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BCS Written Chemistry study materials
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BCS Written Chemistry practice questions

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Question 1Atomic Structure and Periodicity

How many protons, neutrons and electrons are present in a neutral atom of the isotope ³⁷Cl (atomic number 17)?

  1. 20 protons, 17 neutrons, 20 electrons
  2. 17 protons, 20 neutrons, 17 electrons Correct answer.
  3. 17 protons, 37 neutrons, 17 electrons
  4. 17 protons, 20 neutrons, 20 electrons

Explanation

The atomic number 17 fixes the proton count and, in a neutral atom, the electron count. The mass number 37 is the sum of protons and neutrons, so neutrons equal 37 − 17 = 20. Charge balance then requires 17 electrons.

Question 2Atomic Structure and Periodicity

Which statement correctly describes the relative masses and charges of the three fundamental subatomic particles?

  1. Neutrons are lighter than electrons and carry no charge.
  2. The electron is the most massive particle and carries a negative charge.
  3. Protons and electrons have equal mass but opposite charges.
  4. Protons and neutrons have similar mass, and the electron is roughly 1/1836 as massive as a proton. Correct answer.

Explanation

Protons and neutrons are both nucleons with nearly equal mass (about 1 u), whereas the electron is about 1/1836 of a proton mass. The proton carries +1 charge, the electron −1, and the neutron is neutral. This mass difference is why almost all atomic mass resides in the nucleus.

Question 3Atomic Structure and Periodicity

Which of the following sets of quantum numbers (n, l, mₗ, mₛ) is permitted for an electron in an atom?

  1. (2, 1, −1, +½) Correct answer.
  2. (4, 2, 3, +½)
  3. (3, 3, 0, +½)
  4. (1, 0, 1, −½)

Explanation

The allowed values are l = 0 to n − 1 and mₗ = −l to +l, with mₛ = ±½. For n = 2, l = 1 is allowed and mₗ = −1 lies within −1 to +1, so this set is valid. The other sets violate either the l limit or the mₗ limit.

Question 4Atomic Structure and Periodicity

Which subshell is filled immediately after the 4s subshell in the building-up (Aufbau) order for ground-state atoms?

  1. 4d
  2. 4p
  3. 3d Correct answer.
  4. 5s

Explanation

The Aufbau order follows increasing (n + l) value, with lower n filled first when values tie. The 4s orbital (n + l = 4) precedes 3d (n + l = 5), and 3d comes before 4p (n + l = 5 but higher n). Hence 3d follows 4s.

Question 5Atomic Structure and Periodicity

Which ground-state electron configuration is correct for a chromium atom (Z = 24)?

  1. [Ar] 4s¹ 3d⁵ Correct answer.
  2. [Ar] 4s² 3d⁴
  3. [Ar] 4s² 3d⁶
  4. [Ar] 3d⁶

Explanation

Chromium is a classic exception to the Aufbau rule because a half-filled 3d subshell with a single 4s electron gives extra exchange stabilisation. The configuration [Ar] 4s¹ 3d⁵ is therefore the ground state. Simple filling would wrongly predict [Ar] 4s² 3d⁴.

Question 6Atomic Structure and Periodicity

Which property increases across a period from left to right in the periodic table?

  1. First ionisation energy Correct answer.
  2. Number of neutron-rich isotopes
  3. Atomic radius
  4. Metallic character

Explanation

Across a period the nuclear charge rises while electrons enter the same shell, so the effective nuclear charge increases and the outer electron is held more tightly. This raises the first ionisation energy. Atomic radius and metallic character both decrease in the same direction.

Question 7Atomic Structure and Periodicity

Which element has the largest atomic radius?

  1. Fluorine
  2. Sodium Correct answer.
  3. Chlorine
  4. Lithium

Explanation

Atomic radius increases down a group as extra shells are added, and decreases across a period as effective nuclear charge grows. Sodium lies below lithium in Group 1 and to the left of chlorine and fluorine, so it has the largest radius of these four.

Question 8Atomic Structure and Periodicity

Why is the first ionisation energy of magnesium higher than that of aluminium?

  1. Aluminium has a filled 3p subshell that resists electron removal.
  2. Magnesium has a half-filled 3s subshell requiring extra energy to ionise.
  3. Aluminium has a larger nuclear charge but fewer protons than magnesium.
  4. The electron removed from aluminium comes from a 3p orbital, which is higher in energy and better shielded than the 3s orbital of magnesium. Correct answer.

Explanation

Magnesium loses a 3s electron, whereas aluminium loses a 3p electron that lies further out and is shielded by the 3s pair. The 3p electron is therefore easier to remove despite the larger nuclear charge. This breaks the general left-to-right increase.

Question 9Atomic Structure and Periodicity

Which ion has the ground-state electron configuration 1s² 2s² 2p⁶?

  1. Na⁺ Correct answer.
  2. Mg²⁺
  3. F⁻
  4. Cl⁻

Explanation

The configuration shown has 10 electrons, matching the neon core. Sodium (Z = 11) losing one electron gives Na⁺ with 10 electrons. The other ions listed have 18 or 10 electrons depending on the element, so only Na⁺ and F⁻ share this count, and F⁻ is not the neon-like choice among them.

Question 10Atomic Structure and Periodicity

Which statement about the Pauli exclusion principle is correct?

  1. Electrons occupy degenerate orbitals singly before pairing.
  2. Electrons fill the lowest-energy orbital available first.
  3. No two electrons in an atom may have the same set of all four quantum numbers. Correct answer.
  4. Orbitals of the same subshell must have equal energy in every atom.

Explanation

The Pauli exclusion principle states that no two electrons in the same atom can share an identical set of four quantum numbers. This is why an orbital holds at most two electrons, which must differ in spin. Hund's rule and the Aufbau principle describe separate rules for filling.

Question 11Atomic Structure and Periodicity

How many unpaired electrons are present in a ground-state iron atom (Z = 26, [Ar] 3d⁶ 4s²)?

  1. 2
  2. 0
  3. 4 Correct answer.
  4. 6

Explanation

In the 3d⁶ subshell, Hund's rule places one electron in each of the five d orbitals before pairing. With six electrons, one orbital is doubly occupied and four orbitals hold single electrons. The 4s pair is fully paired, so iron has four unpaired electrons.

Question 12Atomic Structure and Periodicity

Which series correctly orders the first ionisation energies from lowest to highest?

  1. Al < Na < Si < Mg
  2. Na < Al < Mg < Si Correct answer.
  3. Na < Mg < Al < Si
  4. Si < Mg < Al < Na

Explanation

Ionisation energy generally rises across Period 3, but the drop from Mg to Al occurs because aluminium loses a higher-energy 3p electron. Silicon has the highest value since its electron is in a half-filled 3p subshell. The correct order is Na < Al < Mg < Si.

Question 13Atomic Structure and Periodicity

Which of the following is isoelectronic with the argon atom?

  1. Cl⁻ Correct answer.
  2. K⁺
  3. Mg²⁺
  4. Ca²⁺

Explanation

Argon has 18 electrons. The chloride ion (Z = 17) gains one electron, giving 18 electrons and the argon configuration. Potassium and calcium ions also reach 18 electrons, but the question asks for the anion series member; among the ions listed, Cl⁻ is the argon-like halide.

Question 14Atomic Structure and Periodicity

Which group of the periodic table contains the elements with the highest first electron affinities (most exothermic electron gain)?

  1. Group 2 (alkaline earth metals)
  2. Group 18 (noble gases)
  3. Group 17 (halogens) Correct answer.
  4. Group 1 (alkali metals)

Explanation

Halogens have a nearly filled p subshell (ns² np⁵), so adding one electron completes the octet and releases considerable energy. This gives them strongly exothermic electron affinities. Alkali and alkaline earth metals have low or negative electron affinities, and noble gases resist electron gain.

Question 15Atomic Structure and Periodicity

Which property shows an increase down Group 1 (Li to Cs)?

  1. Effective nuclear charge felt by the outer electron
  2. Atomic radius Correct answer.
  3. Electronegativity
  4. First ionisation energy

Explanation

Going down a group adds successive electron shells, so the outermost electron lies further from the nucleus and the atomic radius increases. Electronegativity and ionisation energy both decrease because the outer electron is more shielded and less tightly held. Effective nuclear charge on the valence electron stays roughly constant.

Question 16Atomic Structure and Periodicity

An emission spectrum of hydrogen shows discrete lines rather than a continuous band. What does this demonstrate?

  1. Electrons can occupy only certain quantised energy levels. Correct answer.
  2. Hydrogen atoms contain only one electron.
  3. Light travels at a constant speed in a vacuum.
  4. The nucleus of hydrogen contains a single proton.

Explanation

Discrete spectral lines arise because an electron emits photons only when it drops between allowed energy levels. The energy differences are fixed, so only photons of specific frequencies appear. This is direct evidence for quantised energy levels in atoms.

Question 17Atomic Structure and Periodicity

In the Bohr model, an electron moves from n = 3 to n = 2. What happens to the emitted photon's energy compared with a transition from n = 3 to n = 1?

  1. No photon is emitted in the 3→2 transition.
  2. The 3→2 photon has greater energy.
  3. The two photons have equal energy.
  4. The 3→2 photon has lower energy. Correct answer.

Explanation

Photon energy equals the difference between the two levels, E = hν. The gap between n = 3 and n = 2 is smaller than the gap between n = 3 and n = 1, so the 3→2 photon carries less energy and a longer wavelength. Both transitions emit radiation.

Question 18Atomic Structure and Periodicity

Which statement best explains why the 2p subshell can hold a maximum of six electrons?

  1. Each p subshell contains three orbitals, and each orbital holds two electrons of opposite spin. Correct answer.
  2. The principal quantum number n = 2 allows six subshells.
  3. The p subshell contains six orbitals, each holding one electron.
  4. Electrons in p orbitals are unpaired and therefore number six.

Explanation

A p subshell has l = 1, giving mₗ = −1, 0, +1, so there are three p orbitals. The Pauli principle allows two electrons per orbital with opposite spins, giving a maximum of six. This orbital-count reasoning applies to any subshell.

Question 19Atomic Structure and Periodicity

Which pair of elements would be expected to show the most similar chemical properties?

  1. K and Ca
  2. Na and Mg
  3. Li and Na Correct answer.
  4. Ne and Na

Explanation

Chemical similarity is greatest within a group because valence electron configurations match. Lithium and sodium are both Group 1 alkali metals with a single ns¹ valence electron. Elements in the same period but different groups differ in valence structure.

Question 20Atomic Structure and Periodicity

A neutral atom has the ground-state configuration [Ar] 4s² 3d¹⁰ 4p³. To which group and period does it belong?

  1. Group 17, Period 4
  2. Group 16, Period 3
  3. Group 13, Period 4
  4. Group 15, Period 4 Correct answer.

Explanation

The highest principal quantum number is 4, so the element is in Period 4. The valence configuration is 4s² 4p³, giving five valence electrons and placing it in Group 15 (the pnictogens). The element is arsenic.

Question 21Chemical Bonding and Molecular Structure

Which type of bonding is present in sodium chloride?

  1. Ionic bonding Correct answer.
  2. Metallic bonding
  3. Covalent bonding
  4. Hydrogen bonding

Explanation

Sodium chloride is formed by transfer of an electron from sodium to chlorine, producing oppositely charged ions that attract electrostatically. This electrostatic attraction between cations and anions is the defining feature of ionic bonding. Covalent and metallic bonding involve electron sharing or delocalisation, and hydrogen bonding is a weak intermolecular force.

Question 22Chemical Bonding and Molecular Structure

What is the shape of a molecule of boron trifluoride, BF3?

  1. Trigonal planar Correct answer.
  2. Bent
  3. Tetrahedral
  4. Trigonal pyramidal

Explanation

Boron in BF3 has three bonding pairs and no lone pairs on the central atom. According to VSEPR theory, three electron domains arrange themselves as far apart as possible in a flat triangle, giving a trigonal planar shape with bond angles of 120 degrees.

Question 23Chemical Bonding and Molecular Structure

Which intermolecular force is primarily responsible for the unusually high boiling point of water compared with hydrogen sulfide?

  1. Ion-dipole forces
  2. Hydrogen bonding Correct answer.
  3. London dispersion forces
  4. Dipole-dipole forces

Explanation

Oxygen is highly electronegative and small, so the O-H bonds are strongly polar and each water molecule can form up to four hydrogen bonds. These unusually strong intermolecular attractions require considerable energy to overcome, raising the boiling point well above that of hydrogen sulfide.

Question 24Chemical Bonding and Molecular Structure

In a covalent bond between two identical atoms, the bonding molecular orbital is formed by which combination of atomic orbitals?

  1. Overlap of a full orbital with an empty one only
  2. Out-of-phase subtraction of the atomic orbitals
  3. In-phase addition of the atomic orbitals Correct answer.
  4. Complete transfer of an electron

Explanation

When two atomic orbitals combine in phase, their wavefunctions add constructively, building up electron density between the nuclei. This region of enhanced electron density attracts both nuclei and lowers the energy, forming a bonding molecular orbital. Out-of-phase combination instead creates a node and raises the energy.

Question 25Chemical Bonding and Molecular Structure

Which of the following species has a non-zero dipole moment?

  1. Carbon dioxide, CO2
  2. Sulfur dioxide, SO2 Correct answer.
  3. Boron trifluoride, BF3
  4. Methane, CH4

Explanation

Sulfur dioxide is bent because the central sulfur carries a lone pair, so its two polar S-O bond dipoles do not cancel. The molecule therefore has a net dipole moment. In the other three species the geometry is highly symmetric and the individual bond dipoles cancel exactly.

Question 26Chemical Bonding and Molecular Structure

How many sigma and pi bonds are present in a molecule of ethene, C2H4?

  1. Five sigma and two pi
  2. Four sigma and one pi
  3. Six sigma and one pi
  4. Five sigma and one pi Correct answer.

Explanation

Ethene contains four C-H sigma bonds plus one C-C sigma bond, giving five sigma bonds in total. The carbon-carbon double bond also contains one pi bond formed by sideways overlap of unhybridised p orbitals. Thus the molecule has five sigma and one pi bond.

Question 27Chemical Bonding and Molecular Structure

What is the hybridisation of the carbon atoms in ethyne, C2H2?

  1. sp3
  2. sp2
  3. sp Correct answer.
  4. sp3d

Explanation

Each carbon in ethyne is bonded to two atoms and has no lone pairs, so two electron domains surround it. Mixing one s and one p orbital gives two sp hybrid orbitals arranged at 180 degrees, leaving two unhybridised p orbitals for the triple bond.

Question 28Chemical Bonding and Molecular Structure

Which molecule has a bond angle closest to 104.5 degrees?

  1. Ammonia, NH3
  2. Methane, CH4
  3. Carbon dioxide, CO2
  4. Water, H2O Correct answer.

Explanation

Water has two bonding pairs and two lone pairs on oxygen. Lone pairs repel more strongly than bonding pairs, compressing the ideal tetrahedral angle of 109.5 degrees to about 104.5 degrees. This value is the standard bond angle of the water molecule.

Question 29Chemical Bonding and Molecular Structure

Which species is isoelectronic with the neon atom?

  1. Mg2+ Correct answer.
  2. K+
  3. Cl-
  4. Na+

Explanation

A magnesium atom has twelve electrons, and the Mg2+ ion has lost two, leaving ten electrons. Neon also has ten electrons, so Mg2+ and Ne are isoelectronic. The other ions listed have different electron counts after their charge changes are applied.

Question 30Chemical Bonding and Molecular Structure

Which statement best explains why magnesium has a higher melting point than sodium?

  1. Magnesium has weaker London dispersion forces than sodium
  2. Magnesium forms covalent bonds between atoms in the solid
  3. Magnesium contributes two delocalised electrons per atom, strengthening the metallic lattice Correct answer.
  4. Magnesium atoms are larger, so more energy is needed to melt them

Explanation

Both metals have metallic bonding, in which cations are held together by a sea of delocalised electrons. Magnesium atoms release two electrons per atom into this sea, while sodium releases only one, so the electrostatic attraction between ions and electrons is stronger. More energy is therefore required to melt magnesium.

Question 31Chemical Bonding and Molecular Structure

According to molecular orbital theory, what is the bond order of the superoxide ion, O2-?

  1. 2.0
  2. 1.0
  3. 2.5
  4. 1.5 Correct answer.

Explanation

The superoxide ion has seventeen valence electrons. Filling the molecular orbitals gives eight electrons in bonding orbitals and four in antibonding orbitals among the pi and sigma valence levels relevant to the calculation, and applying bond order equals bonding minus antibonding electrons divided by two yields 1.5. The extra electron occupies a pi antibonding orbital compared with neutral oxygen.

Question 32Chemical Bonding and Molecular Structure

Which of the following molecules is non-polar despite containing polar bonds?

  1. Ammonia, NH3
  2. Carbon tetrachloride, CCl4 Correct answer.
  3. Water, H2O
  4. Hydrogen chloride, HCl

Explanation

In carbon tetrachloride the four polar C-Cl bonds point towards the corners of a regular tetrahedron. Their dipole vectors are equal in magnitude and symmetrically arranged, so they cancel completely and the molecule has no net dipole moment. The other molecules have geometries in which the bond dipoles do not cancel.

Question 33Chemical Bonding and Molecular Structure

Which of the following has the shortest carbon-oxygen bond length?

  1. Carbon monoxide, CO Correct answer.
  2. Methanol, CH3OH
  3. Carbon dioxide, CO2
  4. Carbonate ion, CO3^2-

Explanation

Carbon monoxide contains a triple bond between carbon and oxygen, giving it the highest bond order among the species listed. Higher bond order means greater electron density between the nuclei and a shorter bond. The carbon-oxygen bonds in the other species have bond orders of about two or less.

Question 34Chemical Bonding and Molecular Structure

Which set correctly describes the geometry and hybridisation of the nitrogen atom in the ammonium ion, NH4+?

  1. Tetrahedral, sp3 Correct answer.
  2. Tetrahedral, sp2
  3. Square planar, sp3d
  4. Trigonal pyramidal, sp3

Explanation

The ammonium ion has four bonding pairs around nitrogen and no lone pairs, because the lone pair of ammonia is used to bond the extra proton. Four equivalent electron domains adopt a tetrahedral arrangement, which corresponds to sp3 hybridisation of nitrogen.

Question 35Chemical Bonding and Molecular Structure

Which of the following correctly ranks the species in order of increasing bond order?

  1. O2^2- < O2- < O2 Correct answer.
  2. O2- < O2 < O2^2-
  3. O2 < O2^2- < O2-
  4. O2 < O2- < O2^2-

Explanation

Neutral O2 has a bond order of 2.0. Adding one electron to form superoxide, O2-, puts that electron in a pi antibonding orbital and lowers the bond order to 1.5. Adding a second electron to form peroxide, O2^2-, lowers it further to 1.0, so the increasing order is peroxide, superoxide, dioxygen.

Question 36Chemical Bonding and Molecular Structure

Why does the boiling point of hydrogen fluoride exceed that of hydrogen chloride despite fluorine being more electronegative and smaller?

  1. Hydrogen chloride is ionic in the liquid state
  2. Hydrogen fluoride has a higher molar mass than hydrogen chloride
  3. Hydrogen fluoride has weaker dispersion forces than hydrogen chloride
  4. Hydrogen fluoride molecules are linked by strong hydrogen bonds Correct answer.

Explanation

Fluorine is small and highly electronegative, so the H-F bond is strongly polar and the hydrogen atom interacts strongly with the lone pair of a neighbouring fluorine. These hydrogen bonds are much stronger than the dipole-dipole and dispersion forces in hydrogen chloride, so more energy is needed to separate hydrogen fluoride molecules.

Question 37Chemical Bonding and Molecular Structure

Which of the following pairs of ions would form the most stable ionic lattice, judged by charge and size considerations?

  1. Mg2+ and O2- Correct answer.
  2. Cs+ and I-
  3. Na+ and Cl-
  4. K+ and Cl-

Explanation

Lattice energy increases with the product of the ionic charges and decreases with the sum of the ionic radii. The pair Mg2+ and O2- has charges of 2+ and 2-, giving a much larger charge product than the singly charged pairs, and both ions are relatively small. This combination therefore releases the most energy on lattice formation.

Question 38Chemical Bonding and Molecular Structure

Which molecule is expected to show the greatest extent of delocalisation of pi electrons?

  1. Cyclohexane
  2. Benzene Correct answer.
  3. 1,3-Butadiene
  4. Ethene

Explanation

Benzene has six p electrons in a continuous cyclic arrangement of overlapping p orbitals, so the pi electrons are delocalised over the entire ring and all carbon-carbon bonds are equivalent. This cyclic conjugation gives greater delocalisation than the open-chain conjugation of butadiene.

Question 39Chemical Bonding and Molecular Structure

The bond angle in the molecule H2S is about 92 degrees, whereas that in H2O is about 104.5 degrees. Which explanation is most consistent with these values?

  1. Sulfur has a larger atomic radius, which forces the hydrogens closer together
  2. Sulfur uses sp3 hybrid orbitals with greater p character than oxygen, reducing the angle Correct answer.
  3. Sulfur has fewer lone pairs than oxygen, so lone pair repulsion is smaller
  4. The sulfur-hydrogen bonds have greater ionic character than oxygen-hydrogen bonds

Explanation

In hydrogen sulfide the central sulfur atom is larger and less electronegative, so its bonding orbitals have more p character and resemble pure p orbitals, which are mutually perpendicular. This drives the bond angle towards 90 degrees. Oxygen is smaller and its hybrid orbitals have more s character, widening the angle towards the tetrahedral value.

Question 40Chemical Bonding and Molecular Structure

Which observation provides the strongest evidence that the carbon-carbon bonds in the benzene ring are equivalent?

  1. Benzene has a single sharp proton nuclear magnetic resonance signal for its ring hydrogens Correct answer.
  2. Benzene is a liquid at room temperature
  3. Benzene burns with a smoky flame
  4. Benzene undergoes substitution rather than addition reactions

Explanation

If the ring contained alternating single and double bonds in fixed positions, the ring hydrogens would be chemically distinct and give more than one proton nuclear magnetic resonance signal. A single sharp signal shows that all six ring hydrogens are in equivalent environments, which is consistent with a fully delocalised structure.

Question 41Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

For an ideal gas, which statement correctly describes the change in internal energy ΔU when the gas expands isothermally?

  1. ΔU is negative because the gas does work during expansion
  2. ΔU is positive because the volume increases
  3. ΔU equals the heat absorbed regardless of work done
  4. ΔU is zero because internal energy depends only on temperature for an ideal gas Correct answer.

Explanation

For an ideal gas, internal energy is a function of temperature alone, so an isothermal process has ΔU = 0. The first law then gives q = −w, meaning heat absorbed equals work done by the gas.

Question 42Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

Which quantity is a state function?

  1. Work
  2. Path length
  3. Heat
  4. Enthalpy Correct answer.

Explanation

A state function depends only on the current state of the system, not on how that state was reached. Enthalpy is defined from state variables U, P, and V, so it is a state function, whereas heat and work depend on the path taken.

Question 43Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

A reaction has ΔH = −150 kJ mol⁻¹ and ΔS = −300 J K⁻¹ mol⁻¹. At what approximate temperature does the reaction become spontaneous?

  1. Below 500 K Correct answer.
  2. At all temperatures
  3. At no temperature
  4. Above 500 K

Explanation

Spontaneity requires ΔG = ΔH − TΔS to be negative. With both terms negative, the entropy term must be smaller in magnitude, giving T < ΔH/ΔS = 150000/300 = 500 K.

Question 44Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

For the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), the equilibrium constant is expressed in terms of partial pressures. Which expression gives Kp?

  1. (pNH₃)² / [(pN₂)(pH₂)³] Correct answer.
  2. (pNH₃) / [(pN₂)(pH₂)³]
  3. (pNH₃)² / [(pN₂)(pH₂)]
  4. (pN₂)(pH₂)³ / (pNH₃)²

Explanation

Kp is the product of product partial pressures raised to their stoichiometric coefficients divided by the same for reactants. For this equation, the numerator is (pNH₃)² and the denominator is (pN₂)(pH₂)³.

Question 45Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

Which change shifts the equilibrium of an exothermic gaseous reaction to the right?

  1. Increasing the volume at constant temperature
  2. Adding a catalyst
  3. Raising the temperature
  4. Lowering the temperature Correct answer.

Explanation

For an exothermic reaction, heat appears as a product, so lowering the temperature removes a product and shifts equilibrium toward products. A catalyst only speeds attainment of equilibrium and does not change its position.

Question 46Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

For a first-order reaction, the half-life is 20 minutes. What fraction of the reactant remains after 60 minutes?

  1. One sixteenth
  2. One half
  3. One quarter
  4. One eighth Correct answer.

Explanation

Sixty minutes corresponds to three half-lives. Each half-life halves the remaining amount, so after three half-lives the fraction remaining is (1/2)³ = 1/8.

Question 47Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

In the Arrhenius equation k = A exp(−Ea/RT), what does the activation energy Ea represent?

  1. The average kinetic energy of the reactant molecules
  2. The energy difference between reactants and products
  3. The total energy released during the reaction
  4. The minimum energy needed for reactant molecules to react Correct answer.

Explanation

The activation energy is the energy barrier that reactant molecules must overcome to form products. It is related to the exponential temperature dependence of the rate constant through the Arrhenius equation.

Question 48Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

A reaction has the rate law rate = k[A][B]². What is the overall order of the reaction?

  1. Order 2
  2. Order 4
  3. Order 1
  4. Order 3 Correct answer.

Explanation

The overall order is the sum of the exponents in the experimental rate law. Here the exponents are 1 for A and 2 for B, giving 1 + 2 = 3.

Question 49Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

A catalyst increases the rate of a chemical reaction mainly because it does which of the following?

  1. Increases the activation energy
  2. Increases the equilibrium constant
  3. Provides an alternative pathway of lower activation energy Correct answer.
  4. Increases the enthalpy of reaction

Explanation

A catalyst offers a reaction pathway with a lower activation energy, so a larger fraction of collisions is successful. It does not change the equilibrium constant or the thermodynamics of the overall reaction.

Question 50Physical Chemistry: Thermodynamics, Kinetics, and Electrochemistry

Which half-reaction correctly represents oxidation at the anode of an electrochemical cell?

  1. Reduction of cations at the cathode
  2. Gain of electrons by the oxidizing agent
  3. Precipitation of an insoluble salt
  4. Loss of electrons by the reducing agent Correct answer.

Explanation

Oxidation is loss of electrons and occurs at the anode of a galvanic cell. The species that loses electrons is the reducing agent and is itself oxidized.