CONDENSED Flashcards
1
Q
is energy released or absorbed in sublimation
A
absorbd
2
Q
is energy released or absorbed in deposition
A
released
3
Q
density =
A
mass / volume
4
Q
what is the resultant volume
A
the final volume
5
Q
the number of particels in a mole of a substance is numerically equal to the
A
Avogadro’s constant, 6.02 × 10^23 mol-1
6
Q
what is molar mass numerically equal to
A
relative atomic mass
7
Q
percentage composition formula
A
mass of element in compound / molar mass of compound x 100%
8
Q
molar volume of a gas
A
This states that one mole of a gas at STP occupies a volume of 22.7 dm3 (22700 cm3 or 0.0227 m3).
9
Q
relationship between the amount of a gas (in mol) and its volume
A
amount in mol = vol (dm^3) / molar vol (22.7dm^3)
10
Q
an ideal gas is
A
a gas that exhibigts the five postulates of the kinetic molecular theory, as well as obeying gas laws.
11
Q
what are the gas laws
A
Boyle’s law, Charles law, Gay-Lussac’s Law
12
Q
what temperature scale must you use for gas laws
A
kelvin
13
Q
Boyles Law
A
at constant temperature, the pressure and volume of a fixed mass of an ideal gas are inversely proportional to each other.
14
Q
Charles Law
A
at constant pressure, the volume of a fixed mass of an ideal gas, is directly proportional to its absolute temp in kelvin
15
Q
Gay-Lussac’s law
A
at constant volume the pressure of a fixed mass of an ideal gas is directly proportional to its absolute temperature (in kelvin).
16
Q
Combined gas law:
A
PV/T = k
P1V1/T1=P2V2/T2
17
Q
ideal gas equation
A
PV = nRT
18
Q
cm cubed to m cubed
A
divide by 10^6
19
Q
convert dm cubed to m cubed
A
divide by 10^3
20
Q
with real gases, which of the assumptions made about ideal gases no longer apply under certain conditions.
A
At very high pressure the gas particles are closer together. Under these conditions, the actual volume of the particles becomes significant.
At low temperatures, the particles move less rapidly (have lower average kinetic energy). This means that there is a greater opportunity for intermolecular forces between the particles to have an effect.
21
Q
when is the greatest deviation from ideal behaviour seen
A
when the gas is subjected to a low temperature and a high pressure
22
Q
why is ideal behaviour deviated from at high pressure
A
spaces between particles lessen so intermolecular forces operate and particles become attracted to eachother. additionally, the non zero volume of the particles becomes significant in that the volume of a real gas at high pressure is higher than expected and PV/RT > 1.
23
Q
which gases show the most ideal behaviour
A
low molar mass and weakest intermolecular forces
24
Q
what is a standard solution
A
a solution with an accurately known concentration
25
what is a primary standard solution
it has:
High purity (99.9 %).
High molar mass.
Low reactivity.
Does not change composition in contact with air.
26
what is a secondary standard solution
a solution that has been standardised against a primary standard solution.
27
describe a titration
A titration is a method of volumetric analysis in which the concentration of a solution can be determined. In a titration, a burette is filled with a standard solution of known concentration (the titrant). A carefully measured volume of the solution with the unknown concentration (the analyte) is placed in a conical flask below the burette (Figure 2). An indicator is used to determine the end-point of the titration. In acid–base titrations, for example, a suitable indicator is added to the solution in the conical flask and the volume of titrant required to reach the end-point of the titration is recorded. Acid–base titrations are covered in more detail in section 8.2.2.
28
isotopes have the same chemical reactions because
they have the same number of electrons
29
isotopes have differnt physical properties becaues
they have more neutrons, so more mass
30
what is the mass spectrometer used for
to determine the realtive atomic masses of elements. it can also be used to determine the structure of organic compounds
31
how does mass spec work
1. sample vapourised
2. bombarded with high energy electrons
3. cations produced
4. cations accelerated in a electric field
5. cations reach the detector where they produce a mass spec.
the cations are deflected in a magnetic field depending on their mass to charge ration (m/z). ions with a higher mass to charge ratio are deflected less in the mag
field than ion with alower m/z.
32
what is the shape of an s orbital
sphere
33
the main energy levels are split into...
sub levels which are assigned a number and the letter s,p,d or f
34
what is the shape of a p orbital
dumbbell
35
what does the pauli exclusion prinicple state
two electrons cannot have the same quantum number. two electrons can only occupy the same atomic orbital if they have opposite spins.
36
what is heisenberg's uncertainty principle
it is not possible to know, at the same time, the exact position and momentum of an electron
37
n=1
s 2
38
n=2
s 2
p 6
39
n=3
s 2
p 6
d 10
f 14
40
aufbau principle
electrons fill atomic orbitals of lowest energy first
41
which sub level has the lowest energy
the 1s sub level
42
are s orbitals or p lower energy
s
43
degenerate orbitals
equal energy
44
there is an overlap between the 3d and 4s sub levels
this means the 4s sub level is of lower energy and fills before the 3d sub level
45
chromium electron configuration
1s2 2s2 2p6 3s2 3p6 4s1 3d5 or
1s2 2s2 2p6 3s2 3p6 3d5 4s1
46
copper electron configuration
1s2 2s2 2p6 3s2 3p6 4s1 3d10 or 1s2 2s2 2p6 3s2 3p6 3d10 4s1
47
hunds rule
electrons fill orbitals in the saem sub level singly before pairing up.
48
as freq increases
wavelength decreases
49
energy increases alongside
frequency
50
higher energy =
higher frequency = shorter wavelength.
51
lower energy =
lower frequency = longer wavelength
52
differences between spectra
53
when electrons are excited
they jump to higher energy levels
54
rel between energy and frequency formula
E = hv
55
Electron transitions to the n = 1 energy level
UV
56
Electron transitions to the n = 2 energy level
visible light
57
Electron transitions to the n = 3 energy level
IR
58
what is the highest energy end of each series of spectral lines known as
convergence limit
59
Plancks constant, h
6.6.3 x 10^-34 J/s
60
ground state
The ionisation of a hydrogen atom in its ground state corresponds to the electron transition from n = 1 to n = ∞. At this point, the electron is no longer attracted to the nucleus and the atom has been ionised.
61
what is the first ionisation energy of an element
the energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous 1+ ions.
62
why are ionisation energies always positive (endothermic)
energy must be added to overcome the electrostatic attraction between the nucleus and valence electrosn
63
what is second ionisation energy
If an additional mole of electrons is removed from one mole of gaseous 1+ ions
64
why should we expect the ionisation energies to increase progressively as we remove negatively charged electrons from increasingly positive ions
it results in a stronger electrostatic attraction between the nucleus and the remaining electrons.
65
s block
groups 1 and 2
66
d block
groups 3 to 12
67
p block
13 to 18
68
f block
bottom of the periodic table
69
what is a transition element
defined as an element that has an incomplete d sub-level in its atom or one or more of its ions.
70
properties of transition elements
They have variable oxidation states.
They form coloured compounds.
The elements or their compounds show catalytic activity.
They form complex ions in solution.
The metals and their complexes show magnetic properties.
71
why is scandium the only element that cant have a +2 oxidation state
has only one unparied electron
72
what is the reason for these variable oxidation states is
the closeness in energy of the 3d and 4s sub-levels.
73
how to find the coordination number
number of ions in contact with other ions
74
what does the mp of an ionic compound depend on
the ionic charge and the ionic radius of its component ions
75
which exist as diatomic molecules
halogens
hydrogen
oxygen
nitrogen
76
what is different about diatomic oxygen and nirtogen
atoms are bonded via double and tripke bonds
77
coordiante covalent bond
both bonding electrons come from one atom
78
what is a dimer
a larger molecule composed of two identical smaller molecules and can be linekd by coordinate covalent bodns or by hyrdrogen bodns
79
what is bodn order
the number of bonds between a pair of atomso
80
over 1.8 units
ionic
81
1.8 units plus diff in electronegativity
ionic
82
0.5-1.7 units diff in electronegativity
polar covalent
83
0.1-0.4 units difference in electronegativity
non polar or weakly polar covalent
84
0 units electronegativity
pure covalent
85
when do pure covalent bodns occur
between atoms that have no differene in electronegativity, such as those in molecular oxygen chlorine nitgoen
86
when do polar covalent bonds occur
between atoms that have a difference in electronegativity of between 0.5 and 1.7 units
87
what causes a bond dipole
unequal sharing of electrons in a covalent bond
88
what does the octet rule state
the most stable arrangement is 8 electrons
89
exceptions to the octet rule
Hydrogen is stable with only two electrons in its outer shell.
Atoms such as boron, beryllium and aluminium (in compounds) are stable with fewer than eight electrons in their outer shell.
Atoms in period three and higher, such as sulfur, can form expanded octets with up to twelve electrons in their valence shell.
90
benefit of delocalised electrons
give greater stability to a molecule or polyatomic ion
91
silicon and silicon dioxide
giant covalent structure
tetrahedral
109.5 degree bond angle
strong covalent bonds
poor conductor of electricity
92
what are the allotropes of carbon
diamond
graphite
fullerenes
93
graphite
layered structure
planar sheets of hexagonally arranged carbon atoms
The layers are held together by relatively weak London dispersion forces.
Each carbon atom has an electron which becomes delocalised across the plane. The presence of delocalised electrons explains the ability of graphite to conduct electricity along the plane of the crystal when a voltage is applied.
94
fullerene
bond angle is 120
poor conductor of electricity
rings of 5 and 6 carbon atoms
fullerene C60 is by definition a simple molecular substance even though it contains so many bonded carbon atoms
The structure is made up of carbon atoms bonded together in 20 hexagons (six-carbon rings) and 12 pentagons (five-carbon rings), known as a truncated icosahedron.
95
graphene
Its tensile strength is 1000 times greater than steel.
It behaves as a semi-metal, so it is very suitable for electronic devices.
Adding only 1% content of graphene to plastics could allow those plastics to conduct electricity.
graphene is the most chemically reactive. This is because of the reactive edges of the structure, where there are carbon atoms with unoccupied ('dangling') bonds. This reactivity could be used in important ways. For example, membranes of graphene oxide have been shown to be preferentially permeable to water, which could be useful in desalination and water purification.
96
how many atoms is each carbon directly bonded to in these allotropes
diamond
graphite
fullerene
4
3
3
97
temporary dipole
caused by changes in electron density within an atom or molecule
98
waht does the strength of london forces depend on
The ease with which the electrons in an atom or molecule form a temporary or induced dipole (their polarisability).
The surface area of the molecule.
99
when does polarisability increase
alongside molar mass of a molecule. so does melting point and strength of London dispersion forces
100
what does the strength of the metallic bnd depend on
charge ont he metal ion and ionic radius of the metal ion. they also affect density
101
how is a sigma bond formed
direct head-on (axial) overlap of atomic orbitals
The two 1s atomic orbitals overlap head-on, forming a sigma bond.
102
pi bond formation
formed by the sideways overlap of two unhybridised p orbitals
103
double bond
one sigma one pi
104
which is stronger sigma or pi and why
sigma
The extra strength of the sigma bond comes from the greater overlap of the atomic orbitals in the bond. In a pi bond, the atomic orbitals cannot overlap as much which results in a weaker bond.
105
purpose of formal charge
used ot determin ewhich lewis structure is the preferred one when there is more than one possibility
106
formal charge equation
V - 1/2B - NB
107
the sum of formal charges in a neutral molecule must equal
zero
108
CFC's
Chlorofluorocarbons (CFCs) are highly stable compounds. This stability was initially seen as an advantage but it meant that CFC molecules released into the lower atmosphere could remain intact and reach the upper atmosphere. Here, when exposed to UV radiation, compounds such as trichlorofluoromethane (CCl3F) decompose to produce chlorine free radicals (Cl*), as shown in the equation below.
109
standard conditions (Ɵ)
A pressure of 100 kPa.
A temperature of 25°C (298 K).
110
what is the enthalpy change of neutralisation
the enthalpy change when an acid and base react together to form one mole of water
111
waht is the molar enthalpy of combustion, or the standard enthalpy of combustion (ΔHƟc),
the enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions.
112
Percentage error =
((experimental value - theoretical value) ÷ theoretical value) × 100
113
You should be aware of the limitations of calculating enthalpy changes in a school laboratory. These include but are not limited to:
Heat loss to the surroundings and heat absorbed by the calorimeter
Incomplete combustion of the fuel
Assumptions made about the specific heat capacity and density of aqueous solutions.
114
what does hess law state
the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place, as long as the initial and final conditions are the same.
115
what is the standard enthalpy of formation
the enthalpy change when one mole of a compound is formed from the elements in their standard states under standard conditions.
116
Enthalpy of formation values are useful in that they indicate
he stability of compounds in relation to their elements
117
does hgaving a double bond increase or decrease reactivity and why
incease
This region of high electron density is the site of chemical reactivity within the molecule. This means that alkenes undergo addition reactions that take place across the carbon-carbon double bond.
118
enthalpy change of a reaction using combustion
ΔH⦵ = ΣΔH⦵c (reactants) − ΣΔH⦵c (products)
119
bond breaking is
endothermic (releases energy)
120
bond making is
exothermic (absorbs energy)
121
bond enthalpy aka
bond dissociation energy
122
bond enthalpy definition
It is defined as the energy required to break one mole of chemical bonds in the gaseous state
123
average bond enthalpy definitio n
when one mole of bonds are broken in the gaseous state averaged for the same bond in similar compounds
124
The enthalpy change for a reaction can be calculated using average bond enthalpy values.
ΔH = ΣE(bonds broken) − ΣE(bonds formed)
125
which has a higher wavelength oxygen or ozone
ozone
oxygen
λ < 242 nm
higher energy radiation of shorter wavelength
ozone
λ < 330 nm
lower energy radiation of longer wavelength
126
benefit of the decomposition of oxygen
The two reactions in this cycle depend on different wavelengths of UV radiation, and the effect is to remove the higher energy radiation (λ < 242 nm), so that only the longer wavelength, which is less damaging radiation, reaches the Earth's surface.
127
thes strong double covalent bond in oxygen is disrupted by
the suns high energy UV-C raditation to form atoms which are free radicals since they have an unpaired electron. Such oxygen radicals can then react with an oxygen molecule to form ozone. The bonds in ozone, being weaker, can then be broken by the less energetic UV-B radiation (of longer wavelength) to reform oxygen and an oxygen free-radical.
128
the surface of the earth is protected by these breakdown of ozone reactions from the damaging effects of
UV-B and UV-C radiation
129
The lattice enthalpy (ΔH⦵lat) is
the enthalpy change when one mole of a solid ionic compound breaks down to form gaseous ions under standard conditions.
130
The magnitude of the lattice enthalpy of an ionic compound can be thought of as
a measure of the strength of the ionic bonds between the ions.
131
The magnitude of the lattice enthalpy depends on two factors:
the charge of the ions (ionic charge)
the size of the ions (ionic radii).
132
the lattice enthalpy is directly proportional to
the product of the ionic charges and inversely proportional to the distance between the nuclei of the ions.
133
lattice enthalpy increases alongside
ionic charge increasing or ionic radii decreasing
134
lattice enthalpy
The enthalpy change that occurs when one mole of an ionic solid is broken down into its gaseous ions.
135
purpose of born harber
to calculate lattice enthalpy
136
enthalpy of formtation
the enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.
137
enthalpy of atomisation
the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.
138
Bond dissociation energy, E
the energy required to break one mole of bonds in the gaseous state.
139
Ionisation energy, ΔH⦵IE:
the energy required to remove one mole of electrons from one mole of gaseous atoms.
140
Electron affinity, ΔH⦵EA:
the energy released when one mole of electrons are added to one mole of gaseous atoms
141
steps in born harber
Enthalpy of formation, ΔH⦵f
Enthalpy of atomisation, ΔH⦵at
Bond dissociation energy, E
Ionisation energy, ΔH⦵IE
Electron affinity, ΔH⦵EA
142
why do stable ionic compounds tpicall have large negative values for enthalpy of formation
because as more energy is released, the compound becomes more stable
143
waht is enthalpy of hydration
the enthalpy change when one mole of gaseous ions dissolves in water to form a solution of infinite dilution. It can be represented by the general equation:
X+ (g) → X+ (aq)
144
what is a sol of infinite dilution
A solution of infinite dilution is a solution that has a large excess of water and the addition of more water would not cause any more heat to be released or absorbed.
145
The enthalpy change of solution, ΔH⦵sol,
the enthalpy change when one mole of solute dissolves to form a solution of infinite dilution.
146
enthalpy change of sol is the sum of
lattic and hydration enthalpies
147
if a substnace has a high positive value fo the enthalpy of solution
generally insoluble
148
factors that affect enthalpy of hydration
ionic radius and charge on ion
149
the magnitude of enthalpy of hydration decreases whilst
ionic radius increases.
This is due to the weaker ion-dipole forces produced as the size of the ion increases.
charge decreases
150
the charge density of an ion increases with
greater ionic charge and a decrease in ionic radius
151
what does a higher charge density result in
a stronger ion-dipole force between the ion and the water molecule, and a greater, more negative value for the enthalpy of hydration.
152
substances with a high positive (endothermic) value of ΔH⦵sol, are
less soluble
153
is melting a positive or negative entropty change?
increase, positive tnropy
154
is deposition a positive or negative entropty change?
decrease, negative entropy
155
is sublimation a positive or negative entropty change?
increase, positive entropy
156
dissolving a solute to form a solution positive or negative entropty change?
increase, positive entropy
157
spontaneous meaning
A spontaneous process occurs without the addition of energy, other than that required to overcome the initial energy barrier (also known as the activation energy).
158
For a spontaneous process, the total entropy of the system and surroundings (ΔStotal) must increase. This can be represented in equation form as:
ΔStotal = ΔSsystem + ΔSsurroundings ≥ 0
159
how can negative entropy reactions be spontaneous
the entropy of the surroundings increases to a much greater extent, which gives an overall increase in entropy for the process.
160
For the surroundings, the entropy change is given by the following relationship:
ΔSsurroundings=−ΔH/T
161
entropy change for the universe
ΔSuniverse=ΔSsystem−ΔH/T
162
gibbs free energy change formula
ΔG=ΔH−TΔS
163
what is gibbs free energy
the energy associated with a chemical reaction that can be used to do work.
164
This example shows us that there are three factors to be considered when determining the spontaneity of a reaction:
the sign of the ΔH
the sign of the ΔS
the temperature at which the reaction takes place.
165
if gibbs is negative
reaction is spontaneous
166
The Gibbs free energy of formation, ΔG⦵f, is defined as
the change in free energy when one mole of a compound is formed from its elements in their standard states under standard conditions.
167
gibbs free energy change formula
ΔG⦵ = ΣΔG⦵f (products) − ΣΔG⦵f (reactants)
168
requirements fo r a chemical reaction to take place
correct orientation
sufficent energy
169
what is the transitions tate
the highest energy state on a reaction coordinate; it indicates a point at which new bonds are being formed at the same time as old bonds are being broken.
170
what is the maxwell boltzmann distribution
a graph where the x axis is kinetic energy and the y axis fraction of particles
171
the area beneath the curve in the MB curve is directly proportional to
directly proportional to the number of molecules having a value of kinetic energy in that range.
172
as temp increases in the mb curve,
the particles gain ke and the curve flattens out
173
how is the maxwell boltzmann distribution helpful
in understanding how changes in temperature and the use of a catalyst will affect the rate of a reaction.
174
what factors affect rate of reaction
temperature
concentration
pressure
surface area
175
when does particle size affect rate of reaction
when the solid is being reacted directly, not when it is dissolved in solution.
176
The factor by which the concentration of a reactant affects the rate of a reaction is known as
its order of reaction
177
what is the only way the order of a reaction with respect to a particular reactance can be determined
through experimental data
178
what does the rate expression give
the relationship between reactant concentrations and the rate of reaction
179
what is k
the rate constant
180
what is the overall oder of reaction
the sum of indiviual orders of reactions
181
zeroth order
Rate is independent of [A]; any change in the concentration of A does not affect the rate of reaction
182
first order
Rate is directly proportional to [A]; if the concentration of A is doubled, the rate also doubles
183
the rate expression gives the rel between the conc of the reactants and the
overall rate of reaction
184
k is what dependent
temperature
185
zero order units
mol dm-3 s-1
186
first order units
s -1
187
second order units
mol-1 dm3 s-1
188
third order units
mol-2 dm6 s-1
189
CT graphs
0, straight line
1, exp curve
2, exp curve
190
RC graphs
0, horizontal straight line
1, straight line
2, parabolic curbe
191
what is half life
the time it takes for the concentration of a reactant to decrease by half
192
a reaction that takes place in a single step is known as a
elementary reaction
193
what are the steps in a non elementary reaction called
elementary steps
194
slowest elementary step is the
rds
195
the elementary steps must add together to give
the overall balanced equation for the reaction
196
a substance in multiple steps but not the overall equation is known as
a reaction intermediate
197
what does molecularity tell us
the number of reactant particels in an elementary step
198
If only one particle is involved in the elementary step, it is
unimolecular
199
If two reactant particles are involved,
it is bimolecular
200
if three reactant particles are involved,
termolecular
201
the rds has the highest
activation energy as its the slowest
202
why do catalysts appear in the elementary steps of a reaction mechanism and in the rate-determining step but not in the overall balanced equation for the reaction
they are chemically unchanged and can be reused
203
what value constant does a catalyst change
increasing the value of the rate constant k
204
rate constant k is temp dependent, this relatiponsh[ is shown by the arrhenius equation:
k=Ae ^−Ea/RT
205
what is a in the arrhenius equation
the pre-exponential factor or the frequency factor.
206
what is Ea in the arrhenius equation
activation energy (J/mol^-1)
207
what is r in the arrhenius equation
universal gas constant (8.31 J K^1mol^-1)
208
what is e in the arrhenius equation
eulers number
209
what does the arrhenius constant take into account with a
the frequency of collisions and the probability that they have the correct orientation (or geometry)
210
The expression (e−Ea /RT) (the exponential factor) is the fraction of molecules that have
sufficient kinetic energy to react at a certain temperature
211
modified version of arrhenus equation (to find ativation energy)
lnk=−Ea/RT +lnA
212
lnk =−Ea/RT +lnA is a form of y = mx+c
therefore, a graph of ln k against 1/T gives a straight line (Figure 1). The gradient of the line is equal to –Ea/R (where R is the universal gas constant with a value of 8.31 J K−1 mol−1). The intercept on the y-axis is equal to ln A. Once we have determined the activation energy for the reaction, we can calculate the Arrhenius constant by substituting the activation energy value into the Arrhenius equation.
213
Gradient=
−Ea/R
214
−Ea=
gradient×R
215
the rate constant, k, for a reaction increases ???? with increasing temperature
exponentially
216
4 key characteristics of a reaction at equilibrium
1. The forward and reverse reactions occur at the same rate.
2. The concentrations of reactants and products at equilibrium are constant.
3. Equilibrium requires a closed system.
4. There is no change in macroscopic properties at equilibrium.
217
what is Kc
a constant for a given temperature, therefore, it is temperature-dependent. A change in temperature would result in a change in the value of Kc for a reaction (this is discussed in more detail in a later section).
218
high Kc
reaction almost at completion
219
low Kc (less than 10^-10)
reaction barely proceeded
220
what value of Kc means reactants are predominate at eq
0.01
221
what value of Kc means there are equal amounts of reactants and products
1
222
what value of Kc means products are predominate at eq
100
223
what is the value of Kc when the reactioon has almost gone to completion
over 10^10
224
if a reaction is reversed, what happens to the kc
it will be the reciprocal of that value
225
if the reaction coefficents are doubled, the Kc is
squared
226
if the reaction coefficients are halved, the kc is
square rooted
227
what is the reaction quotient Q
a measure of the relative amounts of reactants and products for a reaction that is not yet at equilibrium. The reaction quotient is useful for predicting in which direction a reaction will proceed to reach equilibrium.
228
what is the different in reaction quotient expressiona dn equilibrum constant expression (q and kc)
kc uses equilibrium values
229
Q vs Kc
Q less than Kc, reaction exceeded eq
Q = Kc, reaction at eq
Q bigger than Kc, reaction not at eq yet
230
Le Châtelier’s principle states that:
'When a system at equilibrium is subjected to a change, the system will respond to minimise the effect of the change.'
231
The chemical process that produces ammonia gas involves the following reaction:
N2 (g) + 3H2 (g) ⇌ 2NH3 (g) ΔH = –92 kJ mol–1
232
the haber process: This means that increasing the pressure will shift the equilibrium position to the right, yielding more what
ammonia
233
catalyst of haber process
iron
234
what does contact process create
sulfuric acid
235
main uses of contact process
Fertilisers.
Paints and pigments.
Detergents and soaps.
Dyestuffs.
236
pressure and temp in the contact process
low temp and high pressure according to le chatlier, but this is uneconomic
237
Kc involving the forward and reverse reaction
Kc = Kf / Kr
238
If the rate constant for the forward reaction (kf1) is greater than that of the reverse reaction (kr1), the equilibrium constant will be
large, and the forward reaction is favoured
239
if the rate constant for the backward reaction (kr2) is greater than that of the forward reaction (kf2), the equalibrium constant will be
small and the reverse reaction is favoured
240
You can see that the more negative the value of the standard Gibbs free energy change, the
larger the value of the equilibrium constant, K. This means that the more spontaneous a reaction is, the further the position of equilibrium lies to the right
241
if gibbs is negative, ln k is positive and Kc is less than 1, waht is the pos of e?
Lies to the right – favouring the products
242
the BL theory defines acids and bases as
proton donors and proton acceptors
243
amphiprotic species
able to act as both a bronsted lowry acid and a bronsted lowry base depending on what its reacting with
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amphiprotic only applies to what theory
the BL theory
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amphoteric refers to
substances that can act as any acid or base (not restricted to Bl theory)
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metal + acid →
salt + hydrogen
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metal oxide + acid →
salt and water
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metal hydroxide + acid →
salt and water
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what is a neutralisation reaction
an acid reacts with a base or alkali to produce a salt and water
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ingredients in most antacids
The active ingredients of most antacid tablets are metal carbonates or hydrogen carbonates, NaHCO3, CaCO3, MgCO3, for instance, or insoluble metal hydroxides such as Mg(OH)2 or Al(OH)3. These react with excess stomach acid in neutralisation reactions to relieve the symptoms of heartburn.
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purpose of an acid base titration
to determine the unknown concentration of an acidic or basic solution using a solution of known concentration
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thermometric titration
when heat is released during a titration.
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pH formula
pH = −log[H+(aq)]
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change in one pH unit is equal to
ten times the change in hydrogen ion conc
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acid base indicator is
a weak acid or weak base in which the dissociated and undissociated forms have different colours
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litmus comes in two colours
red and blue
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in an acidic sol, blue litmus urns to
red
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in an alkaline sol red litmus turns to
blue
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universal indicator
red in acid
purple in alkali
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pH probe
more accurate method of measuring pH
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Water molecules do dissociate however, but only to a very small extent; this is known as the auto-ionisation of water. This reaction can be represented by the following equation:
H2O (l) ⇌ H+ (aq) + OH– (aq)
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The equilibrium constant expression (Kc) for the auto ionisatin of water is
Kc=[H+][OH−]/[H2O]
The position of equilibrium for the dissociation of H2O lies very far to the left, so the concentration of the water is effectively constant. Therefore, we can write a new expression which is known as the ionic product of water (Kw):
Kw=[H+][OH−]
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ionic product of water
Kw=[H+][OH−]
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Kw
1 x 10^-14
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The strength of an acid refers to its
degree of dissociation (or ionisation) in aqueous solution.
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three common strong bases
sodium hydroxide, potassium hydroxide and barium hydroxide
metaly hydroxides of group 1
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for the dissociation of weak bases the eq lies very far to the
left
Kc is very small
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remember to use an equilibrium sign (⇌) for
weak acids and bases.
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equimolar solutions
ones that have equal concentrations
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active metals are those
above hydrogen in the activity series
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which have higher elec conductivity
strong bases
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Natural, unpolluted rainwater is acidic, having a pH of approximately
5.6
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acid rain pH
less than 5
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how is acid rain formed
when acidic gases such as sulfur dioxide (SO2) and the oxides of nitrogen (NOx) dissolve in the water in the atmosphere to produce sulfuric acid or nitric acid.
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how can nitrogen monoxide be formed
during lightning storms by the reaction of nitrogen and oxygen, the two predominant atmospheric gases. The second is by the reactions that take place in internal combustion engines. Nitrogen reacts directly with oxygen at high temperatures to produce nitrogen monoxide
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benefit of a catalytic converter
reduces the level of these pollution emissions. This reaction also takes place in the jet engines of aircraft.
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how is sulfur dioxide formed
occurs naturally in the atmosphere as it is released during volcanic eruptions
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sulfur dioxide photochemical oxidation in the atmosphere
2SO2 (g) + O2 (g) → 2SO3 (g)
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sulfur dioxide photochemical oxidation in the atmosphere
2SO2 (g) + O2 (g) → 2SO3 (g)
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what does sulfur dioxide form with water
Sulfur dioxide and sulfur trioxide react with water in the atmosphere to form sulfurous and sulfuric acid:
SO2 (g) + H2O (l) → H2SO3 (aq)
SO3 (g) + H2O (l) → H2SO4 (aq)
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what is dry deposition
acidic particles and gases fall to the ground via dust and smoke in the absence of precipitation. This form of deposition can be washed into streams, lakes and rivers, causing harm to biological systems.
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which are the stronger acids
These four acids (nitric, nitrous, sulfuric and sulfurous acids) are all much stronger than carbonic acid, H2CO3.
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environmental impact of acid deposition
prevent fish eggs from hatching
damage plants
harms biodiversity
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how is acid rain detrimental on statues
calcium carbonate can react with sulfuric acid and nitric acid
CaCO3 (s) + H2SO4 (aq) → CaSO4 (aq) + H2O (l) + CO2 (g)
CaCO3 (s) + 2HNO3 (aq) → Ca(NO3)2 (aq) + CO2 (g) + H2O (l)
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acid rain on humans
can adversely affect the mucous membranes and lungs, causing irritation and possibly exacerbating the symptoms for people with asthma and other respiratory conditions.
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how to reduce acid deopisition
pre combustion methods: removing the sulfur before the coal is combusted
post comubstin methods: removing the sulfur oxides from exhaust gases once they have been formed by reacting with a base.
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hydrodesulfurisation
