Сборник текстов на казахском, русском, английском языках для формирования навыков по видам речевой деятельности обучающихся уровней среднего образования

Hydrogen Bonds

Because of this polarity, the oxygen end of the molecule would repel negative atoms like itself, while attracting positive atoms, like hydrogen. Hydrogen, which becomes slightly positive, would repel positive atoms (like other hydrogen atoms) and attract negative atoms (such as oxygen atoms). This positive and negative attraction system helps water molecules stick together, which is what makes the boiling point of water high (as it takes more energy to break these bonds between water molecules).

In addition to the four types of chemical bonds, there are also three categories bonds fit into: single, double, and triple. Single bonds involve one pair of shared electrons between two atoms. Double bonds involve two pairs of shared electrons between two atoms, and triple bonds involve three pairs of shared electrons between two atoms. These bonds take on different natures due to the differing amounts of electrons needed and able to be given up.

Now, let’s look at determining what types of bonds we see in different compounds. We’ve already looked at the bonds in H₂O, which we determined to be hydrogen bonds. However, now let’s look at a few other types of bonds as examples.

Compound: HNO₃ (also known as Nitric acid)

There are two different determinations we can make as to what these bonds look like; first we can decide whether the bonds are covalent, polar covalent, ionic, or hydrogen. Then, we can determine if the bonds are single, double, or triple.

In order to decide whether the bonds are covalent, polar covalent, ionic or hydrogen, we need to look at the types of elements seen and the electronegativity values. We look at the elements and see hydrogen, nitrogen, and oxygen—no metals. This rules out ionic bonding as a type of bond seen in the compound. Then, we would look at electronegativity values for nitrogen and oxygen. Oftentimes, this information can be found on a periodic table, in a book index, or an educational online resource. The electronegativity value for oxygen is 3.5 and the electronegativity value for nitrogen is 3.0. The way to determine the bond type is by taking the difference between the two numbers (subtraction). 3.5 – 3.0 = 0.5, so we can determine that the bond between nitrogen and oxygen is a covalent bond. We can also determine, from past knowledge, that the bond between oxygen and hydrogen is a hydrogen bond as it was in water.

Now, we need to count the electrons and draw the diagram for HNO₃. For more help counting electrons, please see the page on Electron Configuration. For more help drawing the Lewis structures, please see the page on Lewis Structures. This process combines both of these in order to determine the structure and shape of a molecule of the compound.

Now, we can count how many electrons we have used by counting 2 electrons for each bond placed. We see that we have placed 4 bonds, so we have used 8 electrons. 24 – 8 = 16 electrons that we need to distribute. In order to correctly place the rest of the electrons, we need to determine how many electrons each atom needs to be stable.

The central atom, N, has three bonds attached (equivalent of 6 electrons) so it needs 2 more electrons to be stable. The O to the right has one bond (two electrons) so it needs 6 more to be stable. The O above the N has one bond (two electrons) so it also needs 6 electrons to be stable. The O to the left of the N is bonded both to N and to H, so it has two bonds (4 electrons); therefore, it needs 4 more electrons to be stable. We add up the total amount of electrons needed, 2 + 6 + 6 + 4 = 18, and see that we need 18 electrons to stabilize the compound. We know this is not possible, since we only have 16 available electrons. When this happens, we need to insert a double bond in order to resolve the problem of lack of electrons. This is because, although we count each bond as 2 electrons, the elements joined together in the bond are actually sharing the electrons. Therefore, when we count out the bonds, we are counting some electrons twice because they are shared. This is normal and expected, and resolves not having enough valence electrons. Now, we need to decide where to put the double bond in this compound. We know that the double bond cannot go between O and H, because H does not have enough room to accept another electron. Therefore, we know we must place the bond between N and O. You might be thinking, how do I decide where to put the bond? In this particular example, we can place the bond either between the top O and N, or the right O and N. This is because HNO₃ displays resonance.

Here are the ways you can place the double bond:

We are going to keep the bond between N and the right O in our example. After we add in the bond, we subtract two more electrons from our available electrons (16) and are left with 14 electrons to distribute. Now we need to make sure we have the correct number of electrons. After placing in the double bond, N is now stable because it has 4 bonds (8 electrons) surrounding it. It does not need any additional electrons. The top O (above N) needs 6 electrons, the right O now only needs 4 electrons (because it has a double bond now, which is 4 electrons), and the left O still needs 4 electrons to become stable. We add these numbers together, 6 + 4 + 4 = 14, and we see that 14 is the number of electrons we have, so we can go ahead and distribute them, like this:

Now, our compound is stable with appropriately distributed valence electrons. We can see that there are three single bonds (H—O, N—O, and N—O) and one double bond (N==O). [5]

The early innovation of Periodic law by a Russian chemist, Dmitry I. Mendeleev in the mid-19th century, has been of great value in the growth of chemistry. In chemistry, Periodic law is the arrangement of theelements in order of increasing atomic number, that is, the total number of protons in the atomic nucleus and their physical and chemical properties recur in a pattern with the increasing atomic number. In the periodic table, the horizontal rows, known as the “periods” exemplify these relations. In fact, until the second decade of the 20th century, it was not documented that the order ofelements in the periodic system is that of their atomic numbers.