CHEMICAL BONDING Cocaine Chemistry I – Chapter 8 To play the movies and simulations included, view the presentation in Slide Show Mode.

What is a Chemical Bond? Chemical bonds involve ONLY the valence electrons of atoms. The nucleus is unaffected, so each atom keeps its identity (carbon stays carbon) Chemical bonds seek stability in the octet rule – Most atoms want to have 8 valence electrons. 3 ways to achieve this: Transferring electrons – Ionic bonds Sharing electrons – Covalent bonds Floating electrons – Metallic bonds

Types of Chemical Bonds Ionic Bond: Chemical bonding that results from the bonding between metals and nonmetals. Electrons are transferred from one atom to another, from the metal (electropositive atom) to the nonmetal (electronegative atom.)

Atoms that lose valence electrons become positively charged and are known as CATIONS. Atoms that gain valence electrons become negatively charged and are known as ANIONS. Cations and Anions come together to form a NEUTRAL 3D crystal

Properties of Ionic Compounds Are hard, rigid, brittle crystalline solids which may or may not dissolve in water. Have high melting and boiling points. Good conductors of electricity when melted or when dissolved (aqueous solution); nonconductors when solid.

Metallic Bonds Two metals do not transfer electrons to other metals nor do they share electrons. Their valence electrons are not held to any specific atom in the solid metal and can move freely from one metal atom. Think Hot Potato! (Electron Sea Model).

Properties of Metals Because these delocalized electrons are free to move, metals are excellent conductors of electricity and heat. They have a luster and are malleable, ductile, and usually durable. Melting points are variable but boiling points are all high.

Covalent bond: Chemical bonding that results from the sharing of electrons between two nonmetals Covalent bonds equally (NONPOLAR) or unequally (POLAR) share the electrons between the atoms. A group of covalently bonded atoms is called a molecule

Properties of Covalent Compounds Because covalent bonds are generally weaker than ionic bonds, the melting points and boiling points of covalent compounds are generally lower than those of ionic compounds. Many covalent compounds exist as gases at room temperature.

Review of Chemical Bonds There are 3 forms of bonding: _________—complete transfer of 1 or more electrons from one atom to another (one loses, the other gains) forming oppositely charged ions that attract one another _________—some valence electrons shared between atoms _________ – holds atoms of a metal together Most bonds are somewhere in between ionic and covalent.

Bond and Lone Pairs In Covalent Bonds valence electrons are distributed as shared or BOND PAIRS , and unshared or LONE PAIRS. • •• H Cl shared or bond pair lone pair (LP) This is called a LEWIS structure.

Steps for Building a Dot Structure Ammonia, NH3 1. Decide on the central atom; never H. Why? If there is a choice, the central atom is the atom with lowest electronegativity. Therefore, N is central on this one 2. Add up the number of valence electrons that can be used. H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons

Building a Dot Structure 3. Form a single bond between the central atom and each surrounding atom (each bond takes 2 electrons!) H N 4. Remaining electrons form LONE PAIRS to complete the octet as needed (or duet in the case of H). H •• N 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair (2 electrons).

Building a Dot Structure Check to make sure there are 8 electrons around each atom except H. H should only have 2 electrons. This includes SHARED pairs. H •• N 6. Also, check the number of electrons in your drawing with the number of electrons from step 2. If you have more electrons in the drawing than in step 2, you must make double or triple bonds. If you have less electrons in the drawing than in step 2, you have made a mistake!

Carbon Dioxide, CO2 1. Central atom = 2. Valence electrons = 3. Form bonds. C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons This leaves 12 electrons (6 pair). 4. Place lone pairs on outer atoms. 5. Check to see that all atoms have 8 electrons around it except for H, which can have 2.

Carbon Dioxide, CO2 C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons Do we have octets? 6. Not every atom has its octet. We must form DOUBLE BONDS between C and O. Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond.

Double and even triple bonds are commonly observed for C, N, P, O, and S H2CO SO3 C2F4

Now You Try One! Draw Sulfur Dioxide, SO2

Electronegativity and Polarity The difference in electronegativities between bonding atoms may be used to predict the type of bond that forms. For our purposes we will consider bonds between metals and nonmetals to be ionic and bonds between two nonmetals to be covalent.

Although all covalent bonds involve a sharing of one or more pairs of electrons between bonding atoms, most of the time this sharing is not equal. One of the atoms is more electronegative than the other so it “hogs” the shared bonded pair of electrons more of the time. This type of covalent bond is known as a polar covalent bond.

Electronegativity Difference If the difference in electronegativities is between: 1.7 to 4.0: Ionic 0.5 to 1.7: Polar Covalent 0.0 to 0.4: Non-Polar Covalent Example: NaCl Na = 0.8, Cl = 3.0 Difference is 2.2, so this is an ionic bond!

Bond Polarity Check the polarity of HCl: 3.0 – 2.1 = 0.9 Because it is polar, it has positive and negative ends Cl has a greater share in bonding electrons than does H. Cl has slight negative charge (-d) and H has slight positive charge (+ d) called DIPOLES

Bond Polarity “Like Dissolves Like” Polar dissolves Polar Nonpolar dissolves Nonpolar

Polarity of Whole Molecule Determining the polarity of a bond is easy, but what about the molecule as a whole? Strangely enough, you can have nonpolar molecules that contain polar bonds! SO…. How do you determine if a molecule itself is going to be polar or nonpolar? Look at the Lewis Structure and check out the central atom!

Determining if a Molecule is NonPolar or Polar Look at what is attached to the central atom -If every attached structure is the same -If the attached structures are different -The molecule is symmetrical -The molecule is asymmetrical -There are NO lone pairs on central atom -There are lone pairs on central atom And therefore NONPOLAR And therefore POLAR

Polar Bond δ - δ + H Cl Asymetrical Polar Molecule

O = C = O The molecule is nonpolar So… Polar Bonds δ - δ + δ + δ - O = C = O However, the shape is symmetrical So… The molecule is nonpolar

Nonpolar Bond Br Br Symmetrical Nonpolar Molecule

MOLECULAR GEOMETRY

VSEPR MOLECULAR GEOMETRY Valence Shell Electron Pair Repulsion theory. Molecule adopts the shape that minimizes the electron pair repulsions. VSEPR Valence Shell Electron Pair Repulsion theory. Most important factor in determining geometry is relative repulsion between electron pairs.

VSEPR charts Use the Lewis structure to determine the geometry of the molecule How the electrons are arranged determines the bond angles and shapes. Arrangement focuses on the CENTRAL atom for all data! Think REGIONS WHERE ELECTRONS ARE LOCATED rather than bonds (for instance, a double bond would only be 1 region)

Some Common Geometries Linear Tetrahedral Trigonal Planar

Structure Determination by VSEPR The electron pair geometry is TETRAHEDRAL 1) Looking at the central atom, how many bonding regions are there? 2) How many bonded atoms? How many lone pairs? Water, H2O The molecular geometry is BENT. 2 bond pairs 2 lone pairs

Structure Determination by VSEPR Ammonia, NH3 Draw! The electron pair geometry is Tetrahedral , and the Geometry name is TRIGONAL PYRAMID.

Van der Waal Forces of Attraction In covalent compounds the bonds within a molecule are quite strong but the attraction between different molecules is relatively weak. These weak forces of attraction between individual molecules are known as intermolecular forces (IMFs) or Van der Waal Forces. There are three kinds of Intermolecular forces – London Dispersal Forces, Dipole-Dipole Interactions, and Hydrogen Bonds.

London Dispersal Forces London Dispersal Forces are the weakest of the intermolecular forces and are caused by the temporary shifts of electrons in the electron clouds. These occur when nonpolar molecules are attracted to other nonpolar molecules.

Dipole-Dipole Interactions are attractions between oppositely-charged regions of polar molecules which allows one polar molecule to “stick” to another polar molecule.

Hydrogen Bonding is a special type of dipole-dipole interaction between a hydrogen on one polar molecule and an nitrogen, oxygen, or fluorine on an other polar molecule.

Hydrogen Bonding examples The best example of hydrogen bonding is water The strong IMFs create LARGE amounts of surface tension Explains why you can float a paper clip on the surface of water, but it will sink if it’s already under it. Aquatic bugs use this to walk on water Also explains why water is the only substance where the solid form (ice) floats on top of the liquid form.

Diagram/Summary of IMFs Hydrogen bonding – H, and N or O or F Dipole-Dipole forces – all polar compounds Polar Molecules London Dispersion forces – all compounds

IMFs and Phase Changes IMFs determine if the compound would rather be a solid, liquid, or gas at room temperature Also determine melting and boiling pts of the compound More kinds of IMFs means harder to change phase In addition to phase, intermolecular forces also play a role in compressibility, fluidity, viscosity, surface tension, and capillarity.