ways to represent ammonia:

    1. The molecular formula shows how many atoms of each element are present

    2. The structural formula ALSO shows the arrangement of these atoms!


    Section 8 2 the nature of covalent bonding
    Section 8.2The Nature of Covalent Bonding


    A single covalent bond is
    A Single Covalent Bond is...

    • sharing 2 valence e-

    • Only nonmetals and H

    • Different from ionic bond b/c they actually form molecules.

    • Two specific atoms joined

    • In an ionic solid, you can’t tell which atom e-’s moved from or to


    How to show the formation
    How to show the formation…

    • It’s like a jigsaw puzzle.

    • You put the pieces together to end up with the right formula.

    • C ….. special ex. - can it really share 4 electrons: 1s22s22p2?

      2p

      1s 2s

      C

      • Yes, due to e- promotion!


    How to show the formation1
    How to show the formation…

    • It’s like a jigsaw puzzle.

    • You put the pieces together to end up with the right formula.

    • Carbon is a special example - can it really share 4 electrons: 1s22s22p2?

      2p

      1s 2s


    Water

    H

    O

    Water

    Another example: water formed w/ covalent bonds, by using e- dot structure

    • Each H has 1 valence e-

      - Each H wants 1

    • O has 6 valence e-

      - wants 2 more

    • Sharing completes each octet

    Single bonds


    Water1

    O

    Water

    • Put the pieces together

    • The first H is happy

    • The O still needs 1 more

    H


    Water2

    O

    Water

    • a 2nd H attaches

    • Every atom has full energy levels

    Note the two “unshared” pairs of electrons

    H

    H



    How many bonds do i make
    How many bonds do I make?

    #bonds = # valence e- needed - # valence e- present / 2

    Check out my webpage for a detailed list of rules for electron dot/structural formula rules for covalent bonds


    • Examples:

    • Conceptual Problem 8.1 on page 220

    • We’ll do #7 & 8


    Double and triple covalent bonds
    Double and Triple Covalent Bonds

    • Sometimes atoms share more thanone pair of valence e-’s

    • double bond: atoms share 2 pairs of e-’s (4 total)

    • triple bond: atoms share 3 pairs of e-’s (6 total)

    • Table 8.1, p.222 - Know these 7 elements as diatomic:

      Br2 I2 N2 Cl2 H2 O2 F2


    Dot diagram for carbon dioxide

    O

    Dot diagram for Carbon dioxide

    • CO2 - Carbon is central atom( more metallic )

    • C - 4 valence e-

      • Wants 4 more

    • O- 6 valence e-

      • Wants 2 more

    C

    The chemistry of CO2 6:44


    Carbon dioxide

    O

    Carbon dioxide

    • Attaching 1 oxygen leaves oxygen 1 e- short, and carbon 3 short

    C


    O

    O

    Carbon dioxide

    • Attaching 2nd O

      • both O 1 e- short

      • C 2 e- short

    C


    O

    O

    Carbon dioxide

    • only solution  share more

    C


    O

    O

    Carbon dioxide

    • The only solution is to share more

    C


    O

    Carbon dioxide

    • The only solution is to share more

    O

    C


    O

    Carbon dioxide

    • The only solution is to share more

    O

    C


    O

    Carbon dioxide

    • The only solution is to share more

    O

    C


    Carbon dioxide

    • The only solution is to share more

    O

    C

    O


    Carbon dioxide

    • The only solution is to share more

    • Requires 2 double bonds

    • Each atom counts all e- in bond

    O

    C

    O


    Carbon dioxide

    • The only solution is to share more

    • Requires two double bonds

    • Each atom counts all e- in the bond

    8 valence electrons

    O

    C

    O


    Carbon dioxide

    • The only solution is to share more

    • Requires two double bonds

    • Each atom counts all e- in the bond

    8 valence electrons

    O

    C

    O


    Carbon dioxide

    • The only solution is to share more

    • Requires two double bonds

    • Each atom counts all e- in the bond

    8 valence electrons

    O

    C

    O

    How covalent bonds form - Mark Rosengarden


    How to draw them
    How to draw them?

    • Use the handout guidelines:

    • Add up all valence e-

    • Count total e- needed to make all atoms happy (stable)

    • Subtract; Divide by 2 (tells you how many bonds to draw)

    • Central atom (least electronegative)

    • Start w/ most electronegative atom, fill in remaining valence e- to fill atoms up


    Examples
    Examples

    • NH3, ( ammonia )

    • N – central atom; 5 valence e- (wants 8)

    • H - has 1 (x3) valence electrons, wants 2 (x3)

    • NH3 has 5+3 = 8

    • NH3 wants 8+6 = 14

    • (14-8)/2= 3 bonds

    • 4 atoms with 3 bonds

    N

    H


    Examples1
    Examples

    • Draw in the bonds; start with singles

    • All 8 e- accounted for

    • Everything full – DONE!

    H

    H

    N

    H


    Example hcn
    Example: HCN

    • HCN: C is central atom

    • N - has 5 valence electrons, wants 8

    • C - has 4 valence electrons, wants 8

    • H - has 1 valence electron, wants 2

    • HCNhas 5+4+1 = 10

    • HCNwants 8+8+2 = 18

    • (18-10)/2= 4 bonds

    • 3 atoms with 4 bonds – this will require multiple bonds - not to H however


    HCN

    • Put single bond between each atom

    • Need to add 2 more bonds

    • Must go between C and N (Hydrogen is full)

    H

    C

    N


    HCN

    • Put in single bonds

    • Needs 2 more bonds

    • Must go between C and N, not the H

    • Uses 8 electrons – need 2 more to equal the 10 it has

    H

    C

    N


    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 the N to fill its octet

    H

    C

    N


    Another way of indicating bonds
    Another way of indicating bonds

    • Often use a line to indicate a bond

      • structural formula

    • Each line = 2 valence e-

    H

    O

    H

    H

    O

    H

    =



    A coordinate covalent bond

    C

    O

    A Coordinate Covalent Bond...

    • When one atom donates both electrons in a covalent bond.

    • Carbon monoxide (CO) is a good example:

    Both the carbon and oxygen give another single electron to share


    Coordinate Covalent Bond

    • When one atom donates both e-

      • i.e. Carbon monoxide (CO)

    Oxygen gives both of these electrons, since it has no more singles to share.

    This carbon electron moves to make a pair with the other single.

    C

    O


    Coordinate Covalent Bond

    • When one atom donates both e-

      • Carbon monoxide (CO)

    The coordinate covalent bond is shown with an arrow as:

    C

    O

    C O


    Coordinate covalent bond
    Coordinate covalent bond

    • Most polyatomic cations & anionscontain covalent & coordinate covalent bonds

      • Table 8.2, p.224

    • Sample Problem 8.2, p.225

    • Ammonium ion (NH4+)can be shown as another example


    Bond dissociation energies
    Bond Dissociation Energies...

    • Total energy required to break bond btwn 2 covalently bonded atoms

    • High dissociation energy usually means compound relatively unreactive, b/c high energy needed to break bond


    Resonance is
    Resonance is...

    • When more than one valid dot diagram is possible.

    • Consider the two ways to draw ozone (O3)

    • Which one is it? Does it go back and forth?

    • It’s hybrid of both, shown by double-headed arrow

    • found in double-bond structures!


    Resonance in ozone
    Resonance in Ozone

    Note the different location of the double bond

    Neithersinglestructure is correct, actually a hybrid of the two. To show it, draw all possible structures, and join them with a double-headed arrow.


    Resonance
    Resonance

    • Occurs when more than one valid Lewis structure can be written for particular molecule (due to position of double bond)

    • resonance structures of carbonate ion (used in production of carbonated beverages).

    • The actual structure is an avg (or hybrid) of these structures.


    Polyatomic ions note the different positions of the double bond
    Polyatomic ions – note the different positions of the double bond.

    Resonance in a carbonate ion (CO32-):

    Resonance in an acetate ion (C2H3O21-):


    The 3 exceptions to octet rule
    The 3 Exceptions to Octet rule

    • For some molecules, it is impossible to satisfy the octet rule

      #1. usually when there is an odd numberof valence electrons

      • NO2 has 17 valence electrons, because the N has 5, and each O contributes 6. Note “N” page 228

    • It is impossible to satisfy octet rule, yet the stable molecule does exist


    Exceptions to octet rule
    Exceptions to Octet rule

    • Another exception: Boron

      • Page 228 shows boron trifluoride, and note that one of the fluorides might be able to make a coordinate covalent bond to fulfill the boron

      • #2 -But fluorine has a high electronegativity (it is greedy), so this coordinate bond does not form

    • #3 -Top page 229 examples exist because they are in period 3 or below


    Section 8 3 bonding theories
    Section 8.3Bonding Theories

    • OBJECTIVES:

      • Describe the relationship between atomic and molecular orbitals.


    Section 8 3 bonding theories1
    Section 8.3Bonding Theories

    • OBJECTIVES:

      • Describe how VSEPR theory helps predict the shapes of molecules.


    Molecular orbitals are
    Molecular Orbitals are...

    • The model for covalent bonding assumes orbitals are those of individual atoms = atomic orbital

    • Orbitals that apply to overall molecule, due to atomic orbital overlap are molecular orbitals

      • bonding orbital is molecular orbital occupied by 2 e-’s of covalent bond


    Molecular orbitals definitions
    Molecular Orbitals - definitions

    • Sigma bond- 2 atomic orbitals combine to form molecular orbital symmetricalalong axis connecting nuclei

    • Pi bond- bonding e-’s likely above and below bond axis (weaker than sigma)

    • Note pictures - next slide

    Hybridization video – 1:36


    - Pages 230 and 231

    Sigma bond is symmetrical along axis between 2 nuclei.

    Pi bond is above and below bond axis - weaker than sigma


    Attractive repulsive forces in h 2 bond
    Attractive & repulsive forces in H2 bond

    • + nuclei repel

    • e-’s repel

    • Nuclei and e-’s attract

    • Attractive forces stronger than repulsion

    • As nuclei distances decrease, PE decreases

    • If distance decreases more, PE increases b/c increased repulsion

    • Bond forms w/ bond length = interatomic distance (PE minimum)


    Vsepr stands for
    VSEPR: stands for...

    • Valence Shell Electron Pair Repulsion

    • Predicts 3D shape of molecules

    • The name tells you the theory:

      • Valence shell = outside e-’s

      • Electron Pair repulsion = e- pairs try to get as far awayas possible from each other.

    • determines angles of bonds.


    Vsepr
    VSEPR

    • Based on # of pairs of ve-’s, bonded & unbonded.

    • Unbonded pair called lone pair.

    • CH4 - draw structural formula

    • Has 4 + 4(1) = 8

    • wants 8 + 4(2) = 16

    • (16-8)/2 = 4 bonds


    Vsepr for methane a gas
    VSEPR for methane (a gas):

    • Single bonds fill all atoms.

    • 4 pairs of e-’s pushing away

    • The furthest they can get away is 109.5°


    4 atoms bonded
    4 atoms bonded

    • Basic shape -tetrahedral

    • pyramid w/ triangular base

    • Same shape for everything with 4 pairs

    H

    109.5º

    C

    H

    H

    H


    Other angles pages 232 233
    Other angles, pages 232 - 233

    • Ammonia (NH3) = 107o

    • Water (H2O) = 105o

    • Carbon dioxide (CO2) = 180o


    VSEPR models

    VSEPR theory video 4:52

    - Page 232

    Methane has an angle of 109.5o, called tetrahedral

    Ammonia has an angle of 107o, called pyramidal

    Note the unshared pair that is repulsion for other electrons.



    Hybrid orbitals
    Hybrid Orbitals

    • Provides info for molecular bonding & shape


    Hybridization w single bonds
    Hybridization w/ single bonds

    • Orbitals combine

    • C outer e- configuration 2s2 2p2 but one 2s e- promoted to 2p

      • One 2s e- and 3 2p e-’s

      • Allows bond in methane (CH4)

      • All bonds same due to orbital hybridization

      • Mix to form 4 sp3 hybrid orbitals


    Hybridization w double bonds
    Hybridization w/ double bonds

    • Ethene C2H4 – 1 C-C double bond and 4 C-H single bonds

    • sp2 hybrid orbitals form from one 2s and two 2p atomic orbitals of C


    Section 8 4 polar bonds and molecules
    Section 8.4Polar Bonds and Molecules

    • OBJECTIVES:

      • Describe how electronegativity values determine the distribution of charge in a polar molecule.


    Section 8 4 polar bonds and molecules1
    Section 8.4Polar Bonds and Molecules

    • OBJECTIVES:

      • Describe what happens to polar molecules when they are placed between oppositely charged metal plates.


    Section 8 4 polar bonds and molecules2
    Section 8.4Polar Bonds and Molecules

    • OBJECTIVES:

      • Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


    Section 8 4 polar bonds and molecules3
    Section 8.4Polar Bonds and Molecules

    • OBJECTIVES:

      • Identify the reason why network solids have high melting points.


    Bond polarity
    Bond Polarity

    • do covalent bonds always share equally?

  • e-’s pulled - tug-of-war, btwn nuclei

    • In equal sharing (such as diatomic molecules), the bond is called nonpolar covalent bond


  • Bond polarity1
    Bond Polarity

    • When 2 different atoms bond covalently, unequal sharing

      • more electronegative atomstronger attraction

        • slightly negative charge

    • polar bond

    Lower case delta


    Electronegativity
    Electronegativity?

    The ability of an atom in a molecule to attract shared electrons to itself.

    Linus Pauling

    1901 - 1994


    Bond polarity2
    Bond Polarity

    • Refer to periodic table w/ EN

    • Consider HCl

      H = electronegativity of 2.1

      Cl = electronegativity of 3.0

      • Polar bond

        • Cl: slight - charge

        • H: slight + charge


    Bond polarity3
    Bond Polarity

    • Partial charges, much less than 1+ or 1- in ionic bond

      H Cl

      Partial charges

    d+d-

    d+andd-


    Bond polarity4
    Bond Polarity

    • Can also be:

      • arrow points to more EN atom

    • KNOW Table 8.3, p.238 shows how electronegativity indicates bond type

    H Cl


    Polar molecules
    Polar molecules

    • Sample Problem 8.3, p.239

    • polar bond tends to make entire molecule“polar”

      • areas of “difference”

    • HCl has polar bonds, thus polar molecule

      • molecule w/ 2 poles called dipole, like HCl


    Polar molecules1
    Polar molecules

    • effect of polar bonds on polarity of entire molecule depends on molecule shape

      • CO2 has 2 polar bonds

        • linear

        • nonpolar molecule


    Polar molecules2
    Polar molecules

    • effect of polar bonds on molecule polarity depends on shape

      • water has 2 polar bonds - bent shape

      • highly electronegative O pulls e- away from H

      • very polar!

        polar bond of water molecule

    Chemistry of water 4:46




    Attractions between molecules p 240
    Attractions between molecules p. 240

    • makes solid & liquid molecular cmpds possible

    • weakest called van der Waal’sforces - there are two kinds:

      #1. Dispersion forces: weakest of all, caused by motion of e- - increases as # e- increases (momentarily more on side of molecule closest to neighboring molecule, neighboring molecule’s e-’s move to opp side)

      halogens start as gases; bromine (l); iodine (s) – all in Group 7A


    2 dipole interactions
    #2. Dipole interactions

    • Occurs when polar molecules attracted to each other

    • 2. Dipole interaction happens in water Figure 8.25, page 240

      • +region of one molecule attracts -region of another molecule


    2 dipole interactions1

    d+d-

    d+d-

    H F

    H F

    #2. Dipole interactions

    • Occur when polar molecules attracted to each other

    • Slightly stronger than dispersion forces

    • Opposites attract, but not completely hooked like ionic solids


    2 dipole interactions2

    d+d-

    d+d-

    d+d-

    d+d-

    d+d-

    d+d-

    d+d-

    #2. Dipole Interactions

    d+d-


    3 hydrogen bonding
    #3. Hydrogen bonding

    • …is the attractive force caused by hydrogen bonded to N, O, F, or Cl

    • N, O, F, and Cl very electronegative, so this is very strong dipole

    • And, the hydrogen shares with lone pair in molecule next to it

    • This is strongest of the intermolecular forces


    Order of intermolecular attraction strengths
    Order of Intermolecular attraction strengths

    • Dispersion forces are weakest

    • Little stronger are dipole interactions

    • Strongest is H bonding

    • All are weaker than ionic bonds


    3 hydrogen bonding defined
    #3. Hydrogen bonding defined:

    • When H atom is: a) covalently bonded to a highly EN atom, AND b) is also weakly bonded to unshared e- pair of nearby highly EN atom

      • The H is left very e- deficient (only had 1 to start with!) - it shares with something nearby

      • H is ONLY element with no shieldingfor its nucleus when involved in covalent bond!


    Hydrogen bonding shown in water

    d-

    d+

    O

    d+

    H

    d+

    d-

    H

    H

    O

    H

    d+

    Hydrogen Bonding(Shown in water)

    This H is bonded covalently to: 1) the highly negative O, and 2) a nearby unshared pair.


    H bonding allows h 2 o to be a liquid at room temp

    H

    O

    O

    H

    H

    O

    H

    H

    H

    H

    O

    H

    H

    H

    H

    O

    O

    O

    H

    H

    H

    H bonding allows H2O to be a liquid at room temp


    Attractions and properties
    Attractions and properties

    • Why are some chemicals gases, some liquids, some solids?

      • Depends on type of bonding!

      • Table 8.4, page 244

    • Network solids– solids where all atoms covalently bonded to each other


    Attractions and properties1
    Attractions and properties

    • Figure 8.28, page 243

    • Network solids melt at very high temps, or not at all (decomposes)

      • Diamond does not really melt, but vaporizes to a gas at 3500 oC

      • SiC, used in grinding, has melting pt of 2700 oC


    Covalent network compounds
    Covalent Network Compounds

    Some covalently bonded substances DO NOT form discrete molecules.

    Graphite, a network of covalently bonded carbon atoms

    Diamond, a network of covalently bonded carbon atoms



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