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

Chapter 6 - shapes of molecules + intermolecular forces 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

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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

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3
Q

What is covalent bonding

A

non-metals sharing a pair of electrons

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4
Q

Actual Electron Capacity

A

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

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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

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6
Q

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

A

2.5°

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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

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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

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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

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10
Q

What does a solid line represent (3D bonding)

A

a bond in the plane of the paper

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11
Q

What does a solid wedge represent (3D bonding)

A

a bond coming out of the plane of the paper

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12
Q

What does a dotted wedge represent (3D bonding)

A

a bond going into the plane of the paper

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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

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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

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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

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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)

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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°)

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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

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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)

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20
Q

Factors that affect electronegativity

A

Nuclear charge (proton number)
Electron shielding
Atomic radius

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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)

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22
Q

General trend of electronegativity

A

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

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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

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24
Q

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

A

The protons and electrons are pulled closer together

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25
Why does electronegativity decrease down a group
due to the atomic size increasing, so the ponding pair of electrons are attracted less strongly to the nuclei of the atom
26
Covalent bond type electronegativity difference
0
27
Polar covalent bond type electronegativity difference
0-1.8
28
Ionic bond type electronegativity difference
>1.8
29
Why does one element have a negative dipole and one has a positive
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
30
Why are the noble gases not included in the Pauling Scale
they tend not to form compounds as they are so unreactive (full outer shell)
31
What group has the most electronegative atoms
The non-metals e.g. nitrogen, oxygen, chlorine, fluorine
32
What group has the least electronegative atoms
The group 1 metals e.g. lithium, sodium & potassium
33
What is a non polar bond
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
What is a polar bond
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
What does the δ delta sign mean
Small (e.g δ + is less than a full + charge)
36
The atom with the larger electronegativity has what charge
δ-
37
Why is H2O a polar molecule
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
why is CO2 a non-polar molecule
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
bond polarity arrow in polar molecules etc points towards…?
the more electronegative molecule
40
SbCl3 molecules are polar: explain why
- 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
What are the 3 types of intermolecular forces
- induced dipole-dipole interactions (London forces) - permanent dipole-dipole interactions - hydrogen bonding
42
Intermolecular forces are responsible for what properties?
physical properties e.g melting and boiling points
43
bonding e.g covalent bonds are responsible for what properties of a compound
the identity and chemical reactions
44
Bond Enthalpy of different intermolecular forces versus covalent bonds
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
About london forces (how strong they are etc.)
London forces are weak and exist between Every molecule, whether polar or non polar: they act between induced dipoles
46
How are dipoles induced? (4 steps)
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
the more electrons in each molecule… (to do with london forces)
- the larger the instantaneous and induced dipoles - the greater the induced dipole-dipole interactions - the stronger the attractive forces between molecules
48
what are permanent dipole-dipole forces
the weak intermolecular forces that arise only between permanently polar molecules
49
Why does F2 have a much lower boiling point than HCl even though they have the same amount of electrons and shape
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
What is a simple molecular substance
- made up of simple molecules (small units containing a definite number of atoms) e.g neon (Ne), hydrogen (H2), water (H2O) etc
51
What regular structure do simple molecules form in the solid state
a simple molecular lattice
52
What are the forces and bonding like in a simple molecular lattice
- the molecules are held in place by weak intermolecular forces - the atoms in the molecule are bonded strongly by covalent bonds
53
Ionic compounds have what type of forces of attraction?
Strong electrostatic forces of attraction between the ions (they also form a giant ionic lattice)
54
Simple molecular substances have what type of bonding
Covalent
55
Why do simple molecular substances have weak boiling points
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
Covalent substances with simple molecular structures can be put in what 2 categories
- polar - non polar
57
Why do non-polar simple molecular substances tend to be soluble in non-polar solvents
- 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
Why do simple molecular substances fend fo be insoluble in polar solvents
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
hydrophobic vs hydrophilic parts in biological molecules
- hydrophobic will be non-polar and comprised of a carbon chain - hydrophilic will be polar and contain electronegative atoms (usually oxygen)
60
what does the solubility of polar simple molar substances depend on
the strength of the dipole
61
why do polar covalent substances dissolve in polar solvents
because the polar solute and solvent molecules can attract each other
62
why are simple molecular structures non-conductors of electricity
- there are no mobile charged particles in the structures (with no charged particles to move, there is nothing to complete an electrical circuit)
63
Where are hydrogen bonds found
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
The hydrogen bond is shown by what type of line
A dashed line
65
Why is ice more solid than water
- 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
How many H bonds can each water molecule form
4 hydrogen bonds (as it has 2 lone pairs on the oxygen atom & 2 hydrogen atoms)
67
Why is water denser than ice
- 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
Why does water have such a high melting & boiling point
- 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
Anomalous properties of water
- a relatively high viscosity - relatively high surface tension
70
What bonds hold the double helix structure of DNA
hydrogen bonds
71
Hydrogen bonding in the double helix can only take place between what types of bases
A purine & a pyrimidine base
72
What are van der Waal’s forces
another term to describe induced dipole-dipole interactions
73
What are the 3 electronegative atoms that hydrogen atoms attach to to form H bonds
Oxygen, Nitrogen, Fluorine H-O, H-N, H-F
74
Why does the B.P. of Halogens increase as you go down the group
- number of electrons increases and therefore stronger London forces

Chemistry (8 decks)

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