<p> Covalent Bonds</p><p>Properties  Nonmetal + Nonmetal (usually)  Most common type of bond  Covalent bonds form molecules o A substance made of molecules is called a molecular structure o Molecular formulas are used to describe the composition of a molecule  Form by sharing electrons o The sharing of one pair of electrons is a single bond (X-X) . Another name for single covalent bond is sigma bond symbolized by σ  Sigma bonds form from the overlap of a s orbital with another s orbital, a s orbital with a p orbital, or a p orbital with another p orbital o The sharing of two pairs- double bond (X=X) . Another name for multiple bonds is pi bond symbolized by π  Pi bonds form when parallel orbitals overlap to share electrons o A double covalent bond has one sigma and one pi bond o The sharing of three pairs-triple bond (XΞX) . A triple covalent bond has one sigma and 2 pi bonds  Bond polarity explains the attraction between the sharing o Nonpolar electrons are shared equally Ex. F-F (same electronegativity) o Polar electrons are not shared evenly Ex. H-F (different electronegativity)  Bond dissociation energy is the amount of energy needed to break a bond.  Bond enthalpy (∆H), for the breaking of a particular bond in a mole of gaseous substance  Bond Length is between 2 bonded atoms and is the distance between 2 nuclei o Decreases as the number of bonds between the atoms increase. o The strength of a covalent bond depends on how much distance separates bonded nuclei</p><p>Intramolecular Forces in Bonds Table</p><p>Force Basis of attraction Ionic cations and anions Covalent positive nuclei and shared electrons Metallic metal cations and mobile electrons</p><p>1 Intermolecular Forces</p><p>Intramolecular forces do not account for all attractions between particles. There are forces of attraction called intermolecular forces. o They can hold together identical particles or two different types of particles o Also called van der Waals forces 3 types: 1. Dispersion forces a. Sometimes called London dispersion forces b. The force between oxygen molecules i. Weak forces that result from temporary shifts in the density of electrons in electron clouds: ……… δ- δ+ δ- δ+</p><p>δ+ δδ+ Attraction δ+ δ+ Temporary attraction Temporary attraction δ= δ+ ←|- ←|-</p><p>2. Dipole –dipole: Attraction between oppositely charged regions of polar molecules o Stronger than dispersion forces . The more polar the molecule, the stronger the force 3. Hydrogen bonds: One special type of dipole-dipole dealing with hydrogen bonds  Very strong intermolecular force that is formed with a H end and a F, O, or N atom on the other dipole</p><p>Many physical properties of covalent molecular solids are due to intermolecular forces.  The melting and boiling points are relatively lower than Ionic (that is why salt doesn’t burn when you heat it but sugar will)  Many are gases or vaporized at room temperature  Hardness is also due to the intermolecular forces so covalent solids are soft in comparison to ionic solids</p><p>2 Naming Molecular Compounds Rules for Binary Molecular Compounds are similar to that of naming Ionic compounds except the names include prefixes indicating the number of atoms in the molecule. Numerial Prefixes  Mono-1  Di-2  Tri-3  Tetra-4  Penta-5  Hexa-6  Hepta-7  Octa-8  Nona-9  Deca-10</p><p>Exceptions: </p><p>H2O is water</p><p>NH3 is ammonia</p><p>Examples:</p><p>1. CO2 –Carbon dioxide 2. CO-Carbon monoxide</p><p>3. N2O4-dinitrogen tetroxide</p><p>4. SCl6 –Sulfur hexachloride</p><p>Practice</p><p>1. PBr3</p><p>2. CS2 </p><p>3. N2O</p><p>4. P2O5</p><p>5. SO2</p><p>6. N2O5</p><p>3 Covalent vs. Ionic Examples</p><p>1. CaCl2 –metal and a nonmetal: Ionic</p><p>2. AlPO4-metal and a negative polyatomic ion: Ionic</p><p>3. SO3 –Polyatomic ion with only nonmetals: Covalent</p><p>4. CH4-All organic compounds are covalent</p><p>5. P4-multiple nonmetals: covalent 6. HCl-this is an acid so covalent</p><p>Practice </p><p>1. FeCl3</p><p>2. C2H5OH</p><p>3. H2O</p><p>4. NaOH</p><p>5. NaCl</p><p>6. O2</p><p>7. BaSO4</p><p>8. KI</p><p>Metallic Bonds</p><p>Metallic Bonds are formed when metal cation attract free valance electrons  Often form solid lattices  8 to 12 other metal atoms surround each metal atom  The electrons involved in metallic bonding are called delocalized electrons because they are from to move throughout the metal and are not attached to a particular atom o Delocalized electrons move heat from one place to another much more quickly than the electrons in nonmobile electrons . The movement of mobile electrons around positive metallic cations explains why metals are good conductors  Mobile electrons easily move as a part of an electric current when electrical potential is applied to a metal o Delocalized electrons also interact with light, absorbing and releasing photons, thereby creating the property of luster in metals</p><p>Metal alloys are formed when a metal is mixed with one or more other elements 2 types: </p><p>4 1. Substituational alloy: have atoms of the original metallic solid replaced by other metals of similar size. Ex. Sterling Silver 2. Interstitial alloy: forms when the small holes (interstices) in metallic crystals are filled with smaller atoms. Ex Carbon Steel  Holes found in iron crystals are filled with carbon atoms, and the physical properties of iron are changed o The carbon added make the solid harder, stronger, and less ductile than pure iron, which increases its uses</p><p>Some Commercially Important Alloys Common Name Composition Uses 1. Brass Cu 67-90%, Zn 10-33% Plumbing, hardware, lighting 2. Bronze Cu 70-95%, Zn 1-12 %, Sn 1-18% bearing, bells, medals 3. Pewter Sn 70-95%m Sb 5-15%, Pb 0-15% Tableware 4. Stainless steel Fe 73-79%, Cr 14-18%, Ni7-9% Instruments, sinks 5. Sterling Silver Ag 92.5%, Cu 7.5% Tableware, jewelry 6. 10 carat Gold Au 42%, Ag 35%, Cu 38-46% Jewelry</p><p>Naming Binary Acids</p><p>Binary acids are acids with only two elements.  Prefix –hydro, stem of anion, and suffix –ic</p><p> Exception is HN3: Hydroazoic acid, where the root – azo is used for nitrogen.</p><p>Naming Ternary Acids and Bases</p><p>Ternary acids are acids that contain 3 elements.  Usually no prefix is used and the suffix is –ic.  Exceptions:  One less O than the most common : no prefix and suffix used is –ous  Two less O than the most common: prefix hypo- and suffix –ous  One more O than the most common: prefix per- and suffix –ic</p><p> Ex. HClO3 is the most common: Chloric acid</p><p> HClO2 has one less O so: Chlorous acid  HClO has two less O so: Hypochlorous acid</p><p> HClO4 has one more O than most common so: Perchloric acid</p><p>5 Ternary bases  Arrhenius bases are composed of metallic, or positively charged ions and the negatively charged hydroxide ion. Therefore, these bases are named by adding the word hydroxide to the name of the positive ion.  Ex. Sodium hydroxide is NaOH.</p><p>Naming Organic Acids and Bases</p><p>First you need to determine the functional group. </p><p>Note: You can name organic compounds by their IUPAC name or Common name.</p><p>1. Alkanes: Family of saturated hydrocarbons.  Cn H2n + 2  Molecules that contain only H and C that are connected by single bonds and with –ane ending.</p><p> Ex. Methane: CH4 (1 Carbon and 4 Hydrogens)  Has sp3 hybrid, tetrahedral structure, with bond angles of 109.5 degrees.  Others:2-14 (IUPAC Names) CH3CH3 Ethane</p><p>CH3CH2CH3 Propane </p><p>CH3(CH2)2CH3 Butane</p><p>CH3(CH2)3CH3 Pentane</p><p>CH3(CH2)4CH3 Hexane</p><p>CH3(CH2)5CH3 Heptane</p><p>CH3(CH2)6CH3 Octane</p><p>CH3(CH2)7CH3 Nonane</p><p>CH3(CH2)8CH3 Decane</p><p>CH3(CH2)9CH3 Undecane</p><p>CH3(CH2)10CH3 Dodecane</p><p>CH3(CH2)11CH3 Tridecane</p><p>CH3(CH2)12CH3 Tetradecane</p><p>6 Alkyl Group: If a H atom is removed from an alkane, a partial structure that remains is an alkyl.</p><p> Alkyl groups are named by replacing the –ane ending of the parent alkane with an – yl ending.  Ex. Methyl: -CH3  Others: -CH2CH3 Ethyl -CH2(CH2)5CH3 Heptyl</p><p>-CH2CH2CH3 Propyl -CH2(CH2)6CH3 Octyl</p><p>-CH2(CH2)2CH3 Butyl -CH2(CH2)7CH3 Nonyl</p><p>-CH2(CH2)3CH3 Pentyl -CH2(CH2)8CH3 Decyl</p><p>-CH2(CH2)4CH3 Hexyl  The alkyl groups are often used to determine the common names of compounds.  R represents an alkyl group</p><p>2. Alkenes: Family of unsaturated hydrocarbons with Carbon=Carbon double bonds. Cn H2n  Since fewer H’s than in alkanes they are unsaturated with an –ene ending.  Ex. Ethylene (IUPAC name is ethene): H2C=CH2  Ex. Propylene (IUPAC name is propene): CH3CH=CH2  Has sp2 hybrid, planar structure with bond angles of about 120 degrees.</p><p>3. Alkynes: Family of unsaturated hydrocarbons with CarbonCarbon triple bonds. Cn Hn  alkynes use the suffix: -yne  Line drawing representation would be RCCH Ex. Propyne: CH3CCH  Has sp hybrid, linear structure with bond angles of 180 degrees.  Ex. Acetylene (IUPAC name is ethyne): H-CC-H  Used to prepare acetic acid 4. Amines: Family of organic derivatives of ammonia, NH3.  Amines are named by adding the ending –amine to the name of the hydrocarbon from which it is derived.  Organic bases contain nitrogen with unshared pair of electrons.  Line drawing representation would be R-NH2  Ex. Methylamine H3C-NH2 and Ethanamine CH3CH2-NH2</p><p>5. Carboxylic Acid: Family of the most useful building blocks for synthesizing other molecules. O</p><p>║</p><p>7  Line drawing representation would be C R OH</p><p> They are named by adding the ending –oic acid to the name of the hydrocarbon from which the acid is derived.  Ex. HCOOH is methanoic acid, which is also known as formic acid.  Derivatives of carboxylic acid are: O</p><p>║</p><p>C</p><p>R X Acid halide (X=F,Cl, Br, or I) </p><p>With ending –yl halide</p><p> Ex. Acetyl chloride (from acetic acid) where the R is CH3 and the X is Cl.  Ex. Acetic anhydride where both the R’s are replaced with CH3. O O</p><p>║ ║</p><p>C C</p><p>R O R Acid anhydride: With ending anhydride</p><p>8</p>