Molecular Structure and Shape Covalent Compounds Molecular Structure and Shape

Lewis Structures Covalent compounds share electrons. Goal is to achieve “Noble Gas” configuration. Hydrogen wants to be like Helium (2 electrons) All other nonmetals want to have 8 electrons. This is known as the octet rule. Atoms achieve an octet by sharing 1, 2, or 3 electrons with another atom. Single bond – each atom contributes ONE electron for a pair Double bond – each atom contributes TWO electrons for two pairs Triple bond – each atom contributes THREE electrons for three pairs

Lewis Structures Atoms can also share electrons with more than one atom. Ex. An oxygen atom shares 2 electrons with 2 hydrogen atoms in the formation of water. The oxygen shares 1 electron with each hydrogen atom, forming 2 single bonds. Lewis structures show the electrons that can be shared and therefore can be used to figure out how atoms combine in covalent compounds.

Lewis Structures Remember, a Lewis Structure (or electron dot diagram) is drawn by writing the symbol for an element and then surrounding it with dots that represent the valence electrons. Example: O H

Lewis Structures Electrons are then SHARED between atoms so they can achieve their Noble Gas configuration. Example: H O H

H O H H O H Structural Formulas Lewis Structures are converted to structural formulas by drawing a line in place of a SHARED pair of electrons. H O H H O H

Draw the Lewis Structure for the following elements. Hydrogen Carbon Oxygen Nitrogen Sulfur Chlorine Fluorine Boron Selenium

Identify the number of electrons needed to complete the octet Identify the number of electrons needed to complete the octet. -This is how many electrons will be shared. Hydrogen Carbon Oxygen Nitrogen Sulfur Chlorine Fluorine Boron Selenium

Using the element Lewis Structure, draw the Lewis Structure for the following compounds. Hydrogen, H2 Chlorine, Cl2 Oxygen, O2 Nitrogen, N2 Methane, CH4 Ammonia, NH3 Hydrofluoric acid, HF Carbon dioxide, CO2 Boron trichloride, BCl3 Ethane, C2H6

Molecular Shapes Shapes are determined by the number of shared and unshared pairs of electrons that surround the central atom. Molecules arrange themselves so that their electron pairs are as far apart as possible. This is the VSEPR theory – Valence Shell Electron Pair Repulsion theory http://www.chemmybear.com/shapes.html

Shapes you need to know Linear Tetrahedral, Trigonal Pyramid, Bent Steric Number 2 or higher Tetrahedral, Trigonal Pyramid, Bent Steric Number of 4 Trigonal Planar Steric Number of 3 Octahedral Steric Number of 6

Sigma and Pi Bonds Sigma Bonds are the first bond that forms between two atoms. Sigma bonds are the only bonds that are used to determine the steric number and molecular shape. Sigma bonds form in the space directly between two atomic nuclei. Pi Bonds are the second or third bond that forms between the same two atoms. Pi bonds are NOT used to determine molecular shape. However, they do shorten the length of the bond – pulling the nuclei closer together. Pi bonds form around the sigma bond – either top and bottom or front and back.

Sigma and Pi Bonds

Sigma and Pi Bonds Single bond Double bond Triple bond 1 sigma bond Ethane, C2H6 Double bond 1 pi bond Ethene, C2H4 Triple bond 2 pi bonds Ethyne, C2H2

Polarity Bonds and Molecules are either polar or nonpolar. Polar Bonds form when there is unequal sharing of the electrons. This is due to a difference in electronegativity. The more electronegative nucleus gets “a bigger share”. Polar Molecules form when the polar bonds and unshared pairs surrounding the central atom are asymmetrical in their 3D arrangement. Certain shapes will usually be polar. Trigonal pyramid, bent All shapes can be polar.

Intermolecular Forces Forces BETWEEN molecules Hydrogen Bonds occur when a hydrogen is covalently bonded to a very electronegative atom (F, Cl, O) Strongest of intermolecular forces Dipole Interactions occur with polar molecules. The positive end of one molecule is attracted to the negative end of another molecule. Dispersion forces occur as a result of the motion of electrons within atoms. Weakest of intermolecular forces

Review of Compounds Ionic, Covalent, Metallic

Ions and Ionic Compounds Remember an ion is an atom that has lost or gained electrons Cations – positive – lost electrons Anions – negative – gained electrons Ionic Compounds form when 2 or more atoms are joined by the loss and gain of electrons ALWAYS cation + anion Cation is always first and anion is always second.

Ionic Bonds Cation + Anion Ions are joined by the transfer of electrons Creates Electrostatic forces (attraction of opposite charges) that hold the ions together Ionic Compounds are composed of a continuous arrangement of oppositely charged ions. NOT a single separate unit

Ionic Bonds Bonds between atoms will be ionic when there is a LARGE difference in electronegativity between the atoms. > 1.7 ∆EN Metal + Nonmetal Opposite sides of the periodic table – large differences in electronegativity

Properties of Ionic Compounds Metal + Nonmetal (usually) Crystalline SOLIDS at room temp. (most) Crystal – repeating geometric pattern Brittle HIGH melting points HIGH boiling points Some are so high it takes extreme conditions to get them to change to gas Conduct electrical currents when melted or dissolved in water

Molecular Compounds Compounds that form when atoms SHARE electrons Forms a covalent bond NEVER contain ions NEVER have charges Contain only nonmetals

Properties of Covalent Compounds Nonmetal + Nonmetal Can be solids, liquids, or gases at room temperature State is determined by the bond strength (compounds with stronger bonds tend to be solids) LOW melting points LOW boiling points POOR conductors of electricity under any conditions Can be Polar or Nonpolar Depends on how the electrons are shared Polar compounds are better conductors

Metallic Compounds Metal + Metal Consists of positive metal ions with a “sea of electrons” Electrons are free floating – are not attached to any one atom or ion HIGH melting points HIGH boiling points HIGH conductivity Malleable, Ductile, High luster (shine) All properties are a result of the “free” electrons.