The first law of thermodynamics is based on conservation of
A mass
B energy
C volume
D entropy
First law states energy can neither be created nor destroyed, only converted from one form to another.
Internal energy (ΔU) of an ideal gas depends only on
A pressure
B volume
C temperature
D amount of catalyst
For an ideal gas, internal energy is a function of temperature only because intermolecular forces are negligible.
Work done by a gas during expansion against constant external pressure is
A W = +PΔV
B W = −PΔV
C W = +ΔP•V
D W = −ΔP•V
Chemistry sign convention: expansion work done by system is negative (system loses energy).
Which process has ΔU = 0 for an ideal gas
A isothermal process
B adiabatic process
C isobaric process
D isochoric process
For an ideal gas, ΔU depends only on temperature. In isothermal process, ΔT = 0 → ΔU = 0.
Heat absorbed at constant pressure is equal to change in
A internal energy
B enthalpy
C entropy
D Gibbs energy
At constant pressure, qₚ = ΔH.
Relation between enthalpy and internal energy is
A H = U − PV
B H = U + PV
C H = U + RT
D H = U − RT
Enthalpy is defined as H = U + PV.
For an ideal gas reaction, ΔH and ΔU are related by
A ΔH = ΔU
B ΔH = ΔU + ΔnRT
C ΔH = ΔU − ΔnRT
D ΔH = ΔU + PΔV only
For gaseous reactions, ΔH differs from ΔU due to PV term; Δn = moles(gas products) − moles(gas reactants).
The second law of thermodynamics states that entropy of an isolated system
A always decreases
B remains constant always
C always increases for spontaneous process
D becomes zero always
Entropy increases in spontaneous processes for an isolated system.
For a reversible process, change in entropy is given by
A ΔS = q/T
B ΔS = qrev/T
C ΔS = qirr/T
D ΔS = q×T
Entropy is a state function; its change is calculated using reversible heat flow.
Unit of entropy is
A J/mol
B J/mol•K
C kJ/mol
D J/K only
Entropy change per mole per Kelvin has unit J mol⁻¹ K⁻¹.
Gibbs free energy is defined as
A G = H + TS
B G = H − TS
C G = U + PV
D G = U − PV
Gibbs free energy relates enthalpy and entropy at temperature T.
A process is spontaneous at constant T and P if
A ΔH < 0 only
B ΔS > 0 only
C ΔG < 0
D ΔG > 0
At constant temperature and pressure, spontaneity criterion is ΔG negative.
At equilibrium, Gibbs free energy change is
A positive
B negative
C zero
D infinite
At equilibrium, ΔG = 0 because forward and reverse tendencies balance.
If ΔG is positive, the reaction is
A spontaneous
B non-spontaneous
C at equilibrium
D always fast
ΔG > 0 indicates reaction is not spontaneous in forward direction.
The third law of thermodynamics states that entropy of a perfect crystal at 0 K is
A maximum
B minimum but not zero
C zero
D infinite
Perfectly ordered crystal has only one microstate at 0 K → entropy zero.
The equilibrium constant Kc is defined using
A partial pressures
B concentrations
C masses
D volumes only
Kc is written in terms of molar concentrations for equilibrium mixture.
For a reaction aA + bB ⇌ cC + dD, Kc is
A [A]^a[B]^b / [C]^c[D]^d
B [C]^c[D]^d / [A]^a[B]^b
C [A+B]/[C+D]
D [C+D]/[A+B]
Products in numerator and reactants in denominator, each raised to stoichiometric power.
Kp is expressed in terms of
A concentrations
B partial pressures
C moles only
D mass only
Kp uses partial pressures of gaseous species.
Relation between Kp and Kc is
A Kp = Kc
B Kp = Kc(RT)^Δn
C Kp = Kc/(RT)^Δn
D Kp = Kc + RT
Δn = gaseous moles products − gaseous moles reactants.
If Δn = 0, then
A Kp = Kc
B Kp > Kc always
C Kp < Kc always
D Kp = 0
When Δn = 0, (RT)^Δn = 1.
Reaction quotient Q is
A always equal to K
B evaluated using initial concentrations/pressures
C evaluated only at equilibrium
D independent of concentrations
Q has same expression as K but uses any moment values; compare Q with K to predict direction.
If Q < K, reaction proceeds
A backward
B forward
C stops immediately
D becomes impossible
When Q is smaller than K, more products must form to reach equilibrium.
If Q > K, reaction proceeds
A forward
B backward
C remains at equilibrium
D no effect
If Q exceeds K, system has excess products and shifts backward.
If K is very large, equilibrium mixture contains mainly
A reactants
B products
C equal reactants and products
D catalyst only
Large K means products are favored at equilibrium.
According to Le Chatelier’s principle, increasing pressure shifts equilibrium towards
A side with more moles of gas
B side with fewer moles of gas
C side with more solids
D side with more liquids
System counteracts pressure increase by reducing total gaseous moles.
pH is defined as
A log[H⁺]
B −log[H⁺]
C −log[OH⁻]
D log[OH⁻]
pH = −log₁₀[H⁺].
At 25°C, ionic product of water (Kw) is
A 10⁻⁷
B 10⁻¹⁴
C 10⁻⁵
D 10⁻²
Kw = [H⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C.
At 25°C, neutral water has pH
A 5
B 6
C 7
D 8
In pure water, [H⁺]=[OH⁻]=10⁻⁷ → pH = 7.
A strong acid is one which
A partially ionizes
B completely ionizes
C never ionizes
D depends only on concentration
Strong acids dissociate almost completely in water.
For a weak acid HA, Ka is
A [HA]/[H⁺][A⁻]
B [H⁺][A⁻]/[HA]
C [H⁺]/[HA]
D [A⁻]/[HA]
Ka is equilibrium constant for acid dissociation: HA ⇌ H⁺ + A⁻.
Larger Ka implies acid is
A weaker
B stronger
C neutral
D amphoteric only
Higher Ka indicates greater extent of ionization → stronger acid.
pKa is defined as
A log Ka
B −log Ka
C log Kb
D −log Kb
pKa = −log₁₀(Ka).
A buffer solution resists change in
A volume
B temperature
C pH
D density
Buffer maintains nearly constant pH on addition of small acid/base.
A common buffer consists of
A strong acid and strong base
B weak acid and its salt
C strong acid and its salt
D strong base and water only
Example: CH₃COOH + CH₃COONa.
The pH of acidic buffer is given by Henderson equation
A pH = pKa + log([salt]/[acid])
B pH = pKa − log([salt]/[acid])
C pH = pKb + log([base]/[salt])
D pH = −log Ka
Henderson–Hasselbalch equation for acid buffer.
Solubility product Ksp is defined for
A highly soluble salts
B sparingly soluble salts
C gases
D metals only
Ksp measures equilibrium of slightly soluble ionic compounds.
For AgCl(s) ⇌ Ag⁺ + Cl⁻, Ksp is
A [AgCl]
B [Ag⁺] + [Cl⁻]
C [Ag⁺][Cl⁻]
D [Ag⁺]/[Cl⁻]
Ksp equals product of ion concentrations, each to power of stoichiometry.
Common ion effect causes solubility to
A increase
B decrease
C remain same
D become infinite
Adding common ion shifts equilibrium backward, reducing solubility.
In precipitation, ionic product (Qsp) compared to Ksp must be
A Qsp < Ksp
B Qsp = Ksp
C Qsp > Ksp
D Qsp = 0
When ionic product exceeds Ksp, solution is supersaturated and precipitate forms
Phase is a region which is
A always solid
B homogeneous and physically distinct
C always liquid
D always gas
A phase has uniform composition and properties separated by boundaries.
The number of phases in a system of ice + water is
A 1
B 2
C 3
D 4
Ice (solid) and water (liquid) are two distinct phases.
Gibbs phase rule is
A F = C − P + 2
B F = P − C + 2
C F = C + P + 2
D F = C − P − 2
Phase rule relates degrees of freedom (F), components (C), and phases (P).
For a one-component system with two phases, degrees of freedom is
A 0
B 1
C 2
D 3
For C=1, P=2 → F = 1 − 2 + 2 = 1.
Triple point is the point where
A only solid exists
B only liquid exists
C solid, liquid, and gas coexist in equilibrium
D only gas exists
Triple point corresponds to three-phase equilibrium.
At triple point in a one-component system, degrees of freedom is
A 0
B 1
C 2
D 3
For C=1, P=3 → F = 1 − 3 + 2 = 0 (invariant).
The line separating liquid and vapor region in phase diagram is
A fusion curve
B vaporization curve
C sublimation curve
D critical curve
It represents equilibrium between liquid and vapor.
Critical point is the point where
A solid becomes liquid
B liquid becomes solid
C liquid and gas become indistinguishable
D gas becomes solid directly
At critical point, no boundary between liquid and gas.
In CO₂ phase diagram, solid CO₂ directly changes to gas at 1 atm by
A melting
B vaporization
C sublimation
D deposition
Dry ice sublimes at 1 atm without forming liquid CO₂.
The equilibrium between solid and vapor is represented by
A fusion curve
B sublimation curve
C vaporization curve
D critical curve
Sublimation curve shows solid ⇌ vapor equilibrium.
Standard Gibbs free energy and equilibrium constant relation is
A ΔG° = RT ln K
B ΔG° = −RT ln K
C ΔG° = −RT/K
D ΔG° = RT/K
This fundamental relation connects thermodynamics with equilibrium.