Ionic, Covalent, and Metallic Bonding (Chapter 7&8)

In nature, only the noble gases exists as uncombined atoms (monatomic) All other elements need to lose or gain electrons To form ionic compounds Some elements share electrons To form covalent compounds (molecules) 2

Valance Electrons Electrons in the outermost energy level (shell) of an atom (electrons involved in bonding) To determine the number of valance electrons in an atom, simply look at the group number Except for He = 2 valance electrons

Covalent Bonding

Octet Rule Octet rule: chemical compounds form so each atom (through gaining, losing, or sharing electrons) will have 8 valence electrons Exception: atoms trying to be like Helium (He) Electrons with 8 valance electrons are stable (unreactive) Noble Gases (except He)

DIATOMIC MOLECULES Some molecules are made up of the same element There are 7 naturally occurring diatomic molecules

Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen, Fluorine DIATOMIC MOLECULES Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen, Fluorine Br2. I2. N2. Cl2. H2. O2. F2

Examples: H2O (water), CO2 (carbon dioxide), and N2O (nitrous oxide) Molecules can also be made of different elements known as a molecular compound Examples: H2O (water), CO2 (carbon dioxide), and N2O (nitrous oxide)

Lewis Dot Structures Be C S Diagrams that shows valance electrons for a given element A way to show covalent bonding between atoms Be C S

Rules for Drawing Lewis Dot Structures Determine the TOTAL number of valance electrons in the molecular compound Arrange the atoms around a central atom. The central atom is usually the atom with the lowest number of valance electrons Circle the electrons that will bond with each other (shared pairs) Complete the octets (re-check) you have used the right # of valance electrons 14

Drawing Lewis Structures Rules for Drawing Lewis Structures 1.Count up the total number of valence electrons. First add up the group numbers of all atoms in the molecule 2.Calculate the total number of electrons that would be needed for each atom to have an octet (or doublet for H). 3.Subtract the result of step 1 from the result of step 2. This is the total number of shared or bonding electrons 19

4.Assign two bonding electrons to each bond. 5.If bonding electrons remain, assign them in pairs making some of the bonds double or triple bonds. (Usually, only C,N,O, and S can form double bonds, and only C and N can form triple bonds) 6.Assign remaining electrons as lone pairs, giving octets to all atoms except H 20

Types of Covalent Bonds Single covalent bond Two atoms held together by sharing one pair of electrons 33

A bond that involves two shared pairs of electrons. Double covalent bond A bond that involves two shared pairs of electrons. Triple covalent bond A bond formed by sharing three pairs of electrons 34

In F2, each fluorine atom has 3 unshared pairs of electrons A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair (non-bonding pair) In F2, each fluorine atom has 3 unshared pairs of electrons 35

Lewis Dot Structures Practice 1. HCl 2. H2O 3. SiO2 4. HCN 5.CH3OH 16

How do we describe the way electrons are shared in a covalent bond How do we describe the way electrons are shared in a covalent bond?...because it’s not always equal! 21

Covalent bonds differ in terms of how the bonded atoms share the electrons (equal or unequal) The character/properties of the molecule depends on the kind and number of atoms joined together 22

Nonpolar Covalent Bond (Nonpolar Bond) When the atoms in the bond pull equally the bonding electrons are shared equally Always results when an element bonds with itself 23

Dipole: an arrow which is drawn to represent a polar covalent bond (Polar Bond) Covalent bond between atoms in which the electrons are shared unequally More electronegative atom attracts more strongly and gains a partial negative charge The less electronegative atom has a partial positive charge δ+ δ– H—Cl H—Cl Dipole: an arrow which is drawn to represent a polar covalent bond 24 24

How do you determine whether a (covalent) bond is Polar or Nonpolar? 25

The electronegativity difference between two atoms tells you what kind of bond is likely to form Electronegativity Differences and Bond Types Electronegativity difference range Most probable type of bond Example 0.0–0.40 Nonpolar covalent H—H (0.0) 0.41–2.00 Polar covalent δ+ δ– H—F (1.9) >2.0 Ionic Na+Cl– (2.1) 26

Which type of bond (nonpolar covalent, polar covalent, or ionic) will form between each of the following pairs of atoms? a. N and H b. F and F c. Ca and Cl d. Al and Cl e. C and H f. Li and O 27

Determining if bonds in a molecule are polar or non-polar: CF4 H2O CH4 HCl ** Dipoles are drawn when a polar bond is present. Arrow towards the atom with the most valance electrons (most electronegative)

How do we determine the overall polarity of the molecule (Molecular Polarity)? 28

Draw the correct Lewis dot structure In order to determine whether the molecule as a whole is polar or non-polar you must: Draw the correct Lewis dot structure Determine whether each individual bond is polar or non-polar Assess which dipoles (polar bonds) will cancel each other out (determining the molecular polarity) If all dipoles cancel out, the molecule is non-polar If there are remaining dipoles, the molecule is polar

Basic Rules for Determining Molecular Polarity All non-polar bonds = non-polar molecule 2 atoms with a polar bond = polar molecule 1 or more polar bonds = depends on the shape of the molecule (do the dipoles cancel out?) 31

Determine the molecular polarity of each molecule Practice Problems Determine the molecular polarity of each molecule 1. HF 2. CO2 3. H2O 4. CCl4 5. CFCl3 32

(Valence Shell Electron Pair Repulsion) VSEPR Theory (Valence Shell Electron Pair Repulsion) Used to explain the 3-D shape of molecules The repulsion between electron pairs causes the valence-electron pairs (unshared pairs of electrons) to stay as far apart as possible 39

Common Molecular Shapes Linear Trigonal planar Bent Pyramidal Tetrahedral Trigonal bipyramidal Octahedral Square planar 40

Linear There are 2 atoms bonded to the central atom No unshared pairs of electrons on the central atom The double bonds joining the oxygens to the carbon are farthest apart when the O=C=O bond angle is 180° Carbon dioxide (CO2) No unshared electron pairs on carbon 180° 42

Bent There are 2 atoms bonded to the central atom 1 or 2 unshared pairs of electrons on the central atom With two unshared pairs repelling the bonding pairs, the bond angle is compressed to about 105° 43

Tetrahedral There are 4 atoms bonded to the central atom No unshared pairs of electrons on the central atom In this arrangement, all of the bond angles are 109.5° 41

Trigonal Pyramidal Unshared pair strongly repels the bonding pairs, pushing the bonding electrons together There are 3 atoms bonded to the central atom 1 unshared pair of electrons on the central atom The measured bond angle is only 107° 44

Trigonal Planar In the BF3 molecule, there are 3 atoms bonded to the central atom No unshared pairs of electrons on the central atom The measured bond angle is only 120° 42

MOLECULAR SHAPE SUMMARY 43