CHEMICAL BONDS How can atoms form a molecule? Let s watch the video: Bond types CHEMICAL BONDING Ionic bonding

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1 CHEMICAL BONDS How can atoms form a molecule? Let s watch the video: Bond types CHEMICAL BONDING In 1916, the American chemist Gilbert Newton 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 most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to complete their valence shells. We now know that there are two main 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 basis of the ionic bond. For example, during the reaction of sodium with chlorine: After the reaction takes place, the charged Na + and Cl - ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common: Ionic bonds form between metals and nonmetals. In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride). Ionic compounds dissolve easily in water and other polar solvents. In solution, ionic compounds easily conduct electricity. Ionic compounds tend to form crystalline solids with high melting temperatures.

2 If we consider a solid crystal of sodium chloride, the solid is made up of many positively charged sodium ions (pictured below as small grey spheres) and an equal number of negatively charged chlorine ions (green spheres). Sodium chlorine ions are arranged in an alternating fashion as demonstrated in the schematic. Each sodium ion is attracted equally to all of its neighbouring chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single molecule does not apply to ionic crystals because the solid exists as one continuous system. Ionic solids form crystals with high melting points because of the strong forces between neighbouring ions. 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 two (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 commonly occurs when two nonmetals bond 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 bond 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 up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. The atoms will share each other's single electron, forming one covalent bond. In this way, 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 exist as 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 shell. When two oxygen atoms form the compound O2, they share two pairs of electrons, forming two covalent bonds.

3 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 two periods of the periodic table are shown below. Lewis structures can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H2 and O2 are shown below. Nonpolar and polar covalent bonding - What is electronegativity? There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of the first type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed. Electronegativity is a measure of the attraction of an atom for electrons in a covalent bond. Fluorine, the most reactive non-metal, is assigned the highest value since it has the greatest attraction for the electrons being shared with the other element. Oxygen is also highly electronegative and has a strong attraction for electrons in a covalent bond. When two unlike atoms are covalently bonded, the shared electrons will be more strongly attracted to the atom of greater electronegativity. A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons (electronegativity) than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. The presence or absence of polar bonds within molecules plays a very important role in determining chemical and physical properties of those molecules. Some of these properties are melting points, boiling points, viscosity and solubility in solvents. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule. Water molecules contain two hydrogen atoms (pictured in red) bonded to one oxygen atom (blue). Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares one of its own electrons with each hydrogen, completing the H valence shells.

4 Because oxygen has a stronger pull on the bonding electrons, this leads to unequal sharing and to the formation of a polar covalent bond. The dipole Because the valence electrons in the water 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 partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it and it is called a dipole. Predicting Bond Types The difference in electronegativities of two elements can be used to predict the nature of the chemical bond. Bond type can be described as belonging to one of three classes: 1. nonpolar covalent 2. polar covalent 3. ionic When difference in electronegativity is or greater, the bond is usually ionic. Less than 1.7, the bond is usually covalent and it has some degree of polarity, unless the difference is less than 0.5. When the difference is less than 0.5 bonds are considered to be nonpolar. The most commonly used electronegativity scale is Pauling's. Notice that Pauling didn't assign a value in electronegativity to the Noble gases. Across a period: the electronegativities generally increase from left to right across a period with the Group VII elements having the highest value for the period. Down a group: the electronegativities generally decrease from top to bottom down a group.

5 Metallic bonding Over 3/4 of the elements of the periodic table are metals. Over the years, scientists have tried to come up with different models and theories to explain how metals bond, and how these bonds account for metal's specific properties. Metallic Character Metals have several qualities that are unique, such as the ability to conduct electricity, a lustrous (shiny) appearance, and they are malleable and ductile. Metals have a crystal structure. Metals that are malleable can be beaten into thin sheets, for example: aluminium foil. Metals that are ductile can be drawn into wires, for example: copper wire. The "Sea of Electrons" Theory In this model, the valence electrons are free, delocalized, mobile, and not associated with any particular atom. For example: metallic cations are shown in green surrounded by a "sea" of electrons, shown in purple. This model may account for metal properties, especially the conductivity: since the electrons are free, if electrons from an outside source were pushed into a metal wire at one end, the electrons would move through the wire and come out at the other end (conductivity is the movement of charge).