Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry > Chapter 6 - shapes of molecules + intermolecular forces > Flashcards

Chapter 6 - shapes of molecules + intermolecular forces Flashcards

You may prefer our related Brainscape-certified flashcards:
AP® Chemistry flashcards Periodic Table flashcards Chemistry 101 flashcards Organic Chemistry 101 flashcards
1
Q

What is electron-pair repulsion theory

A

The shape of a molecule is determined by the electron pairs surrounding the central atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What is electron pair repulsion theory based off

A

That pairs of electrons repel all of the other electron pairs - so they move as far as possible to minimise the repulsion and thus holds the bonded atoms in a definite shape

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What is covalent bonding

A

non-metals sharing a pair of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Actual Electron Capacity

A

n = 1 : 2 electrons
n = 2 : 8 electrons
n = 3 : 18 electrons
n = 4 : 32 electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What are is a lone pair

A

a pair of electrons in the outer shell of an atom which aren’t involved in any of the bonding in the atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

For each additional lone pair, the bond angles decrease by…

A

2.5°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Why do lone pairs repel more

A

Because they are closer to the nucleus of the central atom and occupy more space than a bonded pair

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

types of electron pairs in increasing repulsion

A

Lowest repulsion
Bonded pair / bonded pair
Bonded pair / lone pair
Lone pair / lone pair
Highest repulsion

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What are sigma bonds (σ bonds)

A

σ bonds = the first bond between 2 atoms, formed by an overlap of orbitals directly between the bonding atoms

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

What does a solid line represent (3D bonding)

A

a bond in the plane of the paper

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What does a solid wedge represent (3D bonding)

A

a bond coming out of the plane of the paper

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What does a dotted wedge represent (3D bonding)

A

a bond going into the plane of the paper

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Shape and angle when 2BPE / bonding regions and 0LPE e.g C2O

A

Linear, 180°
e.g. O=C=O
Equal distance between 2 BP (all on the same plane & no LP to distort angle

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Shape & angle with 3 bonding regions and 0 LPE e.g. BF3

A

Trigonal Planar
120°
All BP are on the same plane and and there’s no LP’s so the angle isn’t distorted

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Shape & angle with 4 bonding regions and 0 LPE e.g. CH4

A

Tetrahedral
109.5°
1 C-H bond going out of the plane (solid wedge) and 3 bonds on the same plane
All BP angles are equal as there are no lone pairs

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Shape & angle with 6 bonding regions and 0 LPE e.g. SF6

A

Octahedral
90°
6 BP angles all equal (2 on plane, 2 behind, 2 out)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Shape & angle with 3 bonding regions and 1 LPE e.g. NH3

A

Trigonal Pyramid
107°
1 bond going into the plane, one going out and one on the plane + the lone pair at the top of the N (therefore distorting the angle by 2.5°)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Shape & angle with 2 bonding regions and 2 LPE e.g. H2O

A

Non-Linear (bent)
104.5°
Two lone pairs on top of the O and 2 bonds on the same plane going out diagonally

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

What is electronegativity

A

The measure of the tendency of an atom to attract a bonding pair of electrons in a covalent bond (the greater the electronegativity of the atom the more it attracts electrons towards it)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Factors that affect electronegativity

A

Nuclear charge (proton number)
Electron shielding
Atomic radius

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

What scale measures electronegativity

A

The Pauling scale
0-4 (e.g F has an EN of 4 & is there more the most electronegative element)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

General trend of electronegativity

A

Electronegativity increases as you go up and go to the right of the periodic table

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Why does electronegativity increase across a period

A

Because there are more protons (greater nuclear charge) so the bonding pair is attracted more
- Note: The shielding stays the same

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Why does atomic radius get smaller due to an increase nuclear charge

A

The protons and electrons are pulled closer together

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Why does electronegativity decrease down a group

A

due to the atomic size increasing, so the ponding pair of electrons are attracted less strongly to the nuclei of the atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Covalent bond type electronegativity difference

A

0

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

Polar covalent bond type electronegativity difference

A

0-1.8

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Ionic bond type electronegativity difference

A

> 1.8

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

Why does one element have a negative dipole and one has a positive

A

the negative one e.g F in F-H because it has a greater share of electrons, while H has a positive dipole as it has a smaller share of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

Why are the noble gases not included in the Pauling Scale

A

they tend not to form compounds as they are so unreactive (full outer shell)

31
Q

What group has the most electronegative atoms

A

The non-metals e.g. nitrogen, oxygen, chlorine, fluorine

32
Q

What group has the least electronegative atoms

A

The group 1 metals e.g. lithium, sodium & potassium

33
Q

What is a non polar bond

A

when the bonded electron pair is shared equally between the bonded atoms e.g.
- if the bonded atoms are the same
- if the bonded atoms have the same / similar electronegativity
(e.g. H-H, or Cl-Cl)

34
Q

What is a polar bond

A

when the bonded electron pair is shared unequally between the bonded atoms
- when the bonded atoms are different / have different electronegativity values, resulting in a polar covalent bond
e.g. HCl (Cl = δ- and H = δ+)

35
Q

What does the δ delta sign mean

A

Small (e.g δ + is less than a full + charge)

36
Q

The atom with the larger electronegativity has what charge

A

δ-

37
Q

Why is H2O a polar molecule

A

the two O-H bonds each have a permanent dipole
-> these dipoles act in different directions but don’t directly oppose each other as it’s a non-linear molecule (not symmetrical)
so the oxygen end of the molecule = δ- and the hydrogen end = δ+

38
Q

why is CO2 a non-polar molecule

A

the two C=O bonds each have a permanent dipole
-> the dipoles act in opposite directions and exactly oppose each other, so they cancel and the overall dipole is zero
δ- <- O=C=O -> δ-

39
Q

bond polarity arrow in polar molecules etc points towards…?

A

the more electronegative molecule

40
Q

SbCl3 molecules are polar: explain why

A
  • Sb and Cl have a difference in electronegativity so they have dipoles
  • the dipoles are not in opposite directions as the molecule is not symmetrical, so they don’t cancel out
41
Q

What are the 3 types of intermolecular forces

A
  • induced dipole-dipole interactions (London forces)
  • permanent dipole-dipole interactions
  • hydrogen bonding
42
Q

Intermolecular forces are responsible for what properties?

A

physical properties e.g melting and boiling points

43
Q

bonding e.g covalent bonds are responsible for what properties of a compound

A

the identity and chemical reactions

44
Q

Bond Enthalpy of different intermolecular forces versus covalent bonds

A

Covalent bonds are significantly stronger
London forces: 1-10
Permanent dipole-dipole interactions: 3-25
Hydrogen bonds: 10-40
Single covalent bonds: 150-500

45
Q

About london forces (how strong they are etc.)

A

London forces are weak and exist between Every molecule, whether polar or non polar: they act between induced dipoles

46
Q

How are dipoles induced? (4 steps)

A
  1. Movement of electrons produces a changing dipole in a molecule
  2. At any instant, and instantaneous dipole will exist, but it’s position is constantly shifting
  3. The instantaneous dipole induces a dipole on a neighbouring molecule
  4. The induced dipole further induces dipoles on neighbouring molecules, which then attract one another
47
Q

the more electrons in each molecule… (to do with london forces)

A
  • the larger the instantaneous and induced dipoles
  • the greater the induced dipole-dipole interactions
  • the stronger the attractive forces between molecules
48
Q

what are permanent dipole-dipole forces

A

the weak intermolecular forces that arise only between permanently polar molecules

49
Q

Why does F2 have a much lower boiling point than HCl even though they have the same amount of electrons and shape

A

F2 is non polar so only has London forces between the molecules
HCl is polar and has London + permanent dipole-dipole interactions:
-> extra energy is needed to break the additional permanent interactions so the b.p. is much higher

50
Q

What is a simple molecular substance

A
  • made up of simple molecules (small units containing a definite number of atoms) e.g neon (Ne), hydrogen (H2), water (H2O) etc
51
Q

What regular structure do simple molecules form in the solid state

A

a simple molecular lattice

52
Q

What are the forces and bonding like in a simple molecular lattice

A
  • the molecules are held in place by weak intermolecular forces
  • the atoms in the molecule are bonded strongly by covalent bonds
53
Q

Ionic compounds have what type of forces of attraction?

A

Strong electrostatic forces of attraction between the ions (they also form a giant ionic lattice)

54
Q

Simple molecular substances have what type of bonding

A

Covalent

55
Q

Why do simple molecular substances have weak boiling points

A

they have weak intermolecular forces
(when a simple molecular lattice is broken during melting:
- only the weak intermolecular forces break
- the covalent bonds are strong and don’t break)

56
Q

Covalent substances with simple molecular structures can be put in what 2 categories

A
  • polar
  • non polar
57
Q

Why do non-polar simple molecular substances tend to be soluble in non-polar solvents

A
  • when a simple molecular compound is added to a non-polar solvent (e.g. hexane) intermolecular forced form between the molecules and the solvent
  • the interactions weaken the intermolecular forces in the simple molecular lattice, the intermolecular forces break and the compound dissolves
58
Q

Why do simple molecular substances fend fo be insoluble in polar solvents

A

when the two are mixed together there’s little interaction between the molecules in the lattice and the solvent molecules
-> this is as the intermolecular bonding within the polar solvent is too strong to be broken

59
Q

hydrophobic vs hydrophilic parts in biological molecules

A
  • hydrophobic will be non-polar and comprised of a carbon chain
  • hydrophilic will be polar and contain electronegative atoms (usually oxygen)
60
Q

what does the solubility of polar simple molar substances depend on

A

the strength of the dipole

61
Q

why do polar covalent substances dissolve in polar solvents

A

because the polar solute and solvent molecules can attract each other

62
Q

why are simple molecular structures non-conductors of electricity

A
  • there are no mobile charged particles in the structures
    (with no charged particles to move, there is nothing to complete an electrical circuit)
63
Q

Where are hydrogen bonds found

A

between molecules containing
- an electronegative atom with a lone pair of electrons e.g. oxygen, nitrogen, fluorine
- a hydrogen atom attached to an electronegative atom e.g H-O, H-N

64
Q

The hydrogen bond is shown by what type of line

A

A dashed line

65
Q

Why is ice more solid than water

A
  • H bonds hold the H2O molecules apart in an open lattice structure
    -> the water molecules in ice are further apart than in water
    therefore solid H2O is less dense than liquid and floats
66
Q

How many H bonds can each water molecule form

A

4 hydrogen bonds (as it has 2 lone pairs on the oxygen atom & 2 hydrogen atoms)

67
Q

Why is water denser than ice

A
  • H bonds extend outwards, holding the H2O molecules and forming an open tetrahedral lattice full of holes
    -> the holes in the lattice decrease the density of water: when ice melts, the ice lattice collapses and the molecules move close together, so the liquid is denser
68
Q

Why does water have such a high melting & boiling point

A
  • water has H bonds on top of it’s London forces
  • therefore more energy is needed to break the H bonds, so the points are higher than expected
    -> when the H bonds break (when ice boils) the ice lattice breaks to break the hydrogen arrangement
69
Q

Anomalous properties of water

A
  • a relatively high viscosity
  • relatively high surface tension
70
Q

What bonds hold the double helix structure of DNA

A

hydrogen bonds

71
Q

Hydrogen bonding in the double helix can only take place between what types of bases

A

A purine & a pyrimidine base

72
Q

What are van der Waal’s forces

A

another term to describe induced dipole-dipole interactions

73
Q

What are the 3 electronegative atoms that hydrogen atoms attach to to form H bonds

A

Oxygen, Nitrogen, Fluorine
H-O, H-N, H-F

74
Q

Why does the B.P. of Halogens increase as you go down the group

A
  • number of electrons increases and therefore stronger London forces

Chemistry (8 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions