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Chemistry year 12 - speedrun > Topic 2 - Chemical Bonding and Structure > Flashcards

Topic 2 - Chemical Bonding and Structure Flashcards

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1
Q

What is ionic bonding?

A

Ionic bonding is the strong electrostatic attraction between oppositely charged ions.

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2
Q

What affects the strength of ionic bonds?

A
  • product of the charges of the ions

- ionic radii

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3
Q

how does ionic charge affect ionic bonding?

A
  • larger charges means larger lattice energy therefore stronger ionic bond.
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4
Q

How does ionic radii affect ionic bonding?

A
  • the smaller the radii, the smaller the distance between the ions, stronger force of attraction, stronger ionic bonds.
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5
Q

trend in ionic radii down a group?

A
  • ionic radii increases
  • due to increasing number of electron shells.
  • although nuclear charge increases
  • number of electron shells increases
  • shielding effect increases
  • electrons are pulled in less by the nucleus, which also contributes to increasing ionic radii
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6
Q

trend in ionic radii across a period?

A
  • ionic radii decreases
  • number of shells of electrons stays the same.
  • nuclear charge and number of electrons increases.
  • increased nuclear charge with same number of shells means that the attraction between the nucleus and electrons is greater, so the electrons are pulled in closer to the nucleus, therefore ionic radii decreases.
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7
Q

relationship between ionic bonds and product of charges?

A

the electrostatic forces of attraction between oppositely charged ions are directly proportional to the product of the charges.

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8
Q

what is it called if ions of different elements have the same number of electons?

A

isoelectronic (same number of electrons).

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9
Q

How are ionic compounds melted and do they have high melting points?

A
  • when an ionic compound is melted, the giant ionic lattice is broken.
  • in the molten state, the ions are free to move around.
  • ionic compounds / giant ionic lattices consist of many strong electrostatic forces of attraction between the oppositely charged ions which are strong, so lots of energy is required to break these ionic bonds.
  • so, ionic compounds have high melting points.
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10
Q

why are ionic compounds brittle?

A
  • this is because if you move a layer of ions in an ionic compound, you end up with ions with the same charges next to each other.
  • the layers repel each other and the compound/lattice breaks up.
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11
Q

can ionic compounds conduct electricity?

A
  • in solid state, the ions are in fixed positions in the giant lattice, therefore cannot conduct electricity.
  • in molten and aqueous state, the ions are free to move, so can conduct electricity.
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12
Q

how can ionic compounds dissolve in water at room temperature?

A
  • water molecules are polar. Oxygen atoms slightly negatively charged, hydrogen atoms slightly positively charged.
  • the negative (oxygen) end of a water molecule is attracted to the positive ions.
  • the positive (hydrogen) end of a water molecule is attracted to the negative ions.
  • process is called hydration.
  • provides enough energy to separate the ions in the lattice.
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13
Q

relationship between water molecules and charge of ion?

A

The higher the charge of an ion, the more water molecules it attracts.

  • strength between them is measured by enthalpy of hydration.
  • always negative (exothermic).
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14
Q

what is metallic bonding?

A

Metallic bonding is the strong electrostatic attraction between the nuclei of metal cations and delocalised electrons.

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15
Q

what structure do metals have?

A

giant metallic lattices.

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16
Q

where do delocalised electrons come from and what are they?

A
  • the outer shell electrons of each atom leave to join a ‘sea’ of delocalised electrons, which can move freely throughout the structure.
  • the sea of delocalised electrons binds the positive metal cations together, preventing the repulsion between them.
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17
Q

why do metals have high melting points?

A
  • metallic bonds are strong and require a large amount of energy to overcome the strong electrostatic attractions between the nuclei of the metal cations and delocalised electrons.
  • giant lattice. There are many forces to overcome.
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18
Q

what does the strength of a metallic bond depend on?

A
  • charge of cation
  • number of delocalised electrons per cation
  • size of the cation.
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19
Q

how does charge of the metal cation and delocalised electrons affect strength of metallic bond?

A
  • the larger the charge, the larger the number of delocalised electrons.
  • this means that the force of attraction between the cations and delocalised electrons is greater.
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20
Q

how does size of cation affect strength of metallic bond?

A

The smaller the metal cation, the closer the positive nucleus is to the delocalised electrons. (Distance is reduced).
- This results in a greater force of attraction (as they are closer to each other).

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21
Q

Why does magnesium have a higher melting point than sodium and potassium?

A
  • magnesium has two electrons in outer shell and both get delocalised.
  • K and Na only have one electron in outer shell which gets delocalised.
  • so sea of delocalised electrons has twice the electron density in magnesium than K and Na.
  • Magnesium also has a smaller ionic radii than K and Na, so positive nuclei and delocalised electrons are closer.
  • so, Mg has stronger metallic bonds, and hence a higher melting point.
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22
Q

why can metals conduct electricity?

A

In the giant metallic lattice, the delocalised electrons are free to move throughout the structure, and can therefore carry a charge.

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23
Q

What does the electrical conductivity of a metal depend on?

A
  • number of delocalised electrons per unit volume of metal.
  • Potassium has larger cations, so number of delocalised electrons per unit volume is lower than sodium.
  • So, potassium has lower conductivity than sodium.
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24
Q

Why are metals good thermal conductors?

A
  • the delocalised electrons transfer kinetic energy throughout the whole metal structure
  • closely packed cations pass kinetic energy from one to another.
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25
Q

why are metals malleable and ductile?

A
  • the layers of cations are able to slide over each other, so the metal structure does not shatter.
  • the metallic bonds do not break as the delocalised electrons are free to move throughout the whole metal structure, preventing repulsion between the cations when the layers slide over each other.
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26
Q

why do metallic bonds exist in significant extent even in molten state?

A

metallic bonds are non-directional (the delocalised electrons are shared with many neighbouring cations in all directions and not just one).
- So they exist in significant extent, even in molten state.

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27
Q

facts to support that metallic bonds are strong and non-directional?

A
  • high mp and bp temps. Indicates that they are strong.
  • metals are malleable and ductile. This shows that the layers of cations can slide over each other without repelling, which is a consequence of the non-directional nature of the metallic bond.
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28
Q

ionic bonds and covalent bonds. Non directional or directional?

A
  • Ionic bonds: non-directional.

- Covalent: directional. The position of the bonding pair of electrons decides the direction of the covalent bonding.

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29
Q

what is a covalent bond?

A

A covalent bond is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them.

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30
Q

what kind of structure does graphite, diamond, silicon dioxide have?

A

giant covalent lattice.

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31
Q

How is a covalent bond formed?

A

Formed by the overlapping of two atomic orbitals, each containing a single electron.

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32
Q

what types of covalent bonds are there?

A
  • sigma: end on end overlap of two s orbitals.
  • sigma: end on end overlap of two p orbitals.
  • sigma: end on end overlap of a s and p orbital.
  • pi bond: sideways overlap of two p orbitals.
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33
Q

explain a sigma bond

A
  • end on end overlap of orbitals
  • can involve s and p orbitals
  • electron density mainly between the two nuclei of the two atoms.
  • always a single bond.
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34
Q

explain a pi bond

A
  • sideways overlap of p orbitals.

- electron density mostly above and below the two nuclei of the two atoms.

35
Q

what kind of bonds are in a single, double, triple covalent bond?

A
  • single: 1 sigma bond
  • double: 1 sigma and 1 pi bond
  • triple: 1 sigma and 2 pi bonds.
36
Q

Are pi bonds weaker than sigma bonds?

A
  • pi bonds are weaker than sigma bonds since their electron density is further away from the two nuclei. The electrons here are not as effective in attracting the nuclei.
  • there is a greater degree of orbital overlap in a sigma bond than pi bond. Greater orbital overlap means a stronger bond.
37
Q

• Suggest a reason for the following trend in
bond strength.
• C-C > Si-Si > Ge-Ge

A
  • atomic radii increases
  • bond length increases
  • electrostatic attraction between bonding electrons and nuclei decreases as distance between them increases.
  • increase in bond length decreases bond strength.
  • this factor is more significant than nuclear charge increase.
38
Q

How does a dative covalent bond form?

A

A dative covalent bond forms when the bonding pair of electrons in the covalent bond comes from the same atom (one of the bonding atoms).

39
Q

How are dative covalent bonds represented?

A
  • arrow starting from the atom providing the pair of electrons towards the atom accepting it.
40
Q

what are dative covalent bonds also known as?

A

co-ordinate bonds.

41
Q

examples of dative covalent bonding?

A
  • NH4+, H3O-, NH3BF3

2 AlCl3 join to form dimer Al2Cl6 (with dative covalent bonds).

42
Q

linear

A
  • 2 bonding pairs

- 180°

43
Q

trigonal planar

A
  • 3 bonding pairs

- 120°

44
Q

tetrahedral

A
  • 4 bonding pairs

- 109.5°

45
Q

trigonal pyramidal

A
  • 3 bonding pairs
  • 1 lone pair
  • 107°
46
Q

bent

A
  • 2 bonding pairs
  • 2 lone pairs
  • 104.5°
47
Q

trigonal bipyramidal

A
  • 5 bonding pairs

- 120° and 90°

48
Q

octahedral

A
  • 6 bonding pairs

- 90°

49
Q

How to answer explaining the shape of a molecule?

A
  • state number of bonding pairs and number of lone pairs.
  • state that electron pairs repel to get as far apart as possible.
  • if no lone pairs, state that the bonding pairs repel equally.
  • if there are lone pairs, state that lone pairs repel more than bonding pairs.
  • state shape and bond angle.
50
Q

what is electronegativity?

A

Electronegativity is the ability of an atom to attract a bonding pair of electrons in a covalent bond.

51
Q

What are the most electronegative atoms?

A

F, N, O, Cl.

52
Q

Nonpolar covalent bond?

A

Bonding electrons are shared equally between the two atoms.

- occurs between elements in the bond with same/similar electronegativity.

53
Q

Polar covalent bond (Permanent dipole)

A
  • bonding electrons shared unequally between the two atoms.
  • partial charges on the atoms.
  • creates a dipole
  • occurs between elements in a bond with different electronegativities.
    e. g HCl, δ+ δ-
54
Q

Ionic bond in terms of electron transfer?

A

Complete transfer of one or more electrons.

- results in ions with full charges.

55
Q

types of charges in the different types of covalent bonds?

A
  • non polar: no charges
  • polar: partial charges δ+ δ-
  • ionic bonds: full charges on resulting ions.
56
Q

factors affecting electronegativity?

A
  • atomic radii

- nuclear charge

57
Q

electronegativity down a group

A
  • decreases
  • atomic radii increases (more shells).
  • distance between nucleus and bonding pair increases
  • shielding effect increases
58
Q

electronegativity across a period

A
  • increases
  • atomic radii decreases
  • number of shells stays the same (shielding effect doesn’t change)
  • nuclear charge increases.
  • force of attraction between electrons and nucleus increases, so they are pulled in closer towards the nucleus.
  • distance between bonding pair and nucleus decreases.
59
Q

non polar molecule

A
  • symmetrical
  • all bonds are identical
  • no lone pairs
  • non polar molecule, even though the individual bonds are polar (same atoms bonded to central atom).
  • no net dipole moment.
60
Q

polar molecule

A
  • asymmetrical
  • lone pairs
  • different atoms bonded to central atom
  • there is a net dipole moment.
61
Q

why are some molecules non polar even though they have polar bonds?

A
  • the symmetrical molecular shape of CCl4 means that the dipoles on the bonds ‘cancel out’.
  • no net dipole moment
  • CH3Cl is asymmetrical. Partially negative Cl attached, other H are partially positive. There is a net dipole moment, making it a polar molecule.
62
Q

London forces

A
  • also known as instantaneous / induced dipole / dipole interactions.
  • occurs between all molecules
  • in non-polar molecules, London forces are the only intermolecular forces.
  • in any molecule, electrons are moving constantly and electron density can fluctuate. This causes parts of the molecule to become more or less negative forming temporary dipoles.
  • the temporary dipoles can cause dipoles to form in neighbouring molecules (induced dipoles).
63
Q

what do London forces depend on?

A
  • number of electrons: more electrons means a higher chance of formation of temporary dipoles. More electrons, greater London forces.
  • shape of the molecule: Long straight chained alkanes have more points of contact with neighbouring molecules compared to branched alkanes and cycloalkanes. Therefore, long straight chained alkanes have greater London forces.
64
Q

London forces and chain length?

A
  • long chains have greater number of electrons.
  • greater temporary dipoles
  • greater London forces.
65
Q

How does molecular shape affect London forces?

A
  • straight chain alkanes have greater points of contact between neighbouring molecules.
  • more opportunities for induced dipoles to occur.
  • straight chains have a higher boiling point than branched.
66
Q

permanent dipole-dipole forces

A
  • occurs between polar molecules
  • they are stronger than London forces
  • polar molecules have permanent dipole-dipole forces between atoms of different molecules where there is a significant difference in electronegativity.
  • permanent dipole forces occur in addition to London forces.
67
Q

Hydrogen bonding

A
  • occurs in compounds that have a hydrogen attached to a very electronegative atom (F, N, O).
  • when drawing, always draw the lone pairs and the dipoles and signs.
  • hydrogen bonds occur in addition to London forces and permanent dipole-dipole forces.
68
Q

Hydrogen bonding in water

A

Water can form 2 hydrogen bonds per molecule, since the highly electronegative oxygen atom has two lone pairs of electrons on it.

69
Q

Can alcohols form hydrogen bonds?

A
  • yes

- they have higher boiling points and relatively low volatility compared to alkanes with the same number of electrons.

70
Q

Hydrogen bonding in HF

A
  • although there are 3 lone pairs on the fluorine
  • only one hydrogen bonds can occur per molecule with one of the lone pairs on the F.
  • due to lack of hydrogen atoms.
  • the other 2 lone pairs are wasted.
71
Q

Hydrogen bonding in NH3

A
  • although there are 3 hydrogens available for hydrogen bonding,
  • there is a shortage of lone pairs in each molecule (only 1 present).
  • this means that each NH3 molecule can only form 1 hydrogen bond each.
72
Q

strength of hydrogen bonds compared to other intermolecular forces.

A

Hydrogen bonds are stronger than London forces and permanent dipole-dipole forces.

73
Q

Why can’t some ionic compounds dissolve in water?

A
  • If the attraction between the ions or the attraction between the water molecules is greater than the hydration energy, the ionic compound is insoluble.
74
Q

solubility of simple alcohols

A
  • simple alcohols are soluble in water because they can form hydrogen bonds with water molecules.
  • the longer the hydrocarbon chain, the less soluble the alcohol.
75
Q

insolubility of compounds in water

A
  • compounds that cannot form hydrogen bonds with itself cannot form hydrogen bonds with water molecules.
  • these compounds cannot dissolve in water.
76
Q

solubility of different compounds

A
  • molecules with just London forces can dissolve in each other.
  • molecules with London forces and hydrogen bonds can dissolve in each other.
  • molecules with just London forces cannot dissolves in molecules with London forces and hydrogen bonds.
  • substances need the same type of intermolecular force(s) to dissolve in each other.
77
Q

can diamond conduct electricity?

A
  • diamond has a giant covalent structure
  • cannot conduct electricity.
  • all 4 outer electrons per carbon atom is involved in covalent bonding, so there are no delocalised electrons to carry charge.
78
Q

can graphite conduct electricity?

A
  • yes
  • each carbon atom is joined to 3 other carbon atoms.
  • 1 delocalised electron per carbon atom which is free to move between the layers of graphite, can carry charge.
  • weak intermolecular forces between the layers of graphite, so it is brittle.
79
Q

why does diamond and graphite have high melting points?

A
  • giant covalent structure
  • lots of covalent bonds are present
  • covalent bonds are strong and require lots of energy to break.
80
Q

graphene structure and electrical conductivity?

A
  • one layer of graphite
  • each carbon atom is joined to 3 other carbon atoms.
  • one delocalised electron per carbon atom, which can move freely throughout the structure and can carry charge.
  • high tensile strength due to many covalent bonds in the giant covalent lattice.
  • carbon allotrope.
81
Q

nanotubes

A
  • almost same structure and electrical conductivity as graphene.
  • one potential use is to transfer drugs to cells.
82
Q

why is ice less dense than water?

A
  • in ice, water molecules are arranged in rings of six, held by hydrogen bonds. The water molecules are pushed further apart than liquid water.
  • liquid water, hydrogen bonds are constantly formed and broken as the molecules slide past each other, molecules are closer together.
  • since water molecules are pushed further away from each other, ice is less dense than liquid water.
  • ice can float on water.
83
Q

Why does water have a high boiling / melting point?

A
  • between water molecules, there are hydrogen bonds which are relatively strong.
  • therefore, a lot of energy is required to overcome the hydrogen bonds along with the permanent dipole-dipole forces and London forces.
  • hence a high mp / bp.

Chemistry year 12 - speedrun (7 decks)

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