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Chem Module 2 > 2.2.2 Bonding And Structure > Flashcards

2.2.2 Bonding And Structure Flashcards

1
Q

Explain how an induced dipole dipole interaction forms

A

Movement of Electrons produces a Dipole in a molecule creating an instantaneous dipole which induces a dipole onto a neighbouring molecule this further induces more molecules which attract one another’s oppositely charged sides

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2
Q

Why do simple molecular compounds have low melting and boiling points?

A

Small amount of energy is needed to overcome the weak intermolecular forces between molecules

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3
Q

Why do simple molecular compounds not dissolve in water?

A

There is little interaction between the molecules in the lattice and the polar solvent molecules

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4
Q

Why do simple molecular compounds have poor electrical conductivity?

A

There are no free charged particles (delocalised electrons)

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5
Q

What is the structure of ionic compounds?

A

Giant ionic lattice

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6
Q

what are the properties of giant ionic lattices

A
  • high melting points
  • conduct electricity when molten or in an aqueous solution
  • soluble in water
  • hard and brittle
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7
Q

What are examples of giant covalent structures

A

Diamond graphite silicon dioxide SiO2

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8
Q

Why do ionic lattices have high melting points

A

Because a large amount of energies are required to overcome strong electrostatic bonds Between oppositely charged ions

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9
Q

What are the properties of Giant covalent lattices

A
  • high melting and boiling points
  • do not conduct electricity except graphite
  • insoluble in Polar and nonpolar solvents
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10
Q

what type of solvents do Ionic lattices dissolve and why

A

polar solvents eg water Because the partially charged ions in water are attracted to the charged ions in the lattice

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11
Q

What is the structure of simple covalent compounds?

A

Simple molecular

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12
Q

What are the properties of simple molecular structures

A
  • low melting points
  • do not conduct electricity
  • insoluble in water unless they are polar
  • soft
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13
Q

What are planar molecules

A

Can be arrange around central atom in 2D
Eg CO2 HCN H2O SCl2 BF3 SO3

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14
Q

What is Electronegativity

A

The ability of an atom to attract an electron pair in a covalent bond

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15
Q

What is a hydrogen bond

A

an attraction between a lone pair of electrons on an electronegative atom in one molecule and a hydrogen atom in a different molecule

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16
Q

explain why water molecules are polar

A

O is more electronegative than H so the electron pairs in the covalent bond are pulled towards it making it slightly negative So there is a difference in electronegativity and the molecule is not symmetrical, so dipoles do not balance out

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17
Q

how to explain the shape or bond angle

A
  1. state number of lone pair and bond pairs/regions
  2. say that electron pairs repel as far apart as possible
  3. state that lone pairs repel more than bond pairs
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18
Q

what makes a molecule polar

A

if there are polar bonds (when there is a difference in electronegativity) and the molecule is not symmetrical (due to lone pairs so the dipoles don’t cancel out) so there is an overall net charge

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19
Q

how can a molecule have polar bonds but not be a polar molecule

A

When dipoles cancel out in symmetrical molecules

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20
Q

what does electronegativity depend on

A
  • nuclear charge - the more protons, the more electronegative
  • the atomic radius (decreases left to right)
    (the smaller atomic radius, more electronegative)
  • shielding - the less shells, the more electronegative
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21
Q

what are the most electronegative elements

A

F O N Cl

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22
Q

what is the definition of a dipole

A

The separation of opposite charges

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23
Q

why does ice have a higher melting point than expected

A

Hydrogen bonds are the strongest intermolecular forces so more energies needed to break them

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24
Q

Which one is the slightly negative atom of NBr and why

A

N because it is slightly more electronegative so electrons are pulled towards it

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25
What is a covalent bond
the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
26
How can we predict the shape of a molecule from knowing the number of electron pairs surrounding the central atom
Pairs of Electrons surrounding a central atom repel so the shape is determined by the number of bond pairs and lone pairs Where lone pairs repel more
27
what are types of intermolecular forces In order of increasing strength
- Induced dipole-dipole (London forces) - Permanent dipole-dipole interactions - Hydrogen bonding
28
what is a permanent dipole
a polar covalent bond
29
What is a permanent dipole dipole interaction
The intermolecular bond between two polar molecules with permanent dipoles
30
if there are no hydrogen bonds why may boiling points still increase
The more electrons in each molecule, the larger the instantaneous dipoles, the more induced dipole-dipole interactions (London forces), the stronger the attraction between molecules, meaning more energy needed to overcome intermolecular forces/ London forces which increases the boiling point.
31
What is an ionic bond
The strong electrostatic attraction between oppositely charged ions
32
What is metallic bonding
Electrostatic attraction between the positive metal ions and the sea of delocalised electrons
33
Why do giant covalent structures have high melting and boiling points
Lot of energy is required for strong covalent bonds Between shared pair of Electrons and the nuclei to be broken
34
What structures can covalent bonding create
A small simple molecule a giant covalent structure a charged polyatomic ion
35
Why does giant ionic lattices conduct electricity when liquid but not when solid
When solid the ions are fixed in position when they are liquid they are free to carry charge
36
What is average bond enthalpy
A measure of average energy needed to break the bond
37
What is a dative covalent bond
A bond where both of the shared electrons are supplied by the same atom
38
what does the expansion of the octet mean
When a bonded atom has more than eight electrons in the outer shell
39
Describe the bonding in simple molecular structures
Atoms within the same molecule are held by strong covalent bonds and different molecules are held by weak intermolecular forces
40
How does graphite conduct electricity
Delocalised Electrons are present between the layers and are able to move freely carrying the charge Through the layers
41
What is the shape of a molecule depend on
Number of electron pairs in the outer shell and number of Electrons which are bonded and lone pairs
42
Bond angle - 2 bonded pairs 0 lone pairs
180 linear
43
Bond angle - 3 bonded pairs 0 lone pairs
120 Trigonal planar
44
Bond angle - 4 bonded pairs 0 lone pairs
109.5 Tetrahedral
45
Bond angle - 6 bonded pairs 0 lone pairs
90 Octahedral
46
Bond angle - 3 bonded pairs 1 lone pairs
107 Pyramidal
47
Bond angle - 2 bonded pairs 2 lone pairs
104.5 Nonlinear
48
how many degrees does each lone pair reduce the bond angle
2.5
49
How to draw a hydrogen bond
horizontal Dashed lines next to the H and the lone pair on the electronegative element
50
Why might an elements be unreactive
It is non polar the bonds are very strong It has no lone pairs to react
51
how is a polar bond formed
atoms with different electronegativities bond
52
What are London forces
When a molecule with a permanent dipole (Polar bonds) or instantaneous dipole is close to a non-polar molecule it causes it to become slightly polar leading to A weak attraction
53
where can you find the strongest london forces
In bigger molecules as they have more electrons
54
Why is ice less dense than liquid water
In ice water molecules are arranged in a fixed lattice with hydrogen bonds In water molecules are closer together and hydrogens bonds can break and reform so there is no fixed lattice
55
Why does water have a higher melting point than expected
hydrogen bonds are the strongest of the intermolecular forces so extra strength is required to overcome these forces
56
What is the difference between simple molecular structure and giant covalent structure?
Simple molecules are bonded by covalent bonds and joined to each other by weak intermolecular forces Giant covalents lattices are many molecules joined by covalent bonds
57
HBr and HI have both permanent and induced dipole iodine has a lower electronegativity than Br so why does iodine have a higher boiling point?
Because the induced dipole interaction in iodine is higher
58
Four ionic compounds have different melting points why is that?
As ionic charge increases melting point increase, the greater the ionic charge, the greater attraction between ions, and the more energy required to break the ionic bonds
59
Which one has a higher ionic charge NaF Na2O MgO
MgO 2+ and -2
60
What makes stronger hydrogen bonds
Higher difference in electronegativity FH most strong hydrogen bonds H2O mid strength hydrogen bonds NH3 lower strength hydrogen bonds
61
Ice bond 109 Gaseous water bond 105 Why
Ice forms more hydrogen bonds Water has two bond pairs, and two lone pairs Lone pair repel more
62
Why is NCL3 a polar molecule?
This molecule has three bond pairs which repel equally but also one lone pair which causes asymmetry so dipoles don’t cancel out
63
The first 4 alkanes are are gases but the First for alcohols are liquids, explain the difference in terms of intermolecular forces
Alcohols have Hydrogen bonds which are the stronger as well as London forces
64
Explain The difference in boiling points ammonia -33 Fluorine -188 bromine 59
- Ammonia has hydrogen bonding which is stronger than London forces in fluorine - Fluorine and bromine have London forces, but bromines London forces are stronger because bromine has more electrons - London forces in bromine are stronger than hydrogen bonding in ammonia - London forces and hydrogen bonds are intermolecular forces
65
What Is meant by a giant ionic lattice
Repeating patterns of oppositely charged ions
66
How does a increased electronegativity within a molecule affect the boiling point
The more different in electronegativity the higher the boiling point

Chem Module 2 (7 decks)

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