Sodium chloride contains which type of bond

Ionic bonds usually occur between metal and nonmetal ions. For example, sodium Na , a metal, and chloride Cl , a nonmetal, form an ionic bond to make NaCl. In a covalent bond , the atoms bond by sharing electrons.

Covalent bonds usually occur between nonmetals. Is calcium chloride covalent or ionic? Calcium chloride. Calcium chloride is an ionic compound of calcium and chlorine. It is highly soluble in water and it is deliquescent.

Is MgO a covalent bond? Magnesium oxide, or MgO, is a compound that is solid at room temperature. Often used as a mineral supplement, the bonds that hold the compound together are either ionic or covalent. What's the most common type of chemical bond? Covalent bonding is a common type of bonding in which two or more atoms share valence electrons more or less equally.

The simplest and most common type is a single bond in which two atoms share two electrons. How are ionic and covalent bonds different from hydrogen bonds? In summary, hydrogen bonds are relatively weak intermolecular forces, while covalent and ionic bonds are relatively strong intramolecular forces.

Is rubbing alcohol ionic or covalent? Rubbing Alcohol is a Molecular compound. Rubbing alcohol is isopropanol, with the chemical formula C3H8O. Molecular compounds are between non-metals only. Rubbing alcohol contains carbon, hydrogen, and oxygen, which are all non-metals. Why is sodium chloride ionic rather than covalent? Hydrogen can be considered to be in Group 1 or Group 17 because it has properties similar to both groups.

Hydrogen can participate in both ionic and covalent bonding. When participating in covalent bonding, hydrogen only needs two electrons to have a full valence shell. As it has only one electron to start with, it can only make one bond. Hydrogen is shown in Fig 2. In the formation of a covalent hydrogen molecule, therefore, each hydrogen atom forms a single bond, producing a molecule with the formula H 2.

A single bond is defined as one covalent bond, or two shared electrons, between two atoms. A molecule can have multiple single bonds. For example, water, H 2 O, has two single bonds, one between each hydrogen atom and the oxygen atom Fig. Figure 2. Sometimes two covalent bonds are formed between two atoms by each atom sharing two electrons, for a total of four shared electrons. For example, in the formation of the oxygen molecule, each atom of oxygen forms two bonds to the other oxygen atom, producing the molecule O 2.

Similarly, in carbon dioxide CO 2 , two double bonds are formed between the carbon and each of the two oxygen atoms Fig. In some cases, three covalent bonds can be formed between two atoms. The most common gas in the atmosphere, nitrogen, is made of two nitrogen atoms bonded by a triple bond.

Each nitrogen atom is able to share three electrons for a total of six shared electrons in the N 2 molecule Fig. In addition to elemental ions, there are polyatomic ions. Polyatomic ions are ions that are made up of two or more atoms held together by covalent bonds. Polyatomic ions can join with other polyatomic ions or elemental ions to form ionic compounds. It is not easy to predict the name or charge of a polyatomic ion by looking at the formula.

Polyatomic ions found in seawater are given in Table 2. Polyatomic ions bond with other ions in the same way that elemental ions bond, with electrostatic forces caused by oppositely charged ions holding the ions together in an ionic compound bond. Charges must still be balanced. For example, in Fig. In Figure 2. P olyatomic ions can bond with monatomic ions or with other polyatomic ions to form compounds.

In order to form neutral compounds, the total charges must be balanced. A molecule or compound is made when two or more atoms form a chemical bond that links them together.

As we have seen, there are two types of bonds: ionic bonds and covalent bonds. In an ionic bond, the atoms are bound together by the electrostatic forces in the attraction between ions of opposite charge. Ionic bonds usually occur between metal and nonmetal ions. For example, sodium Na , a metal, and chloride Cl , a nonmetal, form an ionic bond to make NaCl. In a covalent bond, the atoms bond by sharing electrons. Covalent bonds usually occur between nonmetals. For example, in water H 2 O each hydrogen H and oxygen O share a pair of electrons to make a molecule of two hydrogen atoms single bonded to a single oxygen atom.

In general, ionic bonds occur between elements that are far apart on the periodic table. Covalent bonds occur between elements that are close together on the periodic table. Similarly, covalent bonding between silicon and oxygen atoms makes strong bonds that form a large group of minerals called silicates more on those later.

Metallic bonds form when the outer shell electrons are shared between neighboring atoms. Unlike covalent bonding however, there are insufficient numbers of electrons in most metal atoms such as copper or silver to form pure covalent bonds. Therefore, the electrons are shared amongst all the nearest neighbor metal ions, forming a metallic bond. This strange arrangement of "metallic ions is a sea of electrons" gives metals their particular physical properties.

Metallic bonds are also explained by band theory. Band theory states that closely packed atoms have overlapping electron energy levels resulting in a conduction "band" wherein the electrons are free to roam between atoms, thus bonding them together. For more information on metallic bonds and band theory, see this web site. Van der Waals bonds are weak bonds that form due to the attraction of the positive nuclei and negative electron clouds of closely packed atoms.

This attraction is opposed by the repulsive force of the electron clouds and the repulsive force of neighboring nuclei. However, the attraction is stronger than the total repulsive forces, leaving a residual, weak attraction.

Van der Waals bonding is important in minerals such as graphite and clay minerals. Test your knowledge with a quiz! Return to main menu. Return to Introductory Geosciences Course Page.