hydrocarbon chemistry lecture notes

8/11/2019 Hydrocarbon Chemistry Lecture Notes

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1.1 Atomic Structure Atom:

Nucleusneutrons (neutral) + protons (+ve), the mass of theatom.

Electron- Negatively charged, negligible mass

Atom is neutral, No. of protons = no. of electrons

Diameter of an atom is about 2 10-10m (200

picometers(pm)) [the unit angstrom() is 10-10m =

100 pm]


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Atomic number(Z)- number of protons in theatom's nucleus

Mass number(A ) - number of protons plusneutrons

Atoms of a given element have the same atomicnumber: 1 for hydrogen, H 6 for carbon, C

Isotopes - atoms of the same element but

different numbers of neutrons

different massnumbers Atom ic mass(atomic weight)- weighted

average mass in atomic mass units (amu) of an

elements naturally occurring isotopes


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1.2 Atomic Structure: Orbitals

4 kinds of orbitals: s, p, d, and f

Orbitals are organised into different layerselectron shells


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Orbitals and Shells Orbitals are grouped in shells of increasing size and energy Different shells contain different numbers and kinds of orbitals Each orbital can be occupied by two electrons First shell contains one sorbital, denoted 1s, holds only two

electrons Second shell contains one sorbital (2s) and threeporbitals (2p),

eight electrons

Third shell contains an sorbital (3s), threeporbitals (3p), and five dorbitals (3d), 18 electrons


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p-Orbitals

3p orbitals within a given shell are oriented inperpendicular directions:px,py, andpz Lobes of aporbital are separated by region of zero

electron density, a node


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1.3 Atomic Structure: Electron

Configurations

Ground-state electron configuration (lowestenergy arrangement)listing of orbitalsoccupied by the electrons.

Rules 1:Lowest-energy orbitalsfirst: 1s2s2p3s3p4s3d(Aufbau (build-up)principle)


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Rules 2:

Electrons act as they are

spinning around axis,2 spin orientations: or

Only 2 electrons occupy

an orbitalPauli exclusion

principle

Rules 3:

If 2 or more empty orbitals

of equal energy areavailable, 1 electronoccupies with parallelspins until orbitals arehalf-full

Hundsrule


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1.4 Development of Chemical Bonding

Theory

Atoms form bonds because the compound thatresults is more stable than the separate atoms

Ionic bonds in salts form as a result of electrontransfers (example: Na+Cl-)

Organic compounds have covalent bonds fromsharing electrons (G. N. Lewis, 1916)


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Kekul and Couper independently observed thatcarbon always has four bonds

van't Hoff and Le Bel proposed that the four bonds ofcarbon have specific spatial directions Atoms surround carbon as corners of a tetrahedron


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Lewis structures(electron dot) show valence electrons of

an atom as dots Hydrogen has one dot, representing its 1selectron Carbon has four dots (2s22p2)

Kekule structures(line-bond structures) have a line drawnbetween two atoms indicating a 2 electron covalent bond.

Method of indicating covalent bonds


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Note that:

Stable molecule = noble gas configuration

completed shell, octet (eight dots) for main-group atoms, twofor hydrogen

Depends on how many additional valence electrons to

reach noble-gas configuration. Carbon has four valence electrons (2s22p2), forming four

bonds (CH4). Nitrogen has five valence electrons (2s22p3) but forms only

three bonds (NH3). Oxygen has six valence electrons (2s22p4) but forms two

bonds (H2O)


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Nonbonding electron

Valence electrons not used in bonding are callednonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia (NH3)

Shares six valence electrons in three covalent bonds and

remaining two valence electrons are nonbonding lone pair


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Drawing Structures

Drawing every bond in organic molecule can

become tedious. Several shorthand methods have been developed to

write structures.

1)Condensed structures dont have C-H or C-C single

bonds shown. They are understood.2)Skeletal structures


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Skeletal structure

3 General Rules:1) Carbon atoms arent usually shown. Instead a

carbon atom is assumed to be at eachintersection of two lines (bonds) and at the end

of each line.

2) Hydrogen atoms bonded to carbon arent

shown.

3) Atoms other than carbon and hydrogen areshown.


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1.5 The Nature of Chemical Bonds:

Valence Bond Theory

2 models to describe covalent bonding: Valence bond theory Molecular orbital theory

Covalent bond forms when two atoms approach eachother closely so that a singly occupied orbital on one

atom overlapsa singly occupied orbital on the other atom

H-H bond is cylindrically symmetrical, bond formed by thehead-on overlap of 2 atomic orbitals are called sigma ()bonds


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Bond Energy

Reaction 2 H

H2releases 436 kJ/mol Product has 436 kJ/mol less energy than two

atoms: HH has bond strengthof 436 kJ/mol.(1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)


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Bond Length

Distance betweennuclei that leadsto maximumstability

If too close, theyrepel becauseboth are positivelycharged

If too far apart,bonding is weak


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1.6 sp3 Orbitals Carbon has 4 valence electrons (2s22p2) In CH4, all CH bonds are identical (tetrahedral) sp3hybrid orbitals:sorbital and threeporbitals

combine to form four equivalent, unsymmetrical,tetrahedral orbitals (sppp = sp3), Pauling (1931)


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Structure of methane

sp3orbitals on C overlap with 1sorbitals on 4 Hatoms to form four identical C-H bonds

Each CH bond has a strength of 436 (438)kJ/mol and length of 109 pm

Bond anglebonds specific geometry,methane: each HCH is 109.5, the tetrahedralangle.


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Structure of ethane

Two Cs bond to each other by overlap of an sp3orbital from each

Three sp3orbitals on each C overlap with H 1sorbitals to form six CH bonds

CH bond strength in ethane 423 kJ/mol CC bond is 154 pm long and strength is 376 kJ/mol

All bond angles of ethane are tetrahedral value of109.5


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1.7 sp2 Orbitals sp2hybrid orbitals: 2sorbital combines with two2p

orbitals, giving 3 orbitals (spp = sp2). This results in adouble bond.sp2orbitals are in a plane with120 angles Remainingporbital is perpendicular to the plane


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Bonds from sp2hybrid orbitals

Two sp2

-hybridized orbitals overlap to form a

bondporbitals overlap side-to-sideto format a pi ()

bond

sp2sp

2

bond and 2p2pbond result insharing four electrons and formation of C-Cdouble bond

Electrons in the bond are centered between

nuclei Electrons in the bond occupy regions are on

either side of a line between nuclei


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Structure of ethylene

H atoms form bonds with four sp2orbitals HCH and HCC bond angles of about 120 CC double bond in ethylene shorter and stronger

than single bond in ethane Ethylene C=C bond length 134 pm (CC 154 pm)


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1.8 spOrbitals C-C a triplebond sharing six electrons Carbon 2sorbital hybridizes with a singlep

orbital giving two sphybrids twoporbitals remain unchanged

sporbitals are linear, 180 apart onx-axis Twoporbitals are perpendicular on the y-axis

and the z-axis


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Orbitals of acetylene

Two sphybrid orbitals from each C form spsp

bondpzorbitals from each C form apzpzbond by

sideways overlap andpyorbitals overlap similarly


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Bonding in acetylene

Sharing of six electrons forms C C

Two sporbitals form bonds with hydrogens


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Comparison of C-C and C-Hbonds

1 9 Hybridization of Nitrogen and


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1.9 Hybridization of Nitrogen and

Oxygen Elements other than C can have hybridized orbitals HNH bond angle in ammonia (NH3) 107.3 C-N-H bond angle is 110.3 Ns orbitals (sppp) hybridize to form four sp3orbitals One sp3orbital is occupied by two nonbonding

electrons, and three sp3orbitals have one electroneach, forming bonds to H and CH3.

1 10 The nature of chemical


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1.10 The nature of chemical

bonds: Molecular Orbital Theory

A molecularorbital(MO): where electrons are mostlikely to be found (specific energy and generalshape) in a molecule

Additive combination (bonding) MO is lower inenergy

Subtractive combination (antibonding) MO is higherenergy


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Molecular Orbitals in Ethylene

The

bonding MO is from combiningporbitallobes with the same algebraic sign

The antibonding MO is from combining lobeswith opposite signs

Only bonding MO is occupied


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Summary

Organic chemistrychemistry of carbon compounds Atom: positively charged nucleus surrounded by

negatively charged electrons Electronic structure of an atom described by wave

equation Electrons occupy orbitalsaround the nucleus. Different orbitals have different energy levels and different

shapes sorbitals are spherical, porbitals are dumbbell-shaped

Covalent bonds- electron pair is shared between atoms


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Valence bond theory - electron sharing occurs byoverlap of two atomic orbitals

Molecular orbital (MO) theory, - bonds result fromcombination of atomic orbitals to give molecular orbitals,which belong to the entire molecule

Sigma () bonds- Circular cross-section and are formedby head-on interaction

Pi () bondsdumbbell shape from sidewaysinteraction ofporbitals


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Carbon uses hybrid orbitals to form bonds in organic

molecules. In single bonds with tetrahedral geometry, carbon has foursp3hybrid orbitals

In double bonds with planar geometry, carbon uses threeequivalent sp2hybrid orbitalsand one unhybridizedp

orbital Carbon uses two equivalent sphybrid orbitalsto form atriple bond with linear geometry, with two unhybridizedporbitals

Atoms such as nitrogen and oxygen hybridize to formstrong, oriented bonds The nitrogen atom in ammonia and the oxygen atom in

water are sp3-hybridized

hydrocarbon chemistry lecture notes

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