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Unit 5: Bonding, Naming, Formulas, Molecular Geometry, VESPER theory

Unit 5: Bonding, Naming, Formulas, Molecular Geometry, VESPER theory. Standards:. C.7.A name ionic compounds containing main group transition metals, covalent compounds, acids, and bases, using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules

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Unit 5: Bonding, Naming, Formulas, Molecular Geometry, VESPER theory

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  1. Unit 5: Bonding, Naming, Formulas, Molecular Geometry, VESPER theory

  2. Standards: • C.7.A name ionic compounds containing main group transition metals, covalent compounds, acids, and bases, using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules • C.7.B write the chemical formulas of common polyatomic ions, ionic compounds containing main group or transition metals, covalent compounds, acids • C.7.C construct electron dot formulas to illustrate ionic and covalent bonds • C.7.D describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability, and ductility • C.7.E predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using Valence Shell Electron Pair Repulsion (VSEPR) theory

  3. Power Point Index: Naming Chemical Compounds… slide 4 Naming Acids and Bases… slide 11 Writing Chemical Formulas… slide 16 Chemical Bonding… slide 22 Molecular Geometry, VSEPR Theory… slide 31

  4. Naming Chemical Compounds Notes C.7.A name ionic compounds containing main group transition metals, covalent compounds, using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules.

  5. Types of Compounds • There are three main types of compounds when working on Naming Compounds. • Metal Binary Compounds – Contain a Metal and a Non-Metal. They form an Ionic Bond. • Non-Metal Binary Compounds – Contain two Non-Metals. They form a Covalent Bond. • Ternary Compounds – Contain Polyatomic ions. The formula will have three or more elements in it.

  6. Metal Binary Compounds • Name the first element. (This will always be the metal.) • Replace the ending on the second element with an “ide” ending. ( This element will be the non-metal) • Example: • NaCl Sodium and Chlorine becomes • Sodium Chloride • MgS Magnesium and Sulfur becomes • Magnesium Sulfide

  7. Naming Compounds with a Transition metal • When some atoms can have more than one possible charge, you name the charge on the atom. • Copper +1 and +2 Iron +2 and +3 Cu +1 is Copper I Fe +2 is Iron II Cu +2 is Copper II Fe +3 is Iron III CuCl is Copper I Chloride FeCl2 is Iron II Chloride CuCl2 is Copper II Chloride FeCl3 is Iron III Chloride

  8. Non-Metal Binary Compounds • Name the first element • Replace the ending on the second element with “ide” • Use Prefixes for the number of atoms in the formula. • CO2 Carbon and Oxygen • Monocarbon Dioxide • N2O Nitrogen and Oxygen • Dinitrogen Monoxide

  9. Pre-fixes • 1 atom = Mono • 2 atoms = Di • 3 atoms = Tri • 4 atoms = Tetra • 5 atoms = Pent • 6 atoms = Hex • 7 atoms = Hept • 8 atoms = Oct • 9 atoms = Non • 10 atoms = Deca

  10. Ternary Compounds: Compounds with Polyatomic Ions • Name the first part of the compound. Element or Polyatomic ion. • Name the second part of the compound. Element or Polyatomic ion. • Example: • MgSO4 NH4OH • Magnesium Sulfate Ammonium Hydroxide • K3PO4 • Potassium Phosphate

  11. Naming Acids and Bases C.7.A name, acids using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules

  12. Naming Acids without Oxygen • Acids without Oxygen are named with the prefix “Hydro” and end in “ic” • Examples: • HCl Hydrochloric Acid • HF Hydrofluoric Acid • HBr Hydrobromic Acid

  13. Naming Acids with Oxygen • For some acids with oxygen have several forms, there are prefixes used with the regular “ic” and “ous” endings. • The “ic” or regular ending for an acid comes from the polyatomic ion with the “ate” ending. This gives the regular count for the oxygen for this type of acid. • Example: • H2SO4 • SO4 is Sulfate so this acid is called Sulfuric Acid

  14. Once you know the “ic” ending you count the number of oxygen in the other forms to find the name for the acid. (REMEMBER: The regular “ic” form comes from the polyatomic ion that ends with “ate”) • Two less oxygen Hypo ________ “ous” Acid • One less oxygen ________ “ous” Acid • Regular “ic” form ________ “ic” Acid • One more oxygen Per ________ “ic” Acid

  15. The other names for the acids will come from the count based from the “regular acid name” • H2SO4 “ate” ending so it is Sulfuric Acid • H2SO3 “ite” ending so it is Sulfurous Acid • H2SO2 two less oxygen will have a prefix and “ous”ending. Hyposulfurous Acid. • H2SO5 one more oxygen will have a prefix “Per” and the regular “ic” ending. Persulfuric Acid

  16. Writing Chemical Formulas C.7.B write the chemical formulas of common polyatomic ions, ionic compounds containing main group or transition metals, covalent compounds, acids

  17. Writing Formulas: Ionic Compounds • Write chemical symbol for each part of the compound. • Write the charge for the element. Do the charges add together and equal zero? • Yes, Stop this is the formula. The number of electrons given away is the same as what is being taken by the second atom. • No, Cross the absolute value of the charge to the opposite element as a subscript. Multiply the new subscript by the charge and see if the new values will add together and equal zero. If yes, Stop you have the formula

  18. Potassium Bromide Formula K +1 Br -1 +1 + -1 = 0 Yes KBr • Magnesium Chloride Mg +2 Cl -1 +2 + -1 = +1 No Mg 1 Cl 2 Mg (1 x +2)= +2 Cl (2 x -1)= -2 Yes MgCl2

  19. Transition Elements • Same rules as normal ionic compounds. The charge for the transition metal will come from the name of the compound. • Iron III Chloride • Fe +3 Cl -1 +3 + -1 = +2 No • Fe1 Cl 3 Fe (1 x +3) +3 Cl (3 x -1) -3 Yes FeCl3

  20. Polyatomic Ions • The rules for polyatomic ions will be the same as ionic compounds. Place the polyatomic ion in parenthesis. • Keep the parenthesis at the end of the process if you have a number greater than one outside of the parenthesis. If you did not cross a number or if you only crossed a one do not keep the parenthesis. • Magnesium Sulfate • Mg +2 (SO4) -2 Yes MgSO4 • Iron III Phosphate • Fe +3 (PO4) -3 Yes FePO4

  21. Sodium Hydroxide • Na +1 (OH) -1 Yes NaOH Do not keep the parenthesis because there is no number crossed. • Calcium Hydroxide • Ca +2 (OH) -1 Ca 1 (OH)2 Ca (1 x +2) +2 (OH)(2 x -1) -2 Yes Ca(OH)2 Keep the parenthesis because there is a number greater than one outside the parenthesis

  22. Chemical Bonding Notes C.7.C construct electron dot formulas to illustrate ionic and covalent bonds

  23. There are three main types of Chemical bonding. Ionic, Covalent, and Metallic. • Ionic Bonding occurs when there is a transfer of electrons. • Covalent Bonding occurs when atoms share electrons. • Metallic Bonding consist of the attraction of free floating valance electrons for positively charged metal ions.

  24. Electro negativities are used to determine what type of bond is formed when atoms come together in a chemical reaction. • To find the type of bond find the difference in the electro negativities. • If the difference is greater than 1.67 an ionic bond is formed. • If the difference is less than 1.67 a covalent bond is formed.

  25. All atoms want to obtain eight electrons in the valence energy level. To do so they will give, take, or share electrons.

  26. NaCl Sodium Chloride • Sodium: (1.01) Chlorine: (2.83) • Na: 1s22s22p63s1Cl: 1s22s22p63s23p5 • Sodium transfers the 3s1 to Chlorine to complete the 3p energy level. • The electronegativity difference is 1.72 • An ionic bond is formed.

  27. Rules for Ionic Bonds • The element with the fewest atoms goes in the center. • The other atoms go around the central atom. • Show the transfer of the electrons with a positive for the atom that lost the electrons and a negative for the atoms that gain the electrons.

  28. AsI3 Arsenic Triiodide Arsenic (2.20) Iodine (2.21) As: 1s22s22p63s23p64s23d104p3 I: 1s22s22p63s23p64s23d104p65s24d105p64d105p5 The electro negativity difference is .01 A covalent bond is formed. The atoms share the electrons.

  29. Rules for showing Covalent Bonds • The element with the fewest atoms goes in the center. • The other elements go around the central atom. • A bonding pair can only form where there is an unpaired electron. • Shared pairs or bonding pairs are shown with a dash. One dash equals two electrons.

  30. Molecular Geometry, VSEPR Theory C.7.E predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using Valence Shell Electron Pair Repulsion (VSEPR) theory

  31. Molecular Geometry • The shape that a covalently bonded substance will take is referred to as its Molecular Geometry. • The shape is determined by the central atom, and the number of shared and unshared electron pairs around the atom. • Electron pairs around the central atom will spread out as far as possible to minimize the repulsive forces. • This gives bond angles depending on the shape.

  32. Linear Trigonal Planar Tetrahedral Trigonal Pyramidal Bent

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