Chapter 8 covalent bonding
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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”

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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


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    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)


    Chapter 8 covalent bonding

    - 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


    Chapter 8 covalent bonding

    This is what happens when you miss Chemistry class

    Do you know this kid?


    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


    Chapter 8 covalent bonding

    • 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


    Chapter 8 covalent bonding

    O

    O

    Carbon dioxide

    • Attaching 2nd O

      • both O 1 e- short

      • C 2 e- short

    C


    Chapter 8 covalent bonding

    O

    O

    Carbon dioxide

    • only solution  share more

    C


    Chapter 8 covalent bonding

    O

    O

    Carbon dioxide

    • The only solution is to share more

    C


    Chapter 8 covalent bonding

    O

    Carbon dioxide

    • The only solution is to share more

    O

    C


    Chapter 8 covalent bonding

    O

    Carbon dioxide

    • The only solution is to share more

    O

    C


    Chapter 8 covalent bonding

    O

    Carbon dioxide

    • The only solution is to share more

    O

    C


    Chapter 8 covalent bonding

    Carbon dioxide

    • The only solution is to share more

    O

    C

    O


    Chapter 8 covalent bonding

    Carbon dioxide

    • The only solution is to share more

    • Requires 2 double bonds

    • Each atom counts all e- in bond

    O

    C

    O


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    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

    =


    Other structural examples

    Other Structural Examples

    H CN

    H

    C 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


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    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


    Chapter 8 covalent bonding

    - 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


    Chapter 8 covalent bonding

    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.


    Vsepr song 4 33

    VSEPR song – 4:33


    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

    HCl


    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


    Chapter 8 covalent bonding

    bonding animations


    Chapter 8 covalent bonding

    Polar and non-polar covalent bonds - Mark Rosengarten


    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-

    HF

    HF

    #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


    Chapter 8 covalent bonding

    Ionic/Covalent Bond Song


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