AP Chemistry 2.1-2.2: Types of Bonds and Intramolecular Force

Cation

A positively charged ion formed when a neutral atom loses one or more electrons, usually from its valence shell.

Anion

A negatively charged ion formed when a neutral atom gains one or more electrons, usually to complete its valence shell.

Ionic bond

The electrostatic attraction between oppositely charged cations and anions. It commonly results from electron transfer between a metal and a nonmetal.

Ionic compound

A neutral compound composed of cations and anions held together by electrostatic attractions; ionic compounds are also called salts.

Why do metals commonly form cations while nonmetals commonly form anions?

Metals generally have relatively low ionization energies and lose valence electrons readily. Nonmetals generally have stronger attractions for additional electrons and tend to gain electrons to fill their valence shells.

How does the reaction of sodium with chlorine illustrate the difference between elements and ionic compounds?

Elemental sodium is a highly reactive metal and elemental chlorine is a poisonous, corrosive gas, but their reaction produces NaCl, a stable crystalline solid with very different properties.

What charge-balance condition must an ionic compound satisfy?

The total positive charge must equal the total negative charge, so the compound has no net charge.

What does the formula of an ionic compound represent?

It represents the simplest whole-number ratio of the ions that produces electrical neutrality, not a discrete molecule.

How do you determine the formula of aluminum oxide from the charges of its ions?

Aluminum forms $\mathrm{Al^{3+}}$ and oxygen forms $\mathrm{O^{2-}}$. The smallest neutral combination is two $\mathrm{Al^{3+}}$ ions and three $\mathrm{O^{2-}}$ ions, so the formula is $\mathrm{Al_2O_3}$ because $(2)(+3)+(3)(-2)=0$.

Why is it usually incorrect to call a unit of solid NaCl a molecule?

NaCl forms a three-dimensional crystal lattice in which each ion interacts with multiple nearby ions. Its formula gives the simplest ratio of $\mathrm{Na^+}$ to $\mathrm{Cl^-}$, rather than identifying one isolated NaCl molecule.

Why are ionic solids rigid and brittle?

Strong electrostatic attractions hold the ions in an ordered lattice, making the solid rigid. When layers shift, like charges can become adjacent, causing strong repulsion and fracture.

Why do ionic compounds generally have high melting and boiling points?

A large amount of energy is required to overcome the strong electrostatic attractions throughout the ionic lattice.

Why does solid NaCl not conduct electricity, but molten or aqueous NaCl does?

In solid NaCl, the ions are fixed in place and cannot move to carry charge. When melted or dissolved, the ions are mobile and can conduct electricity.

What are the typical charges of main-group metal cations in Groups 1, 2, and 13?

Group 1 metals commonly form $+1$ ions, Group 2 metals form $+2$ ions, and Group 13 metals commonly form $+3$ ions.

How can typical main-group cation charges be predicted for Groups 13–17?

For many main-group elements in Groups 13–17, loss of all valence electrons gives a cation charge approximately equal to the group number minus 10. For example, Group 13 commonly gives $+3$.

What electron configuration does a main-group atom usually achieve when it forms a cation?

It generally loses its valence electrons and becomes isoelectronic with the noble gas preceding it in the periodic table.

What does it mean for an ion to be isoelectronic with a noble gas?

The ion has the same number of electrons and the same electron configuration as the noble gas, although it has a different nuclear charge.

Write the electron configuration of $\mathrm{Ca^{2+}}$ and identify the noble gas with which it is isoelectronic.

Calcium loses its two $4s$ electrons: $\mathrm{Ca^{2+}=[Ar]}$. It is isoelectronic with argon.

How are electron configurations changed when transition-metal atoms form cations?

The outermost $s$ electrons are removed before electrons from the $(n-1)d$ subshell, even though the $d$ subshell filled after the $s$ subshell in the neutral atom.

Why do transition metals often form more than one common ion?

Their outer $s$ and nearby $d$ electrons have similar enough energies that different numbers of electrons can be removed, producing multiple oxidation states such as $\mathrm{Fe^{2+}}$ and $\mathrm{Fe^{3+}}$.

Write the electron configurations of $\mathrm{Fe^{2+}}$ and $\mathrm{Fe^{3+}}$.

Neutral Fe is $\mathrm{[Ar]3d^64s^2}$. Thus, $\mathrm{Fe^{2+}=[Ar]3d^6$ and $\mathrm{Fe^{3+}=[Ar]3d^5}$.

What are common charges for inner-transition-metal cations?

Inner-transition metals most commonly form $+3$ ions, typically by losing their outermost $s$ electrons and one additional $d$ or $f$ electron.

What is the inert pair effect, and how does it affect heavy main-group elements?

The inert pair effect is the increased reluctance of a heavy atom's valence $s$-electron pair to participate in bonding or ion formation. It makes lower positive charges more common for some heavy elements, such as $\mathrm{Pb^{2+}}$ versus $\mathrm{Pb^{4+}}$.

What unusual polyatomic cation is formed by mercury in addition to $\mathrm{Hg^{2+}}$?

Mercury can form $\mathrm{Hg_2^{2+}}$, a diatomic cation containing an Hg–Hg bond, in addition to the monatomic $\mathrm{Hg^{2+}}$ ion.

How do most monatomic nonmetal anions achieve stable electron configurations?

They gain enough electrons to fill their outer $s$ and $p$ orbitals, usually becoming isoelectronic with the noble gas that follows them.

Write the electron configuration of the oxide ion, $\mathrm{O^{2-}}$.

Oxygen gains two electrons: $\mathrm{O^{2-}=1s^22s^22p^6=[Ne]}$.

What charge does phosphorus commonly form as a monatomic anion, and what is its electron configuration?

Phosphorus gains three electrons to form $\mathrm{P^{3-}}$. Its configuration is $\mathrm{[Ne]3s^23p^6=[Ar]}$.

Covalent bond

A bond formed when two atoms share a pair of electrons. It is especially common between nonmetal atoms with similar tendencies to attract electrons.

Why does sharing electrons stabilize an $\mathrm{H_2}$ molecule?

Each hydrogen atom contributes one electron to a shared pair. The shared electrons are attracted to both nuclei, giving each hydrogen access to two electrons, like helium.

How does potential energy change as two hydrogen atoms approach to form $\mathrm{H_2}$?

Potential energy decreases as attractive interactions between the shared electrons and both nuclei become important. At very short distances, nucleus–nucleus and electron–electron repulsions increase the potential energy.

Bond length

The internuclear distance at which the potential energy of two bonded atoms is minimized. At this distance, attractive and repulsive interactions produce the most stable arrangement.

What is the energy relationship between breaking and forming a chemical bond?

Breaking a bond requires energy and is endothermic, while forming the same bond releases energy and is exothermic.

What enthalpy change corresponds to dissociating one mole of gaseous $\mathrm{H_2}$ molecules into atoms?

The process $\mathrm{H_2(g)\rightarrow 2H(g)}$ requires $\Delta H=+436\ \mathrm{kJ\,mol^{-1}}$. Forming one mole of $\mathrm{H_2}$ from gaseous H atoms releases $436\ \mathrm{kJ\,mol^{-1}}$.

Pure covalent bond

A covalent bond in which bonding electrons are shared equally, usually because the bonded atoms are identical and have equal electronegativities, as in $\mathrm{H_2}$ or $\mathrm{Cl_2}$.

Electronegativity

A dimensionless measure of an atom's tendency to attract shared electrons toward itself in a chemical bond.

What are the periodic trends in electronegativity?

Electronegativity generally increases from left to right across a period and decreases down a group. Fluorine has the highest electronegativity on the Pauling scale, approximately $4.0$.

How does electronegativity differ from electron affinity?

Electron affinity is the measurable energy change when an isolated gaseous atom gains an electron, expressed in $\mathrm{kJ\,mol^{-1}}$. Electronegativity is a dimensionless relative measure of attraction for shared electrons in a bond.

Polar covalent bond

A covalent bond in which electrons are shared unequally because one atom attracts the shared electrons more strongly. The atoms acquire partial charges, written $\delta^+$ and $\delta^-$.

How are partial charges assigned in a polar covalent bond?

The more electronegative atom attracts more electron density and is assigned $\delta^-$. The less electronegative atom is assigned $\delta^+$.

In an H–Cl bond, which atom is $\delta^-$ and why?

Chlorine is $\delta^-$ because it is more electronegative than hydrogen and therefore attracts the bonding electrons more strongly. Hydrogen is $\delta^+$.

How is the polarity of a bond estimated quantitatively?

Calculate the absolute electronegativity difference: $\Delta\mathrm{EN}=|\mathrm{EN}_A-\mathrm{EN}_B|$. A larger $\Delta\mathrm{EN}$ generally indicates greater bond polarity and more ionic character.

How can the polarity of several covalent bonds be ranked?

Calculate $\Delta\mathrm{EN}$ for each bond and rank them from smallest to largest difference. The more electronegative atom in each bond receives the $\delta^-$ designation.

Rank C–C, C–H, Si–C, and Si–O in order of increasing bond polarity.

The order is C–C ($\Delta\mathrm{EN}=0.0$) < C–H ($0.4$) < Si–C ($0.7$) < Si–O ($1.7$). In the polar bonds, the more electronegative atom is $\delta^-$.

How does increasing $\Delta\mathrm{EN}$ affect bonding character?

As $\Delta\mathrm{EN}$ increases, electron sharing becomes less equal and the bond generally develops greater polar covalent and then ionic character. The boundaries are approximate rather than absolute.

How should the bonds H–H, H–Cl, and Na–Cl be classified using electronegativity differences?

H–H has $\Delta\mathrm{EN}=0$ and is nonpolar covalent; H–Cl has an intermediate difference of about $0.9$ and is polar covalent; Na–Cl has a larger difference of about $2.1$ and is commonly classified as ionic.

Why should electronegativity difference not be used as the only way to classify a bond?

Bond-type categories overlap, and the same approximate difference can occur in compounds with different classifications. The types of elements and their positions in the periodic table are also important.

What types of bonding occur in potassium nitrate, $\mathrm{KNO_3}$?

The attraction between $\mathrm{K^+}$ and $\mathrm{NO_3^-}$ is ionic. The N–O bonds within the nitrate ion are covalent, with polar character.

Why do molecular covalent substances generally have lower melting and boiling points than ionic solids?

Covalent molecules are electrically neutral, so attractions between separate molecules are generally weaker than the extensive electrostatic attractions in an ionic lattice.

Why are most molecular covalent substances poor electrical conductors?

They generally consist of neutral molecules without mobile charged particles. Consequently, they usually do not conduct electricity as solids, liquids, or aqueous solutions.

What is metallic bonding?

Metallic bonding is the electrostatic attraction between a lattice of positively charged metal ions and delocalized valence electrons that move throughout the solid.

How does the delocalized-electron model explain the electrical conductivity of metals?

Metallic solids contain mobile delocalized electrons. When an electric field is applied, these electrons can move through the lattice and carry charge.

Why are metals malleable and ductile?

Metallic bonding is nondirectional because the delocalized electrons attract metal ions throughout the lattice. Layers of ions can therefore slide without completely breaking the bonding network.

Why are metals generally lustrous and good conductors of heat?

Delocalized electrons interact with and reflect light, producing luster, and they transfer kinetic energy efficiently through the metallic lattice, giving metals good thermal conductivity.

What is bond energy?

Bond energy is the energy required to break one mole of a particular type of bond in gaseous molecules. Bond breaking is endothermic, so bond energies are positive.

How are bond strength and bond length related?

Stronger bonds generally have shorter equilibrium bond lengths because the bonded atoms are held more tightly together. Weaker bonds generally have longer bond lengths.

How can average bond energies be used to estimate the enthalpy change of a reaction?

Estimate $\Delta H$ by subtracting the energy released in forming bonds from the energy required to break bonds: $\Delta H_{\mathrm{rxn}}\approx\sum E(\text{bonds broken})-\sum E(\text{bonds formed})$.

What does the minimum on a potential-energy-versus-internuclear-distance diagram represent?

The minimum represents the equilibrium bond length, where attractive and repulsive forces balance and the bonded atoms have their lowest potential energy.

How does the depth of a potential-energy well indicate bond strength?

A deeper potential-energy well indicates that more energy is required to separate the atoms, so it represents a stronger bond. A shallow well represents a weaker bond.

Why does the potential energy of two atoms rise sharply at very short distances?

At very short internuclear distances, electron–electron and nucleus–nucleus repulsions become dominant, causing the potential energy to increase sharply.

What is the potential energy of two noninteracting atoms at infinite separation typically defined as?

It is typically defined as zero. A bonded pair at its equilibrium distance has lower, usually negative, potential energy relative to that separated-atom reference.