Chapter 1: Bonding and Isomerism

"A" group number in periodic table =

number of valence shell e-

coordinate covalent bond

occurs when BOTH e- come from one atom

when are atoms not neutral?

when they acquire or lose an electron

polar covalent bonds

where the e- are shared, but not equally

can molecules be polar?

yes - A molecule whose centers of positive & negative charge do not balance, will be polar.

Molecule

•A particle composed of 2 or more atoms bonded together.

Compound

•A substance composed of identical molecules which are composed of more than one element. Additionally: The word "compound" can be used to describe one molecule of the substance, or a collection of molecules of the substance.

consistent bonding pattern when forming covalent bonds

•Hydrogen has 1 bond and 0 unshared pairs of ve •Carbon has 4 bonds and has 0 unshared pairs of ve •Nitrogen has 3 bonds and 1 unshared pair of ve •Oxygen has 2 bonds and 2 unshared pairs of ve •Halogens have 1 bond and 3 unshared pairs of ve

Electron domain geometry

•arrangement of bonding and nonbonding electron pairs around the central atom - Sometimes just called "electron geometry"

Non-zero formal charges

•must always be included when drawing structural formulas. They must be drawn close to the atom they belong to. - put a circle around formal charges so that they are not mistaken for other + or - signs

covalent bonds (length)

are shorter than the sum of the atomic radii of the two atoms, which is a result of overlap or "merging" of the orbitals - a key feature of covalent bonds!

A functional group

arrangement of atoms with distinctive physical and chemical properties. (ie. They "function" the same way.)

Molecular geometry

arrangement of the bonded atoms around the central atom - ball and stick or space filling models are useful

what is the most electronegative element?

fluorine

regular arrow pointing right

forms or produces (implies 100%)

Because the e- density is symmetric about the internuclear axis in sigma bonds,

groups can rotate about the bond without breaking it

"Condensed" structures

often used to simplify structure drawing. Structures written this way can be conveyed without any special structure drawing software - can also call "abbreviated"

a double bond is composed of

one sigma and pi bond

Hydrocarbons

only contain C & H

d+ or d-

partial charge caused by polar covalent bonding

dipole arrows

points from the positive end of the bond to the negative end - also indicates partial charge

cation

positively charged species

Atoms contain

protons, neutrons and electrons

Electronegativity

the ability of an atom to attract bonding e- to itself.

Curved arrows are used in organic chemistry to show

- the movement of valence electrons. - The rest of the atom (or atoms) gets dragged along. - Curved arrow always show what the electrons are doing.

"Full" structures

- show every atom symbol - do not show the unshared electrons - can be determined by the octet rule

protons

+ 1 charge

electrons

- 1 charge located outside the nucleus

"Constitutional" isomers or "structural" isomers

- Compounds with the same chemical formula, but their atoms are not attached in the same order. (Same molecular formula, but different structural formulas) - have different physical and chemical properties

Where is the split between ionic and polar-covalent?

- Somewhere between 1.7 - 2.1 - subtract electronegativity values of each to find if the bond is polar or not - if the change is zero, its nonpolar - if the change is between 1.7 and 2.1, it's polar - if the change is greater than 2.1 it is ionic

two bonding theories

- Valence Bond and Molecular Orbital (MO) theory - Both use the idea of orbital overlap to create bonds. The main difference is when orbitals are allowed to merge.

Sigma (s) bonds

- are characterized by a region of high e-density along the internuclear axis - can be formed by the overlap of: 2 "s" orbitals, an "s" and "p" (end on), 2 "p" (end on), an "s" and a hybrid, or 2 hybrid orbitals.

covalent bond

- between two non-metals - only way 2 non-metals can both achieve a full octet of valence electrons is to share electrons

ionic bond

- electrostatic attraction between positive & negative ions resulting from an e- transfer - between a metal and nonmetal - resulting compound is an array (lattice or crystal structure) of positive & negative ions packed so attractive forces between ions of opposite charges are maximized, & repulsive forces between like charges are minimized.

hybrid orbital characteristics

- for a given atom, all hybrid orbitals have the same energy - The total number hybrid orbitals equals the total number of atomic orbitals used to make them because you can't create orbitals from nothing - Unhybridized "p" orbitals can be used to form p bonds (or remain empty)

pi (p) bond

- formed by the side to side overlap of 2 "p" orbitals - the e- density in a pi bond is not located on the internuclear axis, but rather on either "side" of it.

valence shell e-

- outer most shell of e- and are the only e-'s involved in chemical bonding - the highest energy "s" and "p" e-'s

element symbol is used to represent

- the nucleus and the core e-'s, and a dot is used for each valence e- - Four regions around the symbol (top, right, bottom and left) correspond to four orbitals, each can hold 2 e-

VSPER Theory

- theory uses the concept that like charges repel to explain the geometry that is actually observed in molecules - Electron pairs (in either bonds or lone pairs) are arranged to be as far apart as possible around the atom. This is because these are areas of high electrondensity, so they repel each other

good points of valence bond theory

1)Hybrid orbitals created by mixing s and p orbitals, which then form localized s & p bonds - simple model. 2)Geometries match VESPR very well.

is a triple bond longer or shorter than a double bond?

A triple bond is shorter than a double bond which is shorter than a single bond.

rules for line structures

1. Carbon and hydrogen atoms are not shown unless needed for special emphasis or clarity. 2. All atoms other than carbon and hydrogen are explicitly shown. 3. A carbon atom is assumed to be at the end of every line segment or at the intersection of two or more lines, which are used to depict bonds. 4. Multiple bonds are shown with multiple lines. 5. The proper number of hydrogen atoms to provide four bonds to each carbon atom is implied including formal charges. Hydrogen atoms on other atoms such as oxygen and nitrogen are explicitly indicated. 6. Lone pair electrons are not normally shown unless needed for special emphasis or clarity.

rules for drawing resonance structures

1.All atoms must be drawn in the same place. 2.Only unshared electrons and pi electrons can be moved. 3.The resonance structures are separated by a resonance arrow. 4.No atom can have more than 8 valence electrons. 5.If the resonance structures are not of equal stability, the more stable resonance structure makes a larger contribution to the hybrid.

Reaction mechanism

: Step by step sequence of events happening during a reaction. Curved Arrow:used to show the movement of electrons in a reaction mechanism

octet rule

Atoms tend to gain, lose, or share electrons to achieve a noble gas e- configuration, (8 valence e- (2 each in 4 orbitals).

degenerate hybrid orbitals (HO)

Atomic orbitals (AO) are combined in various portions to make degenerate hybrid orbitals

Alkanes

C-C single bonds

What is the EN difference for the C-H bond?

C-H bonds are non-polar for all practical purposes

to draw a lewis structure

Determine the number of: 1) nonbonding electron pairs 2) single bonds 3) double bonds 4) triple bonds

chemical properties are determined by

ELECTRONS

resonance structure

Each oxygen has a ½ negative charge, and each C-O bond is ½ the way between a single bond and a double bond.

problems of valence bond theory

Explains localized bonding - but resonance concepts have to be "added" for some molecules, i.e. does not explain delocalized bonding in benzene

are molecule and compound synonymous?

NO For example: A compound has a melting point, where as a molecule does not.

lone pairs or nonbonding electrons

Unshared e- shown as a pair of dots

are unshared electrons dawn in structural formulas?

Unshared electrons are rarely drawn in structural formulas. It is the formal charge that tells you how many electrons each atom has

Are pi or sigma bonds stronger?

The side to side overlap in p bonds is less efficient than for sbonds, so p bonds are weaker than s bonds. it is less efficient to have electrons above and below axis rather than concentrated on the axis like sigma bonds

VSEPR

Valence Shell Electron Pair Repulsion

hybridized atomic orbitals

a kind of blend or combination of two or more standard atomic orbitals

why does a covalent bond form if there isn't a transfer of electrons?

because the e- - nuc attraction is stronger than the nuc-nuc and e- - e-repulsions

why do diatomic molecules form a bond?

because they're. most stable in the diatomic form due to their outer shells being full as compared to when they were singular and not filled

double bonds

carbon can form two double bonds oxygen can form one double bond

nucleus

composed of protons and neutrons is small

dipole moment has units of

debye (D), where 1 D = 3.34x10-30 coul·m

G.N. Lewis

developed one of the earliest successful pictures of chemical bonding, stated as the octet rule

The magnitude of the charge on each atom is dependent on the

difference in EN of the atoms, and the bond length

most of the volume of an atom is due to the

electrons

One arrow pointing up and another pointing left below it parallel

equilibrium (Reaction can go forward and backward)

group 4A

has 4 e- in their valence shell, so form four covalent bonds.

polar molecule

has opposite charges separated by a distance; it is a dipole and has a dipole moment.

Group 5A

have 5 e- in their valence shell, so form three single bonds

Group 6A

have 6 e- in their valence shell, so form two covalent bonds

Group 7A

have 7 e- in their valence shell, so they form one covalent bond

Electrons will shift toward ______ electronegativity atoms

higher

comparing electron domain geometry and molecular geometry

identical in naming but begin to differ because molecular geometry doesn't count lone pairs of electrons as a bond (why there is bent, trigonal pyramidal, etc)

A single headed or "fishhook" arrow

indicates the movement of 1 electron to produce a free radical

A normal double headed curved arrow

indicates the movement of 2 electrons

two major types of bonds

ionic and covalent

what type of bond is a metal and nonmetal?

ionic bond

in order for a molecule to be polar:

it must contain at least one polar bond

dipole moment

magnitude of charge x distance between charges = Qr - The more polar the molecule the stronger the dipole moment - The molecular dipole moment is the vector sum of the bond dipole moments.

most of the mass of an atom is due to the

nucleus

formal charge =

number of valence electrons - ( number of unshared electrons + 1/2 the number of shared electrons )

Anion

negatively charged species

patterns will always be followed when the atoms are

neural That is to say, when they have no formal charges

Neutrons

no charge

Do all covalent bonds have the same length and stability of bonds?

no. it depends on: 1)the degree of orbital overlap and 2)the relative energies of the atomic orbitals which form the bond.

"Line" structures, or "line angle structures"

quickest and clearest way to draw organic structures.

Whether or not the molecule will be polar, depends on the

relative orientations of the polar bonds.

double-headed arrows

resonance

ex of hybridization

s = 1, p = 1, hybrid = 2sp s = 1, p = 2, hybrid = 3sp2 s = 1, p = 3, hybrid = 4sp3

triple bond

shares three pairs of e- between the two atoms as in the case of N2

The result of polar covalent bonding

the bonding e- pair spend more time near the more EN atom which means the more EN atom will aquire a permanent excess negative charge - the other atom acquires a permanent excess positive charge

The more e- pairs that atoms share

the closer the atoms are pulled together

The number of bonds a particular atom forms can be predicted by

the number of e- it needs to complete its octet

bonds are created in valence bond theory when

the orbitals overlap to produce s or p bonds

Ionic "bond lengths"

the sum of the ionic radii of the two ions (this is a "hard sphere" model of atoms)

Most (but not all) of the bonds in organic compounds are between

two non-metals

hybrid orbitals

used to explain many observed molecular geometries, pure "s" and "p" atomic orbitals are combined to produce a set of "hybrid" orbitals on atoms - then form bonds between atoms producing the correct geometry

Can pi bonds rotate?

very rigid, no rotating if it rotates, the bond will break