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A level chemistry > Module 2.2 electrons, bonding and structure > Flashcards

Module 2.2 electrons, bonding and structure Flashcards

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1
Q

Principal quantum numbers

A

The numbers given to shells

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2
Q

shells further away from the nucleus have

A

higher energy (larger principal quantum number) than shells closer to the nucleus

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3
Q

types of sub shell with number of orbitals and maximum electron capacity

A

s sub shell. 1 orbital. max e- = 2
p sub shell. 3 orbitals. max e- –> 3 x 2 = 6
d sub shell 5 orbitals. max e- –> 5 x 2 = 10
f sub shell 7 orbitals max e- –> 7 x 2 = 14

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4
Q

how many electrons can one orbital carry

A

2

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5
Q

shell, what sub-shells it contains and total no,. of electrons

A

1st: 1s total e- = 2
2nd: 2s 2p total e- : 2+(3x2) = 8
3rd: 3s 3p 3d total e- : 2 + (3x2)+ (5x2) = 18
4th: 4s 4p 4d 4f total e-: 2+ (3x2) + (5x2) + (7x2) = 32

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6
Q

1st shell: sub shells and total number of electrons

A

1st
sub-shell: 1s
total e- : 2

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7
Q

2nd shell: sub shells and total number of electrons

A

2nd
sub-shells: 2s 2p
total e- : 2 + (3x2) = 8

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8
Q

3rd shell: sub shells and total no. of electrons

A

3rd
sub-shells: 3s 3p 3d
total e- : 2 + (3x2) + (5x2) = 18

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9
Q

4th shell: sub shells and total no. of electrons

A

4th
subshells: 4s 4p 4d 4f
total e- : 2 + (3x2) + (5x2) + (7x2) = 32

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10
Q

total no. of electrons in 1st shell

A

2

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11
Q

total no of e- in 2nd shell

A

8

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12
Q

total no of e- in 3rd shell

A

18

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13
Q

total no of e- in 4th shell

A

32

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14
Q

sub-shell s: no. of orbitals and max no. of e-

and shape

A

s
orbitals = 1
max e- = 2
shape = spherical

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15
Q

sub shell p: no of orbitals and max no. of e-

and shape

A

p
orbitals = 3
max e- = 6
shape= three dumbbell shaped orbitals at right angles to one another

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16
Q

sub-shell d: no. of orbitals and max no. of e-

A

d
orbitals = 5
max e- = 10

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17
Q

sub-shell f: no. of orbitals and max no. of e-

A

f
orbitals = 7
max e- = 14

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18
Q

no. of e- in sub-shell s

A

2

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19
Q

no. of e- in sub-shell p

A

6

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20
Q

no. of e- in sub shell d

A

10

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21
Q

no. of e- in sub-shell f

A

14

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22
Q

what is an orbital

A

is a bit of space in which the electrons move in. orbitals within the same sub shell have the same energy.

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23
Q

spin of electrons in orbitals

A

orbitals contain 2 electrons each.
they spin in opposite directions
spin pairing

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24
Q

order in which shells fill

A

1s - 2s - 2p - 3s - 3p - 4S- 3D - 4p - 4d - 4f

main point: 4s fills before 3d

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25
Q

why does 4s fill before 3d

A

because 4s has a lower energy level than 3d even though its principle quantum number is bigger.

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26
Q

how do electrons fill orbitals in a sub shell

A

electrons fill orbitals singling before they start sharing

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27
Q

how are noble gas symbols used in sub shell notation

A

noble gas symbols like argon (Ar)
which e- configoration is: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6)
are sometimes used .
e.g calcium: 1s(2) 2s(2) 2p(6) 3s(2) 3s(2) 3p(6) 4s(2) can be written as: [Ar] 4s(2)

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28
Q

define ionic bond

A

is an electrostatic attraction between two oppositely charged ions

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29
Q

electrostatic attraction

A

is the force that holds -ve and +ve ions together

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30
Q

sodium chloride structure:

A

giant ionic lattice
basic repeat unit of Na+ and Cl-
it foms becacuse each ion is electrostatically attracted in all directions to ions of opposite charge.
cine shape

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31
Q

melting point of sodium chloride

A

is extremely high (801 oc) because of the very strong ionic bonds

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32
Q

3 principles of ionic structure that explains behaviour of ionic compounds

A
  1. ionic compounds conduct electricity when they’re
    molten or dissolved.
  2. ionic compounds have high melting and boiling points.
    giant ionic lattices are held together by strong
    electrostatic forces which take a lot of energy to
    overcome.
  3. ionic compounds dissolve in water.
    water molecules are polar and so interact with the
    charges of the ions.
33
Q

define covalent bonding

A

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

34
Q

drawing covalent bonding

A

like ven diagram with shared electrons in the middle

35
Q

boron trifluoride BF(3) is a special case in covalent bonding as

A

boron forms 3 covalent bonds with the 3 fluorine so that each fluorine has a full outer shell. however boron only ends up with 6 electrons in outer shell.

36
Q

sulfur hexafluoride SF(6) is a special case in covalent bonding as

A

sulfur forms a covalent bond with 6 fluorines so that each have full outer shell (8 e-).
to do this sulfur EXPAND THE OCTET
so that it has 12 electrons in its outer shell.

37
Q

average bond enthalpy

A

measures the energy required to break a covalent bond

the stronger a bond is the more energy required and so greater bond enthalpy.

38
Q

double - triple covalent bonds

A

multiple covalent bonds share more than one pairs of electrons between two atoms. e.g
O2 which is a double covalent bond and shares 2 sets of electrons.
N2 has 5 outer electrons so forms tripple bond to complete shell
CO2 carbon has 4 and needs another 4. each oxygen atoms share two electrons.

39
Q

dative covalent bonds

A

where both electrons come from one atom
e.g
ammonium NH4+
nitrogen atom donates a pair of electromns to H+

40
Q

how are dative covalent bonds drawn

A

as an arrow potting away from donor atom

41
Q

shapes of molecules

A

refer to real life flash cards

42
Q

linear molecules

A

2 electron pairs around central atom
no lone pairs
bond angle: 180

43
Q

trigonal planar

A

3 e- pairs around central atom
no lone pairs
bond angle: 120

44
Q

tetrahedral

A

4 e- pairs around central atom
no lone pairs
bond angle: 109.5

45
Q

trigonal pyramidal

A

4 e- pairs around central atom
1 lone pair
bond angle: 107

46
Q

nonlinear / bent

A

4 e- pairs around central atom
2 lone pairs
bond angle: 104.5

47
Q

trigonal bipyramidal

A

5 e- pairs around central atom
no lone pairs
bond angle: 90 / 120

48
Q

octahedral

A

6 e- pairs around central atom
no lone pairs
bond angle: 90

49
Q

angle size according to lone pair to bonding pair

A

lone pair / lone pair = angles are the biggest
lone pair/ bonding pair = second biggest
bonding pair/ bonding pair = angles are the smallest.

50
Q

which type of electron pair repels more

A

lone pairs of electrons repel more producing larger angles

51
Q

shape of a molecule depends on

A

type of electron pairs surrounding central atom

number of electron pairs

52
Q

electron pair repulsion theory of compounds which all have 4 pairs of electrons

A

4 bonding pairs = 109.5 o
1 lone 3 bonding = 107 o lone pair repels the bonding pairs effectively squeezing them together
2 lone 2 bonding = 104.5 o 2 lone pairs reduce bond angle even more

53
Q

electron pair repulsion theory

A

theory that predicts the molecular shape of molecules

54
Q

electronegativity

A

an atoms ability to attract the electron pair in a covalent bond is called electronegativity

55
Q

Pauling scale

A

mesures electronegativity. the greater an element pauling value, the higher its electronegativity.

56
Q

polar molecules

A

if two atoms have different electronegativities in a covalent bond the bonding electrons will be pulled towards the more electronegative atom. making the bond polar. as it causes a permanent dipole.
the greater the difference in electronegativity the more polar the bond is.

57
Q

dipole

A

dipole is the difference in charge between two atoms caused by a shift in electron density in the bond.

58
Q

carbon and hydrogen have similar

A

electronegativities and so are essentially non-polar

59
Q

how does arrangement of bonds change with permanent dipoles

A

the arrangement of polar bonds determines whether the molecule will have an overall dipole.

symmetrical arrangement cause dipoles to cancel out. so non polar

uneven arrangement causes causes molecule to have overall dipole as they do not cancel out and are polar

60
Q

the transition between ionic to convalent bonding is

A

gradual.
only bonds between atoms of the same element can ever be purely covalent e.g diatomic gases. because the electronegativity difference is 0.

very few compounds are completely ionic.

61
Q

how can u predict type of bonding

A

by looking a electrongativity difference
the higher it is the more ionic
the lower it is the more covalent.

62
Q

intermolecular forces are very

A

Weak

63
Q

name the types of intermolecular forces from strongest to weakest

A
  1. hydrogen bonding
  2. peranent dipole- dipole interactions (e.g ionic and
    polar)
  3. induced dipole-dipole or london forces
64
Q

what are induced dipole-dipole forces

A

found between all atoms and molecules
cause attraction.
electrons charge clouds are constantly moving causing temporary dipoles where an electron is nearer one atom than another. this dipole induces dipoles in the opposite direction on a neighbouring atom causing them to be attracted. domino effect.
they are created and destroyed constantly causes overall attraction

65
Q

variations on london forces strength

A
  • larger molecules have larger electron clouds causing
    stronger induced dipole - dipole forces.
    molecules with greater surface area also have stronger - london forces as they have bigger exposed electron
    cloud.
66
Q

boiling points and london forces

A

when boiling liquid you overcome inter molecular forces as they molecules turn to gas.
more energy needed to overcome stronger inter molecular forces the greater the boiling point

67
Q

induced dipole- dipole forces hold lattices

A

responsible for holding iodine in lattice
iodine atoms are held together by strong covanet bonds.
but the molecules are held in a lattice by weak induced dipole -dipole interactions

68
Q

how are permanent dipole - dipole interactions drawn

A

s+ and s- used to label the dipoles of the atom and dotted line used to show the weak electrostatic force of attraction i.e. week bond between s+ of one molecule and s- of another

69
Q

permanent dipole-dipole interactions happen in addition to..

A

induced dipole-dipole interactions (london forces)

70
Q

hydrogen bonding can only happen when hydrogen is covalently bonded to…

A

fluorine, nitrogen or oxygen.

71
Q

how does hydrogen bonding work?

A
  • Hydrogen has a high charge density because it is so
    small
  • fluorine, nitrogen and oxygen are very electronegative
  • the bond is so polarised that a weak bind forms
    between hydrogen of one molecule and a lone pair of
    electrons on the F, N, O in another molecule.
72
Q

molecules with hydrogen bonding usually contain which groups?

A

-OH
-NH
water and ammonia both have hydrogen bonding

73
Q

hydrogen bonding effect on properties of a substance

A
  • they are soluble in water

- higher boiling and freezing points than molecules of similar size

74
Q

why is ice less dense than water

A

because hydrogen bonds are relative long.

75
Q

explain why simple covalent compounds have __ melting and boiling points interms of inter molecular forces

A

they have low melting and boiling points because:
- inter-molecular forces are weak so don’t need much
energy to break.
- they are often liquids or gases at room temperature

76
Q

expain why polar molecules are _______ in water

A

polar molecules are soluble in water because:
- water is polar so only tends to dissolve other polar
substances.
- substances with hydrogen bonds can form hydrogen
bonds with water molecules, so will be soluble
- molecules that only have induced dipole-dipole forces
will be insoluble

77
Q

simple covalent compounds don’t conduct ________ because

A

simple covalent compounds don’t conduct electricity because:

- even though some have permanent dipoles overall covalent molecules are uncharged.

78
Q

when you melt or boil a simple covalent bond you have to overcome the

A

inter molecular forces that hold the molecules together

A level chemistry (5 decks)

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