Question

Rank the bonds a-f below according to increasing bond length.

Short answer

Question: Rank the following bonds in increasing order of bond length: a) C=C, b) C-N, c) O=O, d) O-F, e) C≡C, f) S-Cl. Answer: e < a < c < b < d < f

Step-by-step solution

Identify the bond order of each bond

First, we need to determine the bond order of each bond. Bond order can be calculated as half of the difference between the number of bonding electrons and antibonding electrons. Bond a: Double bond Bond b: Single bond Bond c: Double bond Bond d: Single bond Bond e: Triple bond Bond f: Single bond

Identify the size of the atoms involved in each bond

Next, we need to determine the size of the atoms in each bond. In general, atomic size increases from top to bottom within a group and decreases from left to right across a period on the periodic table.

Use bond order and atomic size to rank the bonds

Now, we can use the bond order and atomic size information to rank the bonds in increasing order of bond length. Here, bonds with lower bond orders (single bonds) will generally be longer than bonds with higher bond orders (double and triple bonds). Since bond e is a triple bond, it should be the shortest. Then, among the double bonds (a and c), the bond with smaller atoms will be shorter. Among the single bonds (b, d, and f), the bond with smaller atoms will also be shorter. With this in mind, we can now rank the bonds in increasing bond length: Bond e < Bond a < Bond c < Bond b < Bond d < Bond f

Key concepts

Bond Order

Understanding bond order is crucial when it comes to determining the strength and length of a chemical bond. Bond order is defined as the net number of chemical bonds between a pair of atoms. For instance, a single bond has a bond order of one, a double bond has a bond order of two, and a triple bond has a bond order of three.

Higher bond orders indicate stronger and shorter bonds, as more electrons are shared between the atoms, pulling them closer together. Conversely, a lower bond order means a bond is longer and weaker as fewer electrons are shared. In the given exercise, the bond order helps us rank the bonds from the shortest to the longest by recognizing that triple bonds are the shortest followed by double bonds, with single bonds being the longest.

Atomic Size

Atomic size, also known as atomic radius, refers to the size of an atom and is usually measured from the nucleus to the outer shell of electrons. An important trend to remember from the periodic table is that atomic size increases as we move down a group and decreases as we move across a period from left to right.

The size of the atoms involved in a bond affects the length of that bond. Larger atoms result in longer bond lengths because their outermost electrons are farther from the nucleus, which means that when they form bonds, the distance between the two nuclei will naturally be greater. Conversely, smaller atoms will form shorter bonds. This concept is utilized in our exercise to further differentiate between bonds of the same type but with different sized atoms.

Periodic Table Trends

The periodic table is an organized chart of elements that reveals important trends in the properties of these elements, including atomic size and electronegativity. For bond length comparisons, the trend in atomic size is particularly significant. As aforementioned, atomic size increases down a group and decreases from left to right across a period. This trend arises due to the number of electron shells increasing as you move down a group and the strength of the positive charge on the nucleus increasing as you move across a period.

These trends allow us to predict the relative sizes of atoms, which is essential when comparing bond lengths between different types of atoms, as seen in the exercise.

Single Bond

A single bond is a covalent bond formed by the sharing of a pair of electrons between two atoms. It is represented by a single line between the atoms (e.g., H-H) and has a bond order of one. Single bonds are the longest and usually the weakest type of covalent bonds because there is less electron density attracting the two atoms to each other compared to double and triple bonds. Understanding single bonds is fundamental for our exercise since several of the bonds being compared are single bonds, which generally contribute to longer bond lengths than multiple bonds.

Double Bond

A double bond involves two pairs of shared electrons between atoms and has a bond order of two, indicated by two lines (e.g., O=O). It is stronger and shorter than a single bond due to the higher number of shared electrons, leading to increased attraction between the atoms. The presence of a double bond can affect the geometry and the length of the bond in comparison to a single bond. In the exercise at hand, comparing double bonds not only with single but also with triple bonds helps us to place them in the correct order when ranking by increasing bond length.

Triple Bond

Triple bonds share three pairs of electrons and have the highest bond order of three, which is denoted by three lines (e.g., N≡N). They are the shortest and strongest type of covalent bonds due to the significant attraction from the shared electron density. In molecules with triple bonds, the atoms are held very close together, resulting in the shortest bond lengths among single, double, and triple bonds. This characteristic was crucial for determining that bond e, which is a triple bond, should have the shortest bond length in our exercise.