Chapter 1: Basics of Organic Chemistry Flashcards
what is an atomic orbital
- orbitals are regions of space around the nucleus where electrons are located
what is the difference between an s orbital and a p orbital
s orbital look like spheres
p orbitals look like dumbells ( more electrons need more space)
what do the colours represent in p orbitals
- the colour of the lobe represents its phase
- the phase is the sign of wavefunction ( not physically important) there is no relationship to charge
what is the Pauli exclusion principle
- no two electrons can have the same four quantum numbers ( cannot have more then 2 orbitals in the same space )
what is the Aufbau principle
- electrons fill orbitals in increasing order of energy ( lowest to highest
are orbital energies similar
- three p orbital energies are degenerate - they have exactly the same energies
- the lower is the most stable
valence vs core electrons
valence - electrons found in valance orbitals ( available for bonding / as lone pairs)
core - electrons found in core orbitals (not available for bonding / as lone pairs )
valence orbitals ve core orbitals
- valence = occupied orbitals of highest energies ( and any accompanying empty orbitals of similar energies)
- core = fully occupied orbitals of lowest energies
what is an ionic bond
- electrostatic attraction between atoms of opposit charges ( opposit charges attract - positive to negative)
what is a covalent bond
- electrons shared between atoms
- a lot stronger and you can mix and match between + - and + +
- more significant bond due to the chemical reaction taking place
what is the octet rule
- atoms create bonds to try to gain electrons
- they want a full shell ( to fill the valence orbitals)
- the elements always have the first orbital filled. the second and third orbital coencide with each element
- an element with fewer than 8 electrons are more reactive
sigma covalent bonds
covalent bonds can be divided into two groups
sigma covalent bonds are a result of direct head on overlap between two orbitals
- p orbitals cannot for covalent bonds
- only s orbital and hybrids orbitals can form covalent bonds
pi bonds result from indirect (side on ) overlap between two p orbitals (two overlapping areas
polarized bonds
- different elements have different affinities for electrons
- if the two atoms in a covalent bond are the same they share electrons equally
- if two atoms in a covelent are different they do not share electrons equally ( there is visibly a much larger end )
- if the electrons in a covalent bond are not shared equally the bond is polar
how are elements assigment electronegativity
- increasing from left to right
- decreasing from top to bottom
what is method 1 of indicating a bond is polar
- the more electronegative atom gets a S-
- the less electronegative atom gets S+
- pros? emphasises partial charges, unique symbols
-cons? works poorly in complex systems
what is method 2 do indidcate a bond is polar
a crossed arrow pointing towards the net parital negative
pros; shows net polarity of the whole molecule
cons ; symbol means something else in other fields
VSEPR theory
- valance shell electron pair repulsion the most stable three dimensional structure for a molecule is the one in which valence electron pairs are as far apart as possible
- amounts to keeping groups (attached atoms or lone pairs)
trigonal planar
- there are three groups then maximizing the distance makes a trigonal planar geometry
tetrahedral
- four groups then maximizing the distance
hybrid orbitals rules
- the number of hybrid orbitals generated is exactly equal to the number of atomic orbitals use to make them ( 1s orbital and 1,2,3 p orbitals)
- hybrid orbitals arise from the combination of an s orbital with one or more p orbitals (they all look the same 1 small and 1 large lobe) but they go off in different directions
lewis structures
- covalent bonds are lines between atoms
- all atoms are shown/labelled (atomic symbols)
formal charges
- some atoms carry formal chargers (either cation or anions)
- usually results from an incomplete octet a different number of bonds than usual (octet rule)
- Formal charge = (group # ) - (# of bonds) - (# of non bonded electrons)
Formal charges example: oxygen
- oxygen is in group ( has six valence electrons)
- ( 6-3) -2 = +1
hash and wedge notation
- regulat lines connect atoms that are on the same plane
- a hashed line connects to the attached atom behind the plane
- a wedged line connects to the attached atom in front of the plane
- should not be more than 90 apart
line angle structures
- just because lone pairs arent there/drawn doesnt mean they arent there
- need to know if there is lone pairs or not
