Chapter 6 Chemical Bonding
Section 6-1 Introduction to Chemical Bonding
Chemical Bonds Valence electrons are attracted to other atoms, and that determines the kind of chemical bonding that occurs between atoms. The mutual (+ to -) attraction between the nuclei and valence e- of different atoms bind atoms together. This is a chemical bond. Why do atoms bond? To become more stable = have less energy To have a stable e- configuration (8ve-)
Types of Bonds Ionic Bonding cations + and anions - are attracted to one another Electrons are given and received Covalent bonding shared electrons are attracted between two atoms (eare attracted to both atoms nuclei) Can be single, double or triple Can be polar or nonpolar Chemical bonds are often a combination of types. The difference in the elements electronegativities determines the degree of ionic or covalent bonding.
Covalent Bonds Difference in electronegativity of 1.7 or less 2 types of covalent bonds Nonpolar - electronegativity difference from 0 to 0.3 The e- are shared equally between atoms Balanced distribution of charge Polar electronegativity difference of 0.3 to 1.7 E- are not shared equally between the two atoms (one atom has a higher electronegativity) Unequal charge distribution, one side of the molecule is partially + and one side is partially -
Ionic Bonds Difference of electronegativity of 1.7 to 3.3 One atoms gives valence e-, one atom receives valence e- Positive cation is attracted to negative anion
Section 6-2 Covalent Bonding and Molecular Compounds
Molecules Molecule Neutral H 2 H 2 O Atoms bonded by covalent bonds Compound with covalent bonds = molecular compound Chemical formula - symbols and subscripts of a compound MgCl 2 Molecular formula symbols and subscripts of a covalently bonded compound CH 3
Formation of a Bond Atoms bond with each other to become more stable. achieve a noble gas configuration Bonding lowers an atom s energy. The atoms will come together and bond at a distance where attractive and repulsive forces are balanced Attractive - electrons and nuclei Repulsive 2 negative electron clouds and 2 positive nuclei
Bond Length Bond length the average distance between two bonded atoms at their minimum potential energy Bond length in an H 2 molecule = 75 pm Varies for different molecules Bond energy energy to break a bond Measured in kilojoules per mole (kj/mol) Varies for different molecules
When two atoms form a covalent bond, their shared electrons form overlapping orbitals.
Bond Energies and Bond Lengths for Single Bonds
The Octet Rule Atoms can fill their outermost s and p orbitals by sharing electrons through covalent bonding. This gives them a noble-gas configuration, where their s and p orbitals are completely filled with 8 e-. Compounds form so that atoms have an octet of electrons in their highest occupied energy level. Atoms can lose, gain or share electrons to do this
Exceptions to the Octet Rule Hydrogen forms bonds in which it is surrounded by only two electrons. Boron has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons. Main-group elements in Periods 3-7 can form bonds with expanded valence, involving more than eight electrons.
Electron-Dot Notation A way to show how many valence electrons an atom has Use chemical symbol and dots that represent ve-
Write the electron dot notation for an atom of hydrogen sodium nitrogen
Lewis Structures Electron-dot notation can also be used to represent molecules. The pair of dots between the two symbols represents the shared electron pair of the hydrogen-hydrogen covalent bond. For a molecule of fluorine, F 2, the electron-dot notations of two fluorine atoms are combined. F F
The pair of dots between the two symbols = the shared pair of a covalent bond each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds. An unshared pair (lone pair) is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.
Lewis Structure: The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash. example: A structural formula indicates the kind, number, and arrangement, and bonds but not the unshared pairs of the atoms in a molecule. example: F F H Cl
Lewis Chapter Structures 6 Lewis structures: symbols represent nuclei and inner-shell e- dot pairs or dashes between symbols represent bonds dots only around one symbol represent lone pairs
Draw the Lewis structure of iodomethane, CH 3 I. 1. Determine the type and number of atoms in the molecule. 2. Write the electron-dot notation for each type of atom in the molecule.
3. Determine the total number of valence electrons available in the atoms to be combined. 4. If carbon is present, it is the central atom. Otherwise, the least-electronegative atom is central. Hydrogen, is never central. 5. Add unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by eight electrons.
Single, Double and Triple Bonds A single bond = a covalent bond in which one pair of electrons (2e-) is shared between two atoms A double bond = a covalent bond in which two pairs of electrons (4e-) are shared between two atoms. often found in molecules carbon, nitrogen, and oxygen molecules shown either by two side-by-side pairs of dots or by two parallel dashes greater bond energies and are shorter than single bonds
Double Covalent Bond
A triple bond = a covalent bond in which three pairs of electrons are shared between two atoms. even stronger and shorter than double bonds
Drawing Lewis Structures with Many Atoms
Draw the Lewis structure for methanal, CH 2 O, which is also known as formaldehyde.
Resonance Structures Resonance = bonding that can t be represented by ONE Lewis structure O 3 ozone