Brainscape's Knowledge GenomeTM

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


See full index

Chemistry Content > 2 - bonding > Flashcards

2 - bonding Flashcards

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

typical properties of metallic bonding

A
  • high melting point
  • good electrical conductivity
  • good thermal conductivity
  • malleable
  • ductile
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

what does metallic bonding look like?

A

rows of metal cations with delocalised electrons from the outer shells of the atoms throughout

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

what is metallic bonding?

A

electrostatic force of attraction between the nuclei of the cations and the delocalised electrons

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

what affects the melting temperature of metals?

A
  • necessary to overcome the forces of attraction between the nuclei of the cations and the delocalised electrons to such an extent that the cations are free to move around the system
  • metals have a giant lattice structure where there are many forces to overcome (requires lots of energy)
  • Group 1 metals have low melting temps whereas Group 2 metals have higher melting temps
  • metals in the d block tend to have higher melting temps since they have more delocalised electrons per ion
  • the smaller the cation the closer the delocalised electrons to the nucleus of the cation - increase in force of attraction between the nuclei and electrons so increased melting temp
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

how come metals conduct electricity?

A

when a potential difference is applied across the ends of a metal the delocalised electrons will be attracted to and move towards the positive terminal of the cell - the movement of electrical charge constitutes an electric current

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

what contributes to the ability of metals to transfer heat energy?

A
  • free moving delocalised electrons pass kinetic energy along the metal
  • cations are closely packed and pass kinetic energy from one cation to another
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

how come metals are malleable and ductile?

A

depend on the ability of the delocalised electrons and the cations to move throughout the structure of the metal
when stress is applied to a metal, the layers of cations may slide over one another however since the delocalised electrons are free moving, they move with the cations and prevent strong forces of repulsion forming between the cations in one layer and the cations in another layer

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

what kind of ‘force’ makes up an ionic bond? (not sure about the question wording here!)

A

strong electrostatic interactions between the ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

when is it hardest to separate two ions?

A

smallest anion and smallest cation such as LiF
a higher charge also makes it more difficult such as Mg^2+

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

what happens to ionic radii as you go down the group?

A

ions have more electron shells and therefore get larger

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

what does isoelectronic mean?

A

same number of electrons and so the same electronic configuration

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

typical properties of ionic compounds?

A
  • high melting temp
  • brittleness
  • poor electrical conductivity when solid but good when molten
  • often soluble in water
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

why do ionic compounds have high melting points?

A

ionic solids consist of a giant lattice network of oppositely charged ions. there are many ions in the lattice and the combined electrostatic forces of attraction among all of the ions is large
so… a large amount of energy is required to overcome the forces of attraction sufficiently for the ions to break free from the lattice and be able to slide past one another

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

what makes ionic solids brittle?

A

if stress is applied to a crystal of an ionic solid, then the layers of ions may slide over one another
ions of the same charge are now side by side and repel one another
the crystals break apart

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

what is the electrical conductivity of ionic compounds like?

A

solid - do not as there are no delocalised electrons and the ions are also not free to move
molten - conduct as the ions are mobile and will migrate to the electrodes when a potential difference is applied
(solid Li3N will conduct electricity though)
aqueous solutions - conduct and undergo electrolysis since the lattice breaks down into separate ions when the compound dissolves

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

what is the solubility of ionic compounds like?

A

many are soluble in water
the energy required to break apart the lattice structure and separate the ions can be supplied by the hydrogen of the separated ions produced
both + and - ions are attracted to the water molecules because of the polarity water molecules possess

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

what evidence is there for the existence of ions?

A

ability of an ionic compound to conduct electricity and undergo electrolysis when either molten or in aqueous solution
e.g. direct electric current is passed through molten sodium chloride, sodium is formed at the negative electrode and chlorine is formed at the positive electrode
explanation:
- positive sodium ions migrate towards the negative electrode where they gain electrons and become sodium atoms
- negative chloride ions migrate towards the positive electrode where they lose electrode and become chlorine molecules

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

when is a covalent bond formed?

A

between 2 atoms when an atomic orbital containing a single electron from one atom overlaps with and atomic orbital, also containing a single electron, of another atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

what leads to the formation of a sigma bond?

A

an end-on overlap e.g. two s orbitals or end on overlap of two p orbitals
leads to formation of a single covalent bond between the two atoms

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

what leads to the formation of a pi bond?

A

sideways overlap of two p orbitals
cannot form until a sigma bond has been formed
pi bonds only exist between atoms that are joined by double or triple bonds

(sigma bonds are usually stronger)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

what relation does bond length have with bond strength?

A

the shorter the bond length the greater the strength
also double, triple, etc bonds are stronger and shorter than single bonds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

what is electronegativity?

A

the ability of an atom to attract a bonding pair of electrons
- decreases down a group
- increases across a period

23
Q

what is a polar covalent bond?

A

bond between two atoms where the bonding electrons are unequally distributed. because of this, one atom carries a slight negative charge and the other a slight positive charge

inbetween pure covalent and pure ionic character

24
Q

when does a dative covalent bond form?

A

when an empty orbital of one atom overlaps with an orbital containing a non-bonding pair (lone pair) of electrons of another atom
represented by an arrow starting from the atom providing the pair of electrons and going towards the atom with the empty orbital

25
Q

what is electron pair repulsion theory?

A
  • shape of a molecule or ion is caused by repulsion between the pairs of electrons, both bond pairs and lone (non-bonding) pairs, that surround the central atom
  • the electron pairs arrange themselves around the central atom so that the repulsion between them is at a minimum
  • lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion
26
Q

what shape is formed from 2 bond pairs? 2 bond pairs and 2 lone pair?

A

linear - 180*
v-shaped - 104.5*

27
Q

what shape is formed from 3 bond pairs? 3 bond pairs and 1 lone pair?

A

trigonal planar - 120*
trigonal pyramidal - 107*

28
Q

what shape is formed from 4 bond pairs?

A

tetrahedral - 109.5*

29
Q

what shape is formed from 5 bond pairs?

A

trigonal bipyramidal - 90* and 120*

30
Q

what shape is formed from 6 bond pairs?

A

octahedral - 90* and 180*

31
Q

what is a dipole?

A

when two charges of equal magnitude but opposite signs are separated by a small distance

32
Q

London forces

A

Fritz London, German physicist, 1930
weaker than covalent and polar bonds
- attractive force increases with increasing number of electrons in the molecule (the more electrons, the greater the fluctuation in electron density and the larger the instantaneous and induced dipoles created)
- the more points of contact between the molecules the greater the London force
- present between all molecules

33
Q

what is an instantaneous dipole?

A

created as electron density fluctuates over time

34
Q

what is an induced dipole?

A

when the electron density is pulled closer to the other molecule generating a partial negative charge on one side and a positive charge on the other

35
Q

permanent dipoles

A

less than the interaction between instantaneous-induced dipoles

possible to induce a dipole in a nearby molecule
permanent dipole-permanent dipole and permanent dipole-induced dipole are permanent dipole-dipole

36
Q

hydrogen bond

A

only with atoms more electronegative than H (O, N, F - usually)
hydrogen bonds form between O and one water molecule and H and one water molecule - extreme dipole-dipole interaction, usually 180*

37
Q

why does the boiling temp increase with increasing molecular mass for unbranched alkanes?

A

1 as molecular mass increases, number of electrons per molecule increases and so the instantaneous and induced dipoles increase
2 length ncreases, number of points of contact between adjacent molecules increase. instantaneous dipole-induced dipole forces exist at each point of contact between the molecules, more points of contact = greater overall London force of attraction

38
Q

why do branched alkanes have lower boiling points than unbranched?

A

fewer points of contact between adjacent molecules i.e. they do not pack as well together therefore decrease in overall intermolecular force of attraction between molecules and a decrease in boiling point

39
Q

why do alcohols have higher boiling points than their equivalent alkane?

A

O-H group can form hydrogen bonds in addition to London forces

40
Q

what is the boiling temp of hydrogen halides like?

A

HF is highest
then increases from HCl to HBr and then HI
on a graph: HF decreases to HCl then constant increase up to HI

41
Q

what anomalous properties does water have?

A

1 high melting and boiling point for a molecule with such few electrons
2 density of ice at 0C is less than that of water at 0C

42
Q

why is the bp of water higher than expected when compared to molecules with the same number of electrons such as HF?

A

usually argument is strong hydrogen bonding however HF has higher hydrogen bond strength
1 HF forms an average of 1 H bond per molecule whereas water forms an average of 2 per molecule
2 not all H bonds in HF are broken on vaporisation since HF is substantially polymerised even in the gas phase

43
Q

why is the density of ice less than that of liquid water?

A

the molecules in ice are arranged in rings of 6 held together by hydrogen bonds
the structure has large areas of open space inside the rings
when ice melts the ring structure is destroyed and the average distance between the molecules decreases causing an increase in density

44
Q

what conditions must be met for a substance to dissolve?

A
  • solute particles must be separated from each other and then become surrounded by solvent particles
  • forces of attraction between the solute and the solvent particles must be strong enough to overcome the solvent-solvent forces and the solute-solute forces
45
Q

when dissolving an ionic solid what interaction is created between the ions and water molecules?

A

ion-dipole interaction

46
Q

why does solubility of alcohols decrease in water?

A

increasing chain length as London forces predominate between the alcohol molecules

47
Q

what compounds do not dissolve in water?

A

non-polar molecules such as alkanes as the attraction is not strong enough to disrupt the hydrogen bonded system between the water molecules
polar molecules can have limited solubility in water because they don’t form hydrogen bonds with water or the hydrogen bonds are weak compared to those in water
haloalkanes are not very soluble in water either

48
Q

what is the rule of thumb for dissolving something?

A

like dissolves like
solvent for a non-polar substance likely liquids that contain similar molecules
e.g. alkanes are soluble in one another

49
Q

what are the 4 most common giant covalent substances?

A

diamond, graphite, graphene, silicon (IV) oxide

50
Q

what is the bonding like in diamond?

A
  • each carbon forms four sigma bonds to 4 other carbon atoms
  • all bond angles are 109.5*
  • extremely hard bc of the strong C-C bonding
  • high melting temp bc a great number of strong C-C have to be broken in order to melt it which requires a large amount of heat energy
51
Q

what is bonding like in graphite?

A
  • each carbon atom is bonded to 3 others by sigma bonds forming interlocking hexagonal rings
  • the fourth electron on each carbon is in a p orbital to overlap with one another to produce a cloud of delocalised electrons above and below the plane of the rings
  • can be used as a solid lubricant - weak London forces between the layers (lubricating ability decreases x5 at high altitude and x8 in a vacuum)
  • fairly good conductor of electricity as delocalised electrons between the layers are free to move under pd (only conducts electricity parallel to its layers)
  • high mp bc a great number of strong C-C have to be broken in order to melt it which requires a large amount of heat energy
52
Q

what is bonding like in graphene?

A
  • pure carbon in the form of a very thin sheet, one atom thick
  • bonded exactly the same way as in graphite and it can therefore be described as a one atom thick layer of graphite
53
Q

examples of molecular solids?

A

ice, iodine (I2), sulfur (S8), buckminster fullerene (C60), sucrose (C12H22O11), solid alkanes, etc

54
Q

what are the physical properties of molecular solids?

A
  • generally have low mp - to melt it is only necessary to overcome the intermolecular forces of attraction not the covalent bonds within the molecule, to boil weak covalent bonds require little energy
  • London forces increase with increased electrons so polyethene has a higher mp than ethene

Chemistry Content (19 decks)

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