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NMAT 2025 > Chemistry: Inorganic Chemistry > Flashcards

Chemistry: Inorganic Chemistry Flashcards

1
Q

A (physical/chemical) property tells how a substance changes into new substances.

A

Chemical

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2
Q

(T/F): A chemical property is the same thing as a chemical change.

A

FALSE

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3
Q

What are five signs of chemical change?

A

Color change, odor, temperature change, evolution of gas (bubbles), precipitate

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4
Q

(T/F): Change in temperature and odor can also be physical changes.

A

TRUE. Heating/painting

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5
Q

(T/F): The ability of a substance to support burning is flammability.

A

FALSE. This is heat of combustion; flammability describes ability of a chemical to burn or ignite

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6
Q

If property doesn’t depend on the amount of matter present, it is an (intensive/extensive).

A

Intensive

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7
Q

Volume, weight, mass, length, and number of things are (intensive/extensive) properties.

A

Extensive

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8
Q

(T/F): Both homogenous mixtures and pure substances have uniform composition and properties throughout.

A

TRUE. Pure substances always have the same composition, but mixtures have different compositions across samples

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9
Q

Material composed of two or more substances that can be separated by physical methods

A

Mixture

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10
Q

Substance that cannot be broken down into chemically simpler components

A

Element

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11
Q

Substance that can be broken down into chemically simpler compounds only by chemical methods

A

Compound

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12
Q

Filtration, distillation, and crystallization are all (physical/chemical) methods of separation.

A

Physical

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13
Q

A sharp melting/boiling point (not a range) indicates a (pure/impure) substance.

A

Pure

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14
Q

The passing of a visible beam through a sample indicates a (homogenous/heterogenous) mixture.

A

Heterogenous as it indicates small suspended particles scatter the light (suspended particles = hetero)

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15
Q

When a certain soft metal is burned in oxygen, lime is produced (with no other products). This indicates lime is a (element/compound).

A

Compound

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16
Q

(Mass number/atomic weight) refers to the number of protons and neurons while (mass number/atomic weight) refers to the average of masses of all isotopes

A

Mass number, atomic weight

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17
Q

Whose model showed that electrons surround a nucleus

A

Rutherford

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18
Q

Whose model showed that farther orbits had higher energy, and that photons are emitted when going down an energy level

A

Bohr

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19
Q

An atom or molecule with an unpaired electron

A

Free radical

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20
Q

If there are no unpaired electrons, the atom is (diamagnetic/paramagnetic).

A

Diamagnetic

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21
Q

Paramagnetic atoms (repel/attract) magnetic fields.

A

Attract

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22
Q

C = 1s^2 2s^2 2p^2 is a (diamagnetic/paramagnetic) atom

A

Paramagnetic

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23
Q

Rule that states that every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied

A

Hund’s rule

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24
Q

(T/F): According to Hund’s rule, the electrons in singly occupied orbitals may have different spins.

A

FALSE; all must have the same spin

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25
Q

Principle that states that for paired electrons in an orbital, one must have a spin of +1/2 and the other -1/2

A

Pauli exclusion principle

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26
Q

Principle that states that it is impossible to know the momentum and position simultaneously

A

Heisenberg uncertainty

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27
Q

Quantum number that describes the main electron energy level or shell number and possible values?

A

n (principal); 1,2,3,…

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28
Q

(T/F): The shell number matches the row of the periodic table

A

FALSE; this does not apply to d and f orbitals (row number-1 for d, 4 & 5 only for f)

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29
Q

Quantum number that describes the 3D shape of the orbital and possible values?

A

l (azimuthal); 0, 1, 2, n-1 (0 = s, 1 = p, 2 = d,…)

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30
Q

Quantum number (orientation) that describes the orbital sub-type and possible values?

A

m_l (orbital sub-type), integers (-L to +L, middle orbital is 0)

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31
Q

Quantum number that labels electron spin and possible values?

A

m_s (spin); +1/2 and -1/2

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32
Q

Formula for maximum e- in shell (in terms of n)

A

2n^2

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33
Q

What is the number of orbitals in the 4th shell

A

16 (1s, 3p, 5d, 7f)

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34
Q

Formula for maximum e- in subshell

A

4(L) + 2

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35
Q

Ions that have (lesser/greater) charge have -ous, and those with (lesser/greater) charge have -ic.

A

Lesser, greater

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36
Q

Polyatomic anions that contain oxygen

A

Oxyanions

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37
Q

Ions that have (less/more) oxygen end with -ate and those with (less/more) oxygen end with -ite

A

Ate, ite

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38
Q

In extended series of oxyanions, prefix hyper/per- is added to the one with (most/least) oxygen, and hypo- is added to the one with the (most/least) oxygen.

A

Most

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39
Q

Polyatomic anions that gain H+ to form anions of lower charge add the word ______ in front.

A

Hydrogen/dihydrogen

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40
Q

Group 1 is called the ______ family.

A

Alkali metals (H, Li, Na, K, Rb, Cs, Fr)

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41
Q

Group 2 is called the ______ family.

A

Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra)

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42
Q

Groups 3-12 are called the _________ family.

A

Transition metals

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43
Q

Under the 5 family system, groups 13-16 consist of which families of elements (3)?

A

Post transition metals, metalloids, non-metals

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44
Q

Under the 9 family system, groups 13-16 are what families (4)?

A

Earth metals (B), tetrels (C), pnictogens (N), chalcogens (O)

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45
Q

Group 17 is called the ______ family.

A

Halogens (F, Cl, Br, I, At, Ts)

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46
Q

Group 18 is called the _______ family.

A

Noble gas (He, Ne, Ar, Kr, Xe, Rn, Og)

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47
Q

The last two periods of the periodic table are called the _________. What elements (Z) are part of these rows?

A

Rare earth metals; 58-71 & 90-103

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48
Q

Actual amount of positive (nuclear) charge experienced by an electron in a polyelectronic atom

A

Zeff = (Z - S where S is shielding constant)

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49
Q

Pull between nucleus and valence electrons

A

Zeff (effective nuclear charge)

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50
Q

Zeff (increases/decreases/is the same) across a period and (increases/decreases/is the same) down a group

A

Increases, is the same

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51
Q

Energy required to remove an electron from a neutral atom in its gaseous phase

A

Ionization energy

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52
Q

The lower the ionization energy, the more readily the atom becomes a (cation/anion)

A

Cation

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53
Q

Ability of an atom’s inner electrons to shield its positively-charged nucleus from its valence electrons; results in decrease of IE down a group

A

Electron shielding

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54
Q

(T/F): Electron affinity is conceptually the opposite of electronegativity

A

FALSE; ionization energy is the opposite of EN

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55
Q

Ability of an atom to accept an electron (quantitative measurement of energy change when electron is added)

A

Electron affinity

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56
Q

∆H rxn for atoms with high electron affinity is (less than/greater than) 0 when gaining e-

A

Less than (note that EA is still reported as positive value)

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57
Q

Atom’s tendency to attract and form bonds with (gain) electrons; force atom exerts on e- in a bond

A

Electronegativity

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58
Q

Of the noble gases, which ones have EN?

A

Kr and Xe

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59
Q

One-half the distance between the nuclei of two atoms

A

Atomic size

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60
Q

Which trend is the only one that decreases across a period and increases down a group?

A

Atomic size

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61
Q

Are anions bigger or smaller than neutral atoms?

A

Bigger (due to e-e repulsion)

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62
Q

Mercury is better suited than water for use in a barometer chiefly because mercury has _______.

A

High density and little evaporation

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63
Q

Hypothetical charge of an atom if all of its bonds to different atoms were fully ionic

A

Oxidation state

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64
Q

Charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally

A

Formal charge

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65
Q

Formula for formal charge

A

Ve - nonbonding e - (total e shared in bonds)/2

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66
Q

Bond formed via the sharing of electrons between two elements of similar EN

A

Covalent

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67
Q

Refers to whether a covalent bond is single, double, or triple

A

Bond order

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68
Q

Higher bond order corresponds to (higher/lower) bond strength, (higher/lower) bond energy, and (higher/lower) bond length.

A

Higher, higher, lower

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69
Q

How many pi bonds does a triple bond have

A

2 (one sigma, 2 pi)

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70
Q

(Nonpolar/polar) bonds have ∆EN of 0.5-1.7

A

Polar

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71
Q

A bond that has an ∆EN higher than 1.7 is?

A

Ionic

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72
Q

Bond wherein a single atom provides both bonding electrons, as most often found in Lewis acid base chemistry

A

Coordinate covalent bonds

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73
Q

H bonds are (strong/weak) attractions between EN atoms of one compound and H of another compound.

A

Weak

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74
Q

Bond formed via the transfer of one or more electrons from an element with relatively low IE to element with high EA

A

Ionic

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75
Q

Large, organized arrays of ions

A

Crystalline lattices

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76
Q

Formula for compounds that shows simplest whole-number ratio of atoms

A

Empirical

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77
Q

Formula for compounds that shows exact # of atoms of each element

A

Molecular

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78
Q

In (electronic geometry/molecular shape), bonded and lone pairs are treated the same, but lone pairs take up more space than bonds in (electronic geometry/molecular shape).

A

Electronic geometry, molecular shape

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79
Q

Possible shapes for an atom with 2 e- groups + bond angle

A

Linear; 180º

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80
Q

Possible shapes for an atom with 3 e- groups + bond angle

A

Trigonal planar, bent; 120º

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81
Q

Possible shapes for an atom with 4 e- groups + bond angle

A

Tetrahedral, trigonal pyramidal, bent; 109.5º

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82
Q

Possible shapes for an atom with 5 e- groups + bond angle

A

Trigonal bipyramidal, sawhorse, t-shaped, linear; 90º & 120º

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83
Q

Possible shapes for an atom with 6 e- groups + bond angle

A

Octahedral, square pyramidal, square planar, t-shaped, linear; 90º

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84
Q

Van de Waals forces includes which intermolecular forces?

A

London dispersion and dipole dipole

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85
Q

Type of force acting between atoms and molecules that are normally electrically symmetric and thus only create temporary dipoles

A

London dispersion forces

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86
Q

Attractive forces between the positive end of one polar molecule and the negative end of another polar molecule

A

Dipole dipole interactions

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87
Q

Strongest intermolecular force in which hydrogen is in proximity to a highly EN element (N,O,F)

A

Hydrogen bond

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88
Q

Why do we assume that gas particles experience no intermolecular forces?

A

Because gas particles are usually a long distance from one another

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89
Q

Ionic compounds conduct electricity when molten (liquid) or in aqueous solution (dissolved in water), because _______________.

A

Their icons are free to move from place to place

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90
Q

Boiling occurs at (one/multiple) temperatures, while evaporation occurs at (one/multiple) temperatures.

A

One, multiple

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91
Q

Which element has the highest melting point?

A

Carbon

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92
Q

Maximum concentration of a solute that can dissolve in a solvent at a given temperature

A

Solubility

93
Q

Units of solubility

A

mol/L; g/L

94
Q

Factors that affect solubility (4)

A

Concentration of solute, temperature of the system, pressure (for gasses), polarity of solute and solvent

95
Q

Describe the curve that represents the relationship between solubility and temperature of most solids

A

Exponential upward

96
Q

Solubility equilibrium is achieved when the rate of ________ is equal to the rate of _________.

A

Dissolution, precipitation

97
Q

Solubility equilibrium equation (how much a salt will dissolve)

A

Ksp = [A+][B-], usually has only one value for a given salt at a given temperature

98
Q

A Ksp that is greater than 1 indicates a (soluble/insoluble) salt.

A

Soluble (same applies when comparing Ksps of salts that produce same number of ions)

99
Q

Given CaF2 dissolved in water, eq. [Ca2+] was found to be 2 x 10^-4. What is the Ksp at 25ºC?

A

Ksp = [Ca2+][F-]^2
Ksp = [2 x 10^-4][4 x 10^-4]^2
Ksp = 3.7 x 10^-11

100
Q

Equilibrium constant for complex formation; is usually (greater/lower) than Ksp

A

Formation or stability constant (Kf), greater

101
Q

Effect of how the solubility of a compound in a solution that already contains one of the ions in the compound

A

Common ion effect

102
Q

When a central cation is bonded to the same ligand in multiple places to sequester toxic metals

A

Chelation

103
Q

Molarity of the solute at saturation

A

Molar solubility

104
Q

Cation bonded to at least one ligand which is the e- pair donor; held together by coordinate covalent bonds

A

Complex ions

105
Q

Formation of complex ions (increase/decrease) solubility

A

Increase

106
Q

Given CaF2 dissolved in water, eq. [Ca2+] was found to be 2 x 10^-4. What is the molar solubility of CaF2

A

CaF2: Ca^2+ = 1:1
also 2 x 10^-4

107
Q

For acids, those with more oxygen are called (-ous/-ic)

A

-Ic

108
Q

Average kinetic energy depends on (temperature/number of moles/both).

A

Temperature only (2 moles at 546K > 1 mol at 293K)

109
Q

(T/F): Average speed of particles depends on molar mass.

A

TRUE

110
Q

Theoretical gas whose molecules occupy negligible space and whose collisions are perfectly elastic

A

Ideal gas

111
Q

Ideal gases have (high/low) temperatures and (high/low) pressures.

A

High, low

112
Q

Volume of 1 mol of gas at STP

A

1 mol gas

113
Q

Ideal gas law formula

A

PV = nRT (R = 8.314 J/mol K)

114
Q

1 atm = ____ mmHg = ___ torr = ____ kPa = ___ psi

A

760, 760, 101.3, 14.7

115
Q

Formula for density of gases in terms of ideal gas law

A

PM/RT

116
Q

Combined gas law formula

A

P1V1/T1 = P2V2/T2

117
Q

In the combined gas law, what variable is constant?

A

n

118
Q

Avogadro’s principle formula

A

n1/V1 = n2/V2 (constant T&P)

119
Q

Boyle’s Law formula

A

P1V1 = P2V2 (constant n&T)

120
Q

Charles’s Law formula

A

V1/T1 = V2/T2 (constant n&P)

121
Q

Gay-Lussac’s Law formula

A

P1/T1 = P2/T2 (constant n&V)

122
Q

A 5 liter bulb and a 3 liter bulb are attached by a closed stopcock. The larger bulb contains helium gas at 540 torr, while the smaller bulb contains neon gas at 320 torr. What is the pressure in the larger bulb after the stopcock is opened and equilibrium is achieved?

A

Use Dalton’s law of partial pressures to get the individual gasses’ pressures at the total volume, then add: 457.5 torr

123
Q

To change a gas into a liquid, (reduce/increase) pressure and (reduce/increase) temperature.

A

Increase, reduce

124
Q

(T/F): Real gases deviate from ideal behavior at lower temperature and high pressure

A

TRUE.

125
Q

At MODERATELY high P, low V, low T, real gases occupy (less/more) volume than predicted by ideal gas law due to intermolecular attractions

A

Less

126
Q

At EXTREMELY high P, low V, low T, real gases will occupy (less/more) volume than predicted by ideal gas law because the particles occupy physical space

A

More

127
Q

Dalton’s Law (total pressure)

A

PT = PA + PB + PC + …

128
Q

Dalton’s Law (partial pressure)

A

PA = XaPT (X = mol fraction)

129
Q

Fruits are canned while hot because ______.

A

The pressure inside will decrease (lower than atmospheric) once the fruit cools.

130
Q

All diatomic gases

A

H2 N2 F2 O2 I2 C2 B2 (Have No Fear Of Ice Cold Beer)

131
Q

Henry’s Law

A

[A] = kH x PA (or [A]1/P1 = [A]2/P2 = kH)

[A] - concentration of A, kH = Henry’s constant, PA = partial pressure of A

132
Q

Loss of water (or a solvent) of crystallization from a hydrated or solvated salt to the atmosphere on exposure to air

A

Efflorescence

133
Q

Formula for root-mean-square speed of a gas

A

u_rms = sqrt(3RT/M)

134
Q

Graham’s law of diffusion formula

A

r1/r2 = √(M2/M1)

where r = rate, M = mass (note inverse rel.)

135
Q

Movements of gas from one compartment to another through a small opening under pressure

A

Effusion

136
Q

Pressure exerted by a vapor in thermodynamic equilibrium with the condensed phases (solid or liquid)

A

Vapor pressure

137
Q

Boiling point denotes the temperature at which ______________.

A

Vapor pressure equals the pressure of the gas above it (e.g. atmospheric)

138
Q

Vapor pressure is the pressure at which there is an equal amount of molecules ______ and ______.

A

Evaporating, condensing back to the liquid

139
Q

Having a low vapor pressure denotes a (lower/higher) BP.

A

Higher; need more energy to equate vapor pressure to atmospheric pressure

140
Q

The freezing point denotes the temperature at which _________.

A

Solid and liquid phases coexist in equilibrium (at 1 atm)

141
Q

In a phase diagram, an exponential curve suggests _________.

A

The transition from liquid to gas

142
Q

If you want to remove carbon dioxide from a solution, you ____ the solution.

A

Boil, carbon dioxide has very low BP

143
Q

Heating of solids to a high temperature for the purpose of removing volatile substances, oxidizing a portion of mass, or rendering them friable

A

Calcination

144
Q

Type of reaction where two or more reactants form one product

A

Combination

145
Q

Type of reaction where a single reactant breaks down

A

Decomposition

146
Q

Type of reaction that involves a fuel (usually an HC) and O2

A

Combustion

147
Q

Products of combustion

A

CO2 and H2O

148
Q

Type of reaction wherein an atom/ion in a compound is replaced by another atom/ion

A

Single-displacement

149
Q

Type of reaction where elements from two compounds swap places

A

Double-displacement (metathesis)

150
Q

Type of reaction wherein acid + base –> salt + H2O

A

Neutralization

151
Q

Neutralization is a type of what reaction

A

Double replacement

152
Q

Avogadro’s number

A

6.022 x 10^23 = 1 mol

153
Q

Represents the fraction of collisions that have enough energy to overcome the activation barrier

A

Arrhenius equation (k = A x e^(-Ea/RT))

154
Q

Rate law

A

rate = k [A]^x[B]^y

155
Q

Kinetic products are (higher/lower) in free energy than thermodynamic products can form in (higher/lower) temperatures

A

Higher, lower

156
Q

Thermodynamic products are (faster/slower) and (more/less) spontaneous

A

Slower, more

157
Q

Formula for heat energy change (∆H)

A

∆H = ∆H (bonds broken in reactants) - ∆H (bonds made in products)

158
Q

Energy required to reach transition state

A

Activation energy

159
Q

The activation energy of an exergonic reaction is (less/more) than its reverse.

A

Less

160
Q

The change in internal energy ∆U is given by

A

∆U = q + w (+w for work done on system, -w for work done on surroundings)

161
Q

Change in Gibbs free energy is defined as

A

∆G = ∆H - T∆S

162
Q

When ∆G is negative, a process will proceed (spontaneously/unspontaneously) and is (exergonic/endergonic)

A

Spontaneously, exergonic

163
Q

When ∆G = ____, system is in equilibrium and the concentrations of products and reactants is constant

A

0

164
Q

Gibbs free energy in terms of activation energy is

A

Ea - Ea(rev)

165
Q

Le Chatelier’s Principle states that if a stress is applied to a system, the system shifts to _______ the applied stress.

A

Relieve

166
Q

A measure of the degree to which energy has been spread throughout a system or between a system and its surroundings.

A

Entropy

167
Q

Increasing pressure shifts the equilibrium towards the side with __________.

A

Fewer total molecules (A molecule is a mole on the equation (# terms x coefficients))

168
Q

Increasing the reactants shifts the equilibrium to the (left/right).

A

Right

169
Q

If the system’s reaction enthalpy is ∆H = 250 kJ, and temperature is decreased in the system, the equilibrium shifts to the (left/right)

A

Left

170
Q

Raising the temperature of a solution of a gas in a liquid (increases/decreases) solubility thereby (increasing/decreasing) equilibrium concentration.

A

Decreases, decreases (because more entropy & freedom)

171
Q

Solubility of a gas in liquid (increases/decreases) with increasing temperature.

A

Decreases

172
Q

According to Henry’s Law, the solubility of gases (increases/decreases) as the partial pressure of the gas above a solution increases.

A

Increases

173
Q

Process where a solute in gaseous, liquid, or solid phase dissolves in a solvent to form a solution

A

Dissolution

174
Q

According to the modern theory, when sodium chloride is dissolved in water, it forms ________ sodium and chloride ions.

A

Hydrated

175
Q

Are ethyl and methyl miscible or immiscible in water? Why?

A

Miscible; attraction of OH to water

176
Q

In dissolution, the breaking of the solute-solute bonds & solvent-solvent bonds is (endothermic/exothermic), and the bonding of the solute with the solvent is (endothermic/exothermic).

A

Endothermic, exothermic

177
Q

In an endothermic dissolution, (less/more) energy is released when the solute-solvent bond compared to the energy needed to break the solute particles.

A

Less

178
Q

In endothermic reactions, solubility is (increased/decreased) when the temperature is increased.

A

Increased; promote dissolution with additional heat

179
Q

In exothermic reactions, solubility is (increased/decreased) when the temperature is increased

A

Decreased; inhibit dissolution to counter more heat loss with additional heat

180
Q

The dissolution of gasses is (exothermic/endothermic).

A

Exothermic; increase in temp –> increase in KE –> weaker bonds

181
Q

Equilibrium constant or reaction quotient formula

A

Keq = ([C]^c [D]^d)/([A]^a [B]^b); exponents are the coefficient

182
Q

Reaction quotient and equilibrium constant have the same formula, but (reaction quotient/equilibrium constant) will fluctuate as the system reacts, whereas the (reaction quotient/equilibrium constant) is based on equilibrium concentrations

A

Reaction quotient, equilibrium constant

183
Q

(T/F): Q < Keq means a ∆G < 0

A

TRUE; whereas an equal Q = Keq means equilibrium

184
Q

Basis of Arrhenius definition of acids & bases

A

H+/OH- in H2O

185
Q

Basis of Bronsted & Lowry definition of acids & bases

A

H+ donor/acceptor

186
Q

Basis of Lewis definition of acids & bases

A

E pair donor/acceptor

187
Q

Species that can behave as an acid or base

A

Amphoteric base

188
Q

Acid with multiple ionizable H atoms

A

Polyprotic acid

189
Q

Strong acids have (high/low) solubility in water.

A

High

190
Q

Formula for pH

A

pH = -log [H+], so [H+] = 10^-PH

191
Q

Formula for acid dissociation constant

A

Ka = [H+][A-]/[HA]

192
Q

What are the seven strong bases?

A

LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2, Mg(OH)2

193
Q

Is ammonia a strong or weak base?

A

Weak; thus, NH4OH is also weak

194
Q

What are the seven strong acids?

A

HCl, HBr, HI, H2SO4, HClO4, HClO3, HNO3

195
Q

Is the bisulfate ion (HSO4-, not found as molecule) a strong or weak acid?

A

Weak

196
Q

A (strong/weak) acid + a (strong/weak) base will give an alkaline solution.

A

Weak, strong

197
Q

In every solution, pOH + pH = 14 because the product of [H+] and [OH-] must always equal the equilibrium constant for the ionization of water, which is ________.

A

1 x 10^-14

198
Q

Formula for pH

A

-log ([H+])

199
Q

Formula for pOH

A

-log ([OH-])

200
Q

If I dilute 5 mL of 0.15 M NaCl to a final volume of 5 L, what is the final concentration of NaCl?

A

0.00015 M

201
Q

Formula for radioactive decay

A

N(t) = N0 (1/2)^(t/t.5)

where N0 - initial amt, Nt = final amt, t = time, t.5 = half life

202
Q

Equivalent mass is the mass of an acid that yields _________ or mass of a base that reacts with ________.

A

1 mole of H+ (thus is equal to #H or OH of the compound)

203
Q

GEW (gram equivalent weight) formula (basically equivalent weight

A

Molar mass / mol H+ or e-

204
Q

Equivalents formula

A

Mass of compound/GEW

205
Q

Normality formula

A

Eq/L (Eq - equivalents of solute, V - volume of solvent in liters) or M*Eq (M - molarity)

For acids, Eq (n) is the # of H+ available from a formula unit

206
Q

Can donate or accept multiple equivalents

A

Polyvalent

207
Q

A student wants to prepare 500 ml of 8 N H2SO4 from a solution which is 8 molar. What volume of the concentrated acid is needed?

A

250 mL (8N H2SO4 = 4 M H2SO4 * 2 Eq)

208
Q

Law of equivalence (used to neutralize the whole solution)

A

N1V1 = N2V2 where 1 is an acid and 2 is a base, or used to know after dilution

209
Q

Molarity formulas

A

mol/L = normality/mol H+ or e-

210
Q

Colligative properties depend upon the _________ of solute molecules or ions, but not upon the identity of the solute.

A

Concentration

211
Q

Adding a solute will (lower/increase) equilibrium vapor pressure. Why?

A

Lower, because less solvent molecules will be able to come to the surface and evaporate

212
Q

Formula for new vapor pressure given addition of a solute

A

Raoult’s Law: Psolvent = Xsolvent P0solvent

213
Q

Adding a solute will (lower/increase) boiling point. Why?

A

Elevate, because it will require more energy to reach the same vapor pressure needed to equate the external surroundings (given more solute = lower vapor pressure)

214
Q

Formula for new boiling/freezing point given addition of a solute

A

∆T = iKbm or ∆T = iKfm

i - van’t Hoff factor (# of particles into which solute dissociates), m - molality, Kb - molal boiling point constant (for water, 0.5121ºC/m), Kf - molal freezing point constant (for water, -1.86ºC/m)

215
Q

Adding a solute will (lower/increase) freezing point. Why?

A

Lower, because the addition of solute makes the mixture more disordered (and thus with higher entropy) and requires lower temperature to make ordered (frozen)

216
Q

Pressure required to achieve osmotic equilibrium

A

Osmotic pressure (by adding this pressure to a solution separated to a pure solvent by a semipermeable membrane, the escaping tendency of the solution in spite of high solute concentration will be raised to the level of the pure solvent)

217
Q

Number of individual particles in solution

A

Osmolarity (1M NaCl = 2 osmol/L)

218
Q

T/F: Osmotic pressure is the pressure required to sustain osmosis.

A

FALSE: stop osmosis

219
Q

Formula for osmotic pressure

A

II = nRT/V or MRT (T is in K, M is total ionic concentration of solute, i.e. 2x molarity for NaCl in sol. as it dissociates)

220
Q

Acidity is increased when the _________.

A

H+ is more available

221
Q

Carboxylic acids have greater acidity compared to alcohols because of __________.

A

Electron delocalization in the carboxylate ion

222
Q

Things that affect acidity include:

A

Availability of proton, stability of conjugate base

223
Q

Higher s orbital character is more (acidic/basic).

A

Acidic because more s character (more unsaturated) means lower energy (stable)

224
Q

Things that affect basicity include:

A

Availability of lone pair, stability of conjugate ion

225
Q

If resonance of the lone pair is possible, it is (less/more) basic.

A

Less

226
Q

EWG increase (basicity/acidity) and decrease (basicity/acidity).

A

Acidity, basicity

227
Q

Higher negative charge indicates more (acidity/basicity) while more positive charge indicates more (acidity/basicity).

A

Basicity, acidity

228
Q

What are the most basic amines? Why?

A

Secondary amines. While primary amines are more solvated due to more hydrogen bonds, it does not have EDG (alkyl groups) which increase basicity.

229
Q
A

NMAT 2025 (11 decks)

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