Chemical Bonding Review

Compounds differ in composition as well as chemical and physical properties. Two major categories of compounds are ionic compounds and molecular (covalent) compounds. Ionic compounds: metal + non-metal or polyatomic ion Covalent compounds: all non-metals

Chemical bonds are formed when valence electrons are: Transferred from one atom to another (ionic) A metal from the left and a non-metal from the right. Shared between atoms (covalent) Can be polar (unequal sharing), non-polar (equal sharing) Mobile within a metal (metallic) The phrase “mobile valence electrons” is often used.

In a multiple covalent bond, more than one pair of electrons is shared between two atoms. Unsaturated organic compounds contain at least one double or triple bond. Single, double, or triple bonds – each bond is a shared pair of electrons Double and triple bonds are usually represented by two or three lines between atoms. For dot diagrams, two or three pairs of dots are between the elements.

Molecular polarity can be determined by the shape and the distribution of charge. Examples of symmetrical (nonpolar) molecules include CO2, CH4, and the diatomic elements. Examples of asymmetrical (polar) molecules include HCl, NH3­­ and H2O. These examples are used year after year. Symmetrical molecules with polar bonds (the C=O double bond is very polar) are non-polar. You should be able to draw or recognize Lewis structures for all of the molecules above.

When an atom gains one or more electrons, it becomes a negative ion and its radius increases. When an atom loses one or more electrons, it becomes a positive ion and its radius decreases.

When a bond is broken, energy is absorbed When a bond is broken, energy is absorbed. When a bond is formed, energy is released. Absorbing energy – endothermic ΔH is (+) Releasing energy – exothermic ΔH is (-)

Atoms tend to bond so that a stable valence electron configuration, like that of a noble gas, is achieved. The octet rule governs much of chemistry. Sometimes they will ask about isoelectronic atoms and ions; they have the same number of electrons; for example, O2- = Ne = Mg2+ = 10 electrons (eight in the second energy level)

Electron dot diagrams (lewis structures) can represent the valence electron arrangement in elements and compounds. You must be able to distinguish ionic and covalent compounds to write Lewis structures successfully! Ionic compounds have brackets and charges. Covalent compounds have shared electron pairs. NaCl right: [Na]+ [: Cl : ]1- wrong: Na:Cl : HCl right: H:Cl: wrong: [H]+ [: Cl : ]1-

Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond. Electronegativity values are assigned according to arbitrary scales. The definition of electronegativity is a common test question. Sometimes they do not use the word “electronegativity” in the question at all – you have to recognize that they are talking about it!. Electronegativity values for the elements are found in Table S.

The electronegativity difference between two bonded atoms is used to assess the degree of polarity in the bond. Little or no difference: non-polar bond; large difference: polar bond.