Bonding (DELETE)

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 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 will be the column number (counting from the left) that the element is in.

Ex: Nitrogen is in the 5th column from the left, so it must have 5 valence shell electrons. Nitrogen is 1s²2s²2p³, n=2 is the outermost energy level. Add all of those electrons (2 from the s, 3 from the p) to get valence = 5.

Question 3

What is the difference in valence electrons for:
Oxygen (O) vs Silicon (Si)?

Answer 3

Oxygen has 6 electrons in the outermost energy level (2 in the 2s and 4 in the 2p).

Silicon has more electrons total, but it only has 4 valence electrons (2 in the 3s and 2 in the 3p).

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.

An atom will usually have to bond with other atoms to aquire this electron structure.

Question 5

According to energy rules, what will cause 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 bond?

Answer 6

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

Question 7

What are the three notable exceptions to the octet rule?

Answer 7

1. odd-electron species. Molecules which have an odd total number of electrons to distribute (hence some atom will come up lacking)
2. incomplete octets. 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 period 3 or higher that expand to fill their d-block. The term is a misnomer, as this is no longer an octet, but will have 10, 12, or 14 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 from incomplete octets due to the improbability of them acquiring enough electrons.

Question 10

Why can phosphorous 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, 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, 12, or 14 electrons. This is due to them expanding into the d-block for greater bonding stability.

Question 11

What are the three major types of chemical 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 metals in a lattice

Question 12

Define and give an example of:

an ionic bond

Answer 12

An ionic bond is formed when a metal transfers one or more electrons to a nonmetal. Often on the MCAT this will be column I and II transfering electrons to a halogen.

Ex: The classic example is an “ionic salt” like NaCl. Sodium transfers its electron to chlorine.

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, hence can be applied to calculate force between the atoms in an ionic bond.

F ∝ q₁*q₂ / r²

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

Question 14

If we know that NaCl has a higher electrostatic force between atoms in its lattice than does KBr, what does that tell us about the strength of the bonds?

Answer 14

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

Question 15

What would be the result in force between ions in an ionic compound if we 1/2 the distance between adjacent ions?

Answer 15

The force would be increased by a factor of 4, or quadrupled from the original value.

E ∝ q₁*q₂ / r²

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

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 refered to as “heat of formation”.

Note: The value of lattice energy is negative, showing that the formation of an ionic compound is exothermic.

Question 17

How could we evaluate the electrostatic energy between ionic bonds?

Answer 17

E ∝ q₁*q₂ / r
This holds for all charged particles, hence 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 will always be negative, due to the charges always being opposite in an ionic compound.

Question 18

What will be the result in energy if we double the distance between two atoms in an ionic bond?

Answer 18

The energy will be decreased by a factor of 2, or halved from the initial value. Notice that doubling r in the denominator will effectively be the same as dividing by two.

E ∝ q₁*q₂ / r

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

Question 19

What type of bond would we expect between atoms in molecules like KBr, CaF₂, and LiCl?

Answer 19

These are all ionic compounds, commonly referred to as salts. We expect highly ionic bonds between metals and nonmetals in general.

Ex: Classic MCAT examples usually include column I and II bonded with the halogen family.

Question 20

Define:

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 (electrons are held more closely by the higher electronegative species)
2. non-polar covalent bond (electrons are perfectly shared between atoms of the same element)
3. coordinate covalent bond (molecules share electrons for greater net stability)

Question 22

What is the difference in bond strength between molecules in an ionic-bonded compound vs a covalent-bonded compound? How will this affect melting point and boiling point?

Answer 22

Ionic bonds create stronger intermolecular forces than pure covalent bonds. This is due to ionic bonds being extremely polar (more so than ANY covalent bond, by definition), hence it will attract its substituent molecules more strongly.

The melting point and boiling point will both be higher for the ionic compount, since it will require more energy to pull the molecules apart.

Question 23

Define and give examples of:

a nonpolar (or pure) covalent bond

Answer 23

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

Note: though some chemistry texts will also include bonds like the C-H bond (since it only measures .4 on the Pauling scale), this is clearly still polar since electrons favor carbon over hydrogen.

Question 24

Define and give examples of:

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 where we get the term “polar”.

The atom that is more electronegative has a partial negative charge and the atom that is less electronegative has a partial positive charge.

Question 25

# Define and give an example of: 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. ## Footnote Ex: Lewis Acid/Base pairs: 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; Hydrogen didn't have any electrons, but now will have a full 1s subshell; both are now more stable. The first equation below is showing all of the actual valence electrons, the second equation below is the Lewis diagram representation.

Question 26

Specifically, 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 is sharing electrons from one molecule to another. Lewis acid/base pairs combine in exactly that fashion. ## Footnote A Lewis acid is a species that will accept an electron pair (ex: Boron in BF₃). A Lewis base is a species that will donate an electron pair (ex: Nitrogen in NH₃).

Question 27

Describe the bond formed between two nonmetallic elements.

Answer 27

Nonmetallic elements form covalent bonds. ## Footnote 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

Metals and nonmetals form ionic bonds in compounds together. ## Footnote 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

A lattice of positively charged nuclei in a background "sea" of free-flowing negatively charged electrons. ## Footnote 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 are able to drift through the electron structure, moving from ion to ion gives these molecules the ability to conduct electricity. ## Footnote 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 vs negative end. ## Footnote Ex: H₂O has polar bonds between O-H, with Oxygen more electronegative = **–** (red in the picture) and Hydrogen = **+** (blue in the picture). We'll have a general region of negative (top) and general positive (bottom). H₂0 must be polar overall!

Question 32

What makes a molecule nonpolar?

Answer 32

When a molecule has a balanced charge distribution, such that there is not any one general positive or negative region, it is nonpolar.

Question 33

Is it possible for a molecule with polar bonds to still be nonpolar?

Answer 33

Yes. Polar covalent bonds can still create a nonpolar molecule if there is symmetry in the geometry, such that the dipoles will cancel and create a balanced molecule. ## Footnote Ex: 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 keeps any one side from being net negative vs net positive.

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 shape of the molecule cancels the two dipoles and creates an even distribution of charge. ## Footnote Ex: in the following diagram Oxygen atoms are red, Carbon atom is black.

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

**Dipole moment** refers to any bond where there is a seperation of positive and negative charge across it. By common convention we use the partial charges +δ and δ- for the positive and negative ends of the bond respectively.

Question 37

Which will have bonds with a higher dipole moment? 1. CH₄ vs CO₂ 2. NaCl vs H₂O 3. O₂ vs NH₃

Answer 37

1. CO₂ will have a higher dipole moment in the polar C=O bonds (due to Oxygen being highly electronegative), than CH₄ will in the less polar C-H bonds. 2. NaCl will have a higher dipole moment in the ionic Na-Cl bonds (due to Chlorine being significantly more electronegative), than H₂O will in the polar O-H bonds. 3. NH₃ will have a higher dipole moment in the polar N-H bonds (due to Nitrogen being slightly more electronegative), than O₂ will in the pure covalent O=O bonds.

Question 38

# Define and give the two most important characteristics of: **a Lewis structure**

Answer 38

A **Lewis structure** is a representation of a molecule that shows how electrons are being shared beween atoms. Characteristically: * every line represents a shared electron pair * every dot represents one electron Ex: The image below (for H₂O) shows the original atoms on the left, then how confusing it would look without a Lewis structure, then on the far right the proper Lewis structure is displayed.

Question 39

# Define and give the ideal value of: **formal charge**

Answer 39

The **formal charge** of an atom in a Lewis structure is the charge it would have if all bonding electrons were shared equally between the bonded atoms. Ideally, an atom should have a FC of zero.

Question 40

# Give the formal equation for: formal charge

Answer 40

Formal charge = (# of valence electrons) - (# of lone pair electrons) - (1/2 # of bonding electrons) ## Footnote FC is essentially a fictious charge assigned to each atom in a Lewis structure just for the sake of helping distinguish the best Lewis structure for a molecule.

Question 41

How would you actually calculate formal charge with an easier, shortcut formula?

Answer 41

In order to calculate formal charge in a given Lewis structure diagram use: FC = # valence electrons - lines - dots. Ex: in CO₂, Carbon has 4 valence, 4 lines and 0 dots, FC = 4-4-0 = 0. Oxygen has 6 valence and 4 dots and 2 lines, FC = 6-4-2 = 0.

Question 42

# Define: **resonance structure**

Answer 42

A **resonance structure** is when two or more valid (stable) Lewis structures can be drawn for the same compound. ## Footnote In resonance structures only the strength of bonds between atoms, not the actual placement of the atoms. If multiple resonance structures exist, the "hybrid" molecule is thought to exist as an average of all of these structures. Ex: the benzene molecule below has two possible structures, so we consider the C-C or C=C bonds to have an average 1.5 strength no matter what position on the hybrid, due to resonance.

Question 43

How do you calculate the average bond strength between any two specific atoms in a resonance hybrid molecule?

Answer 43

Average bond strength is quite simply: the average stregths of all bonds possible in that position. Take any two atoms, count the total number of bonds in that position across all resonance structures, then divide by the total number of structures. Ex: Nitrate (below), taking the top Oxygen and central Nitrogen, we see that there are 2 bonds + 1 bond + 1 bond =4 total in that position. There are three total structures. Hence the 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 of electrons repel each other. Therefore, in order for a molecule to be at its most stable state, electron pairs should be as far from each other as possible.

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). ## Footnote Note: The atoms can be considered to be trying to attain a full octet of valence electrons.

Question 46

What is the trend for strength of electron group repulsion?

Answer 46

When considering electron pair repulsion, remember that it is the negative character of the electrons that forms the basis of the repulsion. 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. ## Footnote So, two lone pairs will repel each other the most, while two bonded pair will suffer the least repulsion. The repulsion between a lone pair and a bonded pair is 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. * atoms are 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), 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 non-bonded electrons. ## Footnote Ex: H₂O has two Hydrogen atoms placed in a "bent" structure, so this is 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. ## Footnote Ex: in H₂O, Oxygen has two Hydrogens bonded and two lone e- pairs skew to those bonds, for 4 total electron positions in a tetrahedral shape around the Oxygen. H₂O has tetrahedral electronic geometry.

Question 54

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

Answer 54

A hybrid orbital is formed when standard atomic orbitals combine, and are given a name that represents which orbitals have overlapped. The most common hybrid orbitals are sp, sp², and sp³ orbitals. * sp: an s orbital and a p orbital hybridize * sp²: an s orbital and two p orbitals hybridize * sp³: an s orbital and three p orbitals hybridize

Question 55

# Describe an: sp hybridization

Answer 55

* One s orbital mixes with one p orbital creating two energetically equilavalent hybrid sp orbitals. * The two remaing unhybridized p orbitals lie at right angles to the plane of the hybrid orbitals. * The molecule will be linear in geometry. Common examples: C₂H₂, 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: H₂CO, 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. * There are 4 single bonds around a central carbon atom. Common examples: CH₄, 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. ## Footnote Sigma bonds generally exist as single bonds, but a sigma bond also makes up the first bond of a double bond or triple bond.

Question 59

# Define: a **pi bond**

Answer 59

**Pi bonds** are parallel regions of electron density that overlap and form orbitals alongside an initial sigma bond. They are generally the result of p orbitals overlaping side by side so the electron density is above and below the internuclear axis. ## Footnote Pi bonds make up the second bond in a double bond, and the second and third 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 ethene, C₂H₂, contain?

Answer 61

* The two C-H bonds are single bonds = 2 sigma bonds * The C≡C form a triple bond = 1 sigma and 2 pi bonds * Therefore, there are a total of 3 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 diagram). 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

Use the shortcut: 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: hybridization does NOT take into account the difference between the double-double bonds in CO₂ vs 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 the lowest electronegaivity. All of these qualities make it likely to create a hybrid in order have more active valence electrons and hence more bonds.