Intermolecular Forces and Molecular Shape

Comprehensive practice flashcards covering intermolecular forces (LDF, dipole-dipole, H-bonding, ion-dipole), phase changes, nonmolecular solids, VSEPR theory, molecular geometries, and polarity.

What is the primary difference between intramolecular forces and intermolecular forces?

Intramolecular forces are forces within molecules (chemical bonds), whereas intermolecular forces are forces between separate molecules.

What causes a dipole-dipole interaction and what factor increases its strength?

Dipole-dipole interactions occur when two polar molecules come into close proximity and align their opposite partial charges ($\delta^+$ and $\delta^-$); the attraction increases with a stronger dipole, which results from a larger difference in electronegativity.

What structural condition is required for hydrogen bonding to occur, and why is it considered a "super dipole"?

Hydrogen bonding occurs in molecules where hydrogen is directly attached to fluorine, oxygen, or nitrogen ($F$, $O$, or $N$). Because of the large electronegativity difference, an exceptionally strong dipole with high charge separation is formed.

Why do DNA sequences rich in cytosine and guanine (G-C) melt at a higher temperature than sequences rich in adenine and thymine (A-T)?

$G-C$ base pairs link together via $3$ hydrogen bonding sites, whereas $A-T$ base pairs have only $2$ hydrogen bonding sites, requiring more thermal energy to separate $G-C$ pairs.

What causes London dispersion forces (LDFs), and in which substances do they occur?

LDFs are caused by the motion of electrons creating temporary, unequal distributions of charge (temporary dipoles) that induce temporary dipoles in neighboring molecules. They occur in all substances and are the only intermolecular force present in nonpolar molecules.

What is polarizability, and what two structural factors cause it to increase?

Polarizability is the ability of an atom or molecule to form a temporary or induced dipole (how malleable its electron cloud is). It increases as the number of electrons increases and as surface area increases.

Between n-pentane and the isomer shown (neopentane), which molecule exhibits weaker London dispersion forces and why?

Neopentane exhibits weaker London dispersion forces because its compact, spherical shape provides less surface area for intermolecular contact and polarizability compared to the straight-chain structure of $n\text{-pentane}$.

What is the relative order of strength among intermolecular forces and chemical bonds from weakest to strongest?

Dispersion ($0\text{--}2\,kJ\,mol^{-1}$) $<$ dipole-dipole ($2\text{--}10\,kJ\,mol^{-1}$) $<$ H-bonding ($10\text{--}20\,kJ\,mol^{-1}$) $<$ ion-dipole ($20\text{--}100\,kJ\,mol^{-1}$) $<$ covalent bonds ($100\text{--}400\,kJ\,mol^{-1}$) $<$ ion-ion interactions ($400\text{--}1000+\,kJ\,mol^{-1}$).

Under what condition is an attraction value of zero possible between particles?

ZERO attraction is only possible with ideal gases.

How do water molecules orient themselves around dissolved sodium cations ($Na^+$) and chlorine anions ($Cl^-$) during hydration?

The partially negative oxygen atoms ($\delta^-$) orient toward the positively charged sodium cations ($Na^+$), while the partially positive hydrogen atoms ($\delta^+$) orient toward the negatively charged chlorine anions ($Cl^-$).

What two factors governed by Coulomb's Law dictate the strength of an ion-dipole interaction?

Charge and distance; the smallest ion with the largest charge experiences the strongest ion-dipole interaction.

When a molecular compound undergoes a phase change (such as ice melting into water), what happens to its chemical bonds and intermolecular forces?

Chemical bonds and chemical composition do not change; only the spacing, movement, and intermolecular forces between molecules change.

Which phase transitions absorb energy (endothermic), and which release energy (exothermic)?

Melting (fusion), vaporization, and sublimation absorb heat energy to weaken/break intermolecular forces; freezing, condensation, and deposition release heat energy as intermolecular forces strengthen/form.

What is the ranking of ethane, dimethyl ether, and ethanol in order of increasing boiling point, and what justifies this trend?

$\text{Ethane} < \text{Dimethyl Ether} < \text{Ethanol}$. Ethane experiences only weak London dispersion forces, dimethyl ether experiences dipole-dipole interactions, and ethanol possesses strong hydrogen bonding.

What are nonmolecular solids, what are their three categories, and why do they have high melting and boiling points?

Nonmolecular solids are substances not held together by intermolecular forces, categorized as ionic, metallic, or covalent network solids; they have high melting and boiling points because they are held together by strong chemical bonds.

What is the atomic-level model of a metallic solid, and what occurs when it melts or vaporizes?

Metal atoms release their valence electrons to a delocalized sea of valence electrons surrounding metal cores; when melted or vaporized, metallic solids separate into individual metal atoms.

How do covalent network solids differ from molecular solids, and what are two main examples?

Molecular solids are held together by intermolecular forces, whereas covalent network solids are held together throughout by covalent bonds; main examples include diamond and quartz ($SiO_2$).

What does the solubility rule "like dissolves like" mean regarding intermolecular forces?

Polar substances dissolve in polar substances via dipole-dipole, hydrogen bonding, or ion-dipole interactions, whereas nonpolar substances dissolve in nonpolar substances via London dispersion forces.

What historical drug crisis in 1957–1961 demonstrates the biological importance of 3D molecular shape?

Thalidomide, marketed in Germany in $1957$ for anxiety and insomnia, was pulled in $1961$ after causing over $10{,}000$ birth defects (with an approximate $40\%$ mortality rate) because its two enantiomeric forms differed only in their three-dimensional shape.

What is the core premise of Valence Shell Electron Pair Repulsion (VSEPR) Theory?

Electrons repel one another (bonded repelling bonded, nonbonded repelling nonbonded, and nonbonded repelling bonded), causing electron pairs around a central atom to adopt positions as far apart as possible.

Why is the bond angle in ammonia ($NH_3$, $107^\circ$) smaller than the bond angle in methane ($CH_4$, $109.5^\circ$)?

Nonbonding electrons (lone pairs) repel more strongly than bonded electrons because they are, on average, closer to the central atom, compressing the remaining bond angles.

How is the total number of electron domains calculated for a central atom?

$\text{Number of Electron Domains} = \text{Number of Lone Pairs} + \text{Number of Bonding Domains}$ on the central atom.

What are the molecular geometries for central atoms with 3 electron domains having 0 and 1 lone pair, respectively?

$0$ lone pairs: Trigonal planar ($120^\circ$, e.g., $BCl_3$, $SO_3$); $1$ lone pair: Bent ($<120^\circ$, e.g., $SO_2$, $NO_2^-$, $O_3$).

What are the molecular geometries for central atoms with 4 electron domains having 0, 1, and 2 lone pairs, respectively?

$0$ lone pairs: Tetrahedral ($109.5^\circ$, e.g., $CH_4$); $1$ lone pair: Trigonal pyramidal ($107^\circ$, e.g., $NH_3$); $2$ lone pairs: Bent ($104.5^\circ$, e.g., $H_2O$).

What are the molecular geometries for central atoms with 5 electron domains having 0, 1, 2, and 3 lone pairs, respectively?

$0$ lone pairs: Trigonal bipyramidal (e.g., $PCl_5$); $1$ lone pair: Bisphenoidal / seesaw (e.g., $SF_4$); $2$ lone pairs: T-shaped (e.g., $ClF_3$); $3$ lone pairs: Linear (e.g., $I_3^-$).

What are the molecular geometries for central atoms with 6 electron domains having 0, 1, and 2 lone pairs, respectively?

$0$ lone pairs: Octahedral (e.g., $SF_6$); $1$ lone pair: Square pyramidal (e.g., $ClF_5$); $2$ lone pairs: Square planar (e.g., $ICl_4^-$).

In the structure of allyl alcohol shown, what are the molecular geometries of carbon A ($CH_2=$), carbon B ($=CH-$), and carbon C ($-CH_2-$)?

Carbon A has $3$ bonding domains and $0$ lone pairs (trigonal planar); Carbon B has $3$ bonding domains and $0$ lone pairs (trigonal planar); Carbon C has $4$ bonding domains and $0$ lone pairs (tetrahedral).

How does vector addition explain why symmetric molecules with polar bonds can be nonpolar overall?

Individual bond dipole moments act as vectors; in symmetric geometries with identical peripheral atoms (such as linear $CO_2$ or tetrahedral $CCl_4$), the opposing dipole vectors cancel each other out completely, resulting in a net dipole moment of zero.

According to the symmetry quick-checks, which geometries are automatically polar, and why?

Asymmetric geometries (such as Bent and Trigonal Pyramidal) are automatically polar because of imbalanced lone pair repulsions that prevent bond dipoles from canceling.

Why is chloromethane ($CH_3Cl$) polar while methane ($CH_4$) is nonpolar, even though both have tetrahedral geometry?

$CH_3Cl$ has different terminal atoms ($H$ and $Cl$) with differing electronegativities, meaning the bond dipoles do not cancel, whereas in $CH_4$ all terminal atoms are identical and the bond dipoles cancel completely.