 A compound is a pure substance composed of more than one atom  A chemical bond is a mutual electrical attraction between atoms in a compound  Compounds can either be molecular or ionic

 There are two types of bonding: 1. Ionic Bonds result from the transfer of electron from metal to nonmetal 2. Covalent Bonds result from the sharing of electrons between two atoms ▪ Polar covalent = unequal sharing ▪ Non-polar covalent = equal sharing

 Composed of oppositely charged ions  Composed of Metal + nonmetal  Metals form cations (+) ▪ Low IE and low EA mean these atoms lose electrons readily  Nonmetals form anions (-) ▪ high IE and high EA mean these atoms gain electrons readily  Electrically neutral  Formula unit: lowest whole # ratio of ions in the compound  EX: NaCl, CaF 2

 Solids at room temperature  High melting point (>400 0 C)  Soluble in water  Conduct electricity when melted or dissolved in water

 Composed of two or more nonmetals  Involve covalent bonding (Electrons are shared)  Sharing “tricks” each element into thinking that their outer shells are filled  Molecule: smallest unit of a molecular compound  Molecular formula: tells the type and number of atoms in a molecule  Ex: H 2 O, CO 2, CO

 low melting point (<400 0 C)  Usually NOT soluble in water  DO NOT conduct electricity

Metal + nonmetal = ionic Nonmetal + nonmetal = molecular NaCl CO LiF H2O MgS More practice in packet

 DUE MONDAY!!!  Ion Flashcards…  Name of ion on one side  Symbol with charge on the other side  6 sets…keep them separate  Quizzes will begin next week…one set at a time

 Chemical formulas show the kinds and numbers of atoms in the smallest representative unit of that compound  For example, CO2, the formula for carbon dioxide tells us that one carbon atom and two oxygen atoms form one molecule of CO2

 Chemical formula for a covalent compound  Show kind and number of atoms in a molecule  Does not give any information about structure

 A variety of diagrams and molecular models can be used to show molecular structure

 Noble gases stand alone, do not combine chemically with any other element  7 elements exist as diatomic molecules  Hydrogen = H2  Oxygen = O2  Nitrogen = N2  Fluorine = F2  Chlorine = Cl2  Bromine = Br2  Iodine = I2

 A formula unit tells the whole-number ratio of ions in an ionic compound  Ionic compounds do not exist as single units  NaCl = 1:1 ratio  MgCl 2 = 1:2 ratio

 Atoms combine in whole number ratios so that the masses of elements in a compound are also in the same proportion

Apply the rule for naming and writing formulas for molecular compounds

 A molecular compound composed of only two non-metallic elements  For example: H2O,  Atoms can often bond in more than one way  For example, CO vs CO 2  Naming conventions are important  Carbon monoxide will kill you, carbon dioxide will not

 Used to describe number of atoms of each element present in one molecule

 Carbon monoxide  Prefix mono = one  CO  Carbon dioxide  Prefix di = two  CO2 Notice…NO mono prefix on the first element

 Identify the prefix and the element that the prefix is attached to  Write the correct symbols for each element with the appropriate subscripts  Example: tetraiodine nonoxide  Tetra = 4 so there are 4 iodine atoms  Non or nona = 9 so there are nine oxygens  Molecular formula = I 4 O 9

 Sulfur trioxide  Phosphorous pentafluoride  Dinitrogen monoxide  Phosphorous trichloride  Dichlorine octaoxide

 Subscript after the symbol for the element will tell the prefix to use for that element  For example, CO 2 the #2 tells you there are two oxygen atoms therefore, oxygen will get the prefix  If first element in formula is only one, no prefix  For example, wouldn’t say monocarbon dioxide, just carbon dioxide  End molecular names with -ide

 CS 2 N2O3N2O3  OBr 2  SO 3  Carbon disulfide  Dinitrogen trioxide  Oxygen dibromide  Sulfur trioxide

 Electrons in the highest occupied energy level  Involved in bonding  Determine the chemical properties of the element

 Look at electron configuration or group # on periodic table  For example, Na  Electron configuration = 1s 2 2s 2 2p 6 3s 1 1 valence e-  All elements in group 1A have 1 valence e-  For example, Cl  Electron configuration = 1s 2 2s 2 2p 6 3s 2 3p 5 7 valence e-  All elements in group 7A have 7 valence e-

 Also called Lewis Dot diagrams  Show valence electrons around element symbol  Do not distinguish between s and p orbitals  Example: chlorine (7 valence electrons)

 all atoms want to achieve “NOBLE” status and be just like the NOBLE gases  Noble gases have 8 valence electrons (octet) ▪ Except helium which has only 2  Atoms will interact in a way to fill their outer energy level so that it contains 8 electrons  Electron dot diagrams can be used to show bonding

 Oxygen (group VIA; atomic#8)  How many valence electrons?  How many more valence electrons needed to be stable?  How many bonds will oxygen form?

Covalent Bonds Rewritten as

Covalent Bonds Rewritten as

1. Identify the elements that the color spheres in your kit will represent  If you have a blue sphere with 5 holes, connect two of the holes with a spring 2. Build the 3-D model, draw it to the best of your ability 3. Structural formula = Lewis dot diagram without unshared electrons represented  Do not try to recreate you ball and stick model

 Electronegativity = tendency of an atom to attract electrons

Differences in electronegativities determine the nature of the bond  If less than 0.4 = nonpolar covalent  equal sharing of electrons between atoms  If between 0.4 and 1.67 = polar covalent  unequal sharing of electrons  If greater than 1.67, IONIC

 Polar covalent?  Nonpolar covalent?  Ionic? 1. Na and Cl (0.9 vs 3.0) 2. carbon and hydrogen (2.5 vs 2.1) 3. hydrogen and oxygen (2.1 vs 3.5)

 Polar covalent?  Nonpolar covalent?  Ionic? 1. Na and Cl (0.9 vs 3.0) 2. carbon and hydrogen (2.5 vs 2.1) 3. hydrogen and oxygen (2.1 vs 3.5)

 Stands for Valence Shell Electron Pair Repulsion  Predicts the shapes of molecules  Depends on the # of electrons or atoms bonded to a central atom

 Bonding groups: 2  Nonbonding pairs: 0  Examples: ▪ BeCl2 ▪ CO2 ▪ HCN

 # of atoms or electron pairs: 2  # of unshared pairs: 1 or 2  Examples: ▪ H2O

 Bonding pairs: 3  Nonbonding pairs: 0  Examples: ▪ BF3 ▪ COCl2

 Bonding pairs: 3  Nonbonding pairs: 1  Examples: ▪ NH3 ▪ NF3 ▪ PCl3

 Bonding pairs: 4  Nonbonding pairs: 0  Examples: ▪ CH4 ▪ CCl4