Exam 3 Flashcards
(172 cards)
1
Q
Our evidence that atoms exist
A
Atomic force Microscopy and scanning tunneling microscopy
2
Q
Do isolated atoms/molecules exist in a state (solid, liquid, or gas)
A
no
3
Q
Do Isolated atoms or molecules have melting points or boiling points
A
no
4
Q
Emergent Properties
A
properties that belong to the collection but not the individual components
5
Q
how do atoms have emergent properties
A
these properties occur when atoms interact to form larger collections
6
Q
ex of emergent properties
A
states of matter, boiling points, and melting points
7
Q
If a solid reacts with a gas, what state of matter will the product be?
A
we can not tell from the information given because of emergent properties that could arise based on how the atoms themselves act as a collective
8
Q
Why does table salt (edible) result from poisonous Cl2 gas and 2Na that reacts violently with water?
A
The existence of emergent properties and the inability to assume the properties of a product solely based on the individual properties of the atoms that create it.
9
Q
What do the properties of a substance depend on
A
the bonding/interactions within that substance
10
Q
atomic interactions
A
LDFs and Covalent bonds
11
Q
what is the difference in strength between LDFs and covalent bonds?
A
covalent bonds are much stronger *remember the potential energy related to distance models
12
Q
how are LDFs and covalent bonds similar
A
they have similar causes; electrostatic attraction between nucleus of one atom and electrons of another
13
Q
Why are covalent bonds and LDFs different in strength
A
the magnitude of the attraction and how electrons are arranged in new species formed from interaction
14
Q
Covalent Bonds
A
- valence electrons from one atom become attracted to the other nucleus
- each nucleus is attracting both electrons (Tug - of - war, not sharing)
- Kinetic energy of electrons is decreasing by bonding
- bond length is the most stable distance between the atoms (ie, lowest potential energy)
- when bonds form, energy is released to the surroundings
15
Q
Bond theories
A
molecular orbital theory, and valence bond theory
16
Q
Wave properties of electrons
A
remember electrons are waves and can therefore combine constructively and destructively
17
Q
Molecular Orbital (MO) Theory
A
n atomic orbitals combine to give n molecular orbitals
18
Q
constructive combination of atomic orbitals
A
form bonding molecular orbitals
19
Q
Bonding MOs
A
- lower in energy than atomic orbitals
- electrons in bonding orbitals make a molecule more stable
20
Q
destructive combination of atomic orbitals
A
anti-bonding molecular orbitals
21
Q
anti-bonding molecular orbitals
A
- higher in energy than atomic orbitals
- electrons in antibonding orbitals make a molecule less stable
22
Q
Molecular Orbital Theory
A
- Atomic orbitals can combine constructively to form a molecular orbital of lower energy than the atomic orbital
- Atomic Orbitals can combine destructively to form a molecular orbital of higher energy than the atomic orbital
23
Q
Positioning of Bonding vs. Antibonding orbitals
A
Both bonding and antibonding orbitals are in the same place in space just at different energies
24
Q
Why do two hydrogen atoms form a bond?
A
- fill molecular orbital from the “bottom-up”, following Pauli Exclusion Principle and Hund’s Rule
- H only has one electron in the 1s atomic orbital therefore when two of these atomic orbitals combine into a molecular orbital, they fill the bottom first because it takes less energy, leaving the antibonding orbital empty
25
Why do 2 helium atoms not form a bond
The molecular bonding orders cancel
26
A bond will form if... what
- there are two electrons in the bonding molecular orbital (these two electrons make up one bond)
- However! if there are also two electrons in an antibonding orbital, this cancels out the stabilization from two electrons in bonding orbital and, therefore, no bond forms
27
bond order equation
(# electrons in bonding orbitals - # electrons in antibonding orbitals) / 2
28
what does a bond order of zero mean
no bond
29
Bonding in Hydrogen
only the bonding molecular orbital is occupied, leading to a net stabilization of the interaction
- to break the bond, enough energy has to be added to raise an electron to the antibonding orbital
30
how can energy be transferred to this system of molecular orbitals?
- transfer of thermal energy into the system through temperature change
- collision of molecules with enough energy to excite the electron and break the bonds (chemical reaction)
31
Bonding in Helium
In He2 both antibonding and bonding molecular orbitals are occupied, making the sum of their energies the same, resulting in no net stabilization of the interaction and therefore no formation of He2
32
Types of bonds
sigma, sigma star (sigma*), Pi, Pi star (pi*)
33
sigma and sigma star orbitals overlap
end-on-end overlap, involving a node between two s and p orbitals for sigma star, and the bond creating one molecular orbital from 2 s atomic orbitals, and bond (creating one MO) of the same domain between two p orbitals, with the unused domain connected by nodes on either side for sigma
34
Pi and Pi star orbitals overlap
sandwiched together, creating a bond between each domain in Pi, and creating a node between two inversely domain p atomic orbitals for Pi star
35
when two O atoms interact, how many molecular orbitals are formed?
3
one for each n ( 1s, 2s, 2p)
36
What is the bond order between two O atoms in O2?
( 10 - 6 ) / 2 = 2
37
What is the significance of a Molecular Orbital having unpaired electrons ?
it makes the molecule paramagnetic
38
what does paramagnetism exemplify?
how molecular level interactions determine the macroscopic properties of a substance
39
Properties of metals
- shiny
- conduct electricity and heat
- malleable and ductile
- most are grey, but some are colored (gold, copper) or colorless (silver)
40
Bonding in metals
- atomic orbitals (lots of them) combine with each other to form molecular orbitals (an equal number)
- as the number of MOs increases, the energy gap between them decreases
- as the number of MOs increases, they form bands of MOs - of almost continuous energy
- electrons can move freely between MOs
- electrons can move freely over whole system
41
why is a property of metals being shiny
- absorption of a photon will promote an electron to a higher energy level
- the electron immediately falls back down emitting a photon and the metal shines
- the metal interacts with light of many wavelengths, so the metal appears silvery
42
why is malleable and ductile a property of metals
because the atoms can move with respect to one another, metals are malleable and ductile - create a sheet of atoms
43
why can metals conduct electricity
because electrons can move freely, metals conduct electricity
44
Band theory of metals
in metals, the valence band (bonding MOs0 overlaps with the conduction band *anti-bonding MOs), meaning electrons can move freely between the two MOs which is why metals conduct electricity
45
Band theory of semi-conductors
there is a small gap between bands meaning it is relatively easy for the electrons to jump back and forth, but not as easy (takes more energy) as the metal overlap
- often used for solar cells and in computer equipment
ex Ge Si
46
Band theory of insulators
ex diamonds
there isa. large energy gap between the valance and conduction band therefore not conducting electricity
47
structure - property relationship
interactions at the atomic/molecular level <-> properties of substance
type of bonding <-> melting point and boiling point
48
phase changes
melting boiling freezing condensing
49
does melting and boiling absorb or release energy
absorb because bonds are being overcome/ broken
50
where does the energy to melt/ boil come from?
the surroundings of they system
51
what attractions are being overcome during melting/ boiling?
between particles
52
does freezing and condensing absorb or release energy
release
53
where does energy go when molecules freeze/ condince
the surroundings
54
where are the attractions that are being formed during freezing/condensing?
between particles
55
what happens to interactions when a substance boils
some interactions between particles must be overcome so they can move relative to each other
56
what happens to interactions when a substance boils
all of the interactions between the particles must be overcome
57
what does the magnitude of the melting and boiling points say about the interactions
provides an estimate of the strength of the attractive interactions between particles
58
difference between covalent bonds and LDFs
covalent bonds:
* Strong (require a lot of
energy to break)
* Caused by attraction of
electrons from one atom to
nucleus of another atom
* Hard to predict bond
strength
* Present only when atomic
orbitals interact
constructively
* Present within molecules or
networks
LDFs:* Relatively weak
* Caused by fluctuating charge
distribution that results in
electrons from one atom
being attracted to nucleus of
another atom
* Increase (predictably) with
size of electron cloud
* Present between all
molecular species
59
discreet materials
substances where atoms are bonded together into individual molecules
60
continuous molecules
substances where atoms are bonded together in a large, interconnected network
ie metallic bonds
61
what does the temperature of the phase change tell you about the strength of interactions that must be overcome between molecules and atoms
H2 and He have low mp because they have weak LDFs between discrete molecules
Li has a high mp because it has strong metallic interactions between metal atoms
62
what kinds of interactions must be overcome between particles in order to melt carbon
LDFs and Covalent
63
Allotropes of Carbon
Diamond, Graphite, Lonsdaleite, Buckminsterfullerene, C540, C70, Amorphous (soot), carbon (nanotube)
- each has a different structure of carbon atoms
64
comparison of Diamond and Graphite
Diamond:
- High MP
- Hard
- Brittle (breaks along planes)
- Translucent (lets light through)
- does not conduct electricity
- stacked cubes molecular structure
Graphite:
- High MP
- Soft
- Slippery
- Grey, shiny
- Conducts electricity
- sheets of hexagons layers on top of one another
65
what would have to happen to "melt" diamond or graphite
break the covalent bonds between carbons
66
why do metals melt while diamonds and graphite do not?
Takes a LOT of energy to break covalent bonds
67
How can diamond and graphite, being made from carbon atoms only, have such different properties?
valence Bond theory
68
Valence Bond theory
assumption that atomic orbitals overlap to form bonds
- singly occupied orbitals overlap and form an energetically-favorable bond (often shown as sticks) ie H + H means the singly occupied orbital of 1s overlaps with that of another H
69
Electrons in MO theory vs VBT
Molecular Orbital Theory:
- describes where all the electrons in the molecule are most likely to be frond over the entire molecule
- Electrons are considered delocalized
Valence Bond Theory:
- Describes where the electrons in a bond between two atoms are most likely to be found
- electrons are considered localized between the two atoms
70
Models of Bonding breakdowns
Molecular Orbital Theory:
* Atomic orbitals combine to
form an equal number of
molecular orbitals.
* Each orbital can contain up
to two electrons.
* Electrons in bonding
orbitals stabilize the system.
* Electrons in anti-bonding
orbitals make it less stable.
* Electrons are delocalized.Valence Bond Theory:
* Atomic orbitals overlap to
form a bond.
* Each bond is made up of
two electrons.
* The greater the overlap, the
stronger the bond.
* Electrons are localized in
the bond.
71
Why is it weird that carbon can form 4 identical bonds in a diamond according to valence bond theory
VBT tells us that for each bond, carbon needs a singly occupied orbital pointing in the direction of the bond but carbon is 1s2 2s2 2p2- meaning each orbital has more than one electron
72
Bonding in diamonds
- each carbon atom forms 4 bonds to 4 identical carbon atoms
- c-c-c bond angle is ~109
73
Tetrahedral
molecular geometry where the bonds arrange themselves towards the corners of a 4-sided figure (a tetrahedron)
74
Hybridization
- valence orbitals mix together to form hybridized orbitals
- hybridized orbitals overlap to form bonds
75
Hybridization for Diamond
- 4 unhybridized orbitals 2s2 2px 2py
- to 4 hybridized orbitals sp^3
(one s orbital plus 3 p orbitals from the given valence orbital shells so that each electron has its own shell)
76
SP^3 hybridized orbitals
s, px, py, and pz combine to form FOUR sp^3 hybrids - no orbitals remain
77
sp^3 overlapping orbitals yield what bond
strong sigma bonds
78
if diamonds were to melt, what does that mean for the bond
these sigma bonds would have to be broken to allow the atoms to move relative to one another, which would require a lot of energy
( diamonds decompose rather than "melt")
79
Explanation of High MP/BP of diamonds
3D network of strong covalent bonds
80
Explantation of diamonds not conducting electricity
Electrons are localized in bonds between atoms - not free to roam. There is a large "band gap" between the bonding and antibonding orbitals
81
Explanation of diamonds being translucent
light passes through or is reflected. To absorb light an electron must be promoted to a higher energy level. There is a large "band gap" between bonding and antibonding orbitals
82
In one sheet of graphite, how many carbon atoms is each carbon bonded to?
3
83
How many orbitals should hybridize to form the bonds needed for each carbon to bond to 3 other carbons
3
84
Hybridization for graphite
2s2 2px 2py have electrons -> when hybridized
3 sp^2 and one 2p unhybridized
85
sp^2 hybridization
s, px, py combine
creating three sp^2 hybrids and leaving a remaining pz "unused"
- these "unused" p orbitals are what connect sheet to sheet
86
Sigma bonds
end-to-end overlap of atomic orbitals
87
Pi bonds
side-to-side overlap of atomic orbitals
88
Explanation of graphite bonding
the VBT says that the carbon sp^2 hybridized orbitals overlap to form a "localized" sigma bond framework
MO theory states the unhybridized p orbitals overlap to form Pi bonds. Since the overlap occurs across the entire sheet of atoms, "delocalized" pi molecular orbitals form over the sheet, allowing electrons to flow more freely
89
Explanation of why Graphite conducts electricity
electrons can move freely over the entire sheet within its delocalized pi MOs
90
Explanation of why graphite is shiny
it can absorb and emit photons of many wavelengths (just like metals)
91
Explanation of why graphite is slippery
sheets can slide over each other - only "held together" by LDFs
92
delocalized bonds
electrons can be placed on one atom, but can also be "shared" between two atoms
93
localized bonds
electrons only belong to only one bond
- like the hybridization
94
what can carbon form bonds to?
C, H, O, N, S, P
95
carbon boning continuous materials vs discrete
continuous = diamond, graphite
discrete = methane, ethane
96
Methane CH4
- contains 4 C-H bonds
-- each bond is the same length
-- bonds are equidistant from each other
97
what type of hybridized orbitals does C use in Methane to form C-H bonds
sp^3 hybridized orbitals
98
Bonding of C and H in CH4
one C sp^3 orbital bonds with the 1s orbital of H
99
what do the colors on an electrostatic potential map mean
red = more e-density and partial negativity
bue = less e-density and partial positivity
100
What type of orbitals does C have in Ethane to from C-H bonds`
sp^3 orbitals
101
Can sigma bonds freely rotate?
yes, the sigma bond does not prevent rotation
102
Constitutional (structural) isomers definition
have the same formula but differ in how the atoms are connected ( movement of electrons)
103
Do constitutional isomers have different properties?
yes
104
Alkenes
- contain a C-C double bond
- each C is sp^2 hybridized
105
what are double bonds made up of
- one sigma bonds (Csp^2 - Csp^2)
- one pi bond (Cp - Cp)
106
Sigma bonds in Ethylene
between H 1s and C sp^2
and Csp^2 and Csp^2
107
Pi bonds in ethylene
one Pi bond between Cpz and Cpz
108
Can Pi bonds freely rotate?
no, because the pi bond would break because the electrons are sticking to each other rather than apart of the same cloud, meaning a twisting of the clouds would remove the electrons from one another- requiring a LOT of energy
109
Drawing Lewis structures steps
1.) Find the total number of valence electrons. Add of subtract electrons for each negative or positive charge, respectively
2.) decide the central atom (least electronegative, except for H) and draw bonds
3.) complete octet rule (8 e- around atom) around the terminal atom
4.) Place leftover electrons on central atom(s). If the central atom doesn't have an octet, create multiple bonds
5.) If polyatomic atoms: add brackets and charge
110
Bonding in Ethylene (C2H4)
- 4 C-H sigma bonds, each formed from overlap of a carbon sp^2 - hybridized orbital and a hydrogen 1s orbital
- 1 C-C sigma bond, formed from overlap of two carbon sp^2-hybridized orbitals
- 1 C-C pi bond, formed from overlap of two carbon unhybridized p orbitals
111
sp Hybridization
1s2, 2px, 2py have electrons, 2pz has none
to
2 sp hybridized orbitals having one electron each and leftover 2py, 2pz orbitals having one electron each
112
Bonding in Ethyne (C2H2)
- 2 C-H sigma bonds, each formed from overlap of a carbon sp-hybridized orbital and a hydrogen 1s orbital
- 1 C-C sigma bond, formed from overlap of two carbon sp-hybridized orbitals
- 2 C-C pi bonds, formed from overlap of carbon unhybridized p orbitals
113
what are triple bonds made of
one sigma bond ie Csp - Csp and two pi bonds ie Cp-Cp
114
Valence electrons and bonding examples
C has 4 valence e-, allowing it to form 4 bonds ( 2 e- per bond) to complete octet rule
N has 5 valence electrons, allowing it to form 3 bonds with a leftover lone pair of electrons to complete octet rule
115
Number of bonds formed by H, C, N, O, X (F, Cl, Br, I)
H - 1 valence e- - typically 1 bond formed
C - 4 valence e- - typically 4 bonds formed
N - 5 valence e- - typically 3 bonds formed
O - 6 valence e- - typically 2 bonds formed
X (F, Cl, Br, I) - 7 valence e- - typically 1 bond formed
116
Formal charge def
a bookkeeping tool to keep track of electrons in Lewis structures
117
Formal Charge equation
= number of valence electrons in neutral atom - [ number of unshared e- + number of bond(s) ]
118
What is the formal charge of oxygen
FC = 6 - [ 6 + 1 ] = -1
119
common exceptions to the octet rule
in some compounds, atoms do not have 8 electrons around them
ex. BH3, compounds with an odd number of valence electrons (i.e., radicals) (commonly very reactive) - ie O-H, Compounds with expanded octets (not covered)
120
Electronegativity definition
ability of an element to attract electrons to itself in a bond
121
what does electronegativity depend on
the effective nuclear charge and size of orbitals
122
effective nuclear charge deff and equation
nuclear charge "felt" by valence e- after core e- shield nuclear charge
Zeff = Z - S
Zeff = # protons - # core electrons
123
what element has the highest Zeff that can form bonds?
F bc Ne doesn't form bonds (no noble gasses do)
124
Trend of electronegativity in periodic table
increases Zeff across table (left to right)
increasing Zeff up a column (bottom up) because of increasing radii down the table
- smaller radii mean a greater attraction, and therefore a higher Zeff and electronegative force
125
Electronegativity overview
* Ability of an element to attract electrons to itself in a bond.
* Depends on effective nuclear charge and size of orbitals.
* Electronegativity increases across the periodic table.
* Electronegativity decreases down the periodic table.
126
Bond order deffinition
of electron-pair bonds connecting two atoms
127
Bond energy deffinition
energy required to break a chemical bond
128
Bond length deffinition
distance between two nuclei
129
comparing bonds (single, double, triple)
single bond - Bond order of 1 - weakest bond energy - longest bond length
double bond - bond order of 2 - mid bond energy - mid bond length
triple bond - bond order of 3 - strongest bond energy - shortest bond length
130
Why can't Ozone be drawn as a lewis structure
its bonds are the same length, even though one is a double bond and the other is a single bond
131
Resonance Structures deff
are used when a single lewis structure does not explain the data (like bond lengths)
132
Resonance hybrid deff
weighted average of all resonance structures
133
why are resonance structures important
important for communicating electrons delocalized in pi molecular orbitals
134
Benzene carbon hybridization
sp^2
135
Benzene resonance structures
there are three pi bonds and three sigma bonds, but they can go in different places, and they have the same length, leading to the use of resonance structures
136
bond order of benzene resonance structure and what is equivalent
- bond order = 1.5
- all C-C bond lengths are equivalent
- All C-C-C bond angles are equivalent
- Benzene has 6 pi electrons delocalized in its conjugated pi system
137
Nitrate ion bond order (NO3-)
(one bond x 1/3 significance) + (two bonds x 1/3 significance) + (one bond x 1/3 significance) = 4/3
create the same atomic structure and rearrange the electrons- focus on the bonds of one atomic pairing, and find the bo
138
Nitrite Ion orbital overlap
the 2p orbitals on all three O atoms overlap with the 2pz orbital on the central N atom
139
localization of Nitrate ion
All six pi electrons are delocalized
140
what happens to lone pairs on a nitrate Ion
Lone pairs on atoms adjacent to a pi system will participate in pi conjugation (communicated vis resonance structures)
141
Pi conjugation
the overlap of adjacent p orbitals across a chain of atoms, allowing for the delocalization of atoms
142
Do resonance structures contribute equally to the structure of a molecule?
no
143
What is the largest resonance contributor
the resonance structure that best explains the data
144
What does the largest resonance contributor tend to do
- minimize formal charges on atoms
- place negative formal charges on more electronegative atoms
145
molecular shape
Molecules adopt 3-D shapes
146
VBT and shape
sp^3 = tetrahedral = ~109.5 degrees
sp^2 = trigonal planar = ~120 degrees
sp = linear = ~180 degrees
147
VSEPR
Valence shell electron-pair repulsion theory
148
General idea of VSEPR modle
shape of AXn molecule is determined by electrostatic repulsion between valence-shell electron pairs of central atom A (central p-block element)
- molecules adopt geometries in which their valence electron pairs are as far apart from each other as possible
149
steps for a VSEPR model
1.) draw lewis structure
2.) Identify the electron domains
3.) determine the electron pair geometry
4.) determine the molecular shape
150
Electron pair geometry and relation to bonds
geometry of the regions of electron density around central atom
- each single, double, triple bond counts as 1 electron domain
- each lone pair of electrons counts as 1 electron domain
151
Molecular geomatry
arrangement of atoms in space around a certain atom
- does NOT include lone pairs of electrons
152
Resonance structure recap
Resonance structures are used when a single Lewis structure does
not adequately communicate the structure of a molecule.
– Resonance structures are different representations, not different
molecules!
– When drawing resonance structures, keep atoms in the same
place but change where the electrons are shown (bonds & lone
pairs)
– Resonance structures are especially important for
communicating electrons delocalized in molecular orbitals →
look for bonds or lone pairs next to bonds* Resonance structures do not necessarily contribute equally to
the structure of a molecule.
– The largest resonance contributor is the one that best
explains the data.
– The largest resonance contributor tends to (but not always!):
* Minimize formal charges on atoms.
* Place negative formal charges on more electronegative
atoms.
153
Polar Bonds definition
When two atoms of different electronegativities bond, the electrons will be more strongly attracted to the more electronegative atom
- resulting in a bond dipole, which we say is a polar
154
What are bond dipoles in molecular polarity
vector quantities
- they have a magnitude and direction
- must take both into account
155
What is needed to be added to bond dipoles to determine polarity
the molecular shape is needed to add together bond dipoles to determine the molecular dipole
156
A molecule will be polar if...
1.) it contains a polar bond
2.) its bond dipoles do not cancel, causing a net dipole moment
a.) there is asymmetry in the atoms or bonds
b.) it is excited in an electric field
157
Steps to determining if a molecule is polar
1. Draw the Lewis structure.
2. Count the electron centers around the central atom to
determine the electron pair geometry.
3. Determine the molecular geometry (and draw the molecule
in that shape).
4. Determine the bond polarities.
5. Add up the bond polarities (take into account direction).
– If there is an overall dipole, the molecule is polar.
– If not, the molecule is nonpolar.
158
What interactions are overcome when Noble gasses are boiled
LDFs
159
Why does Xe have the highest BP of the noble gasses
largest e- cloud -> most polarized -> largest temporary dipoles -> strongest LDFs -> most energy -> highest BP
160
Types of Intermolecular Forces
London Dispersion Forces (all molecules), Dipole-dipole interactions (polar molecules), Hydrogen bonding interactions (only very specific molecules)
161
London Dispersion forces
- caused by temporary fluctuating dipoles
- Depends on size, surface area, and shape of molecule
- Present in ALL substances in solid and liquid states (between discrete molecules or atoms)
- Only intermolecular force present in non-polar molecules
-- eg CH4, CO2, hydrocarbons, Br2
162
Dipole-dipole interactions
- caused by permanent dipoles
- present in polar substances (along with LDFs)
-- eg. HCl, CH3F, CH2O, CH3OCH3
163
Can two isomers have different boiling points because of bonding?
yes, if the bond dipoles cancel in one and are present in another the calncled one only had LDFs whereas the other had dipole-dipole as well
164
Effect of the polarity of the molecule
polar molecules (LDF and Dipole-dipole) have stronger intermolecular interactions than similarly sized non-polar (LDFs only) molecules
- therefore, polar molecules are held together by stronger forces, which require more energy to overcome, resulting in melting and boiling points
165
Hydrogen Bonding interactions definition
Donor/acceptor interaction between the lone electron pair on a heteroatom and a nearly bare proton
- despite neame IMFs are not bonds!!!
166
Presence of Hydrogen bonding interactions
present between two specific types of molecules:
- must contain H covalently bonded to N, O, or F
- Must have N, O, Fatom with lone pair
eg. H2O, CH3OH, CH3CH2OH, HF, NH3
167
Hydrogen bonding interactions overall
extremely strong and specific dipole-dipole interactions
- compounds with H-bonding also have dipole-dipole and LDFs
168
Why is there such a difference between boiling point and substances with similar dipole moment and with the same types of IMFs
IMFs can not be the same strength
- LDFs are similar due to similar size electron clouds
- dipole-dipole interactions are similar because all are highly polar
- but the number of H-bonds changes due to the number of available H atoms with respect to the number of available lone pairs
168
What are strands of DNA held together with
H bonding interactions
169
Unusual properties of H-bonding present in water
- high MP and BP
- Low vapor pressure
- density of Ice < liquid water
170
Why is the density of ice < liquid water
Ice has a hexagonal crystal form because of the 4 hydrogen bonding interactions around each water molecule
- as energy is added the IMFs in ice are overcome causing the rigid cage-like structure to collapse and liquid water to form
171
Structure-property relationship
Molecular formula -> Lewis structure -> Molecular geometry -> Bond Polarities -> Molecular Polarity -> Types and relative strength of IMFs -> Macroscopic properties
