bonding p1,2 Flashcards

(68 cards)

1
Q

what are ions?

A

charged particles that are formed when an atom loses or gains electrons

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

what is the charge of an ion when electrons are gained?

A

negative
positive when lost

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

what are molecular ions?

A

covalently bonded atoms that lose or gain electrons

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

which electrons are lost when as atom becomes a positive ion?

A

electrons in the highest energy levels

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

do metals usually gain or lose electrons?

A

lose electrons

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

which are the 4 elements that don’t tend to form ions and why?

A

Beryllium, boron, carbon and silicon
requires lots of energy to transfer outer shell electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

what are the 3 main types of chemical bonds?

A

ionic, covalent, metallic

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

ionic bonding

A

electrostatic attraction between positive nucleus and negative ions

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

what determines the strength of an ionic bond?

A
  • ionic radius
  • ionic charge
  • ionic bonding is stronger and the melting points higher when the ions are smaller and have higher charges
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

what is the trend in ionic radius down a group?

A

increases down the group
because the ions have more shells of electrons and thus the outermost electron experiences less pull from positive nucleus

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

what are the physical properties of ionic compounds?

A
  • high melting points
  • non conductor of electricity when solid
  • conductor when in solution or molten
  • brittle
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

in a solution of CuCrO4 with connected electrodes which electrode will the 2 ions (Cu2+ & CrO2-) migrate to?

A

Cu2+ - migrates to negative electrode
CrO2- - migrates to positive electrode

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

covalent bonding

A

electrostatic attraction between a shared pair of electrons and the nuclei

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

metallic bonding

A

electrostatic attraction between the positive metal ions and the sea of delocalised electrons

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

why do giant ionic lattices conduct electricity when liquid but not when solid?

A

in solid state the ions are in fixed positions and thus cannot move
when in liquid state the ions are mobile and thus can freely carry charge

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

giant ionic lattices have high melting and boiling point. why?

A

large amount of energy is required to overcome the electrostatic bonds

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

what type of solvents do ionic lattices dissolve in?

A

polar solvents
e.g. water

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

why are ionic compounds soluble in water?

A

water has a polar bond
hydrogen atoms have a 1+ charge and oxygen atoms have a 2- charge
these charges are able to attract charged ions

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

what is it called when atoms are bonded by a single pair of shared electrons?

A

single bond

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

what is the effect of multiple covalent bonds on bond length and strength?

A

double/triple bonds exert greater electron density therefore the attraction between nucleus and electron is greater resulting in a shorter and stronger bond

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

what is a lone pair?

A

electrons in the outer shell that are not involved in the bonding

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

what is formed when atoms share two pairs of electrons?

A

double bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

what is a dative covalent bond?

A

a bond where both of the shared electrons are supplied by one atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

how are oxonium ions formed?

A

formed when acid is added to water
H3O+

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
what does expansion of the octet mean?
when a bonded atom has more than 8 electrons in the outer shell
26
what are the types of covalent structure?
- single molecular lattice - giant covalent lattice
27
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
28
why do simple molecular structures have low melting and boiling points?
small amount of energy is enough to overcome the intermolecular force
29
can simple molecular structures conduct electricity?
no - no free charged particles to move around
30
simple molecular structures dissolve in what type of solvent?
non polar solvents
31
examples of giant covalent structures
diamond, graphite, silicon dioxide
32
properties of giant covalent structures?
- high melting and boiling point - non conductors (except graphite) - insoluble in polar and non polar solvents
33
how does graphite conduct electricity?
delocalised electrons present between the layers are able to move freely carrying the charge
34
why do giant covalent structures have high melting and boiling points?
strong covalent bonds within the molecules need to be broken which requires lots of energy
35
describe the structure of diamond
3D tetrahedral structure of C atoms, with each C atom bonded to 4 others
36
what does the shape of a molecule depend on?
number of electron pairs in outer shell, number of electrons which are bonded/lone pairs
37
2 bonded pairs shape
linear 180*
38
3 bonding pairs shape
trigonal planar 120*
39
4 bonded pairs shape
tetrahedral 109.5*
40
5 bonded pairs shape
trigonal bipyramid 90* and 120*
41
6 bonded pairs
octahedral 90*
42
3 bonded pairs 1 lone pair shape
pyramidal 107*
43
2 boned pairs and 2 lone pairs shape
non linear (bent) 104.5*
44
by how many degrees does each lone pair reduce the bond angle?
2.5*
45
electronegativity
the ability of an atom to attract a (bonding) pair of electrons in a covalent bond
46
what scale is electronegativity measured on?
Pauling scale
47
which direction of the periodic table does electronegativity increase?
top right, towards fluorine
48
what does is mean when a bond is non-polar?
the electron in the bond are evenly distributed
49
how is a polar bond formed?
bonding atoms have different electronegativities
50
why is H2O polar and not CO2?
CO2 is symmetrical so there is no overall dipole
51
intermolecular force
attractive force between neighbouring molecules
52
what are the 3 types of intermolecular forces?
hydrogen bonding, permanent dipoles, London forces
53
permanent dipole-induced dipole interactions
when a molecules with a permanent dipole is close to other non polar molecules it causes the non polar molecule to become slightly polar leading to attraction
54
permanent dipole-permanent dipole interactions
some molecules with polar bonds have permanent dipoles forces of attraction between those dipoles and those of neighbouring molecules
55
London forces
- caused by random movements of electrons - leads to instantaneous dipoles - instantaneous dipole induces a dipole in nearby molecules - induced dipoles attract one another
56
are London forces greater in smaller or larger molecules?
larger due to more electrons
57
does boiling point increase or decrease down the noble gas group? why?
increases the number of electrons increases and hence the strength of London forces also increases
58
what conditions are needed for hydrogen bonding?
between hydrogen atom and a highly electronegative atom - O, N, F
59
why is ice less dense than liquid water?
in ice the water molecules are arranged in an orderly pattern it has an open lattice with hydrogen bonds in water the lattice is collapsed and the molecules are closer together
60
why does water have a higher melting/boiling point than expected?
hydrogen bonds are stronger than other intermolecular forces so extra strength is required to overcome the forces
61
what type of intermolecular forces do alkanes have? why?
London force induced dipole-dipole interaction becuase the bonds are non polar
62
what happens to the boiling point as alkane chain length increases? why?
boiling point increases bc there is more surface area and so higher number of induced dipole-dipole interactions therefore more energy required to overcome the attraction
63
does a branched molecule have lower or higher boiling point compared to equivalent straight chain? why?
branched molecule has a lower boiling point bc they have smaller surface area and hence less induced dipole-dipole interactions
64
are alkanes soluble in water?
insoluble because hydrogen bonds in water are stronger than alkane's London forces
65
what kind of intermolecular forces do alcohols have?
Hydrogen bonding due to the electronegativity difference in the OH bond London forces
66
how do alcohols' melting point and boiling point compare to other hydrocarbons' of similar C chain lengths?
higher due to hydrogen bonding being stronger than London forces
67
are alcohols soluble in water? why does solubility depend on chain length?
soluble when short chain - OH hydrogen bonds to hydrogen bond in water insoluble when long chain - non-polar C-H bond
68
explain the trend of boiling temperatures of hydrogen halides HF to HI
general increase of boiling point from HCl to HI which is caused by increasing London forces becuase of increasing number of electrons there is a big drop from HF to HCl becuase F is very electronegative therefore the hydrogen bonding is much stronger