chem 241 midterm application
Which has the higher first ionization energy: Be or B? Explain.
Be has the higher first ionization energy because Be loses a 2s electron while B loses a higher-energy 2p electron, which is easier to remove.
Which has the higher first ionization energy: N or O? Explain.
N has the higher ionization energy because it has a half-filled 2p subshell with exchange stabilization, while O has a paired 2p electron that increases repulsion and lowers IE.
Which is larger: Na or Mg? Explain.
Na is larger because across a period nuclear charge increases while shielding changes little, so Mg has higher Zeff and a smaller radius.
Which is larger: Na or K? Explain.
K is larger because it is lower in the group, has a higher principal quantum number, and more shielding.
Which has higher electronegativity: C or O? Explain.
Oxygen has higher electronegativity because it is farther right in the same period and has higher Zeff.
Which experiences greater Zeff on the valence electron: Li or Be? Explain.
Be experiences greater Zeff because it has more protons while electrons are added to the same valence shell.
Which 2p orbital fills first in carbon?
All three 2p orbitals fill singly before pairing because they are degenerate and follow Hund’s rule.
Why does ionization energy generally increase left to right?
Protons increase across a period while shielding increases more slowly, so Zeff increases and electrons are held more tightly.
Why does ionization energy decrease down a group?
Valence electrons are farther from the nucleus and more shielded, reducing effective nuclear attraction.
Which bond is more polar: C–H or O–H? Explain.
O–H is more polar because oxygen is much more electronegative than hydrogen compared to carbon.
In trigonal bipyramidal geometry, where does a lone pair go and why?
A lone pair occupies an equatorial position because it minimizes 90° repulsions.
Which bond is stronger: C–C single or C=C double? Explain.
C=C is stronger because it contains one sigma and one pi bond, increasing electron density between nuclei.
Which is shorter: C–C single or C≡C triple? Explain.
C≡C is shorter because it has one sigma and two pi bonds, increasing attraction and pulling nuclei closer.
Why are third-row transition metals not much larger than second-row transition metals?
Lanthanide contraction increases Zeff due to poor f-orbital shielding, reducing expected size increase.
Why is 2s lower in energy than 2p in multi-electron atoms?
The 2s orbital penetrates closer to the nucleus and experiences higher Zeff.
Why does fluorine have very high electronegativity?
It has high nuclear charge and small radius, resulting in strong attraction for bonding electrons.
Which is more metallic: Na or Cl? Explain.
Na is more metallic because it has lower ionization energy and more easily loses electrons.
Why do noble gases have very high ionization energies?
They have full valence shells and strong exchange stabilization.
Which hybridization corresponds to linear geometry?
sp hybridization.
Which hybridization corresponds to trigonal planar geometry?
sp2 hybridization.
Which hybridization corresponds to tetrahedral geometry?
sp3 hybridization.
In acetylene (C2H2), what is the hybridization of carbon and why?
Carbon is sp hybridized because it forms two sigma bonds and has two unhybridized p orbitals for two pi bonds.
In formaldehyde (CH2O), what is the hybridization of carbon and why?
Carbon is sp2 hybridized because it forms three sigma bonds and one unhybridized p orbital for the pi bond.
Why is metallic bonding associated with electrical conductivity?
Delocalized electrons in overlapping bands can move freely through the metal lattice.
Which has stronger orbital overlap: bonds within the same row or between different rows? Explain.
Within the same row, because orbitals are closer in size and energy.
Why does atomic radius decrease across a period?
Increasing Zeff pulls valence electrons closer to the nucleus.
Why is screening by electrons in the same valence shell less effective than core electrons?
Electrons in the same shell do not shield nuclear charge as effectively because they are at similar distances from the nucleus.
Which would have higher ionization energy: Mg or Al? Explain.
Mg has higher ionization energy because Al loses a higher-energy 3p electron compared to Mg losing a 3s electron.
Why is the Si–F bond stronger than expected despite size mismatch?
Strong electrostatic attraction due to high bond polarity compensates for poorer orbital size match.
Which position in trigonal bipyramidal geometry experiences more 90° interactions?
Axial positions experience more 90° interactions and are more crowded.
Why do molecules adopt geometries that minimize electron repulsion?
Electron pairs repel due to electrostatic forces and adopt arrangements that minimize total repulsion.
Explain the following trend in observed bond strengths: C–F (488 kJ/mol) > C–C (348 kJ/mol) > C–Cl (330 kJ/mol) > C–Br (280 kJ/mol).
Bond strength mainly depends on (1) orbital overlap quality and (2) extra electrostatic stabilization from bond polarity. C–F is strongest because C and F are both small (period 2), so overlap is very good and the bond is very polar, adding extra attraction. C–Cl and C–Br are weaker because Cl and Br are much larger (3p/4p) so overlap with carbon’s 2p is worse and the bond is longer; Br is largest so overlap is worst and C–Br is weakest. C–C is in between because overlap is good (2p–2p) but it’s nonpolar so it lacks the extra electrostatic “boost” that makes C–F exceptionally strong.
Define Z effective (effective nuclear charge, Zeff).
Zeff is the net positive charge “felt” by an electron after accounting for shielding by other electrons; conceptually Zeff = Z − S, where Z is the number of protons and S is shielding/screening. Higher Zeff means the electron is held more tightly (smaller radius, harder to remove).
Pick which atom will have the highest Zeff: H vs He vs Li.
He (for its 1s electrons) has the highest Zeff because the electrons are close to a +2 nucleus and there’s minimal shielding (only two electrons total). Li’s valence electron is in 2s and is shielded by the 1s² core, so it feels a smaller Zeff than you might think from Z=3. H is lowest because Z=1.
Pick which atom will have the highest Zeff: C vs N vs O.
O has the highest Zeff (then N, then C). Across the same period, Z increases left→right while shielding doesn’t increase much because added electrons are in the same shell, so Zeff increases.
Pick which atom will have the highest Zeff: Be vs F vs S.
F has the highest Zeff. Be and F are in period 2 and Zeff increases left→right, so F > Be. S is in period 3; its valence electrons are farther from the nucleus and more shielded, so they feel a lower Zeff than F’s period-2 valence electrons.
Define ionization energy.
Ionization energy is the energy required to remove an electron from a gaseous atom (or ion). First ionization energy removes the first electron. Higher Zeff and smaller atomic radius generally increase ionization energy because electrons are held more tightly.
State the Pauli exclusion principle.
No two electrons in the same atom can have the same set of four quantum numbers; therefore, an orbital can hold at most two electrons and they must have opposite spins.
State Hund’s Rule.
Electrons fill degenerate orbitals (same energy) singly with parallel spins before pairing. This is more stable because it spreads electrons out (less electron–electron repulsion) and gives exchange stabilization when parallel-spin electrons occupy degenerate orbitals.
State the Aufbau Principle.
Electrons occupy the lowest-energy orbitals available first, building the electron configuration from low energy to high energy.
What is the relationship between possible angular momentum quantum numbers (ℓ) and the principal quantum number (n)?
For a given n, ℓ can be any integer from 0 up to (n−1). That’s why n=1 only has s (ℓ=0); n=2 has s and p (ℓ=0,1); n=3 has s, p, and d (ℓ=0,1,2).
Why do elements on the left side of the periodic table tend to lose valence electrons in reactions, whereas elements on the right side tend to accept electrons?
Left-side elements (metals) have low ionization energies because their valence electrons feel lower Zeff and are farther from the nucleus, so losing electrons is easier and often leads to a stable noble-gas-like configuration. Right-side elements (nonmetals) have higher Zeff, smaller radii, and higher electronegativity, so they attract electrons strongly and tend to gain electrons to complete a valence shell.
Why do valence orbitals and valence electrons have a large influence on the chemistry of an atom?
Chemical bonding and reactions involve the outermost electrons. Valence electrons are the ones available to be shared or transferred, and valence orbitals determine bond formation, molecular shape, polarity, and reactivity. Core electrons are too low in energy and too close to the nucleus to usually participate.
For H2S: what is the shape around the central atom and its hybridization? Which bond angles are equal, less than, or greater than ideal, and why?
H2S is bent because sulfur has 4 electron groups (2 bonds + 2 lone pairs): tetrahedral electron-group geometry and commonly described as sp³ on S. The H–S–H bond angle is less than the ideal tetrahedral 109.5° because lone pairs repel more strongly than bonding pairs and squeeze the bonding pairs closer together (and in H2S the angle is often even smaller than in NH3/H2O due to different bonding details), so “less than ideal” is the key comparison.
For CO2: what is the shape and hybridization of each atom? Sketch/identify the AOs involved with π bonding. Which bond angles are equal, less than, or greater than ideal (explain).
CO2 is linear, so carbon has 2 electron groups → sp hybridized and the O–C–O angle is 180° (ideal linear). Carbon uses two sp orbitals to form the two C–O σ bonds. Carbon has two unhybridized p orbitals left (perpendicular to the bond axis) that form two π bonds by sideways overlap with p orbitals on the oxygens; the two π systems are perpendicular to each other. Each oxygen is commonly described as sp² (one σ bond + two lone pairs in sp² orbitals) with one unhybridized p orbital used for the π bond.
For (CH3)2C=CH2: what are the shapes/hybridizations, what AOs are involved in π bonding, and how do the angles compare to ideal (explain).
There is a C=C double bond, so each carbon in the double bond is trigonal planar → sp², with bond angles around those carbons near 120° (ideal trigonal planar), with slight deviations depending on substituent crowding. The π bond is formed by sideways overlap of the unhybridized p orbitals on the two sp² carbons. The CH3 carbons are tetrahedral → sp³ with angles near 109.5° (ideal tetrahedral).
For cis-FN=NF: what are the shapes/hybridizations, what AOs are involved in π bonding, and how do the angles compare to ideal (explain).
Each nitrogen is part of an N=N double bond and also bonded to F, and has one lone pair: that’s 3 electron groups around each N → trigonal planar electron geometry → sp² on each nitrogen. The σ framework uses sp² orbitals; the π bond comes from sideways overlap of the remaining unhybridized p orbitals on the two nitrogens. “cis” means the two F atoms are on the same side of the double bond and the system is planar (rotation is restricted by the π bond). Bond angles around each N are near 120° (ideal trigonal planar) but lone pair–bond pair repulsion can compress some angles slightly below 120°.
Why does K often have a 1+ charge and Cl often have a 1− charge? Explain.
K has one valence electron (ns¹) and it is energetically favorable to lose that electron to reach a stable noble-gas configuration, forming K⁺. Cl has seven valence electrons (ns²np⁵) and it is favorable to gain one electron to complete its octet (noble-gas configuration), forming Cl⁻.
Which has a higher ionization energy: Na⁺ or Ne? Explain.
Na⁺ has the higher ionization energy. Na⁺ and Ne are isoelectronic (same number of electrons), but Na⁺ has one more proton, so its electrons feel a higher Zeff and are held more tightly, requiring more energy to remove an electron.