chem 241 midterm study
What is effective nuclear charge (Zeff)?
Zeff is the amount of nuclear charge experienced by an electron in an atom and is given by Zeff = Z − S.
What does Z represent in Zeff = Z − S?
Z represents the actual nuclear charge, equal to the number of protons in the nucleus.
What does S represent in Zeff = Z − S?
S is the screening (shielding) constant that accounts for electron–electron repulsion reducing the nuclear attraction felt by an electron.
Why is hydrogen considered a simple atomic system?
Hydrogen has only one proton and one electron, so there is no electron–electron repulsion and orbital energy depends only on the principal quantum number n.
Why do multi-electron atoms not behave like hydrogen?
Multi-electron atoms experience electron–electron repulsion and screening, so orbital energies depend on both n and Zeff.
What causes screening in atoms?
Inner (core) electrons repel outer electrons and reduce the effective nuclear charge experienced by the outer electrons.
Why do we focus mainly on valence electrons when discussing Zeff?
Valence electrons are the electrons involved in chemical reactions and determine the chemistry of the atom.
How does increasing the number of protons affect Zeff?
Increasing the number of protons increases the nuclear charge and generally increases Zeff if shielding does not increase proportionally.
In lithium, why does the 2s electron experience less nuclear charge than the 1s electrons?
The two 1s core electrons shield the 2s valence electron from the full +3 nuclear charge.
Why is E2s lower than E2p in multi-electron atoms?
The 2s orbital has an inner lobe that penetrates closer to the nucleus, experiences higher Zeff, and is therefore lower in energy than the 2p orbitals.
What is meant by orbital penetration?
Penetration refers to how much an orbital allows electron density to extend close to the nucleus; greater penetration leads to higher Zeff and lower energy.
In hydrogen-like atoms, how are 2s and 2p energies related?
In hydrogen-like atoms, 2s and 2p orbitals have the same energy because energy depends only on n.
In multi-electron atoms, how are 2s and 2p energies related?
In multi-electron atoms, 2s is lower in energy than 2p due to greater penetration and higher Zeff.
What is the electronic configuration of lithium?
1s2 2s1 or [He] 2s1.
Why is the outer 2s electron in lithium not tightly held?
It is shielded by the 1s2 core electrons and experiences a reduced effective nuclear charge.
What determines the order in which atomic orbitals are filled?
The Aufbau principle, Pauli exclusion principle, and Hund’s rule determine orbital filling.
What does the Aufbau principle state?
Electrons fill orbitals in order of increasing energy.
What does the Pauli exclusion principle state?
No two electrons in an atom can have the same four quantum numbers, and each orbital can hold a maximum of two electrons with opposite spins.
What does Hund’s rule state?
Electrons occupy degenerate orbitals singly with parallel spins before pairing.
Why are parallel spins more stable than paired spins in degenerate orbitals?
Parallel spins provide exchange stabilization energy and reduce electron–electron repulsion.
What is exchange stabilization energy?
Additional stabilization that occurs when electrons with parallel spins occupy degenerate orbitals.
What is the electronic configuration of carbon?
1s2 2s2 2p2 or [He] 2s2 2p2.
How should the 2p electrons be arranged in carbon according to Hund’s rule?
They occupy separate 2p orbitals with parallel spins before pairing.
Why is nitrogen particularly stable?
Nitrogen has a half-filled 2p subshell (2p3), maximizing exchange stabilization energy.
What are the numbers of orbitals and maximum electrons in s, p, d, and f subshells?
s: 1 orbital, 2 electrons; p: 3 orbitals, 6 electrons; d: 5 orbitals, 10 electrons; f: 7 orbitals, 14 electrons.
Why are Cr and Cu exceptions to the expected electron configuration pattern?
The 4s and 3d orbitals are very close in energy, and half-filled or fully filled d subshells provide additional stability.
In multi-electron atoms, why can 4s fill before 3d?
The 4s orbital has lower energy than 3d in neutral atoms due to penetration and shielding effects.
Why does orbital energy in hydrogen-like atoms depend only on n?
There is no electron–electron repulsion, so energy depends solely on the principal quantum number.
In multi-electron atoms, what two major factors determine orbital energy?
Principal quantum number and effective nuclear charge (including screening and electron–electron repulsion).
What is the shielding ability order of orbitals?
s > p > d > f.
Why do s orbitals shield better than p orbitals?
s orbitals penetrate closer to the nucleus and have greater electron density near the nucleus.
What is d-orbital contraction?
Poor shielding by filled d orbitals increases Zeff on outer electrons, leading to smaller-than-expected atomic radii.
What is lanthanide contraction?
Poor shielding by f orbitals increases Zeff across the lanthanides, causing atomic radii to be smaller than expected.
Why is lanthanide contraction important in chemistry?
It causes third-row transition metals to have sizes similar to second-row transition metals, affecting bonding and reactivity.
What is first ionization energy?
The energy required to remove the outermost electron from a gaseous atom.
Why does ionization energy generally increase across a period?
Protons increase left to right while electrons are added to the same valence shell, and screening increases slower than nuclear charge, so Zeff increases.
Why does ionization energy decrease down a group?
Principal quantum number increases, orbitals become larger, shielding increases, and valence electrons are held less tightly.
Why is the ionization energy of boron lower than that of beryllium?
Boron removes a 2p electron (higher energy), whereas beryllium removes a 2s electron (lower energy).
Why is the ionization energy of oxygen lower than that of nitrogen?
Oxygen has one paired 2p electron, increasing electron–electron repulsion and lowering ionization energy compared to half-filled nitrogen.
Why are noble gases very stable and have high ionization energies?
They have full valence shells and high exchange stabilization energy.
How does atomic radius change down a group?
Atomic radius increases due to higher principal quantum number and increased shielding.
How does atomic radius change across a period?
Atomic radius decreases because Zeff increases left to right, pulling electrons closer to the nucleus.
What types of atomic radii are commonly discussed?
Covalent radius, van der Waals radius, and ionic radius.
What is electronegativity?
The relative ability of an atom to attract electrons in a bond.
How does electronegativity change across a period?
It increases from left to right.
How does electronegativity change down a group?
It decreases down a group.
What types of bonding exist?
Covalent, polar covalent, ionic, and metallic bonding.
What characterizes covalent bonding?
Atoms share electrons through overlapping valence orbitals.
What characterizes ionic bonding?
Electrons are transferred due to large electronegativity differences, forming oppositely charged ions held by electrostatic attraction.
What characterizes metallic bonding?
Delocalized valence electrons form overlapping bands of electron density across closely packed metal atoms.
What determines bond strength?
Degree of orbital overlap and electrostatic attraction between atoms.
Why is good orbital overlap important?
Greater overlap lowers energy more and strengthens the bond.
What does VSEPR stand for?
Valence Shell Electron Pair Repulsion.
What is the central idea of VSEPR?
Electron pairs repel each other and adopt geometries that minimize electron–electron repulsion.
What is the geometry of AX2?
Linear, 180°.
What is the geometry of AX3?
Trigonal planar, 120°.
What is the geometry of AX4?
Tetrahedral, 109.5°.
What is the geometry of AX5?
Trigonal bipyramidal, with 90° and 120° angles.
In trigonal bipyramidal geometry, what are axial and equatorial positions?
Axial positions are above and below the trigonal plane; equatorial positions lie in the trigonal plane.
Which positions are less crowded in trigonal bipyramidal geometry?
Equatorial positions are less crowded than axial positions.
Where do lone pairs prefer to reside in trigonal bipyramidal geometry?
Lone pairs prefer equatorial positions because they are less crowded.
What is the geometry of AX6?
Octahedral, with all bond angles 90°.
What is hybridization?
The mixing of atomic orbitals to form equivalent hybrid orbitals oriented to match VSEPR geometry.
What is sp hybridization?
Mixing one s and one p orbital to form two equivalent sp orbitals arranged linearly at 180°.
What is sp2 hybridization?
Mixing one s and two p orbitals to form three equivalent sp2 orbitals arranged trigonal planar at 120°.
What is sp3 hybridization?
Mixing one s and three p orbitals to form four equivalent sp3 orbitals arranged tetrahedrally at 109.5°.
What is sp3d hybridization?
Formation of five hybrid orbitals arranged trigonal bipyramidal.
What is sp3d2 hybridization?
Formation of six hybrid orbitals arranged octahedral.
Why must carbon promote an electron before forming four bonds?
Carbon’s ground state is 2s2 2p2, so promotion to 2s1 2p3 provides four unpaired electrons for bonding.
Is electron promotion a physically separate step?
No, it is a model used to visualize bonding; bond formation energy compensates promotion energy.
What bonds are formed by hybrid orbitals?
Sigma (σ) bonds formed by end-to-end orbital overlap.
How many sigma and pi bonds are in a double bond?
One sigma bond and one pi bond.
How many sigma and pi bonds are in a triple bond?
One sigma bond and two pi bonds.
What is Mulliken electronegativity based on?
Mulliken electronegativity is based on measurable quantities related to ionization energy and electron affinity and applies to molecular bonding.
How is electronegativity related to periodic trends?
Electronegativity generally increases from left to right across a period and decreases down a group.
Why are elements to the right and higher in the periodic table generally more electronegative?
They have higher Zeff and smaller atomic radii, allowing them to attract bonding electrons more strongly.
How can hydrogen behave in compounds?
Hydrogen can be either acidic (Hδ+) when bonded to more electronegative atoms or hydridic (Hδ−) when bonded to less electronegative atoms.
What determines whether hydrogen is Hδ+ or Hδ−?
The relative electronegativity of the atom bonded to hydrogen determines the direction of bond polarization.
What does a red region in an electron density diagram represent?
An electron-rich area.
What does a blue region in an electron density diagram represent?
An electron-poor area.
How does electronegativity difference affect bond polarity?
A larger electronegativity difference increases bond polarity and ionic character.
What happens when electronegativity difference is very large?
Orbital overlap becomes poor and electrons are effectively transferred, forming ionic solids.
What is metallic bonding described as?
Delocalization of valence electrons forming overlapping bands of electron density across many atoms.
What experimental method is used to measure energies of electron bands in metals?
Photoelectron spectroscopy.
What is photoelectron spectroscopy used for in atoms?
It measures ionization energies and can eject inner core electrons when using high-energy photons such as X-rays.
How does covalent bonding occur at the orbital level?
Covalent bonds form by overlap of valence atomic orbitals to create molecular orbitals.
What happens to energy as two hydrogen atoms approach each other?
Energy decreases due to electrostatic attraction between electrons and nuclei.
Why does energy increase again at very short internuclear distances?
Electron–electron and nucleus–nucleus repulsion become dominant, increasing energy.
What is bond length?
The internuclear distance at which energy is minimized and stabilization is maximum.
What determines whether a bond is strong or weak?
The extent of orbital overlap and electrostatic attraction between atoms.
Why do orbitals of similar size overlap better?
Similar size orbitals allow more effective spatial overlap, lowering energy more significantly.
Why is overlap generally better within the same row of the periodic table?
Orbitals within the same row are closer in size and energy.
Why must orbitals point toward each other for strong bonding?
Maximum overlap occurs when orbitals are aligned along the internuclear axis.
Why is the Si–F bond extremely strong?
Strong electrostatic attraction due to high bond polarity outweighs the orbital size mismatch between Si and F.
Why is the C–F bond also strong?
Good orbital overlap and high electronegativity of fluorine provide strong covalent and electrostatic contributions.
Why does electrostatic attraction strengthen polar covalent bonds?
Partial charges attract each other, adding ionic character that strengthens the bond.
Why are lone pairs considered larger than bonding pairs in VSEPR?
Lone pairs are localized closer to the central atom and occupy more space, increasing repulsion.
What is steric repulsion?
Electron pair repulsion that influences molecular geometry.
Why do molecules adopt specific geometries according to VSEPR?
To minimize electron–electron repulsion between bonding pairs and lone pairs.
In trigonal bipyramidal geometry, why are equatorial positions less crowded?
Equatorial positions have 120° separation and fewer 90° interactions compared to axial positions.