ORGO CONTENT Exam 1

organic compounds

from LIVING organisms

• compounds that contain CARBON

• with a vital force

inorganic compounds

from minerals

• without a vital force

what makes carbon so special?

carbon is in the middle

• likes to share electrons

how does carbon “act”

carbon acts differently based on the conditions to reach the octet rule

atoms to the right of carbon

like to GAIN electrons to reach the octet

• accept electrons

atoms to the left of carbon

like to GIVE UP electrons to reach octet

• give away electrons

the struture of an atom

1. nucleus → protons + neutrons

2. electron cloud

‘atomic number’ formula

= the # of protons

neutral carbon

• has 6 protons

• has 6 electrons

what do all carbons have

the same atomic number = # of protons

what can carbons atoms have differently

they can have different mass numbers

mass number formula

= (# of protons) + (# of neutrons)

isotopes

• have different mass numbers

• have the same atomic number

4 shells of atomic orbitals

1. s

2. s, p

3. s, p, d

4. s, p, d, f

number of atomic orbitals in each shell

1. 1

2. 1, 3

3. 1, 3, 5

4. 1, 3, 5, 7

max number of electrons in each shell

1. 2 electrons

2. 8 electrons

3. 18 electrons

4. 32 electrons

what is the first shell closest to

the nucleus

the closer the atomic orbital is to the nucleus….

the lower its energy

aubfau principal

an electron goes into the atomic orbital with the LOWEST energy

• 1s < 2s < 2p < 3s < 3p < 3d

pauli exclusion principal

no more than 2 electrons can be in an atomic orbital

hunds rule

an electron goes into an empty degenerate orbial rather than pairing up

when is an atom most stable

if its outer shell is either filled or contains 8 electrons

electrons of atoms in the first column of the periodic table

lose an electron

• lithium & sodium achieve a filled outer shell by losing an electron

electrons in hydrogen

a hydrogen atm achieves an empty shell by…

• losing an electron

• or a filled outer shell by gaining an electron

a bond formed by sharing electrons

a covalent bond

nonpolar covalent bond

bonded atoms are the same or have similar electronegativities

nonpolar covalent bond electronegativity difference

electronegativity difference < 0.5

polar covalent bonds

boded atomas have different electronegativities

polar covalent bond electronegativity difference

electronegativity difference 0.5-1.9

electronegativity difference > 1.9

electrons are not shared

• atoms are held together by the attraction of opposite charges

dipole moment

(size of the charge) x (the distance between the charges)

electronegativity and periodic table

• electronegativity increases across (to right)

• electronegativity decreases down

electronegativitiy and dipole moment

the greater the difference in electronegativity…

• greater the dipole moment

• the more polar the bond

electrostatic potential maps

1. red = most negative electrostatic potential

2. blue = most positive electrostatic potential

formal charge equation

FC = (# VE) - (# LP + # bonds)

how many bonds does carbon form

4 bonds

what happens if carbon DOES NOT form 4 bonds

it has a charge (or its a radical)

how many bonds & LP does nitrogen form

3 bonds & 1 LP

what happens if nitrogen DOES NOT form 3 bonds

it is charged

how many bonds (+ LPs) do halogens form

1 bond & 3 LPs

what should the # of bonds + # of LPs equal

4

lewis structures

show all valence electrons (both bonding and lone pairs) using dots

• provides a complete picture of electron distribution

kekule structures

a simplified version of lewis structures where shared electron pairs (covalent bonds) are represented by solid lines

• lone pairs and formal charges are typically omitted

condensed structures

eliminates most or all of bond lines, relying instead on grouping atoms together to communicate the molecule's connectivity and save space

skeletal structures

show the carbon-carbon bonds as lines, but do not show the carbons or the hydrogens that are bonded to the carbons

atomic orbital

the region of space around the nucleuswhere an electron is most apt to be found

how does an electron behave

like a standing wave

• upward displacement and downward displacement are in opposite phases (+/-)

‘s’ atomic orbitals

a spherical region surrounding an atomic nucleus where electron is most likely to be found

• represents the lowest energy state for any principal energy level

• each s subshell can hold a maximum of two electrons

three ‘p’ atomic orbitals

a region in an atom where there is a high probability of finding an electron

• the lobes of a ‘p’ atomic orbital have opposite phases

forming a sigma bond

the overlapping of ‘s’ orbitals

• formed by the head-to-head (or end-to-end) overlap

• allows electron density to be shared dbetween nuclei of two bonding atoms

constructive combintion (waves)

waves reinforce each other, resulting in bonding

destructive combination (waves)

waves cancel each other, and no bond forms

“orbitals are conserved” meaning

# of molecular orbitals = # of atomic orbitals combined

pi bond

side-to-side overlap of in-phase p orbitals

CH4 name

methane

CH4 bonds & angles

• 4 C-H bonds (same length)

• all bond angles are the same (109.5 degrees)

what must carbon do in order to form 4 bonds (CH4)

carbon must promote an electron

• an electron in the ‘s’ orbital is promoted to the empty ‘p’ orbital

orbital mixing to form hybrid orbitals (CH4)

• 4 orbitals are hybridized

• 4 hybrid orbitals are formed

• sp³ orbital has a large lobe and a small lobe

the hybridization of carbon in CH4

sp³

• carbon is a tetrahedral

• tetrahedral bond angle is 109.5

C2H6

ethane

• sp³ hybridization

bonding in ethane (C2H6)

25% s and 75% p

• C=C bond formed by sp²-sp² overlap

• C-H bond formed by sp³-s overlap

C2H4

ethylene (ethene)

• sp² hybridization

bonding in ethylene (C2H4)

35% s and 67% p

• C=C double bond formed by sp²-sp² overlap

• C-H bonds formed by sp²-s overlap

• double bond has 1 pi bond and 1 sigma bond

C2H2

ethyne (acetylene)

• sp hybridization

bonding in acetylene (C2H2)

50% s and 50% p

• C≡C triple bond formed by sp-sp overlap

• C-H bond formed by sp-s overlap

• triple bond is 2 pi bonds and 1 sigma bond

hybridization of methyl anion (CH3)

sp³ hybridization

• lone-pairs on C are in sp³ orbital

• C-H bond formed by sp³-s overlap

NH3

ammonia

• sp³ hybridization

bonding in ammonia (NH3)

• nitrogen has 3 unpaired valence electrons and forms 3 bonds

• lone-pair electrons on N use sp³

• N-H bonds formed by overlap of an sp³-s orbital

NH4+

ammonium ion

• sp³ hybridization

water

H2O

• sp³ hybridization

bonding of water (H2O)

• oxygen has 2 unpaired valence elections and forms 2 bonds

• oxygen lone-pair electrons ar ein sp³ orbitals

• O-H bonds in sp³-s bonding

halides

group 9

• florine through iodine

length and strength of hydrogen-halide bonds

1. H-F = length of 0.917

2. H-I = length of 1.609

• as you go down the halides, length increases but bond strength decreases

sp³ hybridization

4 electron domains

• One ‘s’ and three ‘p’ orbitals hybridize

sp² hybridization

3 electron domains

• One ‘s’ and two ‘p’ orbitals hybridize

• One unhybridized ‘p’ orbital remains to form a pi bond

sp hybridization

2 electron domains

• One ‘s’ and one ‘p’ orbital hybridize

• Two unhybridized ‘p’ orbitals remain to form two pi bonds

single-bond vs double-bond vs triple-bond orbitals

1. single-bond = one sigma bond

2. double-bond = one sigma bond & one pi bond

3. triple-bond = one sigma bond & two pi bonds

bond strength & bond length

the more bonds holding 2 atoms together, the stronger and shorter it is

• more electron density = stronger & shorter

‘s’ orbital character and length/strength

the more ‘s’ character, the stronger and short the bond

• more ‘s’ charcter = greater bond angle

the shorter the bond….

the stronger the bond

the longer the bond….

the weaker the bond

the greater the electron density in region of orbital overlap…..

the stronger the bond

pi vs sigma bond strength

the pi bond is WEAKER than a sigma bond

canceling dipole moments

when all identical polar bonds are symmetrically arranged around a central atom, with no lone electron pairs