PAP Chapter 6 CHEMICAL BONDING Cocaine

Chemical Bonding  A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

Metallic Bonding  Metallic – holds atoms of metals together  A metallic bond is formed between all metals. Examples include a piece of Copper, Zinc, Sodium, Iron. Any metal  Metallic bonds result from the attraction between metal atoms and the surrounding sea of delocalized electrons  valence electrons can move freely around the whole metal structure—they are not confined to any one atom

Characteristics of Metals O Ductile- draw into thin wire O Malleable-ability to hammer into O Thin sheets O Shiny (luster) O Conducts heat and electricity O Solid at room temperature

Metallic Bonding This model explains many of the properties of metals: the mobile electrons can enter/leave the metal structure, so metals are good conductors of heat and electricity! metal cations are not locked into any crystal structure, so they can slide past each other when stressed. That makes metal malleable and ductile. The de-excitation ( The electron falling back down to a lower energy level) is responsible for the shiny (luster) appearance of metals.

Ionic Bonds O 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

Remember Cations and Anions  An atom which loses an electron becomes positively charged and is called a cation.  Examples: Na +, K +, Ca 2+, Al 3+  Metals usually become cations  An atom which gains an electron becomes negatively charged and is called an anion.  Nonmetals usually become anions.  Examples: Cl -, S 2-, N 3-

Characteristics of Ionic Bonds The electrons align themselves into an orderly arrangement that is known as crystal lattice. This makes the shape that you see into a crystal.

Characteristics of Ionic Compounds  Strong attraction between ions  Soluble in water  Conduct electricity in solution  Conduct electricity when molten  High melting points  High boiling points  Hard but brittle  Solid at room temperature

What are electron dot formulas or Lewis Dot Diagrams  Electron dot formulas are diagrams that show valence electrons in the atoms of an element as dots around the symbol of the element. 1. Make Generalizations 1. Make Generalizations As illustrated by the electron dot formulas in the table, how are the elements in a group similar?

How are electron dot formulas for ionic bonds constructed? Ionic bonds form as a result of electrostatic attractions between cations and anions. Electron dot formulas show the valence electrons and charges of these ions and may be used to illustrate the ionic bonds.

 Examples

Examples O Lithium and Bromine O Magnesium and Chlorine O Sodium and Sulfur O Aluminum and Oxygen

Day 2

Covalent Bonds :  Covalent —valence electrons are shared between nonmetal atoms

Covalent Bond  Nonpolar Covalent bond- Covalent bond in which the bonding electrons are shared equally by the bounded atoms, resulting in a balanced distribution of an electrical charge.  Polar Covalent bond- Covalent bond in which the bonded atoms have an unequal attraction for shared electrons.  Each atom has a charge  Partial positive- (+δ)  Partial negative- (- δ)

Characteristics of Covalent Bonds O Two non-metals bonded together O Relatively weak bonds O Usually a gas or liquid at room temperature O Does not conduct electricity in solution O Low melting point O Low boiling point O Soluble in alcohol and insoluble in water

The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together. Electronegativity Increases Electronegativity Decreases

O If the difference in electronegativities is between: O 1.8 to 4.0: Ionic O 0.4 to 1.7: Polar Covalent O 0.0 to 0.3: Non-Polar Covalent Example: NaCl Na = 0.9, Cl = 3.0 Difference is 2.1, so this is an ionic bond! Electronegativity Difference Example: H 2 O H = 2.1, O = 3.5 Difference is 1.4, so this is a polar covalent bond !

O Diatomic Molecules are nonpolar covalent compounds. O A diatomic molecule is a molecule only containing two atoms O The seven diatomic molecules are O H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 O These diatomic molecules are never by themselves. Diatomic Molecules

Octet Rule Chemical Compounds tend to form so that each atom, by gaining, losing or sharing electrons, has an octet of eight electrons in its highest occupied energy level

Electron dot formulas:covalent bonds O In a covalent bond, no ions form. Instead, nonmetallic atoms share electrons, which results in each atom having a noble- gas configuration. O Single Covalent Bonds -A single covalent bond results when two atoms share one pair of electrons, as in the case of hydrogen gas, which is a diatomic molecule.

O Examples: F 2, HBr, Cl 4 (contd.)

O Double Covalent Bonds- When atoms bond by sharing two pairs of electrons, the result is a double covalent bond, as in a molecule of carbon dioxide, CO 2. The double bond is shown by four dots or two dashes.

Triple Covalent Bonds- When atoms bond by sharing three pairs of electrons, the result is a triple covalent bond, as in a molecule of nitrogen gas, N 2. The triple bond is shown by six dots or three dashes.

Violations of the Octet Rule Usually occurs with B and elements of higher periods. Common exceptions are: Be, B, P, S, and Xe. BF 3 SF 4 Be: 4 B: 6 P: 8 OR 10 S: 8, 10, OR 12 Xe: 8, 10, OR 12

Lewis Structures HF H2OH2O  or H - F  H O H  or H - O - H 

NH 3 CH 4 H N H H  or H - N - H H  H C H H H  or H - C - H H H Lewis Structures

Extra Slides

H 2 O is POLAR because it has a positive end and a negative end. (difference in electronegativity) O has a greater share in bonding electrons than does H. O has slight negative charge (-δ) and H has slight positive charge (+ δ) Bond Polarity

 This is why oil and water will not mix! Oil is nonpolar, and water is polar.  The two will repel each other, and so you can not dissolve one in the other Bond Polarity

 “Like Dissolves Like”  Polar dissolves Polar  Nonpolar dissolves Nonpolar  Polar also dissolved Ionic  Ex: water and salt