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Chemistry 3.1 Physical chemistry > bonding > Flashcards

bonding Flashcards

(86 cards)

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1
Q

what is the formula for a sulphate ion?

A

SO4(2-)

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2
Q

what is the formula for carbonate ions?

A

CO3(2-)

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3
Q

what is the formula for a nitrate ion?

A

NO3(-)

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4
Q

what is the formula for a hydroxide ion?

A

OH-

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5
Q

what is the formula for phosphate ions?

A

PO4(3-)

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6
Q

what is the formula for an ammonium ion?

A

NH4(+)

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7
Q

what is the formula for a hydrogen carbonate ion?

A

HCO3(-)

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8
Q

what is ionic bonding?

A

the electrostatic force of attraction between oppositely charged ions formed by electron transfers

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9
Q

when is ionic bonding stronger and higher melting points?

A

when the ions are smaller an /or have higher charges

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10
Q

why is ionic bonding stronger with smaller and higher charged ions?

A

smaller ions can be packed closer together (shorter distance between oppositely charged ions makes the electrostatic attraction stronger)

bigger charges exert bigger forces between each other

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11
Q

why are positive ions smaller compared to their atoms?

A

it has one less shell of electrons and the ratio of protons to electrons has increased so there is a greater net force on remaining electrons holding them more closely.

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12
Q

where are negative ions larger than the atoms?

A

groups five to seven

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13
Q

why do negative ions form larger ions than the atom?

A

the ion has more electrons than the atom but the same number of protons. so the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger

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14
Q

as you go down the group, what happens to the size of the atomic radi and why?

A

it increases down the group because the ions have more shells of electrons

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15
Q

what is the trend of size of atoms as you go across a period and why?

A

smaller in size because across the period there are more protons so greater attraction between protons and electrons, so electrons are pulled in towards the nucleus, making the atomic radius smaller across the period

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16
Q

why are ionic compounds brittle and shatter when compressed?

A

force applied

ions arranged in layers and they shift, rearranging themselves

same charged ions repel and the compound shatters

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17
Q

why do ionic compounds conduct electricity charge when molten or in aqueous form?

A

ions are not free to move in solid state
ions are free to move in the liquid state(and can therefore carry charge) and the compound conducts electricity

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18
Q

why are ionic compounds solids at room temperature?

A

strong electrostatic attraction between oppositely charged ions, forming a crystal lattice

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19
Q

why do ionic compounds have a high melting point?

A

ions are held together by strong electrostatic forces. it takes a lot of energy to overcome these forces

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20
Q

what are some notes to know about ionic bonding in group 4?

A

carbon and silicon don’t usually form ions

in tin the Sn4+ is more stable but Sn 2+ is also seen

lead compounds generally contain Pb2+ ion but some contain Pb4+

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21
Q

what is a covalent bond?

A

shared pair of electrons

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22
Q

when does a dative covalent bond form?

A

when the shared pair of electrons in teh covalent bond come from only one of the bonding atoms.

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23
Q

what is another name for a dative bond?

A

co-ordinate bonding

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24
Q

what is the lone pairs equation?

A

electrons in outer shell- number of bonded atoms / 2

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25
in dative bonding, what does the atom that donates the pair of electrons have?
a lone pair
26
what is the condition in dative bonding for the atom that accepts the lone pair?
cannot have a full outer shell of electrons
27
what can co-ordinate bonds be shown by?
an arrow- it points from the atom donating the electrons
28
how does the dative covalent bond behave after it forms?
behaves like a normal covalent bond
29
why do simple covalent compounds have low melting/ boiling points?
weak intermolecular forces between the molecules so not much energy is needed to break these.
30
why can simple covalent compounds not conduct electricity in any state of matter?
they have no delocalised electrons or ions to move throughout the structure and carry charge
31
why do simple covalent compounds generally have poor solubility in water ?
water is polar and covalent compounds are non polar
32
describe the structure of diamond?
Tetrahedral each carbon atom bonded to 4 others (covalent bonds between atoms) all 4 outer shell electrons used in bonding (no delocalised electrons)
33
explain why diamond has a really high mp?
giant covalent lattice many strong covalent bonds between atoms 4 bonds per carbon a lot of energy needed to overcome these many strong covalent bonds
34
why are diamond, graphite and graphene insoluble in water?
they are non-polar and water is polar
35
why can't diamond conduct electricity?
no ions and no delocalised electrons to move through the structure and carry charge throughout the structure
36
describe the structure of graphite?
layers of carbon atoms each bonded to 3 others with covalent bonds between atoms between the layers there are weak van der waals forces
37
explain why graphite has a high mp?
giant covalent lattice many strong covalent bonds between atoms 3 bonds per carbon a lot of energy needed to overcome bonds
38
why can graphite/ graphene conduct?
each carbon atom is bonded to 3 others so one electron per carbon is delocalised so can move in the structure, carrying charge
39
why is graphite soft?
between layers there are weak van der waal forces which can break easily so layers can slide over each other
40
since graphite is soft, what can it be used as?
a lubricant
41
describe the structure of graphene?
it is a single layer of graphite each carbon atom is bonded to 3 others
42
why does graphene have a high mp?
giant covalent lattice many strong covalent bonds between atoms each carbon atom is bonded to 3 others a lot of energy needed to overcome the bonds
43
how do we describe structure in metallic structures?
a lattice of positive metal ions existing in a sea of delocalised electrons
44
what kind of attraction is found between positive ions and delocalised electrons in metallic bonding?
electrostatic attraction
45
why does metallic bonding result in high melting and boiling points?
strong electrostatic attraction between positive metal ions and a sea of delocalised electrons within the giant metallic lattice- therefore large amounts of energy needed to overcome strong attraction
46
why do metallic substances conduct as solids and liquids?
electrons are free to move throughout the structure since they are delocalised- they can carry charge through the structure
47
why are metals malleable and ductile?
layers in the giant metallic lattice can slide over each other into new positions, without disrupting the metallic compound (unlike ionic compounds where force applied results in repelling causing shattering)
48
why are metallic structures generally soluble in water?
water is polar- part of its structure is slightly negative and part is slightly positive. Metals have a charge (+ or -) so attraction means it can be dissolved
49
what are the three factors which affect the strength of metallic bonding?
1. number of protons/ strength of nuclear attraction 2. Number of delocalised electrons per atom 3. size of ion (smaller ion= stronger bond)
50
what is electronegativity?
the ability for an atom to attract electrons towards itself it a covalent bond.
51
what scale is electronegativity measured on?
the Pauling scale (0 to 4)
52
what is the most electronegative element?
flourine
53
why does electronegativity increase across a period?
number of protons (nuclear charge) increases and the atomic radius decreases because the electrons in the same shell are pulled in more
54
why does electronegativity decrease down a group?
distance between nucleus and outer electrons increases and the shielding of inner shell electrons increases
55
how will a compound be "purely covalent" due to electronegativity?
a compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent.
56
how will a compound be ionic due to electronegativity?
a compound containing elements of very different electronegativity and hence a very large negativity difference will be ionic
57
when does a polar covalent bond form?
when the elements in the bond have different electronegativities
58
what happens to the electrons in a polar covalent bond?
there is an unequal distribution of electrons and this produces a charge separation (dipole) S+ and S- ends
59
are symmetrical molecules polar? why/
a symmetrical molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds between atoms are polar because the individual dipoles 'cancel out' - there is no net dipole movement so the molecule is non-polar
60
what factors effect electronegativity?
1. nuclear charge 2. shielding (lessons the pull by the nuleus) this is the most important factor to consider 3. closeness to flourine
61
what factors can make a molecule non-polar even if it has polar bonds?
-symmetrical molecules
62
what are the three types of inter molecular forces in order of their decreasing strength?
1. Hydrogen bonding 2. Permanent dipole - permanent dipole forces 3. Van der Waals' forces
63
when do van der waals forces occur / not occur?
occur between all molecular susbstances and noble gases not in ionic substances
64
what is another name for van der waals forces?
transient, induced dipole- dipole interactions
65
where do van der waals (induced dipole-dipole) forces exist?
between atoms and molecules
66
when can atoms form a dipole?
any molecule with electrons can form a dipole when they move near another atom or molecule
67
why do temporary dipoles occur?
occurs as electrons in a molecule or atom can move from one end to another, allowing one side of the atom to be S- and one side to be S+
68
how is a force of attraction created in order to form a temporary dipole-dipole bond?
the S+ on one atom is attracted to the S- on another and a force of attraction is created
69
how do temporary dipole interactions get destroyed?
this temporary dipole only exists when 2 molecules or atoms are near by. When they move away the interaction is destroyed
70
what can van der waals forces do to the structures?
it can hold some molecules in crystal structures e.g. iodine. van der waals forces are weak at holding molecules together
71
what bonds are we breaking when we boil a liquid?
weak van der waals forces and not covalent bonds
72
how are van der waals forces affecting the boiling point of hydrocarbons?
long, straight chain hydrocarbons have more van der waals forces and so more energy is needed to over come these forces- bp increases
73
where do permanent dipole- dipole interactions exist?
in molecules with polarity
74
how is strength different in induced and permanent dipole-dipole interactions?
dipole-dipole interactions involve molecules with a permanent dipole, unlike van der waals, and so are STRONGER
75
molecules that have dipole-dipole interactions also have...
van der waals forces
76
how can polar molecules (like water) be tested?
by placing a charged rod near a steady stream of polar liquid. You should see the liquid bend towards the rod as the molecules align to face the oppositely charged rod.
77
when does hydrogen bonding occur?
when hydrogen on one molecule forms a bond with the lone pair on nitrogen , oxygen or flourine (3 most electronegative elements)
78
how do we show hydrogen bonding?
dotted lines between the lone pairs and hydrogen
79
what do molecules that have hydrogen bonding also have?
van der waals and dipole-dipole too
80
describe the particle arrangement in soilds?
- particles are tightly packed and in regular arrangement which is why they have a high density -particles vibrate on the spot and can't be compressed
81
describe particle arrangement in liquids?
-particles are tightly packed and in random arrangement which is why they have a high density -particles move freely it is very difficult for liquids to be compressed -particles in a liquid have more energy than solid
82
describe how particles are arranged in gases?
-particles spread out in a random arrangement which is why they have a low density - particles move freely so it is relatively easy for gases to be compressed -particles have more energy in gas than solid and liquid
83
what can the shape of molecules be predicted by?
electron pair repulsion theory
84
how do electron pairs and lone pairs repel?
electron pairs repel each other so that they are as far apart as possible lone pairs repel slightly more than bonded pairs because they are more compact
85
what is the order of repulsion?
lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
86
how much do lone pairs reduce the bond angle by?
2.5 degrees per bond pair

Chemistry 3.1 Physical chemistry (6 decks)

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