topic 4/14- bonding and structure Flashcards
1
Q
describe how positively charged ions form
A
metals lose valence electrons to form cations
2
Q
describe how negatively charged ions form
A
non metals gain valence electrons to form anion
3
Q
define an ionic bond
A
the electrostatic attraction between oppositely charged ions
4
Q
state the formula and relative charge of ammonium
A
NH₄⁺
5
Q
state the formula and relative charge of hydroxide
A
OH¯
6
Q
state the formula and relative charge of nitrate
A
NO3-
7
Q
state the formula and relative charge of hydrogencarbonate
A
HCO3-
8
Q
state the formula and relative charge of carbonate
A
CO₃²⁻
9
Q
state the formula and relative charge of sulfate
A
SO₄²-
10
Q
state the formula and relative charge of phosphate
A
PO₄³⁻
11
Q
describe ionic compounds under normal conditions
A
solids with lattice structures
12
Q
describe the physical properties of ionic compounds:
- melting/boiling point
- volatility
- electrical conductivity
- solubility
A
- high melting/boiling points
- low volatility
- do not conduct when solid, but conduct when molten or aqueous
- soluble
13
Q
explain the high melting/boiling points of ionic compounds
A
- strong EFOA between oppositely charged ions
- lot of energy required to overcome them
14
Q
does magnesium oxide or sodium chloride have a higher melting point? why?
A
magnesium oxide; the EFOA between the 2+ and 2- ions in MgO are much stronger than those between the 1+ and 1- ions in NaCl
15
Q
explain the low volatility of ionic compounds
A
the strong EFOA between oppositely charged ions take a lot of energy to overcome
16
Q
explain the electrical conductivity of ionic compounds
A
- in a solid state, the ions are in fixed positions (ionic lattice) and cannot move.
- when aqueous or molten, the ions are able to move as the lattice is broken down
17
Q
describe what happens to the ions in an ionic compound when it dissolves
A
- bonds/attraction between the solid particles are broken
- new bonds form between sps and water molecules
- water molecules surround sps
- due to sliding movement of water molecules, sps move through liquid until evenly distributed
18
Q
why are some substances soluble and others insoluble?
A
there is an energy cost and an energy gain which determines whether a process occurs without an input of energy
19
Q
a substance will be soluble if the energy gain is —– than the energy cost.
A
greater than
20
Q
state the two sources of energy cost in dissolving
A
- breaking hydrogen bonds in water/solvent
- breaking bonds between solute particles
21
Q
state the two sources of energy gain in dissolving
A
- formation of bonds between solute and solvent
- increasing entropy
22
Q
explain the solubility of ionic compounds in polar solvents like water
A
energy released when ions surrounded by H2O molecules (ion-dipole interactions) > energy required to break ionic lattice
23
Q
explain the insolubility of ionic compounds in non polar solvents
A
energy required to break apart ionic lattice > energy released when non-polar solvent forms interactions with the ions (London forces)
24
Q
define a covalent bond
A
the electrostatic attraction between a shared pair of electrons and two positively charged nuclei
25
describe a single covalent bond
one shared pair of electrons
26
describe a double covalent bond
two shared pairs of electrons
27
describe a triple covalent bond
three shared pairs of electrons
28
describe a coordinate covalent bond
a covalent bond where one atom has donated two electrons to the shared pair
29
describe the relationship between bond length and bond strength
as bond length increases, bond strength decreases
30
describe the relationship between number of shared pairs and bond strength/length
as number of shared pairs increases, strength of bond increases but length decreases.
31
explain what happens to the bond strength as you go from HF to HI.
bond strength decreases;
- distance of bond pair from nucleus decreases, shielding of bond pair increases
- this outweighs an increase in atomic number
32
define electronegativity
a measure of the tendency of an atom to attract a bonding pair of electrons
33
what is impure ionic bonding?
when negative ions are polarised by positive ions.
34
why does impure ionic bonding occur?
the positive ion attracts the electrons of the negative ion
35
explain why smaller, more highly charged positive ions have a greater polarising effect
smaller, more highly charged positive ions have a greater charge density, causing them to have a greater attraction to the anion
36
explain why larger negative ions are more easily polarised than smaller negative ions
larger negative ions have a greater distance and shielding, meaning the electrons are less attracted to the nucleus of the anion
37
give the formula for the most impure ionic bonding
small, highly charged positive ion + large negative ion
38
explain why electronegativity increases ACROSS a period
no. of protons is increasing (nucleus more positive) but distance and shielding are staying mostly the same
39
explain why electronegativity decreases DOWN a group
there is more shielding and an increased distance (so less attraction)
40
which element has the highest electronegativity?
fluorine
41
what is a pure covalent bond?
when the electronegativities of the two atoms are the same
42
explain the difference between covalent and ionic bonding in terms of electronegativities.
- in a covalent bond, the difference in electronegativity is zero or very small.
- in an ionic bond the elements have a large difference in electronegativity
43
describe the relationship between difference in electronegativity and type of bond
- <0.4= covalent bond
- 0.4-1.8= polar covalent
- >1.8= ionic bond
44
the more electronegative atom has a slightly ----- charge
the less electronegative atom has a slightly ----- charge
1. negative
2. positive
45
a molecule is polar if....
there is a net dipole
46
what does the 'octet rule' refer to?
the tendency of atoms to gain a valence shell with a total of 8 electrons
47
state two common exceptions two the octet rule
Beryllium and boron; these might form stable compounds with incomplete octets of electrons
48
what are lewis structures?
diagrams that show all the valence electrons in a covalently bonded species
49
when do resonance structures occur?
when there is more than one possible position for a double bond in a molecule
50
draw the resonance structure for C6H6
refer to flashcard
51
draw the resonance structure for CO₃²⁻
refer to flashcard
52
what is a resonance structure?
one of two or more alternative Lewis structures for a molecule or ion that cannot be described fully with one Lewis structure alone
53
what are the shapes of species determined by?
the repulsion of electron pairs according to VSEPR (valence shell electron pair repulsion) theory
54
describe VSEPR theory
pairs of electrons repel each other so that they are as far apart as possible, keeping the force of repulsion to a minimum
55
how many types of covalent shapes are there that ONLY include bond pairs (no lone pairs)?
5
56
name the 5 covalent molecule shapes with only bond pairs
- linear molecules
- trigonal planar molecules
- tetrahedral molecules
- trigonal bipyramidal molecules
- octahedral molecules
57
describe linear molecules (BP only)
- no of bond pairs
- bond angle
- 3 examples
- shape
- 2 bond pairs
- 180
- BeCl2, CO2 and all diatomic molecules
- straight line
58
describe trigonal planar molecules
- no of bond pairs
- bond angle
- 1 example
- shape
- 3 bond pairs
- 120
- BCl3
- flat peace-sign
59
describe tetrahedral molecules
- no of bond pairs
- bond angles
- 1 examples
- shape
- 4 bond pairs
- 109.5
- methane, CH4
- Eiffel Tower
60
describe trigonal bipyramidal molecules
- no of bond pairs
- bond angles
- 1 example
- shape
- 5 bond pairs
- 90 and 120
- PF5
- fidget spinner shot by an arrow
61
describe octahedral molecules
- no of bond pairs
- bond angles
- 1 example
- shape
- 6 bond pairs
- 90
- SF6
- christian cross that has been shot by an arrow
62
why does a lone pair have a greater repulsive effect than a bonding pair?
- lone pair electrons are localised to an atom, so they are closer to each other
63
state the hierarchy of repulsion
lp to lp> lp to bp> bp to bp
64
name the 7 types of lone pair inclusive molecule shapes
- trigonal pyramidal molecules
- v shaped/bent molecules
- square planar molecules
- square pyramidal molecules
- seesaw molecules
- t/arrow shaped molecules
- linear molecules
(two very silly Swiss singers terrify Lav)
65
describe trigonal pyramidal molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like tetrahedral molecules but without top
- 3 bp, 1 lp, 4 ed
- 107
- NH3
- beheaded Eiffel tower
66
describe v shaped/bent molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
type 1 (SO2 DB):
- like trigonal planar but without top
- 2 bp, 1 lp, 3ed
- 104.5
type 2 (H2O):
- like type 1 but extra pair
- 2 bp, 2 lp, 4 ed
- 104.5
67
describe square planar molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like octahedral but without both vertical bits
- 4 bp, 2 lp, 6 ed
- 90
- XeF4
- cross laid on its side + 2lp above and below
68
describe square pyramidal
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like octahedral but without bottom vertical bit
- 5 bp, 1 lp, 6 ed
- 85-87.5
- BrF5
- cross laid on its side + stick up + lp below
69
describe see saw molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- like trigonal bipyramidal but without left side bit
- 4 bp, 1 lp, 5 ed
- 87.5-90 (equatorial-axial), 117 (e-e)
- SF4
70
describe t/arrow shaped molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- like trigonal bipyramidal but only side 3 + 2 lp
- 3 bp, 2 lp, 5 ed
- 87.5-90
- XeOF2
71
describe linear molecules (WITH lone pairs)
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like trigonal bipyramidal but only 2 vertical bits
- 2 bp, 3 lp, 5 ed
- 180
- I3-
- straight line
72
define metallic bonding
the electrostatic attraction between a lattice of positive ions and delocalised electrons
73
why do metals conduct electricity?
the have free-flowing, delocalised electrons which are free to move.
74
why are metals malleable?
they have layers which can slide across each other without disrupting the bonding in place.
75
what would be the ratio of ions to delocalised electrons in lithium? (Li is in group 1)
1:1 (each Li atom loses ONE electron to become an ion)
76
what would be the ratio of ions to delocalised electrons in magnesium? (Mg is in group 2)
1:2 (each Mg atom loses TWO electrons to become an ion)
77
explain why melting and boiling points increase from sodium to magnesium to aluminium.
- slightly smaller radius
- greater charge
- more delocalised electrons
=> greater charge density
=> greater ESFOA so more energy required
78
explain why the melting and boiling points of group 1 metal decrease going down the group
- larger radius
- more shielding
=> lower charge density
=> less ESFOA so less energy required
79
as charge of ion increases, strength of metallic bond ----
increases
80
as radius of ion increases, strength of metallic bond ---
decreases
81
what gives alloys their properties?
the fact that they are mixtures of more than one metal.
82
gives two examples of alloys
brass (copper + zinc)
bronze (copper + tin)
83
explain the properties of alloys
- contains ions of different sizes
- this disrupts the structure/regular lattice arrangement
- the layers slide over each other less easily
=> harder/less malleable
84
what kind of bonding can carbon and silicon form?
giant covalent/network covalent structures
85
Graphite, diamond, buckminsterfullerene and graphene are ? of carbon
allotropes
86
describe the structure of graphite
- hexagonal layers of carbon
- each atom bonded to 3 others
- London forces between layers
- 120
87
describe the structure of diamond
- continuous tetrahedral arrangement of C atoms
- each atom bonded to 4 others
- 109.5
88
describe the structure of buckminsterfullerene
- 60 carbon atoms
- each atom bonded to 3 others (4th ed)
- interlocking hexagons and pentagons
89
describe structure of graphene
- hexagonal sheet
- each atom bonded to 3 others
- 1 layer thick
90
describe structure of silica, or silicon dioxide
- continuous tetrahedral structure
- each Si shared between 4 O
- each O shared between 2 Si
91
explain the properties of diamond
- high melting/boiling point: strong covalent bonds
- non conductor: no de
- hard: each atom forms 4 bonds
92
explain the properties of graphite
- high melting/boiling point: strong covalent bonds within layers
- conductor: each atom forms 3 bonds so there is 1 de
- brittle: weak London forces between layers
93
explain the properties of graphene
- high melting/boiling point: strong covalent bonds
- conductor: each atom forms 3 bonds so there is 1 de
94
explain the properties of buckminsterfullerene
- low melting/boiling point: weak intermolecular forces between molecules
- semi-conductor: each carbon forms 3 bonds so there is 1de
95
explain the properties of silica, or silicon dioxide
- high melting/boiling point: strong covalent bonds
- non-conductor: no de
96
what do covalent bonds arise from?
the overlap of atomic orbitals
97
describe sigma bonds
- formed by any head-on/axial overlap of orbitals
- s-s/s-p/p-p
98
describe pi bonds
- formed by the lateral overlap of two p orbitals
99
5 differences between sigma and pi bonds
- only 1 s bond exists between two atoms; can be more than 1 pi bond between two atoms
- e density is maximum about the bond axis; e density high above and below bond axis
- free rotation about s bond possible; not possible
- s can be independently formed; formed after sigma bond
- relatively strong; weak
100
define formal charge
the charge an atom would have if all the atoms in the molecule had the same electronegativity
101
state the equation for formal charge
FC= no of valence electrons - 0.5(no of bonding electrons) - no of non bonding electrons
102
what Lewis structure is preferred?
the one with the atoms having FC (formal charge) values closest to 0
103
When does delocalisation of electrons in a molecule arise?
when electrons are shared between all atoms in a molecule or ion
104
what wavelength of light can break down ozone?
UV-B light, 280-320nm
105
give the equation for the splitting of ozone
O3--> O2 + O'
106
which type of UV has the highest energy and shortest wavelength?
UV-C
107
why is a specific wavelength of light required to dissociate oxygen and ozone?
as a certain amount of energy is needed in order for the bonds to be broken
108
how does the ozone layer protect the earth from UV radiation?
most of UV-C and UV-B are absorbed by oxygen and ozone molecules. therefore, only UV-A reaches the earth's surface, which is less harmful and lower energy.
109
oxygen (O2) absorbs
UV-C
110
ozone (O3) absorbs
UV-B
111
state the equation for the dissociation of oxygen
O2--> 2O'
112
state the mechanisms by which O3 is broken down by CCl3F
1. CCl3F -(UV)-> 'CCl2F + Cl' (UV causes C-Cl bond to break)
2. Cl' + O3 --> ClO' +O2 (Cl radical attacks O3)
3. ClO' + O' --> O2 + Cl'
113
state the mechanisms by which O3 is broken down by NOx
formation of nitrogen oxide;
- N2O + O' --> 2NO
- NO2 + O' --> NO + O2
breakdown of O3:
- 'No + O3 --> NO2 + O2
- NO2 + O3 --> 2O2 + NO
114
London forces occur between....
all molecules and atoms of noble gases
115
Dipole-dipole interactions occur between...
polar molecules
116
hydrogen bonding occurs when...
a hydrogen atom is bonded to Nitrogen, Fluorine or Oxygen atoms with a lone pair of electrons
117
Explain how London (dispersion forces) arise
London forces are caused by temporary dipoles which arise in atoms due to uneven distribution of electrons
118
Explain the trend in boiling points of Group 7 elements
- bpt increases as you go down the group
- more electrons in atoms; bigger dipole
- London forces stronger
- take more energy to overcome
119
As chain length of alkanes increases, boiling point ?
Why?
increases- more points of contact with each adjacent molecule, increasing London forces
120
Explain how permanent dipole-dipole interactions arise
molecules with permanent dipole:
- e- spend more time with atom that has higher electronegativity, giving it a negative dipole which attracts the positive dipole on the nearby molecule
121
explain how hydrogen bonding arises
1. the polarised H atom has its only electrons withdrawn, exposing the proton in it's nucleus
2. this high positive charge density attracts the lone pairs of electrons on F/N/O atoms of neighbouring molecules.
122
what is the rule with drawing hydrogen bonding?
the covalent bond of the hydrogen's molecule and the hydrogen bond must be parallel.
123
state the relative strengths of the 3 intermolecular forces
London (dispersion) forces< dipole-dipole forces < hydrogen bonds
124
why would ethanol (CH3CH2OH) have a higher melting point than methoxymethane (CH3OCH3)?
ethanol has hydrogen bonding between -O-H bonds, methoxymethane only has London forces.
125
why would pentan-1-ol have a higher boiling point than pentan-3-ol?
- hydrogen bonds in pentan-3-ol are weaker as longer chains have more points of contact
126
when does a hybrid orbital occur?
a hybrid orbital results from the mixing of different types of atomic orbitals on the same atom
127
define hybridisation
intermixing orbitals of slightly different energies so as to redistribute their energies and form new sets of orbitals of equivalent energy and shape.
128
the hybridized orbitals are always equivalent in ——and ——
energy and shape
129
the number of hybridized orbitals formed is equal to —————
the number of orbitals that get hybridized
130
state the 3 types of hybridisation
sp, sp2, sp3
131
describe the formation of sp3 hybrid orbitals in methane
- four valence orbitals of the carbon (one 2s and three 2p orbitals) combine
- form four equivalent hybrid orbitals
- these overlap with the 1s orbitals of hydrogen to form bonds
132
explain why sp3 orbitals result in a tetrahedral arrangement
- each sp3-hybridized orbital bears an electron, and electrons repel each other
- to minimize the repulsion between electrons, the four sp3-hybridized orbitals arrange themselves so that they are as far away as possible from each other
- resulting in the tetrahedral arrangement predicted by VSPER (109.5')
133
describe the formation of sp2 hybrid orbitals in ethene
- three valence orbitals of the carbon (one 2s and two 2p orbitals) combine
- form three equivalent hybrid orbitals
- the remaining electron goes into the unhybridized 2p orbital (different energy=2nd bond)
134
describe the formation of sp hybrid orbitals in ethyne
- two valence orbitals of the carbon (one 2s and one 2p) combine
- form two equivalent hybrid orbitals
- remaining 2 electrons go into the unhybridized 2p orbitals (different energy=2nd and 3rd bond)
135
explain why sp2 orbitals result in a trigonal planar arrangement
- each sp2-hybridized orbital bears an electron, and electrons repel each other
- to minimize the repulsion between electrons, the three sp2-hybridized orbitals arrange themselves so that they are as far away as possible from each other
- resulting in the trigonal planar arrangement predicted by VSPER (120')
136
explain why sp orbitals result in a linear arrangement
- each sp-hybridized orbital bears an electron, and electrons repel each other
- to minimize the repulsion between electrons, the three sp-hybridized orbitals arrange themselves so that they are as far away as possible from each other
- resulting in the linear arrangement predicted by VSPER (180')
137
draw ammonium and ammonia
refer elsewhere
138
draw SO4 2- with lines
refer elsewhere
139
draw dot and cross for OH-
refer elsewhere
140
draw lewis and line for NO3-
refer elsewhere
141
draw CO3 2- with lines
refer elsewhere
142
draw PO4 3- with lines
refer elsewhere
143
draw HCO3- with lines
refer elsewhere
144
145
macroscopic level of gases
properties
146
molecular level of gases
- structure
- dynamics (motion)
- intermolecular forces
147
equation and definition for pressure
P=F/A
= force exerted per area
148
SI units for pressure
Pa = N/m^2
kPa=10^3 Pa
149
1atm =
101.3 kPa, 760mmHg (Torr)
150
1 bar =
100.0 kPa
151
equation for pressure with density, g, and h
p x h x g, where g = 9.8m/s^2
152
what does the manometer lead to?
an equalisation of pressure:
P(atm) = P(Hg) = p(Hg)gh(Hg)
153
if the Hg in a manometer was replaced with water, would the height of the water be greater or less than Hg?
greater, due to water's lower density (P constant, p goes down, g constant, so h goes up)
154
describe kinetic molecular theory - the ideal gas
1. gases made of tiny particles moving completely randomly
2. total volume of particles very small compared to size of container
3. particles do not interact with each other
4. particle collisions are elastic (no energy lost)
5. kinetic energy (KE) increases with temperature
155
describe (5) KE increases with temperature
for a large collection of molecules:
- at a given temperature, all gases have the same distribution of kinetic energy
- each molecule: KE = 1/2mv^2
156
what is the effect of temperature on average kinetic energy?
the average kinetic energy increases with temperature
157
state 2 equations relating kinetic energy with temperature
KE(avg) = 3RT/2
Temperature is in Kelvin
KE is in J/mol
KE (avg) = 3RT/2Na
for a single gas molecule
158
units for R
J/molK
159
3RT/2Na =
1/2mu^2
160
state the 2 equations relating u, R, T, m, Na, and M
161
how is kinetic energy dependent on the mass of the particles?
at the same temperature, more massive molecules move slower
162
what does pressure result from?
gas particles colliding with container walls
163
Boyle's law
at constant T and fixed n, volume is inversely proportional to pressure, or V=k/P
164
low Pext = Pgas
-> high V
165
high Pext = Pgas
-> low V
166
Charles Law
at same P and fixed n, volume is proportional to temperature, or V= kT
167
Avogadro's law
V=kn
168
what equation describes an ideal gas?
PV = nRT
P - atm
V - L
T - K
169
P1V1/n1T1 =
P2V2/n2T2
170
mole fraction of gas A
Xa = na/ntotal = Pa/Ptotal
171
law of partial pressures
Pa + Pb + Pc +.... = PTotal
172
define effusion
escape of a pas through a hole into a vacuum
173
define diffusion
movement of one gas through another
