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Chemistry Year 12 > 2 - Chemical bonding and structure > Flashcards

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 bonding?

A
  • product of the charges of the ions

- ionic radii

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3
Q

How does ionic charge affect ionic bonding?

A
  • the larger the charges, the larger the lattice energy therefore the stronger the ionic bond.
  • the smaller the radii, the smaller the distance between ions, stronger force of attraction (stronger ionic bonds).
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4
Q

trend in ionic radii down a group?

A
  • ionic radii increases.
  • due to increasing number of shells of electrons.
  • increased shielding effect means also means that electrons are pulled in less by nucleus, which also contributes to ionic radii.
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5
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 means that attraction between nucleus and electrons is greater so the electrons are pulled in closer to the nucleus therefore ionic radii decreases.
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6
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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7
Q

What is it called if ions of different elements have the same number of electrons?

A

Isoelectronic (same number of electrons).

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8
Q

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

A
  • when an ionic compound melts, the giant ionic lattice is broken.
  • in the molten state, the ions are free to move around.
  • ionic compounds / giant ionic lattices consist of strong electrostatic forces of attraction between the oppositely charged ions which are strong, so lots of energy is required to break these bonds.
  • so, ionic compounds have high melting points.
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9
Q

Why are ionic compounds brittle?

A
  • this is because if a layer of ions is moved in an ionic compound, you end up with ions with the same charges next to each other.
  • the layers repel each other and the crystal breaks up.
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10
Q

Can ionic compounds conduct electricity?

A
  • ion fixed in lattice and cannot move in solid state. cannot conduct electricity in solid state.
  • in molten, aqueous solution, ions are free to move, can conduct electricity.
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11
Q

How can ionic compounds dissolve in water at room temperature?

A
  • water molecules are polar. Oxygen atoms slightly negatively charged and 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.
  • called hydration.
  • provides enough energy to separate the ions in the lattice.
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12
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 measured by enthalpy of hydration.
  • always negative (exothermic).
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13
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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14
Q

What is the structure are giant metals in?

A

Giant metallic lattices.

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15
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 electrons binds the positive metal cations together and avoids the repulsion between them.
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16
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 forces of attraction between the nuclei of the metal cations and delocalised electrons.
  • giant lattice. There are many forces to overcome.
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17
Q

What does the strength of a metallic bond depend on?

A
  • Charge of the metal cation and number of delocalised electrons per cation.
  • size of the cation.
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18
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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19
Q

How does size of cation affect strength of metallic bond?

A

The smaller the metal ion, the closer the positive nucleus is to the delocalised electrons. This results in a greater force of attraction.

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20
Q

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

A
  • magnesium has two electrons in its outer shell and both of these get delocalised.
  • K and Na only have one electron in their outer shell which gets delocalised.
  • so, the sea of delocalised electrons has twice the electron density in magnesium as it does in K and Na.
  • Magnesium also has a smaller atomic/ionic radii than K and Na, so the positive nuclei and delocalised electrons are closer together.
  • therefore magnesium has stronger metallic bonds and a higher boiling point.
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21
Q

Why can metals conduct electricity?

A

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

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22
Q

What is 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. Hence conductivity is lower.
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23
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 the other.
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24
Q

Why are metals malleable and ductile?

A
  • the layers of cations are able to slide over each other, so the structure does not shatter.
  • The metallic bonds do not break because 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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25
Q

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

A
  • metallic bonds are non-directional (the 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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26
Q

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

A
  • high melting and boiling temps. Indicates that they are strong.
  • metals are malleable and ductile. This shows that the layers of metal 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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27
Q

ionic bonds and covalent bonds. Non directional or directonal?

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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28
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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29
Q

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

A

Giant covalent lattice

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30
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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31
Q

What type of covalent bonds are there?

A
  • sigma bond: end on overlap of two s orbitals.
  • sigma bond: end on overlap of two p orbitals.
  • sigma bond: end on overlap of a s and p orbital;.
  • pi bond: sideways overlap of 2 p orbitals.
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32
Q

Explain a sigma bond

A
  • end on overlap of orbitals
  • can involve s and p orbitals.
  • electron density mostly between the two nuclei of the two atoms.
  • always a single bond.
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33
Q

Explain a pi bond

A
  • sideways overlap of p orbitals.

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

34
Q

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

A
  • Single: 1 sigma bond
  • double: 1 sigma and 1 pi bond
  • triple: 1 sigma and 2 pi bonds.
35
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 positive nuclei. (the electrons here are not as effective in attracting the nuclei).
  • there is a greater degree of orbital overlap with a sigma bond than pi bond. Greater orbital overlap means stronger bond.
36
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 nuclei and the bonding electrons decreases as distance between them increases.
  • this factor is more significant than nuclear charge increase.
37
Q

How does a dative covalent bond form?

A

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

38
Q

How are dative covalent bonds represented?

A
  • by an arrow starting from the atom providing the pair of electrons and going towards the atom accepting it.
39
Q

What are dative covalent bonds also known as?

A

co-ordinate bond.

40
Q

Examples of dative covalent bonding?

A

NH4+, H3O+, NH3BF3)

2 AlCl3 join to form dimer Al2Cl6.

41
Q

Linear

A
  • 2 bonding pairs

- 180°

42
Q

trigonal planar

A
  • 3 bonding pairs

- 120°

43
Q

tetrahedral

A
  • 4 bonding pairs

- 109.5°

44
Q

trigonal pyramidal

A
  • 3 bonding pairs
  • 1 lone pair
  • 107°
45
Q

bent

A
  • 2 bonding pairs
  • two lone pairs
  • 104.5°
46
Q

trigonal bipyradimal

A
  • 5 bonding pairs

- 120° and 90°

47
Q

octahedral

A
  • 6 bonding pairs

- 90°

48
Q

Explain the shape of a molecule (how?)

A
  • state number of bonding pairs and lone pairs of electrons.
  • state that electrons repel to get as far apart as possible.
  • if no lone pairs, state that the electron pairs repel equally.
  • if there are lone pairs, state that lone pairs repel more than bonding pairs.
  • state shape and bond angle.
49
Q

What is electronegativity?

A

The ability of an atom to attract a bonding pair of electrons in a covalent bond.

50
Q

What are the most electronegative atoms?

A

F, N, O, Cl.

51
Q

Nonpolar covalent bond?

A

Bonding electrons are shared equally between the two atoms.

- compounds containing elements of similar electronegativity.

52
Q

Polar covalent bond (permanent dipole)

A

Bonding electrons shared unequally between the two atoms. Partial charges on the atoms. Creates a dipole.

  • when the elements in the bond have different electronegativities.
    e. g HCl, δ+ δ-
53
Q

Ionic bond?

A

Complete transfer of one or more electrons. Full charges on resulting ions.

54
Q

Factors affecting electronegativity

A
  • atomic radii
55
Q

Electronegativity down a group

A
  • decreases
  • atomic radii increases
  • distance between nucleus and bonding pair increases.
  • shielding effect increases.
56
Q

Electronegativity across a period

A
  • atomic radii decreases
  • no of shells stays the same.
  • nuclear charge increases, pulls the outer electrons closer towards the nucleus
  • distance between bonding pair and nucleus decreases.
57
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.
58
Q

polar molecule

A
  • asymmetrical
  • lone pairs
  • different atoms bonded to central atom
  • there is a net dipole moment
59
Q

Why are some molecules non polar even though they have polar molecules?

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.
60
Q

London forces

A
  • also known as instantaneous / induced dipole/dipole interactions.
  • occur 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).
61
Q

What do London forces depend on?

A
  • number of electrons: more electrons means higher chance of formation of temporary dipoles. More electrons, stronger London forces.
  • shape of the molecule: Long straight chained alkanes have more points of contact therefore stronger London forces than branched or cycloalkanes.
62
Q

London forces and chain length?

A
  • long chains have greater number of electrons
  • greater temporary dipoles
  • London forces increase in strength.
63
Q

How does molecular shape affect London forces?

A
  • straight chain alkanes have greater points of contact between the different molecules.
  • more opportunities for induced dipoles to occur.
  • straight chains have a higher boiling point than branched.
64
Q

Permanent dipole-dipole forces

A
  • occurs between polar molecules
  • 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.
65
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.
66
Q

Hydrogen bonding in water

A

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

67
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.

68
Q

Hydrogen bonding in HF

A
  • although there are three lone pairs on the fluorine,
  • only one hydrogen bond can occur per molecule with the one of the lone pairs on the HF.
  • the other 2 lone pairs are wasted.
69
Q

Hydrogen bonding in NH3

A
  • although there are 3 hydrogens available for hydrogen bonding,
  • there are a shortage of lone pairs in each molecule.
  • this means that each NH3 molecule can form only 1 hydrogen bond.
70
Q

Strength of hydrogen bonds

A

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

71
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 in water.

72
Q

Solubility of simple alcohols

A
  • simple alcohols are soluble in water because they can form hydrogen bonds with water.
  • the longer the hydrocarbon chain, the less soluble the alcohol.
73
Q

insolubility of compounds in water

A
  • compounds that cannot form hydrogen bonds with water molecules (basically molecules with no hydrogen bonds) cannot dissolve in water.
74
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 dissolve in Molecules with London forces and Hydrogen bonds.
  • substances need the same type of intermolecular force(s) to dissolve in each other.
75
Q

can diamond conduct electricity?

A

No. All 4 electrons per carbon atom in involved in covalent bonding, so there are no electrons free to move (delocalised electrons).

76
Q

Can graphite conduct electricity?

A

Yes. there is 1 delocalised electron per carbon atom which is free to move along and between the layers of graphite.
- weak intermolecular forces between the layers of graphite (brittle).

77
Q

Why does diamond and graphite have high melting points?

A
  • giant covalent structure

- takes a lot of energy to break the many covalent bonds.

78
Q

Graphene structure and electricity conductivity?

A
  • one layer of graphite
  • high tensile strength, due to many covalent bonds in the giant covalent lattice (allotrope).
  • one delocalised electron per carbon atom which can move throughout the structure, can conduct electricity.
79
Q

Nanotubes

A
  • almost same structure and electrical conductivity reason as graphene.
  • one potential use is to transport drugs to cells.
80
Q

Why is ice less dense than water?

A
  • in ice, the water molecules are arranged in rings of six, held by hydrogen bonds and the water molecules are pushed further apart compared to liquid water.
  • liquid water, hydrogen bonds are constantly formed and broken as the molecules slide past each other, molecules are closer.
  • since the water molecules are pushed further apart from each other, ice is less dense.
81
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 intermolecular forces in water along with London forces and permanent dipole dipole forces. Hence water has a high melting / boiling point.

Chemistry Year 12 (12 decks)

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