MCAT-Gen Chem Flashcards
(178 cards)
1
Q
What is the atomic number of an atom?
A
It is the number of protons in the nucleus of an atom.
2
Q
What is the mass number of an atom?
A
It is the sum of the number of protons and neutrons in the nucleus of an atom.
3
Q
What is the definition of an isotope?
A
When two atoms of the same element have different number of neutrons.
4
Q
Describe Beta - decay
A
Occurs when an unstable nucleus has too many neutrons
Converts a neutron to a proton and an electron
Mass number stays the same
The atomic number increases by one.
5
Q
Describe beta+ decay
A
When an unstable nucleus contains too few neutrons
Converts a proton into a neutron and a positron
Atomic number is reduced by 1
Mass number is not changed
6
Q
Describe electron capture
A
When an unstable nuclei has too few neutrons
converts a proton and an electron into a neutron
Mass number stays the same
Atomic number decreases by one
7
Q
What is unique about beta - decay and alpha decay?
A
Since it changes the atomic number of the element, the identity of the daughter nucleus changes.
8
Q
Describe alpha decay
A
Alpha decay is equivalent to losing a He nucleus, reduce mass of radioactive element by 4 and reduce the atomic number by 2.
9
Q
What is c, the speed of light?
A
3*10^8 m/s
10
Q
What is Avogadro’s number
A
6.023*10^23
11
Q
What is the equation for exponential decay?
A
N=No(1/2)^T/t-half life
N=No*e^-kt
Where k=ln2/t-half life
Note N equals the number of radioactive nuclei remaining after T time. No is the initial amount of radioactive nuclei.
12
Q
What is the definition of mass defect?
A
Delta m= total mass of separate nucleons- mass of nucleus
13
Q
What is the equation for nuclear binding energy?
A
Eb = mass defect* 931.5Mev
Where mass defect is in amu units and binding energy is in units of electron volts
14
Q
What is the value of plancks constant?
A
h= 6.64*10^-34 J•s
15
Q
What is the equation for frequency?
A
f=c/wavelength
16
Q
What is the equation for the energy of a photon?
A
E=h*f
=c/wavelength
17
Q
What is the definition of a Bohr atom?
A
An atom with only one electron
18
Q
What is the Aufbau principle?
A
Electrons occupy the lowest energy orbitals
19
Q
What does Hunds rule state?
A
Electrons in the same sub shell occupy available orbits singly before pairing up
20
Q
What does the Pauli exclusion principle state?
A
There can be no more than two electrons in any given orbital
21
Q
List out all of the electron configurations to 3d
A
1s 2s 2p 3s 3p 4s 3d
22
Q
What is the definition of a diamagnetic atom?
A
It has all of its electrons spin paired in its orbital configurations.
Must have an even number of electrons ( even atomic number)
Is repelled by an external magnetic field
23
Q
What is the definition of a paramagnetic atom?
A
An atom where not all of the electrons are spin paired in the orbital configuration.
Can have an even or odd number of electrons.
Is attracted to an external magnetic field.
24
Q
What are the d block atoms that will have either half filled or filled d orbitals?
A
Cr, Cu, Mo, Ag and Au
They promote electrons from their 4s orbitals to attain these more stable configurations.
25
Transition metals lose electrons from what orbitals when they are ionized?
S orbitals
26
What are isoelectric atoms
An ionized atom that has the same electron configuration as a non ionized atom
27
What about an atom dictate their properties and chemical behavior?
Valence electrons
28
Atomic radius increases how along the periodic table?
Down and to the left.
29
List the order of electro negativity of the 9 most EN atoms
F>O>N=Cl>Br>I>S>C=H
30
What is the trend for ionization energy across the periodic table?
Increasing Up and to the right
31
What is the trend for electron affinity across the periodic table?
More negative up and to the right
32
What is the trend for electronegativity across the periodic table?
Increasing up and to the right
33
What is the trend for acidity across the periodic table?
Increasing down and to the right
34
What is the trend for atomic radius for cations, atoms and anions.
Cation< neutral atom< anion
35
State rules for assigning oxidation states
1) oxidation state of any element in standard state is 0
2) sum of ox stars of a neutral molecule is 0
3) group 1 elements have a +1 ox star and group 2 have a +2 ox state
4) Fl has a -1 ox state
5) H has a +1 ox state when bonded to something more electronegative, -1 ox state when bonded to something less electronegative than carbon and 0 when bonded to carbon.
6) oxygen has a -2 ox state
7) the other halogens have a -1 ox state and the atoms of the oxygen family have an ox state of -2.
Rules higher up have priority over lower rules
36
State the equation for formal charge of an atom in a molecule
FC= valence electrons-1/2(#bonded electrons)- lone paired electrons
37
For similar bonds, state the trend regarding bond order, bond distance and bond strength (bond dissociation energy)
The higher the bond order, the shorter the bond distance and the stronger the bond.
Single bond = bond order of 1
Triple bond has bond order of 2
38
What is a coordinate covalent bond?
Formed when an atom donates both of the shared electrons used to form the covalent bond.
A coordination complex is formed
39
What are the definitions of a Lewis base and Lewis Acid?
Lewis Base- donates a pair of electrons
Is a nucleophile and ligand.
Lewis acid- accepts a pair of electrons
Is an electrophile
40
For an ionic bond to form between a metal and a nonmetal, the difference in what needs to be great?
Electronegativity
41
Ionic bond strength increases and decreases with what?
Increases proportionally with the charges on the ions and decreases with increasing atomic radii
42
How many electron groups are on a bent molecular geometry? How many lone pairs are there?
3 electron groups, one of them is a lone pair
43
How many electron groups are on a trigonal planar molecular geometry? How many lone pairs are there?
3 electron groups, no lone pairs
44
How many electron groups are on a trigonal pyramid molecular geometry? How many lone pairs are there?
4 electron groups, one of the groups is a lone pair
45
How many electron groups are on a tetrahedral molecular geometry? How many lone pairs are there?
4 electron groups
No lone pairs
Dipole moments cancel out
46
What are intermolecular forces?
What are examples?
The relatively weak interactions between neutral molecules ( not ionic or covalent)
Ion dipole
Dipole dipole
Dipole induced dipole
Instantaneous dipole dipole-weakest and transient
47
As the size of a molecule increases so does what intermolecular force?
Dipole induced dipole. As a molecule increases, so does the number of electrons and so does the ability for polarizability. The induced charges of the dipoles become larger and the strength of the force increases.
48
Branching in hydrocarbons does what to intermolecular forces?
What is the impact on melting point, boiling point and vapor pressure?
Branching inhibits intermolecular forces.
Thus a branched isomer will have a lower melting point, lower boiling point and increase the vapor pressure compared to the nonbranched isomer.
49
Substances with stronger intermolecular forces will exhibit?
Greater melting points, boiling points, vapor pressures and viscosities.
50
Larger heavier molecules have lower or higher boiling points and melting points than smaller less heavy molecules?
Higher melting point and boiling point because increased surface area allows for more intermolecular forces.
51
What are the two criteria for H bonding to occur?
1) one molecule must have a covalent bond between H and N, O or F.
2) the other molecule must have a lone pair attached to N, O or F.
52
The weaker a substances intermolecular forces, the it’s vapor pressure?
The higher the vapor pressure
53
The forces of molecular solids are weaker or stronger compared to ionic, network and metallic solids?
Weaker, thus are liquids and gases typically at room temperature whereas the other solids are actually solids.
54
What is the equation for the heat added or removed from a system?
q =mc *change in temp
```
m= mass of substance
c= specific heat of substance
```
Note that m*c = C = heat capacity
55
What is the conversion between calories and joules?
1cal= 4.2 joules
56
What does the specific heat of a substance describe?
How resistant it is to temperature change.
The larger the value, the more resistant it is.
57
Thermodynamics predicts what of a reaction?
What does it not predict?
Spontaneity and the equilibrium of reactions.
It says nothing about the rate of a reaction. That’s kinetics.
58
What are standard conditions?
Pressure is 1 atm
Temp is 298K (25C)
Substances are assumed pure
Concentrations are 1M
This is not standard temperature and pressure*****
59
When a bond is formed, energy is what?
Energy is released, delta H is negative
Exothermic
60
In order to break a bond energy must be?
Energy must be put into a bond to break it.
Delta H is positive. Endothermic
61
What is the equation for the heat of phase changes?
q = n *delta H of phase change
62
What is the equation for the enthalpy of reaction using bond dissociation energies?
Delta H= (summation of BDE bonds broken)- (summation of BDE of bonds formed)
63
What is the definition of the standard heat of formation?
The energy change associated with making 1 mole of a compound from its constituent elements in their natural state or standard state
64
What is the equation for the standard heat of formation for a reaction?
Delta H= (summation of n* heat of formation of products)- (summation of n* heat of formation of reactants)
65
What are the two criteria to evaluate when looking at a chemical reaction and trying to determine the impact on entropy?
Do you create more moles of products versus reactants?
Does the reaction convert a solid to a liquid or gas?
66
What is the equation for Gibbs free energy?
DeltaG= deltaH -T*delta S
67
What directly impacts the state or phase of a substance?
The average kinetic energy of the molecules.
68
Phase changes that bring molecules together release or absorb heat?
They release energy as heat and thus the kinetic energy of the molecules is reduced and intermolecular forces enforce some type of order to the phase.
69
During a phase transition, what happens to the temperature of the substance?
The temperature does not change. Will change once the heat added or removed is equal to q= n* delta H of phase change
70
What has the largest change in heat to the system on a phase transition diagram, heat of fusion or heat of vaporization?
A substances heat of vaporization is always larger than a substances heat of fusion
71
Contrast the axis of a phase transition/heat curve and a phase diagram
Phase transition/heat curve
Y axis is temperature. X axis is heat added (q)
Phase diagram
Y axis is pressure
X axis is temperature
72
How do you determine the normal boiling point and melting point of a substance on a phase diagram?
Draw a horizontal line across the phase diagram at 1 atm. Where the line intersects the phase boundary between liquid and gas is the normal boiling point. Same thing for melting point.
73
Since water is more dense in liquid phase than solid phase, what occurs to the solid liquid line boundary?
It has a slightly negative slope ( tilted towards the left) compared to other substances.
74
Name the phases of a phase diagram chart from left to right
Solid, liquid is in cup, gas
75
What are the units/conversions for volume?
1cm^3= 1cc = 1 ml
1m^3= 1000L
76
What is the conversion from Celsius to Kelvin?
Temp(K)= temp(in Celsius) + 273.15
77
What is standard temperature and pressure?
0C/273.15 K
1atm
78
What are the units and conversions for pressure?
1atm = 760 torr= 760 mmHg= 101.3kPa
79
What is the universal gas constant?
0.0821 L-atm/K-mol
80
What is the standard molar volume of any ideal gas?
22.4 L
81
What is the definition of the standard molar volume?
The volume occupied by 1 mil of an ideal gas at standard temp and pressure
82
What is the relationship between P and V?
What is the relationship between V and T
What is the relationship between P and T
What is the relationship between n and V
P and V are inversely proportional
V and T are directly proportional
P and T are directly proportional
N and V are directly proportional
83
Ideal gases behave the most ideal at what conditions?
High temperature and low pressures
84
What are the assumptions made in the kinetic molecular theory of ideal gases?
1) particles move in straight lines
2) collisions of particles are elastic, maintain the average kinetic energy due to no attraction and repulsion forces between the molecules.
3) The gases take up a negligible amount of volume, volume of gas equals volume of the container
4) the speed of the particles are directly proportional to the temperature and the molecules have a distribution of speeds
85
How do real gases deviate from ideal gases and the kinetic molecular theory?
Ideal gases do not be as ideal at high pressures and low temperatures
Real gases do have intermolecular forces and thus P real< P ideal
Real gases do occupy some volume and reduces the free space available for particle movement.
V real< V ideal
86
What is the Van Der Waals equation?
What does it correct for?
P= nRT/V-nb - n^2*a/V^2
a and b are the van der waal constants and are generally larger for gases with greater intermolecular forces (a) and have larger molecular weights, which increases the volumes (b)
Corrects for intermolecular forces and physical volume of gases at non ideal conditions
87
What are the characteristics of an ideal gas?
Have the weakest intermolecular forces and have the smallest molecular weights/volumes.
Maintaining high temp and low pressures help to minimize interactions and help the volume of gas stay insignificant compared to the volume of the container.
88
2 gases at the same temperature have the same or different average kinetic energy?
Do they have the same or different average speeds and effusion rates?
The same
They have different average speeds and different effusion rates based on their molecular weights
89
What is the equation for kinetic energy?
KE=1/2 mass* velocity^2
90
What is a convenient measurement for average speed of gas molecules?
RmsV
91
What is the relationship between average speeds and molecular weight between 2 gases?
rmsVa/rmsVb=sqrt (mw of b/mw of a)
92
What is the equation for Graham’s law of effusion?
Effusion rate of gas A/ effusion rate of gas B= sqrt(molar mass of gas B/molar mass of gas A)
93
What is the definition of a rate determining step?
It is the slowest step in a process which determines the overall reaction rate.
94
Why is a transition state intermediate short lived in a a regular chemical reaction?
It is high energy
95
The term activated complex is a synonym for?
Transition state intermediate
96
The lower the activation energy, the faster the?
Reaction rate
97
If there are two elementary reactions that form an overall chemical reaction, the reaction with the higher activation energy is the what?
Rate limiting step
98
Two criteria increase the reaction rate of a chemical reaction, what are they?
The greater the concentration of reactants, the faster the rate.
The higher the temperature the faster the rate.
99
A catalyst does what?
It lowers the activation energy of the rate limiting step( thus increasing the reaction rate.
100
Catalysts increase the rate of reaction however they do not impact what?
The thermodynamics or the equilibrium
101
Is the delta G of a catalyze reaction the same as the delta G of an uncatalyzed reaction?
Yes
Catalysts have a kinetic role, not a thermodynamic role
102
What is the sole characteristic of an intermediate formed in a chemical reaction?
Intermediates are formed and the consumed in subsequent elementary reactions.
103
Do catalysts impact the values of change in enthalpies or changes in Gibbs free energy?
No, catalysts serve a kinetic role.
104
What is the rule of thumb regarding the impact of temperature on the rate constant via the Arrhenius equation?
For every 10C increase in temp, the rate constant increases by a factor of 2 to 4.
105
What are the units of the rate constant in a first order reaction?
units are s^-1
106
What are the units for the rate constant in a second order reaction?
M^-1*s^-1
107
What are the units for the rate constant when the overall order of the reaction is third order?
M^-2*s^-1
108
Only reactants involved in the rate limiting step are part of what?
The rate law expression
Rate=k(A)^x*(b)^y
109
What is the expression and definition of the reaction quotient?
Reaction quotient expression is the same as the equilibrium constant except the reaction is not at equilibrium.
If Q=Keq then the reaction is at equilibrium
110
If Q is larger than Keq then what direction will the reaction favor?
The reverse reaction in order to produce more reactants so that Q equals the equilibrium constant
111
If Q is smaller than Keq then what direction will the reaction favor?
The reaction will favor the forward reaction towards products
112
The equilibrium constant is a constant at what?
A given temperature
If you change the temperature you will change the equilibrium constant
113
If a gaseous reaction at equilibrium with unequal number of moles on each side has the volume increased, how will equilibrium be reestablished?
Increasing volume= decreasing pressure
The reaction will favor side of equation of with the largest number of moles to increase pressure back to equilibrium
114
If a gaseous reaction at equilibrium with unequal number of moles on each side has the volume decreased, how will equilibrium be reestablished?
Decreased volume= increased pressure
The reaction will favor the side of the equation with the lower total number of moles in order to decrease pressure back to equilibrium.
115
How do we treat the change in temperature of a reaction at equilibrium?
We treat heat as a reagent.
We must know whether the reaction is exothermic (heat is on products side) or if it is endothermic (heat is on reactants side).
Note that the change in temp of a reaction at equilibrium can impact the value of Keq
116
Adding a catalyst to a reaction that is at equilibrium has what impact?
None. The catalyst increases the forward and reverse reactions equally but does NOT affect the equilibrium.
117
What does the term dissolution mean?
The process of dissolving
118
What is the definition of an aqueous solution?
When a solution has water as the solvent
119
Solutes will dissolve in solvents when?
The intermolecular forces that are being broken in the solute are replaced with equal or greater forces between the solvent and solute.
120
When a saturated solution is formed, what is reached?
The molar solubility of the solute for that particular solvent.
Adding additional solute will result in the reverse process of dissolution which is precipitation at the same rate as dissolving.
121
What does the term solubility refer to?
The amount of solute required to saturate a particular solvent
122
List the three phase solubility rules
1) the solubility of solids in liquids tends to increase with increasing temperature
2) the solubility of gases in liquids tends to decrease with increasing temperature
3) the solubility of gases in liquids tends to increase with increasing pressure
123
The solubility of a gas in a liquid is a function of what?
The partial pressure of the gas above the liquid and Henry’s law constant solubility=k*P
124
List the three salt solubility rules
1) all group 1 and ammonium salts are soluble
2) all nitrate, perchlorate (ClO4-) and acetate (C2H3O2-) are soluble
3) all silver (Ag), lead Pb and mercury Hg salts are insoluble with the exception for their nitrates, perchlorate sand acetates.
125
The extent to which a salt will dissolve in water at equilibrium can be determined from what?
The salts solubility product constant Ksp
You leave solids out of the equilibrium expression
Ksp= ion product concentrations raised to coefficients
126
How does the ion product differ from the salt solubility product?
The expression is the same but Qsp the concentrations don’t have to be at equilibrium.
127
The ion product is what?
Is the reaction quotient (Qsp) for the salt solubility constant Ksp
128
When Qsp< Ksp ?
Qsp =Ksp ?
Qsp >Ksp ?
More salt can be dissolved in solution- forward solubility reaction favored
The solution is saturated- the solubility reaction is at equilibrium
Excess sat will precipitate- the reverse solubility reaction is favored towards salt reactant
129
If you mix two salts in solution what can happen?
The ions can recombine to form new salts and you must consider the new salts Ksp
130
Describe the common ion effect
When an additional salt is added to a concentration where both salts have a common ion. This causes an increase in the ion product concentration. Via le Chatliers principle the reverse solubility reaction is favored and more salt reactant is formed (precipitates). The molar solubilities of the salt reactants is reduced.
131
The addition of ligand a may substantially alter what of a salt?
The solubility via the formation of a complexed metal.
NH3 is a common ligand/Lewis base and forms coordination complex with ions of salts that have low solubilities, thus increasing the solubility of the salt in solution.
132
List the three rules relating deltaG of formation to Keq
When delta G of formation <0, then Keq>1, the forward reaction and products are favored at equilibrium
When deltaG of formation=0, Keq =1, reactants and products are equally present at equilibrium
When delta G of formation >0, Keq<1, the reverse reaction/reactants are favored at equilibrium
133
What is the definition of
Bronsted Lowry Acid
Bronsted Lowry Base
acid are proton donors (H+)
| Bases are proton acceptors (H+)
134
What determines the strength of an acid?
Depends on how completely the acid dissociates/ionize in water
135
Strong acid completely dissociate in water so the reaction does what?
Goes to completion, there is no equilibrium
136
What is the expression for the acid dissociation constant Ka?
Ka= [H+] [A-]/[HA]
137
If Ka>1, what is favored? What is the strength of the acid?
Products are favored
| The acid is strong
138
If Ka<1, what are favored? What is the strength of the acid?
Reactants are favored
| The acid is weak
139
The larger the Ka, the stronger or weaker the acid?
Stronger
140
List the strong acids that have Ka>1
HI (strongest), HBr, HCL, HClO4 (perchloric), H2SO4 (sulfuric), HNO3 (nitric)
Assume any acid other than this list is a weak acid on MCAT
141
What is the expression for the base ionization/dissociation constant?
Kb=[HB] [-OH]/[B]
142
List the common strong bases
Group 1 hydroxides (NaOH)
Group 1 oxides (Li2O)
Some group 2 hydroxides (BaOH2, SrOH2, CaOH2)
Metal amides (NaNH2)
143
The conjugate base of a strong acid has no _______ properties in water?
No basic properties, does not act like a base.
Note** strong acid dissociation reactions go to completion, there is no reverse reaction.
144
The conjugate base of a weak acid is what type of base?
A weak base
The weaker the acid, the stronger the weak conjugate base is, i.e. the more the reverse reaction is favored.
145
Acid 1 has a Ka=5X10^-10
Acid 2 has a Ka= 7X10^-4
Which acid has the stronger conjugate base?
The conjugate weak base of acid 1 is stronger
Both acids are weak due to Ka's<1
The weaker the acid the stronger the conjugate weak base and v.v. for weak bases.
146
The conjugate acid of a strong base has no ______ properties in water?
acidic properties
The ionization reaction goes to completion thus the conjugate acid has no acidic properties
147
What is the definition of an amphoteric species?
It can act as a base or an acid, usually is polyprotic (has more than one proton to donate)
148
Every time a polyprotic acid donates a proton, the resulting species is either a stronger or weaker acid than its predecessor?
It will be a weaker acid and thus smaller Ka value.
149
What is the expression for the ion product constant of water?
Kw=[H3O+] [-OH]= 1.0X10^-14
150
When you add acid to water the equilibrium is disturbed. What reaction is favored and how is equilibrium re-established?
The amount of hydronium ions is increased. The reverse reaction is favored and the concentration of [OH-] is decreased
If you add base, the amount of hydroxide ions is increased and the equilibrium is disturbed. The reverse reaction is favored and the concentration of [H3O+] is reduced to balance the equilibrium and make Kw=1.0X10^-14.
151
pH=-log[H+], what is the expression when you solve for [H+]?
[H+]=10^-pH
Can do the same for pOH
[OH-]=10^-pOH
152
If the [H+] is greater than 10^-7 than the pH will be?
Lower than 7 and the solution will be acidic
A high [H+] means a low pH and v.v.
153
What is the expression relating pOH and pH?
pH + pOH=14
If you know the pH then you know the pOH and V.V
154
what does Log10 (1.0X10^-3) equal?
-3
155
what does -Log10 (1.0X10^-3) equal?
3
156
YX 10^-n, in cases when the H+ concentration does not equal to the whole number power of 10, what is the rule to determine the boundary of the pH?
pH will be between n and (n-1)
EX: [H+]=3.0X10^-4
pH will be between 4 and 3
The closer Y is to 10 the closer the pH is to the smaller value
157
The larger the Ka the ________ the acid?
The smaller the pKa the __________ the acid?
Stronger
Stronger
Same is true for Kb
158
Ka x Kb= ?
Ka x Kb=Kw=1.0X10^-14
If you know the Ka of an acid then you can determine the Kb of the conjugate base
159
pKa + pKb= ?
pKa + pKb= 14
If you know the pKa of an acid you can determine the pKb of the conjugate base
160
What is the expression to use when asked to determine the pH of a weak acid when given the initial acid concentration and it's Ka?
Ka=X^2/acid concentration
Simplify the math as much as possible
Note**
X^2=1x10^-10 , X= 1x 10^-5
Should verify that Ka <10^-4
161
What is the result if equimolar amounts of a strong acid and base react?
The resulting pH will be neutral.
162
What is the expression used to solve how much base (or acid) is needed to neutralize an acid (or base)?
Same expression used to relate the concentration of initial acid or base that has been titrated.
a x [Acid] x Volume of acid= b x [base] x Volume of base
```
a= the number of acidic hydrogen atoms per formula unit
b= # of protons the base can accept
```
163
all neutralization reactions do what?
Go to completion
Note* neutralization reactions=acid base reaction
164
What are the cations of salts that do not react with water?
Group 1 cations
| Larger Group 2 cations
165
What are the cations of salts that are stronger acids than water?
NH4+, Cu, Be, Zn, Al, Cr, Fe
166
What are the anions of salts that are stronger bases than water?
conjugate bases of weak acids
167
What is a buffer made of and how is an ideal buffer designed?
A buffer is made of a weak acid and the salt conjugate base
An ideal buffer is made where [H+]=[A-] so that pH=pKa via Henderson Hasselbach equation
Also, pick a weak acid whose pKa is as close as possible to the pH desired. If no weak acid has exactly the same pKa then adjust initial acid or base concentrations to target the pH.
168
What is the expression for the Henderson Hasselbach Equation for Acids?
pH= pKa + log [A-]/[HA]
or
[H+]= Ka x [HA]/[A-]
Used for determining the pH of a solution containing a buffer when acid or base has been added
Note* if acid is added, [HA] increases with amount of acid added and [A-] decreases concurrently and V.V.
169
What is the expression for the HH equation for base?
pOH= pKb + log [HA]/[A-]
170
log of a fraction is positive or negative?
Negative
171
Log of a positive whole number ex. 1.05 is positive or negative?
Positive
172
What is the definition of a pH indicator?
What is the general rules concerning indicators and their pKa's?
It is a weak acid that changes color when converted to the conjugate base.
Below pKa-1 and the color is the innate color
Above pKa+1 the color is different
Between pKa-1 and pKa+1 the color is a mix between the two.
173
acid base titration is a technique to determine?
The identity of a weak acid/base or determine the concentration of a weak acid or base of a solution.
The titrant is always either a strong acid or base.
174
Where does the half equivalence point occur on an acid base titration graph? What is the expression for this point?
in the middle of the buffer domain
HE point is where the weak acid concentration equals the conjugate base concentration in solution and v.v.
175
What is the definition of the equivalence point on an acid base graph?
Where the moles of titrant equal the initial moles of the acid or base being titrated
For a weak acid being titrated with a strong base the pH > 7 at the eqv point
For a weak base being titrated by a strong acid the pH<7 at the eqv point
For a strong acid being titrated with a strong bases and v.v. the pH =7 at the eqv. point
176
What can be evaluated on an acid base titration graph that allows us to determine whether the ACID (OR BASE) TITRATED WAS STRONG OR WEAK?
look at the pH of the equivalence point. If it equals 7 then it was a strong acid (or base) titrated, if not then it was a weak acid (or base)
177
How can you determine the pKa of the acid or pKb of the base being titrated by evaluating the acid base titration graph?
The pH at the half equivalence point will be the pKa of the acid (or base) titrated
HH equation supports this
178
What is the empirical formula?
How do you convert a molecular formula into an empirical formula?
The chemical formula with the smallest whole numbers with the same ratio as the molecular formula.
Determine the empirical formula from the molecular formula by dividing all subscripts by the greatest common factor.
