Paramagnetic Properties: O2-, NO, Or CN-?
Hey guys! Let's dive into the fascinating world of paramagnetic substances and figure out which of the following molecules—O2-, NO, or CN-—exhibits this cool property. Understanding paramagnetism requires a peek into the molecular orbital configurations of these molecules, so buckle up, and let's get started!
Understanding Paramagnetism
Paramagnetism arises in substances that have one or more unpaired electrons. These unpaired electrons have their own magnetic moment, and when an external magnetic field is applied, these moments align with the field, causing the substance to be attracted. The strength of this attraction is relatively weak, which distinguishes paramagnetism from ferromagnetism (like in iron), where the attraction is much stronger and can be permanent. In simple terms, if a molecule has unpaired electrons, it's like it has its own tiny magnets that can align with a bigger magnet nearby.
To determine whether a molecule is paramagnetic, we usually look at its molecular orbital diagram. This diagram shows how atomic orbitals combine to form molecular orbitals (both bonding and antibonding). We then fill these orbitals with electrons according to certain rules (like Hund's rule, which says we should fill orbitals individually before pairing electrons in the same orbital). If, after filling all the orbitals, there are still unpaired electrons, then the molecule is paramagnetic.
Think of it like this: imagine you're assigning seats on a bus. Each seat represents an orbital, and each person represents an electron. You want to make sure everyone gets a seat before pairing people up. If you end up with some single riders (unpaired electrons), the bus (molecule) is paramagnetic!
Detailed Analysis of Each Option
A) O2-
Let's start with O2-, the superoxide ion. Oxygen (O2) itself is a classic example of a paramagnetic molecule. It has 16 electrons in total. When it gains an extra electron to become O2-, it now has 17 electrons. To figure out if it's paramagnetic, we need to consider its molecular orbital configuration.
The molecular orbital configuration of O2 is (σ2s)2 (σ2s*)2 (σ2p)2 (π2p)4 (π2p*)2. This means there are two electrons in the sigma 2s bonding orbital, two in the sigma 2s antibonding orbital, two in the sigma 2p bonding orbital, four in the pi 2p bonding orbitals, and two in the pi 2p antibonding orbitals. These last two electrons are unpaired in the π2p* orbitals, making O2 paramagnetic.
Now, when O2 becomes O2-, it gains one more electron, which goes into one of the π2p* antibonding orbitals. So, the configuration becomes (σ2s)2 (σ2s*)2 (σ2p)2 (π2p)4 (π2p*)3. Since there are now three electrons in the two π2p* orbitals, Hund's rule dictates that two of these orbitals will have one electron each, and the third will be paired. Thus, O2- still has one unpaired electron, making it paramagnetic.
So, O2- is indeed paramagnetic because it retains an unpaired electron in its molecular orbital configuration. This is a key point to remember when dealing with oxygen-related species.
B) NO
Next up is NO, or nitric oxide. NO has 15 electrons: 7 from nitrogen and 8 from oxygen. Let's examine its molecular orbital configuration to see if it's paramagnetic.
The molecular orbital configuration of NO is (σ2s)2 (σ2s*)2 (σ2p)2 (π2p)4 (π2p*)1. Notice that there's only one electron in the π2p* antibonding orbitals. This single electron is unpaired, which makes NO paramagnetic.
NO is a particularly interesting molecule because of its role in various biological processes. Its paramagnetism contributes to its reactivity, allowing it to participate in radical reactions and act as a signaling molecule in the body. The presence of that one unpaired electron is crucial to its chemical behavior.
Therefore, NO is also paramagnetic due to the presence of an unpaired electron in its molecular orbital configuration. This single electron determines its magnetic properties and influences its chemical interactions.
C) Both A and B
Since both O2- and NO are paramagnetic, this option seems promising. O2- has one unpaired electron, and NO also has one unpaired electron. As we've established, the presence of unpaired electrons is the defining characteristic of paramagnetic substances.
Therefore, option C, which states that both A and B are paramagnetic, is correct. This highlights the importance of analyzing the molecular orbital configurations of molecules to determine their magnetic properties. Both O2- and NO fit the bill, making option C the accurate choice.
D) CN-
Finally, let's consider CN-, the cyanide ion. CN- has 14 electrons: 6 from carbon, 7 from nitrogen, and 1 from the negative charge. To determine if CN- is paramagnetic, we need to look at its molecular orbital configuration.
The molecular orbital configuration of CN- is (σ2s)2 (σ2s*)2 (σ2p)2 (π2p)4. All the electrons are paired in the bonding orbitals. There are no electrons in the antibonding orbitals. Since all the electrons are paired, CN- does not have any unpaired electrons and is therefore diamagnetic, not paramagnetic.
Diamagnetic substances are repelled by magnetic fields, unlike paramagnetic substances, which are attracted. The complete pairing of electrons in CN- results in the absence of a net magnetic moment, leading to its diamagnetic behavior.
So, CN- is not paramagnetic. It's actually diamagnetic, meaning it's repelled by a magnetic field. This is because all its electrons are paired, leaving no unpaired electrons to interact with an external magnetic field.
Conclusion
Alright, guys, after analyzing each option, we can confidently say that the correct answer is C) Both A and B. Both O2- and NO are paramagnetic due to the presence of unpaired electrons in their molecular orbital configurations. O2- retains an unpaired electron even after gaining an extra electron, and NO has a single unpaired electron in its π2p* antibonding orbitals.
Understanding the molecular orbital theory and how electrons are arranged in molecules is crucial for predicting their magnetic properties. Remember, if there are unpaired electrons, the molecule is paramagnetic. If all electrons are paired, the molecule is diamagnetic.
So, the next time you encounter a question about paramagnetism, remember to check for those unpaired electrons! Happy studying!