Chemistry: Hood River Valley High School Unit 4 Note Pack and Goals Name: Period: Unit 4 Bonding, Chemical Names, and Formulas Unit Goals- As you work through this unit, you should be able to: 1. Use the periodic table to infer the number of valence electrons in an atom and draw its electron dot formula. 2. Describe the formation of cations and anions from metals/nonmetals using electron dot formula and electron configuration. 3. Describe the formation of an ionic bond and the characteristics of an ionic bond. 4. Relate the model for metallic bonding to properties of metals. 5. Understand the difference in properties between ionic compounds and covalent compounds. 6. Use electron dot structure to show the formation of single, double, and triple bonds and draw their structural formula. 7. Use VSEPR theory to predict the shapes of simple covalently bonded molecules. 8. Use electronegativity values to classify a bond as nonpolar covalent, polar covalent, or ionic. 9. Name and describe the weak attractive forces that hold groups of molecules together. 10. Distinguish between ionic and molecular compounds. 11. Relate cations and anions to metals and nonmetals. READ: 12. Distinguish between chemical formulas, molecular formulas, and formula units. 13. Know the charges of the formulas for monatomic ions using the periodic table. 14. Be familiar with the charges and the formulas for polyatomic ions. 15. Apply the rules for naming and writing formulas for binary and ternary ionic compounds. 16. Apply the rules for naming and writing formulas for binary molecular compounds. Chapter 7 (pg.) Chapter 8 (pg.) Chapter 9 (pg.) Assignments: Activities, Labs & Test Description Goals 5 4 0 Computer Bonding A1 Ionic Bonding WS (Goals 1-5) Molecular Geometry Modeling Intermolecular Forces A2 A3 A4 A5 A6 A7 Covalent Bonding WS (Goals 5-6) VSEPR Geometry WS (Goals 7-9) Bond Types and Intermolecular Forces WS (Goals 7-9) Chemical Names and Formulas WS (Goals 10-12) Ionic Compounds WS (Goals 13-15) Molecular Compounds WS (Goal 16) Chemical Names and Formulas Unit 4 Test Late Lab Stamp (this stamp means you are not qualified to do lab and test corrections) Key Terms: molecule, molecular compounds ions, cation, anion, ionic compound, chemical formula, formula unit, molecular formula, monatomic ions, polyatomic ions, binary compounds, ternary compounds, valence electron, ion, electron dot structure, octet rule, ionic bond, stability, metallic bonds, ionic compound, single covalent bond, structural formula, unshared pairs, double covalent bond, triple covalent bond, VSEPR theory, non-polar covalent bond, polar covalent bond, dipole, van der Waals forces, dispersion forces, dipole interactions, hydrogen bonds Demo s: Metal vs. Non-metal vs. Metalloids, Electrolytes with light bulbs, Boiling/Melting salt, sugar, coconut oil, ethanol, Intermolecular attractive forces with posters, Fe + O 2, Na in water
Unit 4: Chemical bonding, names and formulas Ch. 7, 8 & 9 NAME Period: Chapter 7: Ions a. Define Ion i. Cations are because they have electrons. form these positive ions b/c they have electronegativity. ii. Anions are because they have electrons. form these negative ions b/c they have electronegativity. iii. In advertising these are called. iv. Charges are written as w/ a symbol followed by a i.e. v. Remember, atoms are. b. How & Why do ions form? They form by or electrons. This is done because all atoms really want full electrons. i. only have a few valence electrons, so they them to expose their full inner shell. ii. are already almost, so they them. iii. Both want to be like the in group. Called the rule. c. Predicting Ion Charges: Look at the e s. Decide how many electrons they want to or in order to have a valence shell. d. Electron Dot Structures: These simple models only show as dots. e. Octet Rule: when forming compounds, atoms tend to seek electrons like which is typically Copy Figure 7.1 here
Draw sodium and chlorine forming ions, and an ionic bond. Q: What type of bond have we formed in the above compound? Called a i. How do we name positive ions? ii. How do we name negative ions? 7.2 Ionic Bonds and Ionic Compounds A. Modeling Ionic Compounds 1. Show the formation of an ionic bond between magnesium and fluorine 2. Show the formation of an ionic bond between sodium and nitrogen 3. Show the formation of an ionic bond between aluminum and sulfur B. Formula Units i. An ionic compound is represented by a. This is the smallest number ratio of to in the compound. This is not a single physical unit, but a repeating grid. Overall the formula must be neutral. A Trick is called crossing charges to find the Formula Unit. Try Magnesium and Chlorine Try Aluminum and Bromine
C. Properties of ionic Compounds: ii. Structure with a repeating pattern of & iii. Solids at. iv. Conduct when dissolved in water 7.3 Metallic Bonds A. Metals bond because there are forces of between the electrons and the charged metal. Metals are described as having a of electrons. 1. This explains whey they conduct electricity. Electrons added to metals are able to among the. 2. This also explain why are so easy to make, since all metals have similar. Alloys are made of. and are great examples. Chapter 8: Covalent Bonding 8.1 - Molecules and Molecular Compounds: a. What type of bond exists between the atoms in a molecule? i. The prefix co- means. Valence means, so means to b. Define molecule c. What type of elements are chemically bonded together in a molecule? d. Create a list of physical properties that molecules share: Page 225 top paragraph e. are the only elements that exist in nature as atoms. These are called, which means they exist as atom. f. There are also seven common elements that exist in nature as two atoms bonded into as a. They make a seven on the periodic table, they are: g. Molecules are electrically. 8.2 The Nature of Covalent Bonding Using Lewis Dot Structures to draw covalent bond. 1. Show the covalent bond(s) in a Fluorine molecule, F 2,
2. Show the covalent bond(s) in an ammonia molecule, NH 3, 3. Show the covalent bond(s) in an Methane molecule, CH 4 A. Define Single covalent bonds: atoms held together by one of shared electrons. B. Define double covalent bonds: atoms held together by two of shared electrons. C. Define triple covalent bonds: atoms held together by three of shared electrons. D. Define an Unshared Pair: A pair of electrons that is shared between atoms. 4. Show the covalent bond(s) in an oxygen molecule, O 2 5. Show the covalent bond(s) in a nitrogen molecule, N 2, 6. Show the covalent bond(s) in a silicon dioxide molecule, SiO 2, 7. Show the covalent bond(s) in a dinitrogen dihydrogen molecule, N 2 H 2, E. Define coordinate covalent Bond: One atom contributes bonding electrons: Try carbon monoxide, CO
F. Define Polyatomic ion: a tightly bound group of atoms that collectively carry a. 8. Hydroxide ion, OH -1 (remember to add or remove electrons as needed) 9. Show the covalent bond(s) Sulfate ion, SO 4-2 10. Show the covalent bond(s) Sulfite ion, SO 3-2 11. Show the covalent bond(s) Ammonium ion, NH 4 +1 8.3 Bonding Theories 1. VSEPR Theory. dimensional models of molecules depend on a. The of covalent bonds b. The types (,, ) of bonds c. The number of electrons on the atom. d. and bonds act like bond each. e. Some molecules have multiple atoms
Shapes of Molecules # of Central Bond areas unshared pairs of electrons on central atom? Diagram Example Polar or Non- Polar? Tetrahedral Trigonal Planer Linear Triatomic Pyramidal Bent Linear Diatomic 8.4 Polar Bonds a. Non-Polar Covalent: The elements have very electronegativity so they share. Look at H 2 b. Non Polar Covalent: Any perfectly molecule is also because there is not possibility of having a region of one vs a region of the opposite. Try CO 2
Copy Figure 8.4 (pg. 238) Electronegativity Differences and Bond Types Electronegativity Difference Range Most probably type of bond Example c. Polar-Covalent: The elements in the bond have electronegative differences between. HCl H 2 O 1. Dipole = a molecule with poles, or opposite is called 8.4 Intermolecular attractions and Polarity A. Intermolecular Forces: Forces that hold together. These determine,, and states. 1. : forces between molecules. A. Forces: Weakest, caused by motion. More electrons = More = More B. : Polar molecules in which attract. 1. : Dipole forces that are strong b/c Hydrogen is bonded to very electronegative atoms (,,or )
Chapter 9: Chemical Names and Formulas 9.1 Naming Ions A. ions: Ions made of single. B. Elements: There is a pattern in predicting how many electrons are lost and gained for the representative elements, can you guess it? Metals Non - Metals B. metals: We cannot determine how many electrons are lost for the metals b/c their in their valence electrons can. Remember, Transition metals are any metal with orbitals. They include all of the groups, and the metals under the. is the exception because it does not have d orbitals. Aluminum is also considered a metal even though it the ladder. Formula: 1. How can we tell someone how many electrons lost for the transition metals? Name: Examples: Name the following transition metal ions: For help, refer to page 144, table 6.3 a. Lead (lost 2 electrons) b. Lead (lost 4 electrons) c. Vanadium (lost 3 electrons) d. Vanadium (lost 2 electrons) C. There are 3 exceptions to this rule: 1. Silver is always a 2. Zinc and Cadmium are always a *Place their charges in their boxes on your periodic table so you don t forget this rule. 3. DO NOT USE A ROMAN NUMERAL WHEN NAMING SILVER, ZINC AND CADMIUM IONS. 4. ALWAYS USE A ROMAN NUMERAL WHEN NAMING ANY OTHER TRANSITION METAL ION.
Write the symbol and charge of the following elements. a. Oxygen O -2 b. Chromium (II) c. Barium d. Neon e. Chlorine f. copper (II) g. Nitrogen h. Zinc i. Potassium Name the ion Cation or Anion? D. Ions: Ions made of. 1. What endings to polyatomic ions receive when naming them? 2. There are 3 important exceptions, they are: 3. Use chart 6.4 (or back of your periodic table) to write formulas of the following polyatomic ions: a. ammonium ion b. sulfate ion c. sulfite ion d. carbonate ion e. nitrate ion 9.2 Naming and Writing Formulas for Ionic Compounds A. Ionic Compounds 1. What are Binary Ionic Compounds? f. permanganate ion g. hypochlorite ion h. phosphate ion i. cyanide ion j. hydroxide ion 2. What are the rules for writing Binary Ionic Compounds? a. Write the (positive) ion first b. Write the (negative) ion last c. The net charge for the compound must add to (positives + negatives = 0) d. Use to indicate how many of each ion you need to balance the charge. *Trick: Crossing charges. Cross the numbers of the charges down as subscripts removing the +/-. Reduce if possible. Ionic Compound Practice: Write the name &formula for the compound formed b/w magnesium and chlorine. Write the name &formula for the compound formed b/w sodium and oxygen. Write the name &formula for the compound formed b/w aluminum and sulfur.
Write the name &formula for the compound formed b/w iron (III) and oxygen. Write the name &formula for the compound formed b/w calcium and sulfur. PRACTICE PROBLEMS: 1. Write the formulas for the compounds formed between these pairs of ions. NAME THEM. a. Be +2, O -2 b. Na +1, N -3 c. Mg +2, P -3 d. Zr +2, Cl -1 2. Write formulas for these compounds. a. silver chloride d. lead (IV) sulfide b. Gallium (III) phosphide e. barium oxide c. potassium nitride f. magnesium carbide 3. Write names for these binary ionic compounds a. ZnS b. KCl c. BaO d. CuBr 2 e. CuO (careful!) f. Ag 2 S g. Al 2 Se 3 B. Writing formulas for Compounds with Compounds A. Remember, they are still just two ions, and all rules from before still apply! B. Write the formula for Sodium Cyanide, a ternary compound: C. Sometimes, we need to take more than one polyatomic ion to balance the charge to 0. If this happens, place the polyatomic ion in parenthesis and the subscript outside of the parentheses. 1. Write the formula for Sodium sulfate: 2. Write the formula for calcium acetate: 3. Write the formula for ammonium phosphite: 4. Write the formula for aluminum dichromate: Practice Problems: 1. Write the name & formulas for ionic compounds formed from these pairs of ions: a. NH 4 +1, CO 3 2- b. Lithium ion, phosphite ion c. Al 3+, SO 4-2
d. Strontium ion, silicate ion 2. Write formulas for these compounds a. Berylium dihydrogen phosphate b. chromium (IV) nitrate c. mercury (IV) oxide d. ammonium oxalate 3. Name these compounds: a. Al(OH) 3 b. LiCN c. InClO 3 d. Hg(SO 4 ) 2 e. Mg(CN) 2 f. (NH 4 ) 2 CO 3 g. Fe(ClO 4 ) 2 h. Ni(H 2 PO4) 3 9.3 Naming Molecular Compounds 1. We use prefixes when naming binary molecular compounds: Refer to Table 6.5 and fill in the following: Prefix Number Prefix Number 1 6 2 7 3 8 4 9 5 10 a. How do we name binary molecular compound?? b. Say the name of the first element, say the name of the second element, ending in IDE, and put the appropriate prefix in to indicate how many of each element there are in the formula: **If the prefix for the first element in a binary molecular compound is, it may be dropped. However, it must be said if it is for the second element. **Don t reduce the subscripts (like you did for binary ionic compounds) Name these binary molecular compounds Write formulas for the following molecules a. H 2 O b. N 2 Cl 4 c. SeCl 5 d. CO e. C 3 F 8 f. S 6 O 9 a. dinitrogen tetrafluoride b. sulfur hexachloride c. trinitrogen octoxide d. dihydrogen dioxide e. nonanitrogen pentasulfide