In unit six, we discussed ionic compounds, which are generally crystalline solids with high melting points. Other compounds, however, have very different properties. Water is a liquid at room temperature. Carbon dioxide is a gas at room temperature. The attractions that hold together the atoms in water and carbon dioxide can not be explained by ionic bonding. Ionic bonding = Covalent bonding = A is a neutral group of atoms joined together by covalent bonds. In nature, only noble gas elements exist as uncombined atoms. They are monoatomic; that is they consist of. Diatomic molecules = Molecular Compounds = Recall: Molecular Formulas ** A molecular formula does not tell you about a molecule's structure (the arrangement of the various atoms in space or which atoms are covalently bonded to one another).
Both ionic compounds and molecular compounds contain atoms of different elements that are combined chemically. However, the formulas describe different representative units. Ionic Compounds Molecular Compounds Properties of molecular compounds: Characteristics Ionic Compound Molecular Compound Representative Unit Bond Formation Type of Elements Physical State Melting Point Solubility in Water Electrical Conductivity of Aqueous Solution Stop and Review: 1. What information does a molecular formula provide? 2. How is the representative unit of a molecular compound different from the representative unit of an ionic compound?
3. What are the only elements that exist in nature as uncombined atoms? What term is used to describe these elements? 4. Describe how the molecule whose formula is NO is different from the molecule whose formula is N 2 O. 5. Give an example of a diatomic molecule. Remember, that when ionic compounds form, electrons tend to be transferred so that each ion acquires a noble gas configuration. A similar rule applies to covalent bonding. Single Covalent Bonds Two atoms held together by sharing one pair of electrons are joined by a single covalent bond. We can use electron dot structures to represent covalent bonds.
A structural formula represents the covalent bonds as dashes and shows the arrangement of covalently bonded atoms. Double and Triple Covalent Bonds Sometimes atoms bond by sharing more than one pair of electrons. Strength in Covalent Bonds Bond dissociation energy = Stop and Review 1. What electron configuration do atoms usually achieve by sharing electrons to form covalent bonds? 2. How is the strength of a covalent bond related to its bond dissociation energy?
3. How is an electron dot structure used to represent a covalent bond? 4. When are two atoms likely to form a double bond between them? A triple bond? 5. What kinds of information does a structural formula reveal about the compound it represents? 6. Draw electron dot structures for the following molecules. H 2 S PH 3 ClF VSEPR THEORY The electron dot structure and structural formula of methane (CH 4 ) show the molecule as if it were flat and merely two dimensional. In reality, methane molecules are three dimensional! In order to explain the three dimensional shape of molecules scientists use valence shell electron pair repulsion theory (VESPR theory). VESPR THEORY
Name # of Bonds # or Unshared Pairs Shape Stop and Review 1. What do scientists use VSPER theory for? 2. What are the different three dimensional shapes that molecules can take according to the VESPR theory? Describe each one.
POLAR BONDS AND MOLECULES Covalent bonds involve between atoms. However, covalent bonds differ in terms of how the bonded atoms share the electrons. The bonding pair of electrons in covalent bonds are pulled between the nuclei of the atoms sharing the electrons When the atoms in the bond pull equally, the bonding electrons are shared equally, and each bond formed is a. Polar Covalent Bond (Polar Bond) **The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. Electronegativity = Electronegativity difference =
Polar Molecule = Dipole = Practice:
Identify the bonds between atoms of each pair of elements as nonpolar covalent, polar covalent, or ionic. H and Br K and Cl C and O Cl and F Li and O Br and Br ATTRACTIONS BETWEEN MOLECULES Molecules can be attracted to each other by a variety of different forces. Intermolecular attractions are weaker than either ionic or covalent bonds. Nevertheless, you should not underestimate the importance of these forces. These attractions are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature. Dipole Interactions Dispersion Forces Hydrogen Bonding