Chemical Bonding I: Basic Concepts Chapter 9 – (Topic 4 and 14)

Chemical Bonding I:Basic Concepts

Chapter 9 –

(Topic 4 and 14)

4.1 – Ionic Bonding

9.1

Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons thatparticpate in chemical bonding.

1A 1ns1

2A 2ns2

3A 3ns2np1

4A 4ns2np2

5A 5ns2np3

6A 6ns2np4

7A 7ns2np5

Group # of valence e-e- configuration

9.1

Lewis Dot Symbols for the Representative Elements &Noble Gases

• Why do substances bond? – More stability – Atoms want to achieve a lower energy state

Ionic Bonding

• Between a metal and a non-metal.

• Metals lose electrons becoming a cations, while non-metals gain electrons becoming anions.

• An ionic bond is an electrostatic attraction between the oppositely charged ions.

9.2

Li + F Li+ F -

The Ionic Bond

1s22s11s22s22p5 1s21s22s22p6[He][Ne]

Li Li+ + e-

e- + F F -

F -Li+ + Li+ F -

9.3

Lattice energy (E) increases as Q increases and/or

as r decreases.

cmpd lattice energyMgF2

MgO

LiF

LiCl

2957

3938

1036

853

Q= +2,-1

Q= +2,-2

r F- < r Cl-

Electrostatic (Lattice) Energy

E = kQ+Q-r

Q+ is the charge on the cation

Q- is the charge on the anionr is the distance between the ions

Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.

Ionic Structures• In an ionic compound (solid), the ions are packed

together into a repeating array called a crystal lattice.

• The simplest arrangement is one in which the spheres in the base are packed side by side. Opposite charges are attracted to each other.

• Its called simple cubic packing (NaCl is an example)

http://www.avogadro.co.uk/structure/chemstruc/ionic/g-ionic.htm

4.5 Physical Properties

General physical properties

• Depend on the forces between the particles

• The stronger the bonding between the particles, the higher the M.P and BP– MP tends to depend on the existence of a

regular lattice structure

Impurities and Melting points

• An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens the bonding. – They always LOWER melting points– Its often used to check purity of a known

molecular covalent compound because its MP will be off, proving its contamination

4.2 – Covalent Bonding

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

9.4

8e-

H HO+ + OH H O HHor

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

8e- 8e-8e-double bonds double bonds

Triple bond – two atoms share three pairs of electrons

N N8e-8e-

N N

triple bondtriple bond

or

9.4

Lengths of Covalent Bonds

Bond Lengths

Triple bond < Double Bond < Single Bond 9.4

H F FH

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms (electrons are shared unequally)

electron richregion

electron poorregion e- riche- poor

+ -

9.5

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.

Electronegativity - relative, F is highest

9.5

H F

electron poorregion

electron richregion

9.5

The Electronegativities of Common Elements

Nonpolar Covalent

share e- equally

Polar Covalent

partial transfer of e-

Unequal sharing

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Difference Bond Type

0 Nonpolar Covalent

2 Ionic

0 < and <2 Polar Covalent

9.5

Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; andthe NN bond in H2NNH2.

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0 NonPolar Covalent

9.5

The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.

H2 (g) H (g) + H (g) H0 = 436.4 kJ

Cl2 (g) Cl (g)+ Cl (g) H0 = 242.7 kJ

HCl (g) H (g) + Cl (g) H0 = 431.9 kJ

O2 (g) O (g) + O (g) H0 = 498.7 kJ O O

N2 (g) N (g) + N (g) H0 = 941.4 kJ N N

Bond Energy

Bond Energies

Single bond < Double bond < Triple bond

9.10

Coordinate Covalent or Dative Bond

• A covalent bond in which one of the atoms donates both electrons.

• Properties do not differ from those of a normal covalent bond.

Coordinate covalent bonds (dative)

• A covalent bond that occurs between two atoms in which both electrons shared in the bond come from the same atom.

• Both electrons from the nitrogen are shared with the upper hydrogen

• Ammonium has 3 polar covalent bonds and 1 coordinate (dative) covalent bond.

Examples

• Hydronium (H3O+)

• Carbon monoxide (CO)

1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.

2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

3. Complete an octet for all atoms except hydrogen

4. If structure contains too many electrons, form double and triple bonds on central atom as needed.

Writing Lewis Structures

9.6

Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6

Write the Lewis structure of the carbonate ion (CO32-).

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6

Step 5 - Too many electrons, form double bond and re-check # of e-

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.

O O O+ -

OOO+-

O C O

O

- -O C O

O

-

-

OCO

O

-

- 9.8

What are the resonance structures of the carbonate (CO3

2-) ion?

• Resonance is possible whenever a Lewis structure has a multiple bond and an adjacent atom with at least one lone pair.

• The following is the general form for resonance in a structure of this type.

Exceptions to the Octet Rule

The Incomplete Octet

H HBeBe – 2e-

2H – 2x1e-

4e-

BeH2

BF3

B – 3e-

3F – 3x7e-

24e-

F B F

F

3 single bonds (3x2) = 69 lone pairs (9x2) = 18

Total = 24

9.9

Exceptions to the Octet Rule

Odd-Electron Molecules (radicals -very reactive)

N – 5e-

O – 6e-

11e-

NO N O

The Expanded Octet (central atom with principal quantum number n > 2)

SF6

S – 6e-

6F – 42e-

48e-

S

F

F

F

FF

F

6 single bonds (6x2) = 1218 lone pairs (18x2) = 36

Total = 48

9.9

Valence shell electron pair repulsion (VSEPR) model:

Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.

AB2 2 0

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

10.1

linear linear

B B

Cl ClBe

2 atoms bonded to central atom

0 lone pairs on central atom

10.1

AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB3 3 0trigonal planar

trigonal planar

10.1


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

10.1

AB4 4 0 tetrahedral tetrahedral

14.1 - Molecules with more than 4 electron pairs

• Molecules with more than 8 valence electrons [expanded valence shell]

• Form when an atom can ‘promote’ one of more electron from a doubly filled s- or p-orbital into an unfilled low energy d-orbital

• Only in period 3 or higher because that is where unused d-orbitals begin

Why does this ‘promotion’ occur?

• When atoms absorb energy (heat, electricity, etc…)their electrons become excited and move from a lower energy level orbital to a slightly higher one.

• How many new bonding sites formed depends on how many valence electrons are excited.


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

10.1


AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

10.1


AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB6 6 0 octahedraloctahedral

bonding-pair vs. bondingpair repulsion

lone-pair vs. lone pairrepulsion

lone-pair vs. bondingpair repulsion> >

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

AB2E 2 1trigonal planar

bent

10.1

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB3E 3 1


tetrahedraltrigonal

pyramidal

10.1

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


10.1

AB3E 3 1 tetrahedraltrigonal

pyramidal

AB2E2 2 2 tetrahedral bent

H

O

H

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

10.1

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal

bipyramidaldistorted

tetrahedron

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

10.1

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal


tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

ClF

F

F

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

10.1

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal


tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

AB2E3 2 3trigonal

bipyramidallinear

I

I

I

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

10.1


AB5E 5 1 octahedral square pyramidal

Br

F F

FF

F

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

10.1


AB5E 5 1 octahedral square pyramidal

AB4E2 4 2 octahedral square planar

Xe

F F

FF

Predicting Molecular Geometry1. Draw Lewis structure for molecule.

2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.

3. Use VSEPR to predict the geometry of the molecule.

What are the molecular geometries of SO2 and SF4?

SO O

AB2E

bent

S

F

F

F F

AB4E

distortedtetrahedron

10.1

Polarity and shape

• The shape of the molecule directly influences the overall polarity of the molecule.

• If there is symmetry the charges cancel each other out, making the molecule non-polar

• If there is no symmetry, then its polar

http://www.mhhe.com/physsci/chemistry/chang7/esp/folder_structure/bo/m4/s2/index.htm

• Polar bonds do not guarantee a polar molecule

• Ex: CCl4 and CO2 both have polar bonds, but both are NON-POLAR molecules. They have a dipole moment of zero

• The greater the dipole moment, the more polar the molecule

Why is molecular polarity important?

• Polar molecules have higher melting and boiling points (for example the BP of HF is 19.5° C, and the BP of F2 is –188° C).

• Polar solvents dissolve ionic and polar molecules more efficiently than non-polar solvents

Dipole Moments and Polar Molecules

10.2

H F

electron richregion

electron poorregion

10.2

Which of the following molecules have a dipole moment?H2O, CO2, SO2, and CH4

O HH

dipole momentpolar molecule

SO

O

CO O

no dipole momentnonpolar molecule

dipole momentpolar molecule

C

H

H

HH

no dipole momentnonpolar molecule

Does BF3 have a dipole moment?

10.2

The bent shape creates an overall positive end and negative endof the molecule = POLAR

The symetry of the molecule Cancels out the “charges” Making this NON-POLARNo overall DIPOLE

Examples to Try

• Determine whether the following molecules will be polar or non-polar

–SI2

–CH3F

–AsI3

–H2O2

Summary of Polarity of Molecules

• Linear:– When two atoms attached to central atom are the

same, the molecule will be Non-Polar (CO2)

– When the two atoms are different the dipoles will not cancel, and the molecule will be Polar (HCN)

• Bent:– The dipoles created from this molecule will not

cancel creating a net dipole moment and the molecule will be Polar (H2O)

Summary of Polarity of Molecules

• Pyramidal:– The dipoles created from this molecule will not

cancel creating a net dipole and the molecule will be Polar (NH3)

• Trigonal Planar:– When the three atoms attached to central atom

are the same, the molecule will be Non-Polar (BF3)

– When the three atoms are different the dipoles will not cancel, resulting in a net dipole, and the molecule will be Polar (CH2O)

Tetrahedral

• When the four atoms attached to the central atom are the same the molecule will be Non-Polar

• When three atoms are the same, and one is different, the dipoles will not cancel, and the molecule will be Polar

Allotropes of Carbon

Allotropes

• Carbon can bond with itself in at least three different ways giving us 3 different materials– Diamond– Graphite– Buckyballs and nanotubes

Diamond

• Carbons are bonded via sp3 hybridization to 4 other carbon atoms forming a giant network covalent compound.

Properties of Diamond

• High melting point due to strong directional covalent bonds (3550 C)

• Extremely hard because it is difficult to break atoms apart or move them in relation to one another

• No electrical conductivity because electrons are localized in specific bonds

• Insoluble in polar and non-polar solvents because molecular bonds are stronger than any intermolecular forces

Graphite

• Carbon atoms are bonded via sp2 hybridization.

• Carbon atoms form sheets of six sided rings with p-orbitals perpendicular from plane of ring.

Graphite Structure

• Carbon has 4 valence e- to bond with. 3 are used for closest atoms in rings. 1 is delocalized in p-orbitals

• The presence of p-orbitals allows for strong van der waals forces that hold the sheets together

Properties of Graphite

• Different from Diamond– Conducts electricity because of delocalized

electrons– Slippery can be used as lubricant, sheets can

easily slip past each other (think of a deck of cards)

• Same as Diamond– High melting point (higher actually because of

delocalized e-, 3653C)– Insoluble (same reason)

Fullerenes

• Buckyballs: spherical • Nanotubes: tube

shaped• Both have very

interesting properties – Super strong– Conduct electricity and

heat with low resistance

– Free radical scavenger

Buckyballs

• Carbon atoms bond in units of 60 atoms (C-60) forming a structure similar to a soccerball with interlocking six sided and five sided rings.

• sp2 hybridization • Extra p-orbitals form pi

bonds resulting in– Electrical conductivity– Stronger covalent bonds,

therefore stronger materials

4.3 - Intermolecular Forces

11.2

Intermolecular forces are attractive forces between molecules.

Intramolecular forces hold atoms together in a molecule.

Intermolecular vs Intramolecular

• 41 kJ to vaporize 1 mole of water (inter)

• 930 kJ to break all O-H bonds in 1 mole of water (intra)

Generally, intermolecular forces are much weaker than intramolecular forces.

“Measure” of intermolecular force

boiling point

melting point

Hvap

Hfus

Hsub

Intermolecular Forces

Dipole-Dipole Forces

Attractive forces between polar molecules

Orientation of Polar Molecules in a Solid

11.2

Intermolecular Forces

Ion-Dipole Forces

Attractive forces between an ion and a polar molecule

11.2

Ion-Dipole Interaction

Intermolecular ForcesDispersion Forces

Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules

11.2

ion-induced dipole interaction

dipole-induced dipole interaction

Induced Dipoles Interacting With Each Other

11.2

Intermolecular ForcesDispersion Forces Continued

11.2

Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted.

Polarizability increases with:

• greater number of electrons

• more diffuse electron cloud

Dispersion forces usually increase with molar mass.

SO

O

What type(s) of intermolecular forces exist between each of the following molecules?

HBrHBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.

CH4

CH4 is nonpolar: dispersion forces.

SO2

SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

11.2

Intermolecular ForcesHydrogen Bond

11.2

The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom.

A H…B A H…Aor

A & B are N, O, or F

Why is the hydrogen bond considered a “special” dipole-dipole interaction?

Decreasing molar massDecreasing boiling point

11.2

4.4 Metallic bonding

Types of CrystalsMetallic Bonds- electron sea model - metal cations in a sea of valence electrons

• Lattice points in crystal are occupied by metal atoms• the valence electrons do not “belong” to a single cation, but

are delocalized and may move about • Good conductors of heat and electricity

11.6

Cross Section of a Metallic Crystalnucleus &

inner shell e-

mobile “sea”of e-

Metallic bond

• Occurs between atoms with low electronegativities

• Metal atoms pack close together in 3-D, like oranges in a box.

• Close-packed lattice formation

• Many metals have an unfilled outer orbital

• In an effort to be energy stable, their outer electrons become delocalised amongst all atoms

• No electron belongs to one atom

• They move around throughout the piece of metal.

• Metallic bonds are not ions, but nuclei with moving electrons

Physical Properties

Conductivity• Delocalised electrons are

free to move so when a potential difference is applied they can carry the current along

• Mobile electrons also mean they can transfer heat well

• Their interaction with light makes them shiny (lustre)

Malleability

• The electrons are attracted the nuclei and are moving around constantly.

• The layers of the metal atoms can easily slide past each other without the need to break the bonds in the metal

• Gold is extremely malleable that 1 gram can be hammered into a sheet that is only 230 atoms thick (70 nm)

Melting points

• Related to the energy required to deform (MP) or break (BP) the metallic bond

• BP requires the cations and its electrons to break away from the others so BP are very high.

• The greater the amount of valence electrons, the stronger the metallic bond.

• Gallium can melt in your hand at 29.8 oC, but it boils at 2400 oC!

4.5 Physical Properties

General physical properties

• Depend on the forces between the particles

• The stronger the bonding between the particles, the higher the M.P and BP– MP tends to depend on the existence of a

regular lattice structure

Impurities and Melting points

• An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens the bonding. – They always LOWER melting points– Its often used to check purity of a known

molecular covalent compound because its MP will be off, proving its contamination

How would this ideal heat curve look different if the substance was contaminated?

Volatility

• A qualitative measure of how readily a liquid or solid is vaporised upon heating or evaporation– It is a measure of the tendency of molecules and

atoms to escape from a liquid or a solid. – Relationship between vapour pressure and

temperature (B.P)

• Mostly dealing with liquids to gas, however can occur from solid directly to gas (dry ice).

• The weaker the intermolecular bonds, the more volatile

Conductivity

• Generally molecules have poor solubility in polar solvents like water, but if they do dissolve they do not for ions

• There are no charged particles to carry the electrical charge across the solution.

• Example: sugar dissolves in water

• C12H22O11(s) C12H22O11(aq)

Dissolving sugar (covalent compound)

• It takes energy to break the bonds between the C12H22O11 molecules in sucrose crystal structure.

• It also takes energy to break the hydrogen bonds in water so that one of these sucrose molecules can fit into solution.

• In order for sugar to dissolve, there must be a greater release of energy when the dissolution occurs than when the breaking of bonds occur.

Ionic compounds

• The energy needed to break the ionic bond must be less than the energy that is released when ions interact with water.

• The intermolecular ion-dipole force is stronger than the electrostatic ionic bond

• Breaks up the compound into its ions in solution.

• Soluble salt in water breaks up as

NaCl (s) Na+ (aq) + Cl- (aq)

• http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

Ionic compounds

• Held together by strong 3-d electrostatic forces.

• They are solid at room temperature and pressure

• If one layer moves a fraction, the ions charges are off and now repulsion occurs. This is the reason they are strong, yet brittle.

• Molten or dissolved ionic compounds conduct electricity

• Insoluble in most solvents, yet H2O is polar and attracts both the + and – ions from salts

Covalent bonding properties

Giant covalent • Ex: diamond, silicon

dioxide• Very hard• Very high MP

(>1000oC)• Does not conduct• Insoluble in all

solvents

Molecular covalent• Ex: CO2, alcohols, I2

• Usually soft, malleable

• Low MP (<200oC)• Does not conduct• More soluble in non-

aqueous solvents, unless they can h-bond

Solubility of methanol in water

• http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/clm2s3_4.swf

• Alcohols generally become less soluble, the longer the carbon chain due to the decreasing tendency for hydrogen bonding to occur intermolecularly.

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/clm2s3_4.swf

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/clm2s3_4.swf

States of matter

• Physical state depends on intermolecular forces

• The weaker the attraction, the more likely it’s a gas, while stronger attractions indicate solid.

14.2 - Hybridization

• the concept of mixing atomic orbitals to form new hybrid orbitals

• Used to help explain some atomic bonding properties and the shape of molecular orbitals for molecules.

• The valence orbitals (outermost s and p orbitals) are hybridised (mathematically mixed) before bonding, converting some of the dissimilar s and p orbitals into identical hybrid spn orbitals

• We must know sp, sp2, and sp3 hydrid orbitals

Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals.

1. Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals.

2. Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process.

3. Covalent bonds are formed by:

a. Overlap of hybrid orbitals with atomic orbitals

b. Overlap of hybrid orbitals with other hybrid orbitals

10.4

Hybrid orbitals

• Carbon has 4 valence electrons.

• 2 electrons paired up in the s-orbital, and 2 electrons unpaired in the p-orbital.

• So why does it commonly make 4 bonding sites?

• One of carbon’s paired s-orbital electrons is ‘promoted’ to the empty p-orbital

• This produces a carbon in an excited state which has 4 unpaired electrons (4 equivalent bonding sites)

10.4

Formation of sp3 Hybrid Orbitals

Formation of sp Hybrid Orbitals

10.4

Formation of sp2 Hybrid Orbitals

10.4

# of Lone Pairs+

# of Bonded Atoms Hybridization Examples

2

3

4

5

6

sp

sp2

sp3

sp3d

sp3d2

BeCl2

BF3

CH4, NH3, H2O

PCl5

SF6

How do I predict the hybridization of the central atom?

1. Draw the Lewis structure of the molecule.

2. Count the number of lone pairs AND the number of atoms bonded to the central atom

10.4

Sigma bond () – electron density between the 2 atomsPi bond () – electron density above and below plane of nuclei

of the bonding atoms 10.5

Describe the bonding in CH2O.

CH

OH

C – 3 bonded atoms, 0 lone pairsC – sp2

10.5

Sigma () and Pi Bonds ()

Single bond 1 sigma bond

Double bond 1 sigma bond and 1 pi bond

Triple bond 1 sigma bond and 2 pi bonds

How many and bonds are in the acetic acid(vinegar) molecule CH3COOH?

C

H

H

CH

O

O H bonds = 6 + 1 = 7

bonds = 1

10.5

Chemical Bonding I: Basic Concepts Chapter 9 – (Topic 4 and 14)

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Transcript of Chemical Bonding I: Basic Concepts Chapter 9 – (Topic 4 and 14)