Chapter 1: Basics of Organic Chemistry Flashcards

1
Q

what is an atomic orbital

A
  • orbitals are regions of space around the nucleus where electrons are located
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2
Q

what is the difference between an s orbital and a p orbital

A

s orbital look like spheres
p orbitals look like dumbells ( more electrons need more space)

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3
Q

what do the colours represent in p orbitals

A
  • the colour of the lobe represents its phase
  • the phase is the sign of wavefunction ( not physically important) there is no relationship to charge
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4
Q

what is the Pauli exclusion principle

A
  • no two electrons can have the same four quantum numbers ( cannot have more then 2 orbitals in the same space )
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5
Q

what is the Aufbau principle

A
  • electrons fill orbitals in increasing order of energy ( lowest to highest
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6
Q

are orbital energies similar

A
  • three p orbital energies are degenerate - they have exactly the same energies
  • the lower is the most stable
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7
Q

valence vs core electrons

A

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 )

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8
Q

valence orbitals ve core orbitals

A
  • valence = occupied orbitals of highest energies ( and any accompanying empty orbitals of similar energies)
  • core = fully occupied orbitals of lowest energies
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9
Q

what is an ionic bond

A
  • electrostatic attraction between atoms of opposit charges ( opposit charges attract - positive to negative)
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10
Q

what is a covalent bond

A
  • 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
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11
Q

what is the octet rule

A
  • 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
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12
Q

sigma covalent bonds

A

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

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13
Q

polarized bonds

A
  • 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
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14
Q

how are elements assigment electronegativity

A
  • increasing from left to right
  • decreasing from top to bottom
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15
Q

what is method 1 of indicating a bond is polar

A
  • 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
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16
Q

what is method 2 do indidcate a bond is polar

A

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

17
Q

VSEPR theory

A
  • 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)
18
Q

trigonal planar

A
  • there are three groups then maximizing the distance makes a trigonal planar geometry
19
Q

tetrahedral

A
  • four groups then maximizing the distance
20
Q

hybrid orbitals rules

A
  1. 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)
  2. 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
21
Q

lewis structures

A
  • covalent bonds are lines between atoms
  • all atoms are shown/labelled (atomic symbols)
22
Q

formal charges

A
  • 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)
23
Q

Formal charges example: oxygen

A
  • oxygen is in group ( has six valence electrons)
  • ( 6-3) -2 = +1
24
Q

hash and wedge notation

A
  • 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
25
Q

line angle structures

A
  • just because lone pairs arent there/drawn doesnt mean they arent there
  • need to know if there is lone pairs or not