Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

CHEMISTRY FOR MOCKS 2022 > 2 - Chemical bonding and structure > Flashcards

2 - Chemical bonding and structure Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

What is ionic bonding?

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What affects strength of ionic bonding?

A
  • ionic radii (smaller means stronger)

- ionic charge (higher charges means stronger).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Why does MgO have a higher boiling point than NaCl?

A

The ions involved in MgO (Mg2+ and O2-) have a smaller ionic radii and also have higher charges than the ions involved in NaCl (Na+ and Cl-).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Why are positive ions smaller compared to their atoms?

A
  • one less shell of electrons
  • the force of attraction of the nucleus is shared over a smaller number of electrons, so the attraction per electron is greater.
  • so the electrons are pulled in closer towards the nucleus.
  • so the positive ions are smaller than their atoms.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Why are negative ion larger compared to their atoms?

A
  • number of shells does not change.
  • negative ion has more electrons than the corresponding atom, but the same number of protons.
  • the force of attraction of the nucleus is shared over a greater number of electrons, so the attraction per electron is less.
  • so the electrons are not held as close towards the nucleus.
  • so the ion is bigger.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Why does ionic radii decrease across a period?

A
  • these are ions, not atoms.
  • number of electrons stays the same (same electronic structure).
  • proton number increases.
  • force of attraction of nucleus on the electrons increases, attraction per electron increases.
  • the electrons are pulled in closer towards the nucleus.
  • therefore ionic radii decreases.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Why does ionic radii increase down a group?

A
  • the number of electron shells increases down a group.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Physical properties of ionic compounds?

A
  • high melting points (ionic bonds are strong and require a lot of energy to overcome).
  • non conductor of electricity WHEN SOLID (ions are in fixed positions in a lattice structure, and the ions are not free to move).
  • conductor of electricity when aqueous (in solution) or molten (ions are free to move).
  • brittle (a little force can push ions along and cause similar ions to be next to each other. There will be repulsion between the ions with the same charge and cause the layer of ions to be pushed apart.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

In simple electrolysis, where do positive ions (cations) move to?

A

cathode (negatively charged electrode).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

In simple electrolysis, where do negative ions (anions) move to?

A

anode (positively charged electrode).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What is a covalent bond?

A

A covalent bond is the strong electrostatic force of attraction between two nuclei and their bonding pair of electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

why do giant covalent structures like diamond and graphite have high melting points?

A
  • giant covalent structures contain many covalent bonds.

- covalent bonds are strong and require a lot of energy to overcome.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

describe the electron density in a covalent compound?

A

In a covalent compound, there is significant electron density between the nuclei of the atoms.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Effect of multiple bonds on bond strength and length?

A

Nuclei joined by multiple bonds (double and triple) have a greater electron density between them.

This causes a greater force of attraction between the nuclei and the electrons between them.

This results in a shorter bond length and a greater bond strength.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is a dative covalent bond?

A

A dative covalent bond is when the shared pair of electrons in a covalent bond comes from only one of the bonding atoms.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

How do you represent a dative covalent bond?

A

Draw an arrow going from the atom that is providing the lone pair to the atom that is deficient

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Linear shape

A
  • 2 bonding pairs
  • no lone pairs
  • 180º
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Trigonal planar

A
  • 3 bonding pairs
  • no lone pairs
  • 120º
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Tetrahedral

A
  • 4 bonding pairs
  • no lone pairs
  • 109.5º
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Trigonal pyramidal

A
  • 3 bonding pairs
  • 1 lone pair
  • 107º
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Bent

A
  • 2 bonding pairs
  • 2 lone pairs
  • 104.5º
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Trigonal bypyramidal

A
  • 5 bonding pairs
  • no lone pairs
  • 90º and 120º
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Octahedral

A
  • 6 bonding pairs
  • no lone pairs
  • 90º
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

What are steps in explaining the shape of a molecule?

A
  • state number of bonding pairs
  • state number of lone pairs
  • state that electron pairs repel to get as far apart as possible.
  • if no lone pairs, state that the bonding pair of electrons repel equally.
  • if there are lone pairs, state that lone pairs repel more than bonding pairs.
  • state the shape of the molecule
  • state the bond angle(s).
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

lone pairs and bonding pairs repulsion comparison?

A

Lone pairs repel more than bonding pairs.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Define electronegativity

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

How is electronegativity measured?

A
  • Pauling scale

- ranges from 0 to 4 (4.0 is most electronegative, F).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

What factors affect electronegativity?

A
  • number of protons in nucleus
  • atomic radius
  • distance between nucleus and outer electrons.
  • shielding effect
29
Q

Why does electronegativity increase across a period?

A
  • number of protons increases
  • atomic radius decreases as electrons in the same shell are pulled in more.
  • shielding effect remains similar.
30
Q

Why does electronegativity decrease down a group?

A
  • although number of protons increases
  • number of shells increases
  • distance between nucleus and outer electrons increases
  • shielding effect increases.
  • atomic radius increases.
31
Q

Intermediate bonding (covalent)

A

A compound containing elements of similar electronegativity and hence a small electronegativity difference will be covalent (since the bonding pair of electrons aren’t really pulled more towards one side).

32
Q

Intermediate bonding (ionic)

A

A compound containing elements of very different electronegativity and hence a large electronegativity difference will be ionic (bonding pair of electrons are much more attracted towards one of the elements than the other, maybe to the point where the electrons are fully transferred to that element, causing the formation of ions and ionic bonding).

33
Q

What is a polar covalent bond?

A

A polar covalent bond forms when the elements in the covalent bond have different electronegativities.

  • In a polar covalent bond, there is an unequal distribution of the bonding pair of electrons and creates a dipole.
  • e.g H-Cl, Cl is much more electronegative than H, therefore the bonding pair of electrons are attracted more towards the Cl, creating a dipole.
    δ+ H
    δ- Cl
34
Q

What are symmetrical molecules?

A
  • all bonds are identical
  • no lone pairs
  • the molecule won’t be polar even if individual bonds are polar.
  • the individual dipoles on the bonds cancel out.
  • there is no net dipole moment (the molecule is non-polar).
35
Q

What are unsymmetrical molecules?

A
  • bonds are not identical
  • lone pairs may be present
  • there is a net dipole moment
  • the molecule is polar.
36
Q

London forces

A
  • instantaneous induced dipole-dipole interactions.

- they occur between all simple covalent molecules and separate atoms of noble gases.

37
Q

What affects London forces?

A
  • More electrons means stronger London forces. More energy is required to break them, so b.p will be greater.
  • Molecules with more points of contact with adjacent molecules have stronger London forces (e.g straight-chained alkanes have stronger London forces than branched).
38
Q

Why is Cl2 a gas and I2 a solid?

A
  • both exist as diatomic molecules
  • I2 has more electrons than Cl2.
  • I2 has stronger London forces between adjacent molecules than Cl2.
  • therefore more energy is required to break the London forces between I2 molecules than Cl2.
  • so I2 is solid, Cl2 is gas.
39
Q

Permanent dipole dipole interactions

A
  • occurs between polar molecules
  • stronger than London forces, so the compounds have higher boiling points.
  • polar molecules have a permanent dipole (e.g compounds with H-Cl or C-Br bonds).
  • polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
  • permanent dipole-dipole interactions occur in addition to London forces.
40
Q

Hydrogen bonding

A
  • occurs in compounds that have a hydrogen atom attached to a one of these highly electronegative atoms: F, N, O.
  • these atoms must have an available lone pair of electrons.
  • Hydrogen bonding occurs in addition to London forces and permanent dipole-dipole interactions.
  • 180º
41
Q

Tips when drawing hydrogen bonding

A
  • Always show the lone pair of electrons on F, N, O

- include and show all the dipoles.

42
Q

Examples of compounds that can form hydrogen bonds?

A
  • alcohols
  • carboxylic acids
  • proteins
  • amines
  • amides
43
Q

Why does ethanol have a much higher boiling point than propane even though they have the same number of electrons?

A
  • alcohols can form hydrogen bonds
  • alkanes can only form London forces
  • hydrogen bonds are the strongest intermolecular forces and require a lot of energy to overcome compared to London forces (weakest type).
  • therefore alcohols have higher boiling points and relatively lower volatility compared to alkanes.
44
Q

How many hydrogen bonds can a water molecule form?

A

2 per molecule

45
Q

Why does ice float on top of water?

A
  • water molecules in ice are arranged in rings of six, held together by hydrogen bonds.
  • there are large area of open space inside the rings.
  • when ice melts, the ring structure is destroyed and the average distance between the molecules decrease, increasing the density.
  • this means that solid ice is less dense than liquid water, because the water molecules are further apart from each other compared to in liquid water.
46
Q

Why do ionic substances dissolve in water?

A
  • the δ - end of water molecules attract the positive ions sufficiently to remove them from the lattice.
  • The positive ions become surrounded by water molecules.
  • the δ + end of water molecules attract negative ions sufficiently to remove them from the lattice.
  • the negative ions become surrounded by water molecules.
  • this process is called hydration and the energy released is known as the ‘hydration energy’.
47
Q

Solubility of alcohols in water

A
  • shorter chained alcohols are soluble in water because they are able to form hydrogen bonds with water molecules.
  • as chain length increases, solubility decreases.
48
Q

What compounds are insoluble in water?

A
  • e.g polar molecules like halogenoalkanes
  • e.g non polar substances like hexane.

They are insoluble in water because they are unable to form hydrogen bonds with water molecules.

49
Q

Solubility in non-aqueous solvents

A

Compounds that have similar intermolecular forces to those in the solvent will generally dissolve.

50
Q

Non-polar substances will dissolve other non-polar substances

A
  • Iodine only forms London forces between its molecules.
  • Non polar solvent such as hexane also only forms London forces between its molecules.
  • Therefore iodine will be able to dissolve in hexane.
51
Q

What is metallic bonding?

A

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

52
Q

What factors affect strength of metallic bonding?

A
  • number of protons in the nucleus (more protons, stronger the metallic bond).
  • Number of delocalised electrons per atom (outer shell electrons become delocalised) (more delocalised electrons, stronger the bond).
  • size of ion (smaller the ion, the stronger the bond).
53
Q

Why do metals have high melting temperatures?

A

Metallic bonds (strong electrostatic forces of attraction between positive metal cations and the delocalised electrons surrounding them) are strong, so they require a lot of energy to break.

54
Q

Why does Mg have stronger metallic bonding than Na?

A
  • More electrons donated to the ‘sea’ of delocalised electrons per Mg atom (more delocalised electrons per unit volume).
  • Mg2+ ions have a greater charge than Na+ ions.
  • Mg ions have a smaller ionic radii than Na ions (distance between positive nuclei and delocalised electrons is shorter, stronger bond).
  • Mg ions have one more proton than Na ions.
  • There is a stronger electrostatic force of attraction between the positive Mg cations and the delocalised electrons
  • therefore higher melting point.
55
Q

Can metals conduct electricity and why?

A
  • Yes

- The delocalised electrons are able to carry a charge and move freely throughout the metal structure.

56
Q

Why are metals malleable?

A
  • the delocalised electrons can move freely throughout the structure.
  • when the layers of cations slide over each other, the delocalised electrons prevent strong forces of repulsion forming between the cations in one layer and another layer.
57
Q

properties of metals?

A
  • conductor of electricity
  • conductor of heat
  • malleable
  • ductile
58
Q

How are metals able to conduct heat?

A
  • delocalised electrons are free to move throughout the structure and able to pass kinetic energy along the material.
  • The metal cations are closely packed and pass kinetic energy from one cation to another.
59
Q

Why does diamond have such a high melting temperature?

A
  • diamond is a carbon allotrope and is a giant covalent lattice structure.
  • In order to melt diamond, the strong covalent bonds between the carbon atoms must break.
  • A lot of energy is required to break the many strong covalent bonds.
  • Hence high melting temperature.
60
Q

Why can’t diamond conduct electricity?

A
  • All of the 4 outer electrons of each carbon atom is involved in covalent bonding.
  • They are localised and cannot move.
  • therefore there are no delocalised electrons to carry a charge throughout the whole structure.
61
Q

Why does graphite have a high melting temperature?

A
  • Graphite is a carbon allotrope and is a giant covalent lattice structure.
  • In order to melt, the strong covalent bonds between the carbon atoms must break.
  • A lot of energy is required to break the many strong covalent bonds.
  • Hence high melting temperature.
62
Q

Why can graphite conduct electricity?

A
  • 3 of the 4 outer electrons of each carbon atom is involved in covalent bonding, forming hexagonal rings.
  • the fourth outer electron per carbon atom is delocalised. These delocalised electrons are free to move throughout the individual graphite layers.
  • they can carry a charge.
  • Graphite is able to conduct electricity parallel to its layers.
63
Q

Why can graphite be used as a solid lubricant?

A
  • The graphite layers are able to slide easily over each other.
  • There are only weak London forces of attraction between the layers which can be easily broken.
64
Q

Graphene

A
  • carbon allotrope
  • basically a single layer of graphite
  • 3 of 4 outer electrons of each carbon atom involved in covalent bonding.
  • 4th outer electron per atom is delocalised and can carry a charge, able to freely move throughout the structure, hence graphene can conduct electricity.
  • high melting temperature due to many covalent bonds that have to be broken.
65
Q

What is a use of carbon nanotubes?

A
  • They can be used to deliver drugs to cells.
66
Q

Why are ionic substances brittle?

A
  • when layers of ions slide over one another
  • ions of the same charge are now next to each other
  • they repel one another and the crystal breaks apart.
67
Q

properties of metals

A
  • high melting temperatures
  • conductor of electricity (delocalised electrons)
  • thermally conductive
  • malleable
  • ductile
68
Q

properties of ionic compounds

A
  • high melting temperatures
  • brittle
  • solid ionic compounds cannot conduct electricity (ions fixed in lattice).
  • aqueous solutions of ionic compounds and molten ionic compounds are able to conduct electricity (ions are free to move)
  • soluble in water

CHEMISTRY FOR MOCKS 2022 (2 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions