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BSSE Chemistry 1429 study materials
BSSE Chemistry 1429 study resources

BSSE Chemistry 1429 practice questions

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Question 1Atomic Structure and the Periodic Table

Which subatomic particle carries a relative charge of −1 and has negligible mass compared with a proton?

  1. Electron Correct answer.
  2. Alpha particle
  3. Neutron
  4. Proton

Explanation

The electron is the negatively charged particle of the atom, with a relative charge of −1 and a mass roughly 1/1836 that of a proton, so its mass is treated as negligible in most calculations. Protons are positive and comparatively heavy, neutrons are neutral and heavy, and an alpha particle is a helium nucleus containing two protons and two neutrons.

Question 2Atomic Structure and the Periodic Table

In the notation ³⁵₁₇Cl, what does the subscript 17 represent?

  1. The mass number
  2. The atomic number Correct answer.
  3. The number of nucleons
  4. The number of neutrons

Explanation

The subscript placed before the symbol gives the atomic number, which equals the number of protons in the nucleus and determines the element's identity. The superscript gives the mass number, the total count of protons plus neutrons, which for this isotope is 35.

Question 3Atomic Structure and the Periodic Table

Which pair of species are isotopes of the same element?

  1. ¹⁶O and ¹⁸O Correct answer.
  2. ²³Na⁺ and ²³Na
  3. ³⁵Cl⁻ and ³⁷Cl⁻
  4. ¹⁴N and ¹⁴C

Explanation

Isotopes are atoms of the same element with identical proton number but different neutron numbers, so their mass numbers differ. Oxygen-16 and oxygen-18 both have eight protons but ten and twelve neutrons respectively, making them isotopes.

Question 4Atomic Structure and the Periodic Table

An atom of an element has 16 protons, 18 neutrons and 18 electrons. Which description is correct?

  1. An anion with mass number 36
  2. An anion with mass number 34 Correct answer.
  3. A cation with mass number 34
  4. A neutral atom with mass number 36

Explanation

The mass number is the sum of protons and neutrons, 16 + 18 = 34. Because there are two more electrons than protons, the species carries a 2− charge and is an anion, so it is an anion of mass number 34.

Question 5Atomic Structure and the Periodic Table

Bromine consists of ⁷⁹Br (abundance 50.7%) and ⁸¹Br (abundance 49.3%). Which value is closest to its relative atomic mass?

  1. 80.50
  2. 78.92
  3. 81.00
  4. 79.99 Correct answer.

Explanation

Relative atomic mass is the abundance-weighted average of the isotope masses: (79 × 0.507) + (81 × 0.493) ≈ 40.05 + 39.93 = 79.98. Rounding gives about 79.99, close to bromine's accepted value.

Question 6Atomic Structure and the Periodic Table

How many neutrons are present in one atom of ⁵⁶Fe?

  1. 82
  2. 26
  3. 30 Correct answer.
  4. 56

Explanation

Iron has atomic number 26, so each atom contains 26 protons. The neutron number equals mass number minus atomic number, 56 − 26 = 30.

Question 7Atomic Structure and the Periodic Table

Which transition in the hydrogen atom emits a photon of the highest energy?

  1. n = 4 to n = 2
  2. n = 2 to n = 1 Correct answer.
  3. n = 3 to n = 2
  4. n = 4 to n = 3

Explanation

Photon energy equals the energy gap between the two levels, and the gaps grow larger as n decreases toward the nucleus. The 2→1 jump spans the widest interval among those listed, so it releases the most energetic photon, in the ultraviolet Lyman series.

Question 8Atomic Structure and the Periodic Table

According to the Bohr model, what happens when an electron moves from a lower to a higher energy level?

  1. It emits a photon of energy equal to the gap
  2. It splits into two electrons
  3. Its mass increases slightly
  4. It absorbs energy equal to the level difference Correct answer.

Explanation

In the Bohr model an electron occupies discrete circular orbits, and promotion to a higher level requires absorbing energy exactly matching the difference between the two levels, usually as a photon. Emission occurs in the reverse process, when the electron drops to a lower level.

Question 9Atomic Structure and the Periodic Table

Which statement best describes an atomic orbital?

  1. A fixed shell containing exactly two electrons
  2. A region where an electron is most likely to be found Correct answer.
  3. A circular path followed by an electron
  4. The nucleus-centred volume holding protons

Explanation

Quantum mechanics replaces definite orbits with orbitals, mathematical functions whose squared magnitude gives the probability of finding an electron in a region of space. An orbital is therefore a probability distribution, not a precise trajectory.

Question 10Atomic Structure and the Periodic Table

What is the maximum number of electrons that can occupy the 3d subshell?

  1. 6
  2. 14
  3. 10 Correct answer.
  4. 2

Explanation

A d subshell contains five orbitals, and each orbital holds at most two electrons by the Pauli exclusion principle, giving a maximum of ten. The values 2, 6 and 14 correspond to the capacities of a single orbital and of s, p and f subshells respectively.

Question 11Atomic Structure and the Periodic Table

Which set of quantum numbers is not permitted for an electron in an atom?

  1. n = 1, l = 1, mₗ = 0 Correct answer.
  2. n = 3, l = 2, mₗ = −2
  3. n = 4, l = 0, mₗ = 0
  4. n = 2, l = 1, mₗ = 0

Explanation

The angular momentum quantum number l may take only integer values from 0 up to n − 1. For n = 1 the only allowed value is l = 0, so l = 1 is impossible for that shell.

Question 12Atomic Structure and the Periodic Table

What is the ground-state electron configuration of a chlorine atom (Z = 17)?

  1. 1s² 2s² 2p⁶ 3s² 3p⁵ Correct answer.
  2. 1s² 2s² 2p⁶ 3s² 3p⁶
  3. 1s² 2s² 2p⁶ 3s² 3p⁴ 3d¹
  4. 1s² 2s² 2p⁵ 3s² 3p⁶

Explanation

Seventeen electrons fill orbitals in order of increasing energy: 1s² 2s² 2p⁶ accounts for ten, then 3s² brings the total to twelve and 3p⁵ adds the remaining five. This gives a total of seventeen electrons with a half-filled-plus-one p subshell.

Question 13Atomic Structure and the Periodic Table

Which element has the ground-state configuration [Ar] 4s² 3d¹⁰ 4p³?

  1. Arsenic Correct answer.
  2. Selenium
  3. Antimony
  4. Phosphorus

Explanation

The argon core supplies 18 electrons and the 4s² 3d¹⁰ 4p³ portion adds 15 more, giving Z = 33. Counting to atomic number 33 in period 4 places the element in group 15, which is arsenic.

Question 14Atomic Structure and the Periodic Table

Which explanation correctly accounts for the trend of decreasing atomic radius across period 3 from sodium to chlorine?

  1. New electron shells are added at each step
  2. Nuclear charge increases while electrons enter the same shell Correct answer.
  3. The number of neutrons rises steadily
  4. Effective nuclear charge decreases as shielding increases

Explanation

Across a period, protons are added to the nucleus while the extra electrons enter the same principal shell, so shielding stays roughly constant. The growing effective nuclear charge pulls the electron cloud inward, shrinking the radius.

Question 15Atomic Structure and the Periodic Table

Which element in group 2 has the largest atomic radius?

  1. Magnesium
  2. Beryllium
  3. Barium Correct answer.
  4. Calcium

Explanation

Atomic radius increases down a group because each successive element places its valence electrons in a higher principal shell, farther from the nucleus and more shielded. Barium lies lowest in group 2 among the choices, so it has the largest radius.

Question 16Atomic Structure and the Periodic Table

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

  1. Magnesium has a greater effective nuclear charge for the electron removed Correct answer.
  2. Magnesium has a larger atomic radius
  3. Sodium has a full outer shell
  4. Magnesium removes an electron from a p orbital

Explanation

Magnesium has one more proton than sodium while the electron being removed comes from the same 3s subshell, so it experiences a stronger effective nuclear attraction. That greater hold raises the energy needed to remove the electron.

Question 17Atomic Structure and the Periodic Table

The first ionisation energies of the elements increase across period 2 but aluminium's is lower than magnesium's. Which factor best explains this dip?

  1. Aluminium has a lower nuclear charge than magnesium
  2. The electron removed from aluminium is in a 3p orbital, which is higher in energy and better shielded Correct answer.
  3. Aluminium has a larger number of neutrons
  4. The electron removed from aluminium comes from a filled 3s subshell

Explanation

Across period 2 the trend is generally upward, but the electron removed from aluminium occupies the 3p subshell, which lies slightly higher in energy and is shielded by the 3s pair. This makes that electron easier to remove than one from magnesium's filled 3s subshell.

Question 18Atomic Structure and the Periodic Table

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

  1. Electronegativity Correct answer.
  2. Number of occupied shells
  3. Metallic character
  4. Atomic radius

Explanation

Electronegativity rises across a period because the effective nuclear charge grows while the valence shell stays the same, so atoms attract bonding electrons more strongly. Atomic radius, metallic character and the number of shells all decrease or stay constant in that direction.

Question 19Atomic Structure and the Periodic Table

Which group of the periodic table contains the most reactive non-metals?

  1. Group 18
  2. Group 2
  3. Group 1
  4. Group 17 Correct answer.

Explanation

Group 17 halogens have seven valence electrons and a strong tendency to gain one more, giving them high electronegativity and strong oxidising power. Fluorine, at the top of the group, is the most reactive non-metal known.

Question 20Atomic Structure and the Periodic Table

Which ion has the electron configuration 1s² 2s² 2p⁶?

  1. Ca²⁺
  2. Mg
  3. Na⁺ Correct answer.
  4. Cl⁻

Explanation

The configuration given contains ten electrons, matching the neon core. Sodium has eleven electrons, so losing one to form Na⁺ leaves exactly the neon configuration.

Question 21Atomic Structure and the Periodic Table

Which species are isoelectronic with one another?

  1. Mg²⁺ and Ar
  2. Na⁺ and Ne Correct answer.
  3. F⁻ and Cl⁻
  4. Na⁺ and K⁺

Explanation

Isoelectronic species have the same total number of electrons. Na⁺ has 11 − 1 = 10 electrons and neon has 10, so the pair is isoelectronic.

Question 22Atomic Structure and the Periodic Table

Which block of the periodic table contains an element whose atoms fill the f subshells?

  1. p block
  2. d block
  3. f block Correct answer.
  4. s block

Explanation

In the f block, the differentiating electron enters an f subshell, which corresponds to the inner transition series of lanthanides and actinides. The s, p and d blocks fill their respective subshells instead.

Question 23Atomic Structure and the Periodic Table

Why do transition metals such as iron show variable oxidation states?

  1. They have only one valence electron
  2. Their nuclei contain variable numbers of protons
  3. They always have full d subshells
  4. Their 4s and 3d electrons are close in energy and can both be lost Correct answer.

Explanation

In transition metals the 4s and 3d orbitals lie close in energy, so electrons from both can participate in bonding and ionisation. This allows several oxidation states to be accessible, as in Fe²⁺ and Fe³⁺.

Question 24Atomic Structure and the Periodic Table

Which statement about the shielding of outer electrons is correct?

  1. Core electrons exert no shielding effect at all
  2. Shielding increases the effective nuclear charge felt by an outer electron
  3. Inner-shell electrons shield outer electrons more effectively than electrons in the same shell Correct answer.
  4. Electrons in the same shell shield outer electrons most effectively

Explanation

Electrons in inner shells lie between the nucleus and an outer electron and screen the nuclear charge strongly, whereas electrons in the same shell are less effective at screening each other. Greater shielding reduces the effective nuclear charge felt by the outer electron.

Question 25Chemical Bonding and Molecular Structure

Which type of bonding is best described as the electrostatic attraction between a lattice of positive ions and a sea of delocalized electrons?

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

Explanation

Metallic bonding arises when metal atoms release valence electrons into a shared delocalized pool, leaving positive cores that attract the mobile electron sea. This delocalization explains the characteristic conductivity and malleability of metals. Ionic and covalent bonds instead involve discrete electron transfer or shared electron pairs between specific atoms.

Question 26Chemical Bonding and Molecular Structure

According to VSEPR theory, what is the molecular geometry of methane, CH4?

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

Explanation

Methane has four bonding pairs and no lone pairs on the central carbon, giving an electron-pair geometry of four regions. Four equivalent bonding pairs arrange themselves as far apart as possible, producing a tetrahedral shape with bond angles of about 109.5 degrees.

Question 27Chemical Bonding and Molecular Structure

Which molecule contains a polar covalent bond but is overall nonpolar because its bond dipoles cancel?

  1. Ammonia, NH3
  2. Water, H2O
  3. Hydrogen chloride, HCl
  4. Carbon dioxide, CO2 Correct answer.

Explanation

In carbon dioxide the two polar carbon-oxygen bonds point in exactly opposite directions along a linear axis, so their dipole moments are equal and cancel. The result is a molecule with polar bonds but zero net dipole moment. The other molecules have geometries that leave a nonzero resultant dipole.

Question 28Chemical Bonding and Molecular Structure

Which of the following is an example of an ionic compound?

  1. Iodine, I2
  2. Carbon dioxide, CO2
  3. Sodium chloride, NaCl Correct answer.
  4. Methane, CH4

Explanation

Ionic compounds form when a metal transfers electrons to a nonmetal, creating oppositely charged ions that pack into a lattice. Sodium loses one electron to chlorine, giving Na+ and Cl- ions held by electrostatic attraction. The other examples consist of nonmetals bonded covalently.

Question 29Chemical Bonding and Molecular Structure

What is the electron-pair geometry around the central nitrogen in the ammonium ion, NH4+?

  1. Trigonal planar
  2. Linear
  3. Tetrahedral Correct answer.
  4. Octahedral

Explanation

The ammonium ion has four bonding pairs around nitrogen and no lone pairs, because the positive charge reflects loss of the lone pair present in ammonia. Four electron regions repel each other into a tetrahedral arrangement. The molecular shape therefore matches the electron-pair geometry.

Question 30Chemical Bonding and Molecular Structure

Which pair of elements would most likely form an ionic bond?

  1. Potassium and fluorine Correct answer.
  2. Carbon and oxygen
  3. Nitrogen and oxygen
  4. Hydrogen and sulfur

Explanation

Ionic bonding is favored when a low-electronegativity metal reacts with a high-electronegativity nonmetal, giving a large electronegativity difference. Potassium is an alkali metal that readily loses an electron, while fluorine strongly attracts electrons. The other pairs are nonmetal combinations that tend to form covalent bonds.

Question 31Chemical Bonding and Molecular Structure

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

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

Explanation

Each carbon in ethene is bonded to two hydrogens and to the other carbon, giving four C-H sigma bonds plus one C-C sigma bond, for five sigma bonds. The carbon-carbon double bond also contains one pi bond formed by sideways overlap of unhybridized p orbitals.

Question 32Chemical Bonding and Molecular Structure

Which hybridization is assigned to the central carbon atom in a molecule with a triple bond, such as ethyne, C2H2?

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

Explanation

Each carbon in ethyne is bonded to two other atoms and has two electron regions, which corresponds to sp hybridization. The two unhybridized p orbitals on each carbon overlap sideways to form the two pi bonds of the triple bond. Linear geometry with 180 degree bond angles results.

Question 33Chemical Bonding and Molecular Structure

Which of the following molecules has a trigonal pyramidal molecular geometry?

  1. Carbon tetrachloride, CCl4
  2. Sulfur dioxide, SO2
  3. Phosphorus trichloride, PCl3 Correct answer.
  4. Boron trifluoride, BF3

Explanation

Phosphorus in PCl3 has three bonding pairs and one lone pair, giving four electron regions in a tetrahedral arrangement. The lone pair occupies one vertex, so the molecular shape is trigonal pyramidal with slightly compressed bond angles. Boron trifluoride has no lone pair and is trigonal planar.

Question 34Chemical Bonding and Molecular Structure

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

  1. Metallic bonding
  2. Hydrogen bonding Correct answer.
  3. Ion-dipole interactions
  4. London dispersion 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 relatively strong directional interactions require substantial energy to overcome, raising the boiling point. Sulfur is larger and less electronegative, so H2S molecules interact mainly through weaker dipole and dispersion forces.

Question 35Chemical Bonding and Molecular Structure

Which species has a bond angle closest to 120 degrees?

  1. NH3
  2. CH4
  3. BF3 Correct answer.
  4. H2O

Explanation

Boron trifluoride has three bonding pairs and no lone pair, so the electron regions adopt a trigonal planar arrangement with ideal bond angles of 120 degrees. The other species have four electron regions and tetrahedral-based angles near 109.5 degrees or less. Lone pairs further compress those angles below the ideal tetrahedral value.

Question 36Chemical Bonding and Molecular Structure

Which molecule is nonpolar despite containing polar bonds, owing to its symmetrical trigonal planar geometry?

  1. Hydrogen cyanide, HCN
  2. Chloromethane, CH3Cl
  3. Nitrogen trichloride, NCl3
  4. Boron trifluoride, BF3 Correct answer.

Explanation

In boron trifluoride the three polar B-F bonds are arranged symmetrically at 120 degrees in a plane, so the individual bond dipoles sum to zero. The molecule is therefore nonpolar overall. In the other species the geometry or the presence of a lone pair prevents complete cancellation of dipoles.

Question 37Chemical Bonding and Molecular Structure

Which statement correctly describes the difference between a sigma bond and a pi bond?

  1. Sigma bonds can exist only in diatomic molecules, while pi bonds occur in polyatomic molecules.
  2. Pi bonds are always stronger than sigma bonds because they involve greater orbital overlap.
  3. Sigma bonds form only between p orbitals, while pi bonds form only between s orbitals.
  4. Sigma bonds form by head-on overlap along the internuclear axis, while pi bonds form by sideways overlap of parallel p orbitals. Correct answer.

Explanation

A sigma bond results from direct, head-on overlap of orbitals along the line joining the nuclei, giving maximum overlap and free rotation. A pi bond arises from side-by-side overlap of parallel p orbitals above and below the internuclear axis, with less overlap and restricted rotation. Pi bonds are therefore generally weaker than sigma bonds.

Question 38Chemical Bonding and Molecular Structure

What is the formal charge on the central oxygen atom in the hydronium ion, H3O+, when it is drawn with three single bonds and one lone pair?

  1. +1 Correct answer.
  2. -1
  3. +2
  4. 0

Explanation

Formal charge equals valence electrons minus nonbonding electrons minus half the bonding electrons. Oxygen has six valence electrons; with one lone pair (two nonbonding electrons) and three bonds (six bonding electrons), the calculation gives 6 - 2 - 3 = +1. This positive formal charge is consistent with the overall charge of the ion.

Question 39Chemical Bonding and Molecular Structure

Which of the following correctly ranks the bond lengths from shortest to longest?

  1. C-O < C≡O < C=O
  2. C-O < C=O < C≡O
  3. C=O < C≡O < C-O
  4. C≡O < C=O < C-O Correct answer.

Explanation

Bond length decreases as bond order increases because additional shared electron pairs pull the nuclei closer together. A triple bond is therefore shorter than a double bond, which is shorter than a single bond between the same atoms. The correct order runs from the triple bond through the double bond to the single bond.

Question 40Chemical Bonding and Molecular Structure

Which molecule contains a central atom that exceeds the octet and is described as hypervalent?

  1. Methane, CH4
  2. Sulfur hexafluoride, SF6 Correct answer.
  3. Water, H2O
  4. Carbon dioxide, CO2

Explanation

Sulfur hexafluoride has six bonding pairs around sulfur, giving twelve electrons in the valence shell, which exceeds the octet. This is possible for period 3 elements that can use d orbitals or expanded bonding descriptions. Carbon and oxygen in the other molecules obey the octet rule.

Question 41Chemical Bonding and Molecular Structure

Which of the following has the greatest lattice energy, assuming comparable ion arrangements?

  1. Cesium chloride, CsCl
  2. Potassium chloride, KCl
  3. Magnesium oxide, MgO Correct answer.
  4. Sodium chloride, NaCl

Explanation

Lattice energy increases with the product of ionic charges and decreases with the sum of ionic radii. Magnesium oxide contains doubly charged Mg2+ and O2- ions, so the charge product is four times that of a 1:1 salt, and the ions are also relatively small. This combination gives MgO the largest lattice energy among the choices.

Question 42Chemical Bonding and Molecular Structure

What is the molecular geometry of the sulfur dioxide molecule, SO2?

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

Explanation

Sulfur in SO2 has two bonding regions and one lone pair, giving three electron regions in a trigonal planar electron geometry. The lone pair occupies one vertex, so the molecular shape is bent with an O-S-O angle below 120 degrees. Only the positions of atoms define the molecular geometry.

Question 43Chemical Bonding and Molecular Structure

Which hybridization best describes the carbon atoms in benzene, C6H6?

  1. sp
  2. sp3d2
  3. sp2 Correct answer.
  4. sp3

Explanation

Each carbon in benzene is bonded to two other carbons and one hydrogen, giving three electron regions and sp2 hybridization. The remaining unhybridized p orbital on each carbon overlaps with its neighbors to form a delocalized pi system above and below the ring. This delocalization gives all carbon-carbon bonds an intermediate length.

Question 44Chemical Bonding and Molecular Structure

Which of the following molecules exhibits only London dispersion forces as its intermolecular attraction?

  1. Hydrogen bromide, HBr
  2. Tetrachloromethane, CCl4 Correct answer.
  3. Hydrogen fluoride, HF
  4. Ammonia, NH3

Explanation

Tetrachloromethane is a symmetrical tetrahedral molecule whose polar C-Cl bond dipoles cancel, leaving no net dipole moment. Without a permanent dipole or hydrogen bonding capability, its molecules interact only through London dispersion forces. The other molecules have permanent dipoles and, in some cases, hydrogen bonding.

Question 45Chemical Bonding and Molecular Structure

Which of the following best explains why the first ionization energy of magnesium is higher than that of sodium?

  1. Magnesium has a larger atomic radius, so its outer electron is farther from the nucleus.
  2. Magnesium has one more electron shell than sodium, increasing shielding.
  3. Magnesium has a higher effective nuclear charge, so its outer electron is held more tightly. Correct answer.
  4. Magnesium has a half-filled p subshell that provides extra stability.

Explanation

Across a period, the number of protons increases while electrons enter the same shell, so effective nuclear charge rises and atomic radius decreases. Magnesium's outer 3s electrons therefore experience a stronger attraction than sodium's single 3s electron removed from a larger atom. Ionization energy consequently increases from sodium to magnesium.

Question 46Chemical Bonding and Molecular Structure

Resonance structures are used to describe a molecule when which condition applies?

  1. The molecule contains only single bonds and no lone pairs on the central atom.
  2. The molecule rapidly interconverts between distinct isomers that can be isolated.
  3. A single Lewis structure cannot represent the delocalized electron distribution, so multiple equivalent structures are averaged. Correct answer.
  4. The central atom exceeds the octet and requires an expanded valence shell.

Explanation

Resonance applies when two or more valid Lewis structures differ only in the placement of electrons, such as pi bonds and lone pairs, while the atomic positions stay fixed. The true structure is a weighted average or hybrid of these contributors, giving bond orders intermediate between single and double. The contributing structures are not real, separable species.

Question 47Chemical Bonding and Molecular Structure

Using molecular orbital theory, why is the dioxygen molecule, O2, paramagnetic?

  1. It has a completely filled set of pi antibonding orbitals.
  2. Its bond order is zero, leaving all electrons unpaired.
  3. It contains an unpaired electron in a sigma bonding orbital.
  4. Its two highest-energy electrons occupy degenerate pi antibonding orbitals singly with parallel spins. Correct answer.

Explanation

Filling the molecular orbitals of O2 places the last two electrons in the degenerate pi antibonding orbitals, and Hund's rule requires them to occupy separate orbitals with parallel spins. This gives two unpaired electrons and the observed paramagnetism. The bond order of O2 is two, consistent with a double bond.

Question 48Chemical Bonding and Molecular Structure

Which factor best explains why the N-N single bond in hydrazine, N2H4, is weaker than the N-N bond in nitrogen, N2?

  1. The single bond in hydrazine has lower bond order and experiences greater lone-pair repulsion between adjacent nitrogen atoms. Correct answer.
  2. Hydrazine has a longer bond because nitrogen uses d orbitals to strengthen the bond in N2.
  3. Nitrogen molecules contain hydrogen bonding that strengthens the N-N bond.
  4. The triple bond in nitrogen has greater bond order but less orbital overlap than the single bond.

Explanation

Bond strength generally increases with bond order, so the triple bond in N2 is much stronger than the single bond in hydrazine. In hydrazine the adjacent nitrogen lone pairs also repel each other, further weakening the N-N single bond. Bond order and lone-pair repulsion together account for the difference.

Question 49States of Matter and Intermolecular Forces

Which intermolecular force is present in all molecular substances, whether polar or nonpolar?

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

Explanation

London dispersion forces arise from instantaneous fluctuations in electron distribution, which create temporary dipoles in every molecule. Because all molecules have electrons, this force operates universally, whereas the other listed forces require permanent dipoles or ions.

Question 50States of Matter and Intermolecular Forces

A liquid is heated at constant pressure until it becomes a gas. Which statement describes the temperature at the boiling point during the phase change?

  1. It falls because vaporization absorbs energy.
  2. It rises steadily as more heat is added.
  3. It oscillates between the liquid and gas values.
  4. It remains constant while the liquid converts to vapor. Correct answer.

Explanation

At the boiling point, added heat is consumed as latent heat of vaporization to overcome intermolecular attractions rather than to increase kinetic energy. Since average kinetic energy stays fixed, the temperature remains constant until all liquid has vaporized.