Chemical Bonding

Question 1

What are valence electrons, and how do they affect an atom’s chemistry?

Answer 1

Valence electrons are the electrons in the atom’s highest (outermost) energy subshells. They are chemically relevant because they are the electrons that form chemical bonds.

Question 2

How would you quickly find the number of valence electrons by using the periodic table?

Answer 2

Valence electron number can quickly be read off of the periodic table, as it equals the column number (counting from the left) that the element is in.

For example, nitrogen is in the 5th column from the left, so it must have 5 valence shell electrons. Nitrogen is 1s²2s²2p³, and n = 2 is the outermost energy level. Add all of those electrons (s² and p³) to get valence = 5.

Question 3

What is the difference in the valence electron configuration of oxygen (O) and silicon (Si)?

Answer 3

Oxygen has 6 electrons in the outermost energy level (2s² and 2p⁴).

Silicon has more electrons total, but it only has 4 valence electrons (3s² and 3p²).

Question 4

Define:

the octet rule

Answer 4

The octet rule states that atoms desire eight electrons in their valence shells, as this gives them the electron configuration of a noble gas.

Most atoms have to bond with other atoms to acquire this electron configuration.

Question 5

What energy principle causes two atoms to form a bond between them?

Answer 5

Chemical bonds form because they lower the potential energy of the electron clouds. Electrons are shared between atoms across the bond, allowing the final state to be more stable than the two were alone.

In general, bonding proceeds so that as many atoms as possible gain a full octet of valence shell electrons.

Question 6

What are the three possible types of interatomic bonds?

Answer 6

1. single bond: one pair of electrons is shared between two atoms
   For example, an O-H bond in H₂O, or a C-H bond in CH₄.
2. double bond: two pairs of electrons are shared between atoms
   For example, O=O bonds in O₂, or C=O bonds in CO₂.
3. triple bond: three pairs of electrons are shared between atoms
   For example, N≡N bonds in N₂, or C≡N bonds in HCN.

Question 7

What are the three notable exceptions to the octet rule?

Answer 7

1. odd-electron species
   Molecules like NO must break the octet rule since they have an odd total number of valence electrons to distribute.
2. incomplete octets
   These are atoms where attaining a full octet would require too many bonds. H, He, Li, Be, B are all considered incomplete octet species due to the high number of electrons they would need.
3. expanded octets
   Atoms in row 3 or higher, like P or Cl, can hold electrons in their d orbitals. These atoms can hold between 10 and 18 electrons depending on the central atom being bonded.

Question 8

Why does nitrogen not succeed in getting a full octet in nitric oxide (NO)?

Answer 8

The most electronegative atom will always gain a full octet first.

Nitric oxide has 11 valence electrons. Oxygen is more electronegative, so will end up with a full octet (two sets of lone e-pairs and a double bond with N). Nitrogen will be left with 3 lone electrons (and two bonds), as it is less electronegative, for a total of 7 electrons.

Question 9

Why will it be unlikely for Li to ever attain a full octet?

Answer 9

Lithium has only one valence electron. It would need to acquire 7 additional electrons in order to have a full octet. This is statistically (and structurally) unlikely.

H, He, Li, Be, B are all elements that will form incomplete octets due to the improbability of them acquiring enough electrons.

Question 10

Why can phosphorus expand its octet to form 5 bonds to chlorine in PCl₅?

Answer 10

Phosphorus can expand its five valence electrons into the 3d block, allowing it to bond to five chlorine atoms and completing those chlorine atoms’ octets.

Elements in the third row of the periodic table and beyond (such as P, Cl, S, Xe, and Ar) often exhibit expanded octets of 10 to 18 electrons. This is due to them expanding into the d-block for greater bonding stability.

Question 11

What are the three major types of intramolecular bonds?

Answer 11

1. ionic bonds: electrons are transferred from a metal to nonmetal
2. covalent bonds: electrons are shared between nonmetals
3. metallic bonds: electrons “float” between a lattice of metallic nuclei

Note that intramolecular bonds connect atoms of the same molecule, unlike intermolecular forces (like hydrogen bonds).

Question 12

Define:

an ionic bond

Answer 12

Ionic bonds form when a metal transfers one or more electrons to a nonmetal.

On the MCAT, often this will be a metal from Group I or II transferring electrons to a halogen. The classic example is an “ionic salt” like NaCl. Sodium transfers its one valence electron to chlorine, leaving each with an octet of valence electrons.

Question 13

How does Coulomb’s law apply to ionic bonds?

Answer 13

Coulomb’s law states that the magnitude of the electrostatic force between two charged particles is directly proportional to the product of the magnitude of each of the charges, and inversely proportional to the square of the distance between the two particles.

This holds for all charged particles, and can be applied to calculate force between the atoms in an ionic bond.

F ∝ q₁*q₂ / r²

• q₁ and q₂ = charge magnitude of the ions
• r = distance between the ions

Question 14

Knowing that Na and Cl have a larger electronegativity difference than K and Br, what does Coulomb’s law predict about the strength of the NaCl bond vs. that of KBr?

Answer 14

NaCl will have stronger bonds between adjacent atoms than KBr, due to the stronger electrostatic force between ions.

Question 15

How would the force between ions in an ionic compound change if the distance between the adjacent ions was decreased by one-half?

Answer 15

The electrostatic force would be 4 times the original value.

F_bond ∝ q₁*q₂ / r²

Since new R = r/2:

F_new ∝ q₁*q₂ / R²
= q₁*q₂ / (r/2)²
= q₁*q2 / (r²/4)
= F_orig/(1/4) = F_orig*4

Question 16

Define:

lattice energy

Answer 16

The lattice energy of an ionic compound is the energy associated with forming a crystalline lattice of the compound from the gaseous ions.

This may also be referred to as “heat of formation.” Note that the value of lattice energy is negative, showing that the formation of an ionic compound is exothermic.

Question 17

What is the electrostatic energy (E_es) of an ionic bond?

Answer 17

E_es ∝ q₁*q₂ / r

This holds for all charged particles, and thus can be applied to calculate energy between the atoms in an ionic bond.

• q₁ and q₂ = charge magnitude of the atom
• r = distance

Note: the value of E_es will always be negative, due to the charges always being opposite in an ionic compound.

Question 18

How does the energy of an ionic bond change if the distance between the two atoms in the bond doubles?

Answer 18

The final energy will be one-half the initial value.

E_orig ∝ q₁*q₂ / r

Since new R = 2r:

E_new ∝ q₁*q₂ / R

q₁*q₂ / 2r = E_orig/2

Question 19

What type of bond will exist between atoms in molecules like KBr, CaF₂, and LiCl?

Answer 19

These are all ionic compounds, commonly referred to as salts. Ionic bonds will generally form between metals and nonmetals.

Classic examples of ionic compounds will usually include a cation from Group I or II bonded to one or two anions from the halogen family.

Question 20

Define:

a covalent bond

Answer 20

A covalent bond forms when a nonmetal bonds with another nonmetal by sharing electrons between them, resulting in an overlap of their electron orbitals.

The image below of methane shows the polar covalent C-H bonds.

Question 21

What are the three types of covalent bonds?

Answer 21

1. polar covalent bond: electrons are held more closely by the higher electronegative species
2. nonpolar covalent bond: electrons are perfectly shared between atoms of the same element
3. coordinate covalent bond: entire molecules share electrons for greater net stability

Question 22

What is the difference in bond strength between molecules in an ionic compound vs. a covalent compound, and how will the compounds’ melting points and boiling points compare?

Answer 22

Ionic bonds create stronger intermolecular forces than pure covalent bonds. Since ionic bonds are extremely polar (more so than any covalent bond, by definition), a strong electrostatic interaction exists between the ions.

The melting point and boiling point will be higher for ionic compounds, as more energy is required to pull the atoms in the bonds apart.

Question 23

Define:

a nonpolar covalent bond

Answer 23

In a nonpolar covalent bond, both atoms are of the same element. Hence, the electron pair is shared equally. On the MCAT, the only commom pure covalent bonds are the following diatomics: Br₂, I₂, N₂, Cl₂, H₂, O₂, F₂ (BrINClHOF).

Note that although some chemistry texts will also include bonds like the C-H bond (since the electronegativity difference is only 0.4 on the Pauling scale), that bond is still slightly polar since electrons favor carbon over hydrogen.

Question 24

Define:

a polar covalent bond

Answer 24

In a polar covalent bond, the electron pair is pulled closer to the more electronegative atom. The result of this is a bond dipole (one end positive, one end negative), hence the term “polar.” The atom that is more electronegative will carry a partial negative charge and the atom that is less electronegative will carry a partial positive charge.

Examples of polar covalent bonds include O-H bonds, N-H bonds, and C-N bonds.

Question 25

Define:

a coordinate covalent bond

Answer 25

In a coordinate covalent bond between molecules, one atom from one molecule will contribute both electrons to the bond pair. This creates better stability between both molecules.

A common example is a Lewis acid / base pair: NH₃ (N donates the electron pair) and H⁺ (H⁺ accepts the electron pair). Nitrogen had a full octet to start with, but was very polar until donating the electrons to the bond with hydrogen. H⁺ didn’t have any electrons, but now will have a full 1s subshell. Both are now more stable.

The first equation below shows all of the actual valence electrons. The second equation is the Lewis diagram representation.

Question 26

Why do Lewis acids and Lewis bases generally combine to form coordinate-covalent bonded molecules, instead of common ionic salts?

Answer 26

By definition, a coordinate-covalent bond occurs when one molecule donates both electrons to the covalent bond. Lewis acid/base pairs combine in exactly that fashion.

A Lewis acid is a species that will accept an electron pair (like boron in BF₃).
A Lewis base is a species that will donate an electron pair (like nitrogen in NH₃).

Question 27

Describe the bond formed between two nonmetallic elements.

Answer 27

Nonmetallic elements form covalent bonds.

Nonmetals tend to have similar values of electronegativity, and thus share electron density fairly evenly when bound together.

Question 28

Describe the bond formed between a metal and a nonmetal in a compound.

Answer 28

Metal-nonmetal bonds will be ionic.

The electronegative nonmetal withdraws one or more electrons from the electropositive metal, leaving each species ionically charged.

Question 29

Describe the bonding for metals in their standard states.

Answer 29

Metals form a lattice of positively charged nuclei in a background “sea” of free-flowing negatively charged electrons.

The valence electrons of metals are loosely bound to the atomic cores, and can be considered to be effectively unattached.

Question 30

Explain why metals, or metallic bonded elements, are good conductors of electricity.

Answer 30

The “sea” of electrons is able to drift through the electron structure, moving from ion to ion, giving these structures the ability to conduct electricity.

In metallic bonding, each metal atom in the crystal structure contributes valence electrons to form a “sea” of delocalized electrons.

Question 31

What makes an entire molecule polar?

Answer 31

A molecule is polar when the polar covalent bonds in the molecule add via the molecular geometry to give the entire molecule a positive and a negative end.

For example, H₂O has polar bonds between O-H, with oxygen more electronegative = – (red in the picture) and hydrogen = + (blue in the picture). There is a negative region (top) and a positive region (bottom). H₂O is therefore a polar molecule.

Question 32

What makes a molecule nonpolar?

Answer 32

When a molecule has a balanced charge distribution, such that no comparably positive or negative region exists, it is nonpolar.

Question 33

When will a molecule with polar bonds still be nonpolar?

Answer 33

Polar covalent bonds can still create a nonpolar molecule if the geometry is symmetrical, such that the dipoles will cancel and create a balanced molecule.

For example, In BF₃ the B-F bond will pull electrons towards the more electronegative F atoms = – (red in the picture) and away from the B atom = + (blue in the picture). Though these are polar bonds, the trigonal planar symmetry means the molecule has no net dipole.

Question 34

Is CO₂ a polar or nonpolar molecule?

Answer 34

CO₂ is a nonpolar molecule.

Although the C=O bond is polar, the symmetrical, linear shape of the molecule (shown below) results in cancelled dipoles and an even distribution of charge.

Question 35

What is the difference between a bonding pair and a lone pair of electrons?

Answer 35

• bonding pair: a pair of electrons that is shared between two atoms across a bond (represented by a line in the Lewis diagram)
• lone pair: a pair that is associated with only one atom and stays on just that atom (represented by two close dots in the Lewis diagram)

Question 36

Define:

dipole moment

Answer 36

A dipole moment refers to any polar bond with a separation of partial positive and negative charge. By common convention, the partial charges δ+ and δ- are used to indicate the positive and negative ends of the bond, respectively.

Question 37

Which compound contains bonds with a higher dipole moment?

1. CH₄ vs, CO₂
2. NaCl vs. H₂O
3. O₂ vs. NH₃

Answer 37

1. The C=O bonds in CO₂ have a higher dipole moment, due to oxygen’s high electronegativity, than the less polar C-H bonds in CH_4.
2. The ionic NaCl bond has a higher dipole moment than the covalent O-H bonds in H₂O.
3. The polar covalent N-H bonds in NH₃ have a higher dipole moment than the pure covalent O=O bond in O₂.

Question 38

Define:

Lewis structure

Answer 38

A Lewis structure is a representation of a molecule that shows how electrons are being shared between atoms.

Characteristically:

• every line represents a shared electron pair
• every dot represents one electron

For example, the image below (for H₂O) shows the original atoms on the left, then how confusing it would look without a Lewis structure; on the far right, the proper Lewis structure is displayed.

Question 39

Define:

the formal charge of an atom in a molecule

Answer 39

The formal charge of an atom is the charge it would have if all bonding electrons were shared equally between the bonded atoms.

Typically, a molecule’s ideal Lewis structure will have zero formal charge on each atom.

Question 40

How is the formal charge on an atom calculated?

Answer 40

Formal charge = (# of valence electrons) - (# of lone pair electrons) - (1/2 # of bonding electrons)

Formal charge is essentially a fictitious charge assigned to each atom in a Lewis structure for the sake of helping distinguish the best Lewis structure for a molecule.

Question 41

What is an easier method for calculating formal charge of an atom in a compound, given the compound’s Lewis structure?

Answer 41

Formal charge = # valence electrons - lines - dots

For example, in CO₂, carbon has 4 valence electrons, 4 lines, and 0 dots. FC = 4-4-0 = 0. Oxygen has 6 valence electrons, 4 dots, and 2 lines. FC = 6-4-2 = 0.

Question 42

Define:

resonance structures

Answer 42

A resonance structure exists when two or more valid (stable) Lewis structures can be drawn for the same compound.

In resonance structures, only the strength of bonds between atoms varies, not the actual placement of the atoms. If multiple resonance structures exist, the actual molecule’s structure is an average of all of these.

For example, the benzene molecule below has two possible structures, so the C-C bonds have an average bond order of 1.5 in any given position due to resonance.

Question 43

How is the average bond strength between any two specific atoms in a molecule with multiple resonance structures calculated?

Answer 43

Average bond strength is simply the average strengths of the bond between the atoms in each resonance structure.

In short, total the bonds between the two atoms in each resonance structure, then divide by the total number of structures.

For example, to calculate the N-O bond order in nitrate (below), consider the upper oxygen. There are 2 bonds + 1 bond + 1 bond = 4 total bonds possible in that position across the 3 unique structures. Hence, the average bond strength is 4/3.

Question 44

Define:

valence shell electron pair repulsion theory

Answer 44

The valence shell electron pair repulsion (VSEPR) theory states that electron pairs repel each other. Therefore, for a molecule to be at its most stable state, electron pairs should be as far from each other as possible in three-dimensional space.

Question 45

According to the valence bond theory, what does a chemical bond result from?

Answer 45

A chemical bond occurs when one partially-filled orbital of one atom overlaps with a partially-filled orbital of another atom, allowing both to attain a more stable state by sharing electrons. A single bond has one shared pair, a double bond has two pairs (4 total), and a triple bond has three pairs (6 total).

Note that atoms typically form enough bonds to attain a full octet of valence electrons.

Question 46

Which is more electronically repulsive, a pair of electrons in a bond or a nonbonded pair of electrons?

Answer 46

A pair of nonbonded electrons has more negative character than a pair of bonding electrons, since the nonbonded electrons are not shared between two positive nuclei.

So, two lone pairs will repel each other the most, while two bonded pairs will suffer the least repulsion. The repulsion between a lone pair and a bonded pair falls somewhere in between.

Question 47

Describe the characteristics and give an example of:

a linear molecule

Answer 47

A linear molecule:

• has 3 atoms bonded
• has atoms arranged 180 degrees apart
• usually contains two double bonds, or one triple and one single bond

Classic examples include CO₂ (double-double) and HCN (single-triple).

Question 48

Describe the characteristics and give an example of:

a bent molecule

Answer 48

A bent molecule:

• has 3 atoms bonded
• has a 104.5 degree bond-bond angle
• contains two single bonds and two lone e^- pairs

Classic examples include H₂O and SCl₂.

Question 49

Describe the characteristics and give an example of:

a trigonal planar molecule

Answer 49

A trigonal planar molecule:

• generally has 4 atoms bonded (one central atom and 3 peripheral)
• has a 120 degrees bond-bond angle

Classic examples include SO₃ (3 double bonds), BF₃ (3 single bonds), and H₂CO (two single bonds, one double bond).

Question 50

Describe the characteristics and give an example of:

a trigonal pyramidal molecule

Answer 50

A trigonal pyramidal molecule:

• generally has 4 atoms bonded (1 central and 3 peripheral)
• has a 107 degree bond-bond angle
• contains 3 single bonds, and one lone e^- pair

Classic examples include NH₃ and PCl₃.

Question 51

Describe the characteristics and give an example of:

a tetrahedral molecule

Answer 51

A tetrahedral molecule:

• generally has 5 atoms bonded (1 central and 4 peripheral)
• has a 109.5 degrees bond-bond angle
• contains 4 single bonds

Classic examples include CH₄ and CCl₄.

Question 52

Define:

molecular geometry

Answer 52

Molecular geometry is the actual placement of atoms in a molecule, ignoring any nonbonded electrons.

For example, H₂O has two hydrogen atoms placed in a “bent” structure, so it has a bent molecular geometry.

Question 53

Define:

electronic geometry

Answer 53

Electronic geometry is the position of all bonding electrons and lone e^- pairs around one central atom.

For example, in H₂O, oxygen has two hydrogens bonded and two lone e- pairs in a skew position to those bonds, for 4 total electron positions in a tetrahedral shape around the oxygen. H₂O has tetrahedral electronic geometry, while its molecular geometry is bent.

Question 54

What is a hybrid orbital, and what are the three most common hybrid orbital types?

Answer 54

A hybrid orbital is formed when standard atomic orbitals combine. They are given a name representing which orbitals have overlapped.

The most common hybrid orbitals are sp, sp², and sp³ orbitals.

• sp: an s orbital and a p orbital
• sp²: an s orbital and two p orbitals
• sp³: an s orbital and three p orbitals

Question 55

Describe an:

sp hybridization

Answer 55

• One s orbital mixes with one p orbital, creating two energetically equivalent hybrid sp orbitals.
• The two remaining unhybridized p orbitals lie at right angles to the plane of the hybrid orbitals.
• The molecule will be linear in geometry.

Common examples include C₂H₂ and BeCl₂.

Question 56

Describe an:

sp² hybridization

Answer 56

• One s orbital mixes with two p orbitals.
• The one remaining unhybridized p orbital lies at a right angle to the plane of the three hybrid orbitals.
• The resulting molecule will be trigonal planar in geometry.

Common examples include H₂CO and BF₃.

Question 57

Describe an:

sp³ hybridization

Answer 57

• One s orbital mixes with three p orbitals.
• The resulting molecule will be tetrahedral in geometry.
• Four single bonds are oriented around a central atom, usually carbon.

Common examples include CH₄ and CCl₄.

Question 58

Define:

a sigma bond

Answer 58

A sigma bond occurs when atomic orbitals overlap end-to-end and result in an accumulation of electron density directly between the nuclei.

Sigma bonds generally exist as single bonds, but a sigma bond also makes up one bond of any double or triple bond.

Question 59

Define:

a pi bond

Answer 59

Pi bonds are parallel regions of electron density that form orbitals alongside an initial sigma bond. They result from p orbitals overlapping side by side so the electron density is above and below the internuclear axis.

Pi bonds make up one bond in a double bond and two bonds in a triple bond.

Question 60

How many sigma and/or pi bonds are there in the following?

1. single bonds
2. double bonds
3. triple bonds

Answer 60

1. A single bond contains one sigma bond and no pi bonds.
2. A double bond contains one sigma bond and one pi bond.
3. A triple bond contains one sigma bond and two pi bonds.

Question 61

How many sigma and pi bonds does acetylene, C₂H₂, contain?

Answer 61

C₂H₂ contains a total of 3 sigma and 2 pi bonds.

• The two C-H bonds are single bonds, meaning there are 2 sigma bonds.
• The C≡C is a triple bond, yielding an additional sigma and 2 pi bonds.

Question 62

How is the hybridization of a carbon atom determined?

Answer 62

1. Draw the molecule (or look at the Lewis structure).
2. Count the number of atoms bonded around the carbon.
3. The hybridization of carbon = sp^(x-1), where x is the number of atoms bound to the carbon.

Ex: CH₄ has 4 atoms bonded around carbon, so its hybridization must be sp^(4-1)=sp³

Question 63

What is the hybridization of carbon in the following?

1. H₂CO
2. CCl₄
3. CO₂
4. HCN

Answer 63

Remember that if x total atoms are bonded to the carbon, then the carbon’s hybridization is sp^(x-1).

1. H₂CO: sp² (three atoms bonded)
2. CCl₄: sp³ (four atoms bonded)
3. CO₂: sp (two atoms bonded)
4. HCN: sp (two atoms bonded)

Note that sp² hybridization can involve either the double-double bonds in CO₂ or the single-triple bonds in HCN.

Question 64

Why do central or interior atoms have the highest tendency to hybridize?

Answer 64

A central atom by definition is the atom in a molecule that is bonded to more than one other atom, has the highest number of valence electrons, and has the lowest electronegativity.

All of these qualities make it likely to hybridize in order to have more active valence electrons. Central atoms, then, tend to form more bonds.