What Type of Bond Results When Electrons Are Completely Transferred From One Atom to Another?
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Though the periodic table has but 118 or then elements, there are apparently more substances in nature than 118 pure elements. This is because atoms tin can react with one another to course new substances called compounds (meet our Chemic Reactions module). Formed when two or more atoms chemically bond together, the resulting chemical compound is unique both chemically and physically from its parent atoms.
Let's look at an example. The chemical element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet. The element chlorine is a light-green-colored gas that is then poisonous that information technology was used every bit a weapon in Globe War I. When chemically bonded together, these two dangerous substances form the chemical compound sodium chloride, a chemical compound so prophylactic that we eat it every day - mutual table table salt!
In 1916, the American chemist Gilbert N. Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are well-nigh stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than 8 valence electrons bond together to share electrons and consummate their valence shells.
While some of Lewis' predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the footing of what is known today most chemical bonding. We now know that there are two principal types of chemical bonding; ionic bonding and covalent bonding.
Ionic bonding
In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the footing of the ionic bond.
For example, during the reaction of sodium with chlorine:
Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an boosted valence electron. This is typical of the relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. Afterward the reaction takes identify, the charged Na+ and Cl- ions are held together past electrostatic forces, thus forming an ionic bail. Ionic compounds share many features in common:
- Ionic bonds form between metals and nonmetals.
- In naming uncomplicated ionic compounds, the metal is always kickoff, the nonmetal second (eastward.g., sodium chloride).
- Ionic compounds dissolve easily in water and other polar solvents.
- In solution, ionic compounds hands conduct electricity.
- Ionic compounds tend to class crystalline solids with loftier melting temperatures.
This last feature, the fact that ionic compounds are solids, results from the intermolecular forces (forces betwixt molecules) in ionic solids. If we consider a solid crystal of sodium chloride, the solid is made upward of many positively charged sodium ions (pictured below as small gray spheres) and an equal number of negatively charged chlorine ions (light-green spheres). Due to the interaction of the charged ions, the sodium and chlorine ions are bundled in an alternate fashion every bit demonstrated in the schematic. Each sodium ion is attracted every bit to all of its neighboring chlorine ions, and too for the chlorine to sodium attraction. The concept of a single molecule does non apply to ionic crystals because the solid exists as one continuous arrangement. Ionic solids course crystals with high melting points considering of the strong forces betwixt neighboring ions.
| Sodium Chloride Crystal | NaCl Crystal Schematic | ||||
 | |||||
| Cl-one | Na+1 | Cl-1 | Na+i | Cl-1 | |
| Na+ane | Cl-one | Na+1 | Cl-ane | Na+1 | |
| Cl-1 | Na+1 | Cl-1 | Na+1 | Cl-1 | |
| Na+1 | Cl-1 | Na+1 | Cl-1 | Na+1 | |
Covalent bonding
The second major type of atomic bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when 2 (or more) elements share electrons. Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). This most normally occurs when two nonmetals bail together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bail is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick upward a 2nd electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. Considering the hydrogen compound is a combination of equally matched atoms, the atoms volition share each other'due south single electron, forming ane covalent bail. In this mode, both atoms share the stability of a full valence shell.
Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed. Thus covalent molecules are not strongly attracted to one another. As a result, covalent molecules move about freely and tend to be equally liquids or gases at room temperature.
Multiple Bonds: For every pair of electrons shared between two atoms, a single covalent bond is formed. Some atoms can share multiple pairs of electrons, forming multiple covalent bonds. For example, oxygen (which has six valence electrons) needs two electrons to complete its valence crush. When two oxygen atoms form the compound O2, they share 2 pairs of electrons, forming ii covalent bonds.
Lewis Dot Structures: Lewis dot structures are a shorthand to represent the valence electrons of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons. The Lewis structures for the elements in the first 2 periods of the periodic table are shown below.
Lewis structures tin can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and tin can be represented by a pair of dots or a dash (each nuance represents one pair of electrons, or one bail). Lewis structures for H2 and O2 are shown below.
Polar and nonpolar covalent bonding
There are, in fact, two subtypes of covalent bonds. The Hii molecule is a good example of the starting time type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule accept an equal allure (or affinity) for electrons, the bonding electrons are equally shared by the ii atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed.
A polar bond is formed when electrons are unequally shared between 2 atoms. Polar covalent bonding occurs because one cantlet has a stronger affinity for electrons than the other (yet not enough to pull the electrons abroad completely and course an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the cantlet that has the stronger affinity for electrons. A proficient example of a polar covalent bail is the hydrogen-oxygen bond in the water molecule.
Water molecules incorporate two hydrogen atoms (pictured in blue) bonded to 1 oxygen atom (red). Oxygen, with six valence electrons, needs ii boosted electrons to consummate its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from ii hydrogen atoms to complete its own valence shell, and in render shares two of its ain electrons with each hydrogen, completing the H valence shells.
The principal difference betwixt the H-O bail in water and the H-H bail is the degree of electron sharing. The large oxygen atom has a stronger analogousness for electrons than the minor hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond.
The dipole
Because the valence electrons in the h2o molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a fractional positive accuse. Ions are not formed; however, the molecule develops a fractional electrical charge across it called a dipole. The water dipole is represented by the arrow in the pop-upwards animation (to a higher place) in which the caput of the arrow points toward the electron dense (negative) end of the dipole and the cross resides nigh the electron poor (positive) cease of the molecule.
Summary
Chemical bonding betwixt atoms results in compounds that tin be very different from the parent atoms. This module, the 2nd in a series on chemical reactions, describes how atoms proceeds, lose, or share electrons to course ionic or covalent bonds. The module lists features of ionic and covalent compounds. Lewis dot structures and dipoles are introduced.
Source: https://www.visionlearning.com/en/library/Chemistry/1/Chemical-Bonding-previous-version/246