properties of carbon element

Properties of Carbon Element


• We have learnt carbon element as the basis of Organic Chemistry.

• Why does carbon can form so many different compounds?

• Carbon is found the second period group IVA of the periodic table.

EXCEPTIONS OF CARBON COMPUDS WHICH ARE NOT ORGANIC

• oxides of carbon (CO2, CO)

• carbonates,bicarbonates(NaHCO3,CaCO3)• cyanides (NaCN, etc)


• The Lewis structure for carbon shows 4 unpaired valence electrons.

• To fulfill the octet rule, a carbon atom needs 4 more electrons.

• A carbon atom may form 4 covalent bonds and is capable of forming long chains with single, double or triple bonds between carbon atoms.

• These chains may be continuous (straight) or branched.

• The 2 ends of a chain can bond together to form a ring.

• Carbon compounds are divided into classes based on their chemical similarity.

Hydrocarbons• Hydrocarbons are compounds containing hydrogen and

carbon. Hydrocarbons may have different numbers of bonds between carbon atoms.

• The four hydrocarbon classes are: alkane (single bond), alkene, (double bond), alkyne (triple bond), aromatic (benzene ring).

• Alkanes contain only single C-C bonds. They contain as many hydrogen atoms as possible, and are said to be saturated.

• Hydrocarbons containing double or triple bonds are unsaturated.

• A homologous series is series of compounds that differ by a

constant increment. Aromatic hydrocarbons include a benzene ring- 6 carbon atoms with all the bonds alternating between a single and a double bond.


Carbon is unique– It has 6 electrons in its outer

shell arranges 1s22s2sp2

– It has room for 4 bonds to 4 other atoms.

– Carbon-to-carbon bonds can be single (A),

– double (B), or – triple (C).

– Note that in each example,

each carbon atom has four dashes, which represent four bonding pairs of electrons, satisfying the octet rule.

Alkanes Alkenes Alkynes SATURATED means that each carbon is bonded to four other atoms through single covalent bonds. Hydrogen atoms usually occupy all available bonding positions after the carbons have bonded to each other.

UNSATURATED hydrocarbons contain either double or triple bonds. Since the compound is unsaturated with respect to hydrogen atoms, the extra electrons are shared between 2 carbon atoms forming double or triple bonds.

PARAFFINS which is derived from a Latin word meaning "little activity", and means that the compounds are very unreactive.

Alkenes are also called OLEFINS because they form oily liquids on reaction with chlorine gas.

Alkynes are also generally known as ACETYLENES from the first compound in the series.

HYDROCARBONS

HydrocarbonsHydrocarbons

C C C C

C C

C

C

C

C

C

C

H

H

H

H

H

H

C C C C C

H

H

H

H

H H

H

H

H

H

H

H

Alkanes Alkenes

Alkynes Aromatics

C C C C C

H

H

H

H

H

H

H

H

H

H

C C C C CH

H

H

H

H

H

H

H


• A)The carbon atom forms bonds in a tetrahedral structure with a bond angle of 109.5O.

• (B) Carbon-to-carbon bond angles are 109.5O, so a chain of carbon atoms makes a zigzag pattern.

• (C) The unbranched chain of carbon atoms is usually simplified in a way that looks like a straight chain, but it is actually a zigzag, as shown in (B).


Carbon-to-carbon chains can be

• (A) straight, • (B) branched, or • (C) in a closed ring.

• (Some carbon bonds are drawn longer, but are actually the same length.)

Why does carbon can form so many different compounds?

• There are now more than ten million organic compounds known by chemists.

• Many more undoubtedly exist in nature, and organic chemists are continually creating (synthesizing) new ones.

• Carbon is the only element that can form so many different compounds because each carbon atom can form four chemical bonds to other atoms, and because the carbon atom is just the right, small size to fit in comfortably as parts of very large molecules.

• Having the atomic number 6, every carbon atom has a total of six electrons.

• Two are in a completed inner shell, while the other four are valence electrons—outer electrons that are available for forming bonds with other atoms.


• The carbon atom's four valence electrons can be shared by other atoms that have electrons to share, thus forming covalent (shared-electron) bonds.

• They can even be shared by other carbon atoms, which in turn can share electrons with other carbon atoms and so on, forming long strings of carbon atoms, bonded to each other like links in a chain.

• Silicon (Si), another element in group 4A of the periodic table, also has four valence electrons and can make large molecules called silicones, but its atoms are too large to fit together into as great a variety of molecules as carbon atoms can.


• Carbon's ability to form long carbon-to-carbon chains is the first of five reasons that there can be so many different carbon compounds; a molecule that differs by even one atom is, of course, a molecule of a different compound.

• The second reason for carbon's astounding compound-forming ability is that carbon atoms can bind to each other not only in straight chains, but in complex branchings, like the branches of a tree.

• They can even join "head-to-tail" to make rings of carbon atoms. • There is practically no limit to the number or complexity of the branches

or the number of rings that can be attached to them, and hence no limit to the number of different molecules that can be formed.

• The third reason is that carbon atoms can share not only a single electron with another atom to form a single bond, but it can also share two or three electrons, forming a double or triple bond.

• This makes for a huge number of possible bond combinations at different places, making a huge number of different possible molecules.

• And a molecule that differs by even one atom or one bond position is a molecule of a different compound.

Why does carbon can form so many different compounds

• The fourth reason is that the same collection of atoms and bonds, but in a different geometrical arrangement within the molecule, makes a molecule with a different shape and hence different properties.

• These different molecules are called isomers.• The fifth reason is that all of the electrons that are not being

used to bond carbon atoms together into chains and rings can be used to form bonds with atoms of several other elements.

• The most common other element is hydrogen, which makes the family of compounds known as hydrocarbons.

• But nitrogen, oxygen, phosphorus, sulfur, halogens, and several other kinds of atoms can also be attached as part of an organic molecule.

• There is a huge number of ways in which they can be attached to the carbon-atom branches, and each variation makes a molecule of a different compound.

The Greater Stability of C-C Bonds

• Since the average bond dissociation energy of C-C is greater than the average bond energies between different atoms.

• Thus the energy released when carbon atom bonds to another carbon atom is greater than the energy released when the other atoms like B,N,O,Si,P and S bonds to each other.

• Thus C-C bond is more stable than the others like B-B,N-N,

O-O,Si-Si,P-P and S-S.

Bonding Atoms

Bond Energy (kJ mol-1)

B-B 293 C-C 343 N-N 163 O-O 157 Si-Si 222 P-P 201 S-S 266

Ability to Form Chains Between Their Atoms

• The atoms closer to C in the periodic table are B,N,O,Si,P and S.

• The ability of these atoms to bond each other to form chains is lower than C.

• For examle Si can produce chains made of at most 11 atoms of it and N at most three atoms it.

• Although the ability to form chains between their atoms for P and S is greater than Si and N but it is very much smaller compared to C.

Ability to Form Chains Between Their Atoms

The greater ability of carbon to form chains compared to atoms closer to it in the periodic table can be explained by two reasons:

1. The average bond dissociation energies of them is lower than that of carbon.

2. The electronegativity values B,Si and P lower than that of C.atoms.Thus the attraction forces between these atoms are smaller than that of carbon.This is also true when these atoms are bonded to the other atoms like hydrogen or halogens.

Li Be B C N O F 1,0 1,5 2,0 2,5 3,0 3,5 4,0 Na Mg Al Si P S Cl 0,9 1,3 1,6 1,9 2,2 2,5 3,2

Electronegativity values of some elements ording to Pauling’s Scale

Electronegativity

• Electronegativity:Electronegativity: – a measure of an atom’s attraction for the

electrons it shares with another atom in a chemical bond

• Pauling scalePauling scale– generally increases left to right in a row– generally increases bottom to top in a

column

Greater Bonding Capacity of C compared to N and O

The electronegativity values of N and O are greater than that of C. But their bonding capacities are smaller than that of C since they have lower number of unpaired electrons.

Lewis Dot Diagrams of Selected Elements

Summary…

• Compared to C atom B,Si,P,N and O atoms can not be expected to form greater number of compounds and unbrached and branched chains and cyclic compounds.

• Carbon compounds are more stable than Si4,P4,O3,S8 and B4 molecules.

Electron Configuration of Elements

Lewis Dot Structures…• Gilbert N. Lewis• Valence shell:Valence shell:

– the outermost occupied electron shell of an atom• Valence electrons:Valence electrons:

– electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions

• Lewis dot structure:Lewis dot structure: – the symbol of an element represents the nucleus and

all inner shell electrons– dots represent valence electrons

Lewis Dot Structures

• Table 1.4 Lewis Dot Structures for Elements 1-18

N OB

H

Li Be

Na

He

Cl

F

S

Ne

Ar

C

SiAl P

1A 2A 3A 4A 5A 6A 7A 8A

Mg ::

::

::

.

.

.

.

.

.

.

..

..

. .

.

.

.

:

:

:

::::::::

::::::.

:::

:

Lewis Model of Bonding…

• Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number– an atom that gains electrons becomes an anionanion– an atom that loses electrons becomes a cationcation– the attraction of anions and cations leads to the

formation of ionic solidsionic solids– an atom may share electrons with one or more atoms

to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bondcovalent bond

– bonds may be partially ionic or partially covalent; these bonds are called polar covalent bondspolar covalent bonds

Covalent Bonds!• The simplest covalent bond is that in H2

– the single electrons from each atom combine to form an electron pair

– the shared pair functions in two ways

– simultaneously; it is shared by the two atoms and fills the valence shell of each atom

• The number of shared pairs

– one shared pair forms a single bond

– two shared pairs form a double bond

– three shared pairs form a triple bond

H H H-H+ • H0 = -435 kJ (-104 kcal)/mol•

Hydrogen Molecule Formation

0.74 A

- 436

0

H – H distance

En

erg

y (K

J/m

ol)

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318

no interaction

increasedattraction

balanced attraction& repulsion

increasedrepulsion

Potential Energy Diagram - Attraction vs. Repulsion

(internuclear distance)

Lewis Structures!• To write a Lewis structure

– determine the number of valence electrons

– determine the arrangement of atoms

– connect the atoms by single bonds

– arrange the remaining electrons so that each atom has a complete valence shell

– show a bonding pair of electrons as a single line

– show a nonbonding pair of electrons as a pair of dots

– in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

Table of Lewis Structures!

• In neutral molecules– hydrogen has one bond– carbon has 4 bonds and no lone pairs– nitrogen has 3 bonds and 1 lone pair– oxygen has 2 bonds and 2 lone pairs– halogens have 1 bond and 3 lone pairs

H2O (8) NH3 (8) CH4 (8) HCl (8)

C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)

H-O-H H-N-HH

H-C-HH

HH-Cl

H-C C-HH

HC O

H

HC C

H

HO O

CH H

O

Ethylene

Hydrogen chlorideMethaneAmmoniaWater

Carbonic acidFormaldehydeAcetylene

Resonance!• In chemistry, resonance is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. • A molecule or ion with such delocalized electrons is represented by several structures called resonance structures.

Tautomerization• Tautomerization usually involves the movement of a hydrogen atom between a different location on the molecule, resulting in two or more molecular structures.• These structures are called tautomers, which exist in dynamic equilibrium with each other.

Enol form Keto form

Molecular Geometry and Bonding Theories

molecular formula

structural formula

molecular shape

ball-and-stick model

CH4 C

H

H

HH

H

H

H

H

109.5o

C

tetrahedrontetrahedralshape ofmethane

CH

H

H

H

Methane & Carbon Tetrachloride

molecular formula

structural formula

molecular shape

ball-and-stick model

CH4 C

H

H

HH

H

H

H

H

109.5o

C

CCl4

space-filling model

C

Cl

Cl

ClCl

Molecular Geometry

Linear Trigonal planar

Tetrahedral

Trigonal pyramidalBent

107.3o104.5o

H2O CH4 AsCl3 AsF5 BeH2 BF3 CO2

180o

H

H

H

H

109.5o

C

C109.5o

H

HHH

N107o HH

H

..

O104.5o H

H

..

..

CH4, methane NH3, ammonia H2O, water

Molecular ShapesThree atoms (AB2) Four atoms (AB3)

Five atoms (AB4)

•Linear (180o)•Bent

•Trigonal planar (120o)•Trigonal pyramidal

•Tetrahedral (109.47o)

B BA

B

B

A

B

linear trigonal planar

B

A

BB

B

tetrahedral

Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 313.

Bonding and Shape of Molecules

Number of Bonds

Number of Unshared Pairs Shape Examples

2

3

4

3

2

0

0

0

1

2

Linear

Trigonal planar

Tetrahedral

Pyramidal

Bent

BeCl2

BF3

CH4, SiCl4

NH3, PCl3

H2O, H2S, SCl2

-Be-

B

C

N

:

O

:

:

CovalentStructure

Molecular Shapes

AB2

Linear

AB3

Trigonal planar

AB4

Tetrahedral

AB2EAngular or Bent

AB3ETrigonal

pyramidal

AB2E2

Angular or Bent

Molecular Polarity

Molecular Structure

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Dipole Moment

• Direction of the polar bond in a molecule.

• Arrow points toward the more electronegative atom.

H Cl+ -

Determining Molecular Polarity

• Depends on:– dipole moments– molecular shape



• Nonpolar Molecules– Dipole moments are symmetrical and

cancel out.

BF3

F

F F

B



• Polar Molecules– Dipole moments are asymmetrical and

don’t cancel .

netdipolemoment

H2OH H

O


CHCl3

H

Cl ClCl


• Therefore, polar molecules have...– asymmetrical shape (lone pairs) or – asymmetrical atoms

netdipolemoment


Dipole Moment

Nonpolar

Polar

....

H H

O

C OO

Bond dipoles

Overall dipole moment = 0

Bond dipoles

Overall dipole moment

The overall dipole moment of a moleculeis the sum of its bond dipoles. In CO2 thebond dipoles are equal in magnitude butexactly opposite each other. The overall dipole moment is zero.

In H2O the bond dipoles are also equal inmagnitude but do not exactly oppose eachother. The molecule has a nonzero overall dipole moment.

221

dqqk

F Coulomb’s law = Q r Dipole moment,


......

Polar and Nonpolar Molecules

H Cl

Polar

A molecule has a zero dipole moment because their dipoles cancel one another.

H HO

PolarF F

B

F

Nonpolar

HH

H

N

Polar

Cl

ClC

Cl

Nonpolar Polar

Cl

HC

Cl

H

H

Be

H

H

BeH2

s p

Formation of BeH2 using pure s and p orbitals

The formation of BeH2 using hybridized orbitals

atomic orbitals atomic orbitals

Be

s p

Be H

H

s p

atomic orbitals

hybrid orbitals

No overlap = no bond!

sp p

Be HH

All hybridized bonds have equal strength and have orbitals with identical energies.

BeH2Be

Be = 1s22s2

Hybridization - The Blending of OrbitalsHybridization - The Blending of Orbitals

Poodle

+

+Cocker Spaniel

=

=

=

=

+

+s orbital p orbital

Cockapoo

sp orbital

sp Hybrid Orbitals

Ground-state Be atom

1s 2s 2p

1s 2s 2p

Be atom with one electron “promoted”

s

px py pz

sp

hybrid orbitals

En

erg

y

hybridize

s orbital p orbital

two sp hybrid orbitals sp hybrid orbitals shown together(large lobes only)

1s sp 2p

Be atom of BeH2 orbital diagram

H HBe

n = 1

n = 2

sp Animation

sp2 Hybrid Orbitals

2s 2p

Ground-state B atom

s

px py pzEne

rgy

sp2 2p

B atom of BH3 orbital diagram

hybridize

s orbital

2s 2p

B atom with one electron “promoted”

sp2

hybrid orbitals

p orbitals sp2 hybrid orbitals shown together

(large lobes only)three sps hybrid orbitals

H

H

HB

Sp2 Animation

s

px py pz

Carbon 1s22s22p2

Carbon could only make two bondsif no hybridization occurs. However,carbon can make four equivalent bonds.

sp3

hybrid orbitals

En

erg

y

sp3

C atom of CH4 orbital diagram

B

A

BB

B


sp3 Hybrid Orbitals

Sp3 Animation

Multiple Bonds

2s 2p 2s 2p sp2 2p

promote hybridize

C C

H

H H

H

C2H4, ethene

one bond and one bond

H

H

CC

H

H

H

H

CC

H

H

Two lobes ofone bond

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326

Multiple Bonds

2s 2p 2s 2p sp2 2p

promote hybridize

C C

H

H H

HC2H4, ethene (ethylene)

one bond and one bond

H

H

CC

H

H

H

H

CC

H

H

Two lobes ofone bond

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326

C C

H

H

sp2

sp2

sp2

H

H

sp2

sp2

sp2

p p

p p

Sigma and pi Bonds Animation

bond

Internuclear axis

p p

Two atomic p orbitals Pi (p) Bond

3D view of Pi (p) Bond

(pi) Bond – overlap of two p orbitals oriented perpendicular to the line connecting the nuclei.

bond

H s

Orbital Picture of Ethylene

C C

H s 1

sp2

sp2

sp2

sp2

sp2

sp2

H s

H s

px px

CH

HC

H

H

Ethylene Animation

Bonding in Formaldehyde

bonds in Benzene

H

CHC

H

C

H

C

H C

H

C


C6H6 = benzene

2p atomic orbitals in Benzene


bonds and bonds

H

C

HCH

C

H

C

H C

H

C


bonds in BenzeneH

CH

C

HC

H

C

H

C

H

C

HC

H

C


bondsH

C

H

C

HC

H

C

H

C

H

C

HC

H

C


Ethyne sp Hybridization

Orbital Picture of Acetylene

C C

spsp

px

py

px

py

H s H ssp sp

1

2

C C HH

Acetylene Animation

Number of electron domains

Electron-domain geometry

Predicted bond angles

TetrahedralTrigonalplanar Tetrahedral

109.5o 120o 109.5o

C C OH H

H

H O

4 3 4

Acetic Acid, CH3COOH

Hybridization of central atom sp3 sp2


sp3

VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory

• Electron pairs orient themselves in order to minimize repulsive forces.


......

The Shapes of Some Simple ABn Molecules

H HB

Linear

O OS

BentF F

B

F

Trigonalplanar

FF

F

N

Trigonalpyramidal

SF6

H2O


VSEPR Theory

• Types of e- Pairs– Bonding pairs - form bonds– Lone pairs - nonbonding electrons

Lone pairs repel

more strongly than

bonding pairs!!!


VSEPR Theory

• Lone pairs reduce the bond angle between atoms.

Bond Angle


• Draw the Lewis Diagram.

• Tally up e- pairs on central atom.– double/triple bonds = ONE pair

• Shape is determined by the # of bonding pairs and lone pairs.

Know the 5 common shapes

& their bond angles!

Determining Molecular Shape


Common Molecular Shapes

2 total

2 bond

0 lone

LINEAR180°BeH2


B BA

3 total

3 bond

0 lone

TRIGONAL PLANAR

120°

BF3



B

B

A

B

4 total

4 bond

0 lone

TETRAHEDRAL

109.5°

CH4



B

A

BB

B

4 total

3 bond

1 lone

TRIGONAL PYRAMIDAL

107°

NH3



4 total

2 bond

2 lone

BENT

104.5°H2O



• CO2

O C O2 total

2 bond

0 lone LINEAR

180°

Examples


• PF3

4 total

3 bond

1 lone

TRIGONAL PYRAMIDAL

107°

F P FF

Examples


No Loners Animation

Loners Animation

References

• http://science.jrank.org/pages/1202/Carbon-Why-carbon-special.html

• http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/perlewis.html

• www.nisd.net/communicationsarts/pages/chem

properties of carbon element

Documents

carbon bonds

carbon compounds

b carbon

carbon chains

carbon bond angles

exceptions of carbon

organicoxides of carbon

triple bonds