Unit 4 Notes and In-Class Problems

Transcription

1 Unit 4 Notes and In-Class Problems There are three main types of bonding: Ionic, Metallic, and Covalent Ionic bonding occurs between Metallic bonding occurs between Covalent bonding occurs between Positive ion: Negative ion: Metallic Bonding Metallic bonds consist of positively charged metallic cations that donate electrons to the The sea of electrons are shared by all atoms and can move throughout the structure. Properties: o Thermal conductivity o Electrical conductivity o Malleability the ability to be hammered into thin sheets o Ductility the ability to be drawn into a wire Ionic Bonding When a metal bonds with a nonmetal, an bond is formed An ionic bond always involves the TRANSFER of electrons from the to the An ionic compound does not consist of individual molecules. Instead, there is a huge network of positive and negative ions that are packed together in a Because their bonds are so strong, ionic compounds tend to have very melting points Ionic compounds are, which means they can conduct electricity When forming compounds, the positive and negative charges must balance! Covalent Bonding Covalent bonds are between 1

2 Covalent bonds are formed when electrons are between two atoms. o If two atoms share 2 electrons, they form a o If two atoms share 4 electrons, they form a o If two atoms share 6 electrons, they form a There are two types of covalent bonds: polar and nonpolar. o Polar bonds usually involve o In polar bonds the electrons are shared o Nonpolar bonds usually involve o In nonpolar bonds the electrons are shared Covalent compounds can exist in any state (solid, liquid, gas). They have melting and boiling points. Determining the Type of Bond: Determine the type of bond (Ionic, Covalent or Metallic) in the following compounds: Compound Bond Type Compound Bond Type NaCl NCl3 CO FeNi SiS2 PF3 CaCl2 Fe2O3 Covalent Naming: Binary covalent compounds are characterized by having two nonmetals. Naming these compounds involves the use of numerical prefixes: Prefix Number Prefix Number

3 Rules: 1. The subscript tells you how atoms of that particular element are in the compound. 2. The prefix that goes IN FRONT OF the element s name is based on how many atoms there are of that element. 3. The name of the first element DOES NOT change, except by adding a prefix for how many there are. 4. The name of the second element DOES change, by changing the ending to ide and adding a prefix. 5. Do not use the prefix mono- for the first element if there is only one atom of that element. Practice: o N2O4 o Diarsenic pentoxide o XeF4 o Phospohorus pentabromide o N2O5 o Carbon tetraiodide o CO o Trisilicon tetranitride o CBr4 o Tetraphosphorus decoxide Ionic Naming: Writing formulas: Criss-Cross method- 1. The charge on the metal crosses down to become the subscript on the nonmetal. 2. The charge on the nonmetal crosses down to become the subscript on the metal. 3. Drop the signs. 4. Reduce if necessary. 3

4 Examples: 1. Barium and Oxygen 2. Lithium Iodide 3. Strontium Chloride 4. Aluminum Nitride 5. Sodium Sulfide Basic Ionic Naming: 1. The first element listed is the metal. The name of the metal does not change. 2. The second element listed is the nonmetal. The ending of the nonmetal changes to ide 3. If the metal has more than one possible charge (the transition metals), use a Roman Numeral to indicate which charge the metal has. Examples: 1. Al2O3 2. BaCl2 3. Ca3N2 4. KF 4

5 Ionic Compounds containing metals that can have more than 1 charge: Transition metals and some p-block metals can have multiple oxidation states (charges). Exceptions: Silver (Ag) is ALWAYS +1 Zinc (Zn) is ALWAYS +2 Cadmium (Cd) is ALWAYS +2 Aluminum (Al) is ALWAYS +3 *Remember, Group 1 metals always have a charge of +1 and group 2 metals always have a charge of +2. The transition metals and p-block metals that are not exceptions form two or more cations with different charges. Roman Numerals are used to indicate the charge on the metal. Example: Fe +2 would be named Iron (II) How would you name Fe +3? Roman Numerals: Number Roman Numeral 1 (I) 2 (II) 3 (III) 4 (IV) 5 (V) 6 (VI) 7 (VII) 5

6 Rules: Criss-cross subscripts back up to determine charges First element is always positive Second element is always negative o Nonmetals in Ionic compounds can only have 1 possible charge use those to help you un-reduce if needed Nonmetal charges in ionic compounds: Group 15: Group 16: Group 17: Hydrogen (listed second): Examples: 1. NiCl3 2. PbO2 3. Fe2S3 When writing the formula from the name, write the ions with their appropriate charge then criss-cross down. Reduce if necessary. 1. Iron (III) oxide 2. Tin (IV) chloride 3. Chromium (II) sulfide 6

7 Naming Ionic Compounds Containing Polyatomic Ions: Ions formed from a single atom are known as monatomic ions You wrote formulas for ionic compounds using monatomic ions Some ionic compounds contain more than 2 elements, where there are 2 or more atoms combined covalently, but have a charge, known as a polyatomic ion The name of the polyatomic ions must be memorized Rules: Follow the rules for all other ionic naming. The name of the polyatomic ion will not change When criss-crossing down, place parentheses around the polyatomic ion to be sure you don t put a subscript where it doesn t belong When criss-crossing back up, subscripts that are part of the polyatomic ion s formula STAY with the polyatomic ion! Do not make it a charge! Examples: 1. Sodium sulfate 2. Ammonium oxide 3. Zinc phosphate 4. BaCO3 5. NaNO3 6. Al(CN)3 7

8 Acid Nomenclature If the compound begins with Hydrogen, it is an acid. Steps: 1. Name the compound ionically. 2. Look at the ending of the ionic name. a. If the compound ends in -ide, change the ending to -ic and add hydro- to the beginning. Add "acid" to the end. b. If the compound ends in -ate, change the ending to -ic and drop hydrogen completely. Add "acid" to the end. (c. If the compound ends in -ite, change the ending to -ous and drop hydrogen completely. Add "acid" to the end.) Examples: 1. HCl 2. H2SO4 3. H3PO4 4. Hydrobromic acid 5. Nitric acid 6. Chloric acid 8

9 Lewis Structures Lewis structures (also called dot diagrams) represent how atoms bond in ionic or covalent compounds. Need to know: o Elements symbols o Number of valence electrons in each element in the compound Draw dot diagrams for the following elements: 1. Carbon 2. Nitrogen 3. Oxygen Drawing dot diagrams for covalent compounds: The atom listed first in the compound is typically the central atom (exceptions: H2O, H2S) Use the Needed, Available, Shared method for determining how many bonds will be drawn o Needed how many electrons are needed for each atom (most atoms want 8 electrons; some are fine with less ) o Available how many total valence electrons are available for bonding; sum of the valence electrons in the compound o Shared = needed available Examples: 1. PCl3 2. CBr4 9

10 Ionic Dot Diagrams: For Ionic Compounds, the electrons are gained and lost, NOT shared, so the diagram looks a little different. Each atom is in brackets with the valence electrons around it and the charge outside of the bracket. Examples: 1. Sodium chloride 2. Aluminum sulfide Polarity (Covalent only!) The bonding pairs of electrons in covalent bonds are pulled, as in the tug of war between the nuclei of the atoms sharing the electrons. Polarity is when a molecule acts like a little magnet. This is because one side of the molecule has more electrons than the other side. For BOND POLARITY look at differences in electronegativity! Difference in Electronegativity Most probable type of bond Nonpolar covalent Polar covalent 1.8 or greater Ionic **You do not need to memorize any specific electronegativity values. You will be provided with a chart containing the values for any questions pertaining to bond polarity. 10

11 Examples: Bond Electronegativity Difference Type of bond C N Li F N Cl Na Cl O F B H Ba F C H Al S For MOLECULAR POLARITY LOOK AT SYMMETRY! If the molecule is symmetrical, then it is non-polar If the molecule is asymmetrical, then it is polar o Lone pairs on the central atom make the molecule polar! Examples (you have to DRAW!): 1. NH3 4. CO2 2. SCl2 5. CHCl3 3. CCl4 6. BF3 11

12 VSEPR: Valence Shell Electron Pair Repulsion VSPER is used to describe the of molecules Single, double or triple bonds act the same Electron densities (lone pairs and bonds) will arrange themselves equidistant around an atom to repulsive forces Unbonded electrons (lone pairs of e-) take up space than bonded pairs 12