Download
1 / 39
390 likes | 465 Views
CHAPTER 6. CHEMICAL BONDING. 6.1. Introduction to Chemical Bonding. 6.1 NOTES. Seldom do atoms exist as independent particles, they are almost always bonded to another atom.
E N D
CHAPTER 6 CHEMICAL BONDING
6.1 Introduction to Chemical Bonding
6.1 NOTES Seldom do atoms exist as independent particles, they are almost always bonded to another atom. A chemical bond is an electrical attraction between the nucleus of one atom and the valence electron(s) of a differentatom. These electrons can be lost, gained, or shared.
6.1 NOTES 3 Types of bonding: • Ionic • Covalent • Metallic • What is it based on? • Use table of electronegativities to determine if the bond is covalent or ionic. • Covalent = 0.0 – 1.9 • Ionic = 2.0 and greater
6.1 NOTES How to use the chart of electronegativites: Think of bonding as a continuum with no distinct lines separating ionic and covalent.
6.1 NOTES 3 Types of bonding: • Ionic – bond in which electron(s) are transferred from one atom to another. Ions are formed in the process.
6.1 NOTES 3 Types of bonding: • Ionic – bond in which electron(s) are transferred from one atom to another. Ions are formed in the process. Usually between a metal (loses e-) and a nonmetal (gains e-).
6.1 NOTES 3 Types of bonding: • Ionic – bond in which electron(s) are transferred from one atom to another. Ions are formed in the process.
6.1 NOTES 3 Types of bonding: • Covalent - bond in which electron(s) are shared between atoms. Also based on the table of electronegativities. • What are the two types of covalent bonds?
6.1 NOTES 3 Types of bonding: • Covalent – bond in which electron(s) are shared between atoms. Also based on table of electronegativities. • Nonpolar covalent bond – electron(s) are shared evenly. Electronegativity difference is between 0 and 0.4 • Polar covalent bond – electron(s) are not shared evenly. Electronegativity difference is between 0.5 and 1.9 A polar bond is shown by adding an arrow pointed toward the atom that has the higher electronegativity lower higher
6.1 NOTES 3 Types of bonding: • Covalent – bond in which electron(s) are shared between atoms. Also based on table of electronegativities.
6.1 NOTES 3 Types of bonding: • Metallic – unique bonding in metals which accounts for their properties (magnetism, electrical and heat conductors)
6.1 NOTES What type of bonding do you expect between the following atoms? If an arrow is needed, which way should it be pointing? • C and Cl • C and O • Na and F • H and O • O and O • Ca and O • H and Br • Li and I
6.2 Covalent Bonding the Molecular Compounds
6.2 NOTES What is the particle formed in a covalent bond?
6.2 NOTES A molecule is the particle formed when two atoms bond covalently. In forming a covalent bond, the atoms become more stable and energy is released (-kJ/mol) If that same bond is broken, the same amount of energy has to be added. (+kJ/mol) The length of the bond (distance between nucleii) depends on 2 things: • the strength of the bond • the size of the atoms in the bond
6.2 NOTES Each bond between two atoms has a particular distance it establishes between the two atoms for the lowest energy state.
6.2 NOTES In forming covalent bonds, atoms share electrons, so at least some of the time they attain a noble gas configuration. (octet rule) These electrons are shared throughout the entire molecule. p. 182 & 183
6.2 NOTES What are each of the following: • Single covalent bond – • Double covalent bond – • Triple covalent bond -
6.2 NOTES How to draw covalently bonded molecules: Use electron dot notation to represent valence electrons (p. 184)
6.2 NOTES How to draw covalently bonded molecules: Use electron dot notation to represent valence electrons (p. 184) Lewis Structures show atoms and shared electrons
6.2 NOTES How to draw covalently bonded molecules: Use electron dot notation to represent valence electrons (p. 184) Lewis Structures show atoms and shared electrons
6.2 NOTES How to draw covalently bonded molecules: Use electron dot notation to represent valence electrons (p. 184) Lewis Structures show atoms and shared electrons
6.2 NOTES How to draw covalently bonded molecules: • Determine the type and number of atoms in the molecule. • Determine the total number of valence electrons in the molecule. (Polyatomic ions are covalently bonded. If drawing one of these either add electron(s) if it has a negative charge or subtract electron(s) if it has a positive charge.) • Arrange the atoms by putting carbon in the middle if it is present, otherwise put the least electronegative element in the middle. Hydrogen never goes in the middle. • First add electrons in pairs between atoms (bonds), second add electrons on the outer atoms (unbonded). • If the center atom has less than an octet, form multiple bonds (double or triple) to satisfy the octet rule. • If there are extra electrons, place them on the center atom.
6.2 NOTES Structural formulas are the same as Lewis, except the electrons involved in the bonding have been replaced by lines. p. 186 & 187
6.2 NOTES Let’s draw some molecules! HCl
6.2 NOTES Let’s draw some molecules! PF3
6.2 NOTES Let’s draw some molecules! SiH4
6.2 NOTES Let’s draw some molecules! O2
6.2 NOTES Let’s draw some molecules! N2
6.2 NOTES Let’s draw some molecules! CCl4
6.2 NOTES Let’s draw some molecules! CO2
6.2 NOTES Let’s draw some molecules! HCN
6.2 NOTES Let’s draw a polyatomic ion! NO31-
6.2 NOTES Let’s draw a polyatomic ion! NH41+
6.2 NOTES Exceptions to the octet rule: • Molecules with an odd number of electrons- ClO2, NO, NO2
6.2 NOTES Exceptions to the octet rule: • Molecules with an odd number of electrons- ClO2, NO, NO2 • Molecules in which an atom needs less than an octet – B only needs 6 electrons to become stable and Be only needs 4 electrons.
6.2 NOTES Exceptions to the octet rule: • Molecules with an odd number of electrons- ClO2, NO, NO2 • Molecules in which an atom needs less than an octet – B only needs 6 electrons to become stable • Molecules in which an atom can have more than an octet. This typically occurs in period 3 atoms and beyond. P, S, and Xe are common examples.
6.2 NOTES Resonance structures are molecules which can’t be accurately represented by only one structural formula.
