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organic chemistry Flashcards

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1
Q

what is an atomic orbital

A

a region around the nucleus that can hold up to two electrons with opposite spins

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2
Q

what are the different types of orbital

A

s f d p

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3
Q

how many s orbitals are there in one sub level

A

1

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4
Q

how many p orbitals are there in one sub level

A

3

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5
Q

how many d orbitals are there in one sub level

A

5

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6
Q

how many f orbitals are there in one sub level

A

7

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7
Q

what is Aufbau’s principle

A

electrons enter at the lowest energy orbital available

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8
Q

what is Hund’s rule?

A

electrons prefer to occupy orbitals on their own and only pair up when no empty orbitals of the same energy level are available

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9
Q

explain how the electronic configuration for an ion would work

A

highest energy electrons are lost when and ion is formed
4s lost before 3d

  1. remove outer shell electrons first
  2. remove p-electrons, then s , then d
  3. remove paired electrons before unpaired electrons in the same sub lvl
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10
Q

what is an ionic bond?

A

electrostatic attraction between oppositely charged ions which are formed from the transfer of electrons

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11
Q

what is a covalent bond

A

strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

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12
Q

what is a dative covalent bond?

A

a pair of electron shared between two atoms, one of which provides both electrons to the bond

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13
Q

what is a sigma bond

A

a bond resulting from the direct overlap of two orbitals along the internuclear axis or between the two bonding nuclei

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14
Q

what is a pi bond

A

a bond resulting from overlap of adjacent atomic orbitals above and below the two bonding nuclei

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15
Q

define electronegativity

A

relative ability of an atom to attract the two electrons in a covalent bond

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16
Q

what are the factors which affect electronegativity

A
  • nuclear charge
  • atomic radi
  • shielding
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17
Q

how does nuclear charge impact electronegativity

A

more protons=stronger attraction between nucleus and bonding pair of electronsh

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18
Q

how does atomic radi impact electronegativity

A

closer to nucleus=stronger attraction between nucleus and bonding pair of electrons

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19
Q

how does shielding impact electronegativity

A

less shells of electrons between nucleus and electrons=less shielding=stronger attraction between nucleus and bonding pair of electrons

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20
Q

how does electronegativity change down a group and why

A

it decreases

atomic radi increases=more shielding=less attraction between nucleus and bonding pair electrons

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21
Q

how does electronegativity change across a period and why

A

it increases

atomic radi decreases=increased nuclear charge = stronger attraction between nucleus and bonding pair of electrons

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22
Q

describe non polar covalent bonds

A
  • two atom in a covalent bond have the same electronegativity (less than 0.4 difference in electronegativity)
  • two electrons are shared equally as both atoms attract with a similar power therefore the electrons are midway in the bond
  • eg cl2
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23
Q

describe polar covalent bonds

A
  • two atoms in a covalent bond have different electronegativity

-covalent bond where the two electrons are not shared equally (one has greater attraction)

-the more electronegative atom will have a greater share and become slightly negative, the other, slightly positive

  • eg hcl
24
Q

what is a bond dipole moment

A

measure of the strength and direction of the polarity in the bond

the bigger the difference in electronegativity, the bigger the bond dipole moment

25
Q

describe electron pair repulsion

A

bonding pairs
-involved in boning and repel e/o equally

lone pairs
-not involved in bonding
- repel e/o more than bonding pairs (more electron dense)
- each lone pair reduces the bond angle by 2.5 degrees

26
Q

what is meant by ‘shrinking the octet’

A

in the outer shell neither beryllium nor boron have enough electrons to form an octet
- unpaired electrons will pair up and so the element will not achieve an octet

27
Q

what is meant by the ‘expansion of octet’

A

phosphorus can have 5 bonds
- another expansion: sulphur

28
Q

what are london forces

A

electrons in a molecule are in a state of constant motion
- the distribution of electrons will not be symmetrical there is likely to be a slight surplus on one of the atoms
- this is known as temporary dipole (lasts for a short time as electrons constantly move)

29
Q

what is an induced dipole

A

electrons on adjacent molecules are repelled by the negative parts of the dipole and attracted to the positive part and move accordingly —> this results in an attraction between the 2 molecules (london forces)

30
Q

which two factors affect the strength of london forces

A
  • number of electrons in the molecule
  • surface are of the molecule
31
Q

what is a permanent dipole

A
  • polar molecules have polar bonds
  • in polar molecules the dipoles cancel out
32
Q

explain how hydrogen’s lack of electron density allows it to form dipole bonds

A

as hydrogen has almost no electron density it is able to approach electronegative atoms on adjacent molecules very closely and form strong intermolecular dipole bonds

33
Q

how do you draw hydrogen bonds

A
  1. relevant dipoles
  2. dashed hydrogen bond between the lone pair and hydrogen
  3. h bond and covalent bond have an angle of 180
34
Q

describe the arrangement of water molecules in ice

A
  • water molecules arrange themselves in order to maximise the amount hydrogen bonding between the molecules
  • the results in a very open hexagonal structure with large spaces within crystals, this accounts for its low density
35
Q

describe the arrangement of water molecules as ice melts

A

the structure collapses into the open spaces , this causes water to occupy less space and thus becoming more dense

36
Q

describe how the surface area of molecules can impact the strength of london forces

A
  • the larger the surface area of a molecule, the more contact it will have with adjacent molecules
  • the more contact, the greater the ability to induce a dipole in an adjacent molecule and the greater the london force and therefore greater melting and boiling points
37
Q

describe how the number of electrons in a molecule can impact the strength london forces

A
  • the greater the number of electrons in a molecule, the greater the likelihood of
    a distortion
  • the greater the frequency and magnitude of the temporary
    dipoles
  • the London Forces between the molecules are stronger and
    the melting and boiling points are larger
38
Q

how can you decide if a bond is polar or non polar

A

if the difference in electronegativity is:
- 0.4 or less = non polar
- 0.4 - 1.8 = polar

39
Q

what is a polar molecule

A
  • has polar bonds
  • the dipoles from the polar bond do not cancel out
40
Q

what is dipole-dipole bonding

A

an attraction between a permanent dipole on one molecule and permanent dipole on another molecule

41
Q

define first ionisation energy

A

the energy required to remove one electron from each atom in one mole of gaseous atoms to produce one mole of gaseous 1+ ions

42
Q

how does ionisation energy differ across a period

A

it increases because NC increases, AR decreases and shielding stays the same , increase attraction and so increased energy to remove

43
Q

why does the ionisation energy of aluminium dip

A

group 2 = s block, group 3= p block , this means group 2 is closer to the nucleus than group 3, therefore 3 requires less energy to remove

44
Q

Why does the ionisation energy of sulphur dip?

A

group 5 electrons lose from orbital with 2 electrons, rest lose from orbital with 1 and due to electron-electron repulsion, group 5 is easier to lose

45
Q

explain the difference in boiling points from silicon to argon

A

si- macromolecular, giant covalent structure, covalent bonds, lots of energy to break bonds

ar- monoatomic, very weak london forces, noble gas

46
Q

name the shape and angle of the molecule with 2 electron pairs

A

linear, 180

47
Q

name the shape and angle of the molecule with 3 bonding pairs and 0 lone pairs

A

trigonal planar, 120

48
Q

name the shape and angle of the molecule with 4 bonding pairs and 0 lone pairs

A

tetrahedral, 109.5

49
Q

name the shape and angle of the molecule with 5 bonding pairs and 0 lone pairs

A

trigonal bipyramidal, 120 and 90

50
Q

name the shape and angle of the molecule with 6 bonding pairs and 0 lone pairs

A

octahedral, 90

51
Q

name the shape and angle of the molecule with 2 bonding pairs and 1 lone pairs

A

bent v shape, 118

52
Q

name the shape and angle of the molecule with 3 bonding pairs and 1 lone pairs

A

trigonal pyramidal, 107

53
Q

name the shape and angle of the molecule with 2 bonding pairs and 2 lone pairs

A

bent v shape, 104.5

54
Q

name the shape and angle of the molecule with 4 bonding pairs and 1 lone pair

A

trigonal pyramidal, 119 and 89

55
Q

name the shape and angle of the molecule with 3 bonding pairs and 2 lone pairs

A

trigonal planar, 120 or 89

56
Q

name the shape and angle of the molecule with 5 bonding pairs and 1 lone pair

A

square pyramidal, 89

57
Q
A

chemistry (2 decks)

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