Bonding and Structure - Topic 2 Flashcards
1
Q
What are London forces
A
The attractive forces between all atoms and molecules
2
Q
Why does an increase in pi bonds increase the London forces
A
As the bigger the electron cloud the more uneven electron distribution there is so there would be a much stronger imbalance in charge causing there to be a stronger dipole, so a higher polarizability so stronger LDN forces
3
Q
covalent bonding
A
the electrostatic attraction between the bonding shared pair of electrons and the two nuclei
4
Q
What is expansion of the octet
A
when an atom is able to have more than 8 valence electrons
5
Q
What are molecular orbitals
A
The overlap of atomic orbitals from separate atoms
6
Q
How many electrons can molecular orbitals hold
A
2
7
Q
Types of molecular orbitals
A
Sigma
Pi
8
Q
Where are sigma bonds found
A
Between atoms
9
Q
Where are pi bonds found
A
Below and above atoms
10
Q
what are sigma bonds
A
the direct overlap of two atomic orbitals
11
Q
How are pi bonds formed
A
When p orbitals in separate atoms overlap sidewards forming pi bonds above and below the atoms
12
Q
Properties of sigma bonds
A
- all single bonds are sigma bonds
- sigma bonds are free to rotate as rotating each nuclei has no impact on the bonding orbital
- have a high level of attraction between nuclei and shared electrons so are very strong
13
Q
How many pi and sigma bonds are there in double bonds
A
x1 sigma
X1 pi
14
Q
How many pi and sigma bonds are there in triple bonds
A
x1 sigma
x2 pi
15
Q
Why are pi bonds weaker than sigma bonds
A
As the electrons in the pi bond are further away from the nuclei compared to sigma bonds so has there is a weaker level of attraction making the bonds easier to break
16
Q
What is hybridisation
A
The concept of mixing w atomic orbitals to create a new type of hybridised orbital which has a different shape and energy level
17
Q
Number of hybridised and unhybridised orbitals and what is the hybridised orbital called in single bonds
A
0 unhybridised orbitals
4 hybridised orbitals
Name: sp3 orbitals
18
Q
Number of hybridised and unhybridised orbitals and what is the hybridised orbital called in double bonds
A
1 unhybridised p orbital
3 hybridised orbital
Name: sp2
19
Q
Number of hybridised and unhybridised orbitals and what is the hybridised orbital called in triple bonds
A
2 unhybridised p orbital
2 hybridised orbitals
Name: sp
20
Q
In terms of hybridised orbitals explain why carbon has a tetrahedral shape
A
Each sp3 orbital has the same energy so the electrons will have equal repulsion between all 4 orbitals causing the orbitals to point away from each other at maximum distances at the angle of 109.5 degrees between then
21
Q
No lone pairs 2 electron pairs geometric shape name
A
Linear
22
Q
Bond angle of linear molecule
A
180 degrees
23
Q
3 electron pairs no lone pairs name of molecule
A
Triagonal planar
24
Q
What are electron domains
A
Things attached to the central atoms e.g other atoms or lone pairs
25
What makes up a linear geometric shape
2 electron pairs no lone pairs
26
What makes up a triagonal planar shape
3 electron pairs no lone pairs
27
2 bonding pairs 1 lone pair geometrical shape
V shaped
28
2 bonding pairs 2 lone pairs shape
V shaped
29
Bond angle for Trigonal planar
120 degrees
30
Bond angle of v shape
104.5 degrees
31
What makes up a v shape
2 bonding pairs of electrons and 2 lone pairs
OR
2 bonding pairs and 1 lone pairs of electrons
32
4 bonding pairs of electrons and no lone pairs geometric shape
Tetrahedral
33
Tetrahedral bond angle
109.5 degrees
34
What makes up a tetrahedral shape
4 bonding pairs no lone pairs
35
3 bonding pairs 1 lone pair geometric shape
Trigonal pyramidal
36
trigonal pyramidal bond angle
107 degrees
37
3 bonding pairs 1 lone pairs geometric shape
Trigonal pyramidal shape
38
6 bonding pairs no lone pairs
Octahedral
39
What makes up an octahedral
6 bonding pairs
40
What makes up a triganol bipyramidal shape
5 bonding pairs no lone pairs
41
Triagonal bipyramidal bond angles
90 and 120 and 180
42
Octahedral bond angles
90 and 180
43
why can polar substances
i) molecules containing electronegative atoms
ii) NH3 HF and alcohols
dissolve in water
i) as the positive part of water attracts the electronegative part of the molecule and the negative part of the water attracts to the positive part of the molecule
2) as they all form hydrogen bonds between water molecules
44
trends of electronegativity
- increases across a period
- decreases down a group
44
why does electronegativity increase across a period
as the nuclear charge increases causing the atomic radius to decrease but the number of shielding remains constant so there is a stronger attraction between the nucleus and bonding pairs of electrons
45
why does electronegativity decrease down a group
as the number of shielding electrons increases the atomic radius increases so this means that the shared electron pair is further from the nucleus so less attraction between nucleus and bonding pairs of electrons
46
why cant non polar substances dissolve in water
These molecules don't have regions of partial positive or partial negative charge, so they aren't electrostatically attracted to water molecules and their weak intermolecular forces cannot disturb the strong hydrogen bonding between water molecules
47
why does solubility decrease as HC chain length increases in alcohols
- As when the length of the hydrocarbon chain increases, the London forces predominate
- the non-polar carbon chain cannot form hydrogen bonds, so as we increase the length a greater part of the molecule is unable to hydrogen bond with water, causing long chain alcohols to be less soluble than short-chain
48
what is required for a substance to dissolve
- soluble particles must be separated from each other and become surrounded by solvent molecules
- the forces of attraction between solute and solvent molecule must be strong enough to overcome the solvent-to-solvent forces and the solute-to-solute forces
49
definition for electronegativity
the ability of an atom to attract the bonding pairs of electrons in a covalent bond
50
what is a saturated solution
a solution that contains as much solute as possible at a particular temperature
51
definition for solubility
a measure of the concentration of a saturated solution of a solute at a specific temperature
52
factors affecting electronegativity
- nuclear charge
- atomic radius
- shielding
53
how does nuclear charge affect electronegativity
more protons stronger attraction between nucleus and bonding pairs of electrons
54
how does atomic radius affect electronegativity
closer to the nucleus the stronger the attraction between nucleus and bonding pairs of electrons
55
how does shielding affect electronegativity
the less shielding so the less repulsion and a stronger attraction between the nucleus and bonding pairs of electrons
56
what are the electronegative atoms
chlorine bromine and iodine
57
what atoms form hydrogen bonds
O F N
58
3 types of intermolecular forces
- induced dipole dipole interactions (London forces)
- permanent dipole dipole interactions
- hydrogen bonding
59
What causes induced dipoles to form
By the random movement of electrons causing an uneven distribution of electrons on one side of the atom. This forms an instantaneous dipole. The negative charge of the instantaneous dipoles repels the electrons in the other atom causing the electrons to move away from the atoms forming an induced dipole. Each atom them experiences a force of attraction between the positive and negative end of the atom, called a London force
60
How does an increase in electrons affect London forces and the boiling point of an atom
The more electrons there are the stronger the London forces so more energy is required to over come them, so the boiling point increases
61
how does the shape of a molecule affect the london forces
- the interactions between long thin molecules are stronger than those between short fat molecules because the interactions between long thin molecules can take effect over a larger surface area
- so the more branched the molecule the lower the surface area of contact and therefore the weaker London forces
62
What causes permanent dipole dipole interactions
This forms when there is an atom with a greater electronegativity so attracts pair of electrons in the covalent bond more strongly causing the electronegative atom to be slightly negative forming a permanent dipole. When two permanent dipoles attract each other they form forming a permanent dipole interaction
63
Why does tetrachloromethane have no overall permanent dipole
As the molecule is completely symmetrical so the bond polarities cancel out so it has no overall permanent dipole
64
Why does tetrachloromethane have a higher boiling point than trichloromethane even though it has no overall permanent dipole
As it has more electrons so stronger London forces so more energy is required to overcome the London forces
65
What causes hydrogen bonding
The electronegative element (nitrogen oxygen or fluorine) strongly attracts the pairs of electrons in the covalent bond causing it to be slightly negative and the hydrogen atom to have a slightly positive charge. The positive hydrogen atom of the other polar molecule (that forms hydrogen bonds) attracts to the lone pair of electrons in oxygen ,fluorine or nitrogen molecule forming a hydrogen bond
66
How many lone pairs of electrons in fluoride nitrogen and oxygen
3 - fluoride
2 - oxygen
1 - nitrogen
67
Conditions for hydrogen bonding
- we need a hydrogen atom bonded to a strongly electronegative element
- the electronegative atom has to have at least one lone pair of electrons
68
Water properties due to hydrogen bonding
- it has a high boiling and melting point due to the hydrogen bonds so requires a lot of energy to overcome hydrogen bonds
- solid form of water is less dense than liquid form
69
Why is ice less dense than water
As when water reaches its freezing point (0 degrees Celsius) water molecules are arranged in a lattice / hexagon
or hydrogen bonds are longer than covalent bonds so there are gaps between water molecules rather in liquid wate, water moelcules are closer
70
How does ice having a low density help organisms
Due to its low density it floats on water, insulating the water below preventing the water from freezing completely
71
definition of hydrogen bonding
a strong intermolecular force between delta-positive hydrogen atom covalently bonded to a F N or O and a lone pair of electrons on the delta negative of the O F or N atom of a nearby molecule
72
What is dative bonding
When an atom uses a lone pair of electrons to form a covalent bond
73
How is a dative bond shown in a displayed formula
Using an arrow
74
Why isn’t a carbon to carbon double bond stronger than a carbon to carbon single bond
As the double bond contains 1x sigma and 1x pi whilst the single bond contains 1x sigma bonds, and its easier to break the pi bond rather than a sigma bond as sigma bonds are stronger
75
what is ionic bonding
the strong electrostatic attraction between two oppositely charged ions
76
How are ionic bonds formed
When a metal transfers its electrons to a non metal so that they both have the same electron configuration of a noble gas causing them to be oppositely charged. And the electrostatic attraction between oppositely charged ions is called an ionic bond
77
What do the square brackets represent in ionic bonding drawings
That the charge is spread over the whole ion
78
How do ionic compounds/ionic crystals form
When a metal and non metal react the ions produced arrange them self into a giant ionic lattice forming an ionic compound
79
Properties of ionic compounds
- they have high mp and bp - due to the strong electrostatic forces of attraction so requires a lot of energy to overcome electrostatic forces of attraction
- they are soluble in polar substances such as water - as the water cluster around the ions and bind to them and the energy released when water molecules bind to the ions is enough to overcome the electrostatic forces of attractions holding the ions lattice together
- they don’t conduct electricity when solid - as the ions are in a fixed in place by the electrostatic forces of attraction so they cannot carry a charge
- hard brittle crystalline substances
80
Why does magnesium oxide have a higher mp than sodium chloride
- magnesium oxide is in group 2 so has a higher charge and a higher charge density so has stronger electrostatic forces of attraction so more energy is needed to overcome so has a higher mp
81
Why does solubility decrease as the charge of ionic compounds increases
As the hydration energy released when the ions are hydrated (water molecules binding to them) is not large enough to overcome the electrostatic forces of attraction holding the lattice together
82
isoelectronic definition
molecules and ions with the exact same number and arrangement of electrons
83
Properties for simple molecular substances
- they have low melting and boiling points - due to the weak London forces between molecules so require little energy to overcome
- do not conduct electricity as they don’t have any delocalised electrons or freely moving ions
- only polar simple substances are soluble in water
84
Why do giant covalent structures have high mp and bp
- have high melting and boiling points due to the strong covalent bonds between atoms so requires a lot of energy to overcome
85
How are the atoms arranged in diamond and what is its bond angle
Tetrahedral structure
Bond angle: 109.5
86
Properties of diamond
- high mp and bp
- does not conduct any electricity as every electron in its covalent bond are fixed between pairs of atoms so no delocalised electrons to act as charge carriers
- insoluble as solvents cannot disrupt the larger number of covalent bonds
- good thermal conductor due to the fixed covalent bonds meaning that the atoms close to the heat get hotter and move faster so the vibrations move quickly throughout the structure heating the whole structure
- very hard due to the strong covalent bonds operating in 3D dimensions
- sublimes at ordinary pressures due to strong covalent bonds
87
How are the carbon atoms in graphite arrange and what’s the bond angle
Planar hexagonal structure
Bond angle: 120
88
Properties of graphite
- conducts electricity due to the delocalised electrons from each carbon atom being able to freely move THROUGHOUT the sheets
- high mp and bp (same as diamond)
- can act as a lubricant due to ADSORBED gases on the surface of carbon atoms (gas molecules sticking to the surface) - but lubricating properties are not there in a vacuum
- insoluble in water (same reason as diamond)
- lower density as diamond due to the relatively amount of large space between layers
89
What is graphene
A single layer of graphite
90
Suggest why boron nitrate can act as a lubricant in a vacuum whilst graphite cannot
As they don’t need gaseous atoms to be ADSORBED (not absorbed, ADSORBED) onto the surface of layers
91
Properties of graphene
- conducts electricity
- can be rolled into a carbon nanotube
- one atom thick
92
Buckminister fullerene formula
C60
93
Properties of silicone dioxide
- high mp and bp due to strong silicone - carbon covalent bonds
- hard
- doesn’t conduct electricity
- insoluble - as there are no attractions which could occur between solvent molecules and the silicone and oxygen atoms which could overcome the covalent bonds in the structure
94
what is metallic bonding
The strong electrostatic attraction between positive metal ions and the sea of delocalised electrons
95
Properties if metallic bonding
- high mp and bp - due to the strong electrostatic forces of attraction between cations and delocalised electrons so requires a lot of energy to overcome these forces and move the metal ions away from their positions in the lattice
- high density - as the atoms are closely packed with little space between them
- can conduct electricity - due to the freely moving delocalised electrons so are attracted to the positive electrode and flow through the metal and this flow is an electrical current
- can conduct heat - when the metal is heated the delocalised electrons move around and conduct heat to the other parts of the metal
- malleable due to metal ions being able to slide over each other when a force is applied (which is known as a slip)
- insoluble
96
What is the electron pair repulsion theory
That the shape of a molecule is determined by the electron pairs surrounding the central atom as the electron pairs repel other electron pairs so move as far as possible to minimise repulsion
97
Why do lone pairs reduce the bond angle by 2.5 degrees
As they repel more strongly
98
why do branched molecules have a lower bp than unbranched molecules
as the molecules dont pack as close together so London forces are reduced as there are fewer contact points
