Chapter 6.2 (Partial), 6.4 and 7.1 (Partial) Covalent Compounds Naming Covalent Compounds Metallic Bonds

Covalent Bonds Molecules are formed as elements make covalent compounds. Covalent bonding decreases potential energy because each atom achieves an electron configuration like a noble gas. Individual atomic orbitals of each element overlap where the covalent bond occurs. This results in a concept called molecular orbitals.

Covalent Bonding

Molecules Carbon Dioxide CO2 Adenine

Covalent Bonds Potential energy determines the length of the bond. At minimum potential energy, the distance between two bonded atoms is called bond length. The bonded atoms vibrate about the bond, therefore, bond length is an average length.

Covalent Bonds Polar molecules have partially positive and partially negative ends. Such molecules are called “dipoles”. The positive end is designated as δ+ and the negative end as δ- example: δ+H--Fδ-

Naming Covalent Compounds First name: name of first element in formula usually least electronegative requires a prefix if more than one of them Second name: ends in –ide the prefix “mono” is sometimes omitted completely, but we will use it on the second element in the compound.

Prefixes 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca

Naming Covalent Compounds

Metallic Bonding Section 6.4

Properties of Substances with Different Types of Bonds

Introduction Delocalized: refers to valence electrons that are free to move throughout the structure In metals, valence electrons are delocalized and can move throughout the metal (shared by all atoms). Metals are therefore composed of cations surrounded by delocalized electrons.

Continued Sometimes referred to as the electron-sea model Metallic bonding is the electrostatic attraction between the metal ions and the delocalized electrons Delocalized electrons = stabilized structure

Visual of the Electron Sea

Physical Properties of Metals Ductile: can be drawn into a wire Malleable: can be hammered into thin sheets Both of these properties are attributed to the metal's ability to have layers pushed so they slide over one another without disrupting the metallic bonding Some metals are better at this than others

Thermal and Electrical Conductivity When a voltage is applied, the delocalized electrons are free to move This also aids in conducting heat as the movement of the electrons carries kinetic energy

Comparison of Properties Chapter 6

Introduction The physical properties of a substance depend on the forces between the particles of the chemical species of which it is composed. The stronger the bonding and the intermolecular forces, the harder the substances, and the higher the boiling point

Continued The presence of impurities always lowers the melting point as the regular lattice is disrupted. Volatility: a qualitative measure of how readily a liquid or solid is vaporized upon heating or evaporation

Metallic Structure Variable hardness, malleable rather than brittle Melting points and boiling points vary, depending on the number of valence electrons, but is generally high. Good electrical and thermal conductivity as solids and liquids Generally insoluble, except in other metals to form alloys (mixture of metals). Electron sea model: mass of positive nuclei in a sea of electrons

Ionic Structure Hard, but brittle High melting and boiling points DO NOT conduct heat and electricity as solids, but DO when molten or in solution. Often quite soluble in water compared to other solvents. (+) + (-) [From PT, Metal + Nonmetal]

Covalent Structure Usually soft and malleable unless H-bonded Low melting and boiling points. Liquids and gases at room temperature are usually molecular covalent. Do not conduct electricity or heat in any state More soluble in non-aqueous solvents, unless they can H-bond to water or react with it Examples: CO2, CH3CH2OH, I2 Shared e-; [From PT: nonmetal + nonmetal]