Bonding

= Topic 6 - Bonding =
 * ** Why do elements bond? **
 * Every element except for the noble gases (because they already have a happy full outer level) want to have a full outer level. Therefore they bond with each other to have a full outer level and be like the noble gases.


 * ** Bonding Theories **
 * **Sigma Bonds**
 * two atoms are held together by one electron pair.
 * two atomic orbitals such that they are symmetrical about the axis they form between the two nuclei.[[image:http://upload.wikimedia.org/wikipedia/commons/thumb/0/06/Sigma_bond.svg/743px-Sigma_bond.svg.png width="292" height="130" caption="Sigma Bond"]]
 * **Pi Bonds**
 * Two lobes of one involved electron orbital overlap two lobes of the other involved electron orbital
 * When orbitals do not overlap end on end
 * This bond is weaker than sigma bonds
 * [[image:http://library.thinkquest.org/10429/media/geometry/pi.gif width="278" height="139"]]
 * **Single Bond: Sigma Bond**
 * **Double Bond: 1 Sigma Bond, 1 Pi Bond**
 * **Triple Bond: 1 Sigma bond, 2 Pi Bond**
 * **VSE****PR**
 * Valence-Shell Electron-Pair Repulsion theory describes repulsion between electron pair which causes molecular shapes to adjust so that the valence electron pairs stay as for apart as possible.
 * Draw the Lewis structure of the molecule.
 * Count the total number of bonding and non-bonding electron pairs in the valence shell of the central atom.
 * Arrange the electron pairs around the central atom so that they are as far apart from each other as possible. It is important not to forget to take into consideration non-bonding pairs.
 * [[image:vinstan:image37.gif width="329" height="318" caption="VSEPR"]]
 * **Ionic Bond**
 * An ionic bond is __a type of chemical bond that involves a metal and a nonmetal ion__ (or polyatomic ions such as ammonium) through electrostatic attraction. The metal donates one or more electrons, forming a positively charged ion or cation with a stable electron configuration. These electrons then enter the non metal, causing it to form a negatively charged ion or anion which also has a stable electron configuration. The electrostatic attraction between the oppositely charged ions causes them to come together and form a bond.


 * **Covalent Bond**
 * When two or more atoms with similar electronegativities interact, they achieve a noble gas electron configuration by sharing electrons in what is known as a **covalent bond**.


 * || Ionic || Covalent  ||
 * Representative Unit || Formula unit – loves whole # ratio of atoms  || Molecule  ||
 * Bond Formation || Transfer of electron between atoms  || Electron shared  ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Type of atoms || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Metals & Nonmetals  || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Nonmetals  ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"> <span style="font-family: 'Calibri','sans-serif'; font-size: 12pt;">Physical <span style="font-family: 'Calibri','sans-serif';">State  || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Solid  || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Solid, liquid, gas  ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Melting Point || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Very high (> 300 °C)  || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Lower than 300 °C  ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Solubility || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Usually high  || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Variable(low -> high)  ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Conductivity in H2O || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Good conductors  || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;"><span style="display: block; font-family: Calibri,sans-serif; font-size: 12pt; text-align: center;">Poor to conduct, except for strong acids  ||

> **
 * ** Octet Rule **
 * <span style="margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0.5em; padding-bottom: 0px; padding-left: 3em; padding-right: 0px; padding-top: 0px;">An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons.
 * <span style="margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0.5em; padding-bottom: 0px; padding-left: 3em; padding-right: 0px; padding-top: 0px;">[[image:http://www.uoregon.edu/~ch111/images/o2dot3.gif height="155"]]
 * <span style="color: #000000; font-family: arial,helvetica,sans-serif; font-size: 13px; line-height: 19px; margin: 0px; padding: 0px;">Exceptions to the Octet Rule
 * <span style="font-family: Arial,Helvetica,sans-serif; font-weight: normal;">There are three types of ions or mole molecules that do not follow the Octet Rule:
 * <span style="font-family: Arial,Helvetica,sans-serif; font-weight: normal;">Ions or molecules with an odd number of electrons
 * <span style="font-family: Arial,Helvetica,sans-serif; font-weight: normal;">Ions or molecules with less than an octet.
 * <span style="font-family: Arial,Helvetica,sans-serif; font-weight: normal;">Ions or molecules with more than eight valence electrons (an expanded octet).

> <span style="display: block; margin-bottom: 0pt; margin-left: 0.25in; margin-right: 0in; margin-top: 0in; text-indent: -0.25in;"> A coordinate covalent bond (dative bond) forms when 2 atoms share a pair of electrons but one of the atoms provides both electrons for the bonding pair || Group II (alkali earth) Metals, eg, Mg Transition Metals, eg, Fe ||< Group I Metal + Non-metal, eg, NaCl Group II Metal + a non-metal, eg, MgCl2 ||< Molecular Substances: Group VII elements, eg, Cl2, Group VI non-metals, eg, O2 hydrogen + non-metal, eg, H2O Coordinate covalent bond examples: NH4+, H3O+ Three dimensional covalent networks: Si, C (graphite and diamond), B, SiO2(quartz) || Three dimensional covalent networks: high || Three dimensional covalent network solids are insoluble. || = REVIEW QUESTIONS = = should be in terms of principles of molecular structure and intermolecular forces. =
 * ** Lewis Electron-dot Diagram **
 * A representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms.
 * Writing Lewis Structures
 * 1) Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
 * If it is an anion, add one electron for each negative charge.
 * If it is a cation, subtract one electron for each positive charge.
 * 1) The central atom is the least electronegative element that isn't hydrogen.
 * 2) Connect the outer atoms to it by single bonds.
 * 3) Fill the octets of the outer atoms
 * 4) Fill the octet of the central atom.
 * 5) If you run out of electrons before the central atom has an octet, form multiple bonds until it does.
 * Only valence electrons are shown in a Lewis structure.
 * [[image:http://www.roymech.co.uk/images14/lewis_elements.gif width="496" height="217" caption="Lewis Structure"]]
 * ** Polarization **
 * Polarization is the separation of charges in a chemical particle (Eg. molecules, ions).
 * Big particles with relatively small charge are easily polarizable, eg. halide ions.
 * Small particles with greater charge polarize others easily, eg. Al3+.
 * ** Electronegativity **
 * The ability of an atom to attract toward itself the electrons in a chemical bond.
 * Elements with high electronegativity have a greater tendency to attract electrons than do elements with low electronegativity.
 * The electronegativity is related to electron affinity and ionization energy.
 * Within each group, electronegativity decreases with increasing atomic number, and increasing metallic character.[[image:http://www.grandinetti.org/Teaching/Chem121/Lectures/Electronegativity/assets/ElectronegativityTrends.gif width="618" height="184"]]
 * ** Intermolecular Forces **
 * An attractive forces between molecules.
 * This force holds atoms together in a molecule.
 * **Van der Waals forces**
 * **Dipole-Dipole interactions**
 * An attractive forces between polar molecules between molecules that possess dipole moments.
 * It happens when polar molecules are attracted to one another.
 * [[image:http://www.ccs.k12.in.us/chsBS/kons/kons/wonderful_world_of_files/image006.gif width="176" height="258" caption="Dipole Dipole interaction"]]
 * **Dispersion forces**
 * It caused by electron movement.
 * It happens in non-polar molecules.
 * It increases with molar mass because molecules with larger molar mass tend to have more electrons, and it increases in strength with the number of electrons.
 * **Hydrogen Bonding**
 * A special type of dipole-dipole interaction between the hydrogen atom in a polar bond.
 * H+ covalently bond to a highly electron element is weakly associated to an unshared pair of electron on another.
 * This bonding is the strongest of all intermolecular forces.
 * ** Intramolecular Forces **
 * A force that holds the atoms or ions together in a compound.
 * There are three types of Intramolecular Forces
 * 1) Metallic Bonding
 * 2) Ionic Bonding
 * 3) Covalent Bonding
 * Intramolecular forces are much stronger than intermolecular forces.
 * The physical properties of metals and ionic substances are dependent ONLY on strong intramolecular forces (Metallic Bonding and Ionic Bonding).
 * The physical properties of three dimensional covalent network substances are determined by strong intramolecular forces (Covalent Bonding).
 * The physical properties of molecular covalent substances are determined by weaker intermolecular forces.
 * ** Comparison of Intramolecular Forces **
 * <  ||~ Metallic Bonding ||~ Ionic Bonding ||~ Covalent Bonding ||~   ||
 * ~ occurs when ||< metal atoms bond to each other ||< cations and anions bond ||< 2 atoms share a pair of electrons and each atom provides 1 electron for the bonding pair.
 * ~ occurs between ||< metal atoms ||< metal and non-metal ions ||< non-metal atoms ||
 * ~ bond characteristics ||< delocalised electrons shared between atoms ||< cations and anions are held together by electrostatic attraction or forces ||< electrons are shared between two atoms ||
 * ~ typical example ||< Group I (alkali) Metals, eg, Na
 * ~ melting/boiling point ||< high ||< high ||< Molecular Substances: low
 * ~ solubility in water ||< insoluble ||< soluble ||< Molecular Substances: dependent on the intermolecular forces
 * ~ conductivity of solid ||< good ||< poor ||< usually poor ||
 * ~ conductivity of liquid ||< good ||< good ||< usually poor ||
 * ~ conductivity of aqueous solution ||< N/A ||< good ||< usually poor unless the substance reacts with water to form ions (eg, HCl reacts with water to form hydrogen ions and chloride ions) ||  ||
 * # Free Response Question from 2010 **
 * = Use the information in the table of the link [|EX.1](Pg. 11 #5) respond to the statements and questions that follow. Your answers =
 * 1) Draw the complete Lewis electron-dot diagram for ethyne in the appropriate cell.
 * 2) Which of the four molecules contains the shortest carbon-to-carbon bond? Explain.
 * 3) Identify a compound from the table of the link [|EX.1] (Pg. 11 #5) that is non-polar. Justify your answer.
 * 4) Ethanol is completely soluble in water, whereas ethanethiol has limited solubility in water. Account for the difference in solubilities between the two compounds in terms of intermolecular forces. ANS EX.1 (Pg. 12~ Pg.14)


 * # Free Response Question from 2008 **
 * Answer the following questions about atomic fluorine, oxygen, and xenon, as well as some of their compounds.
 * 1) Xenon can react with oxygen and fluorine to form compounds such as XeO3 and XeF4 . Draw the complete Lewis electron-dot diagram for each of the molecules.
 * 2) On the basis of the Lewis electron-dot diagrams you drew for part (1), predict the following:
 * 3) The geometric shape of the XeO3 molecule
 * 4) The hybridization of the valence orbitals of xenon in XeF4
 * 5) Predict whether the XeO3 molecule is polar or nonpolar. Justify your prediction. ANS EX.2 (Pg. 13)