2. Types of Compounds Metal Binary Compounds – Metal and Non-Metal, forms an Ionic Bond. Non-Metal Binary Compounds – two Non-Metals, forms a Covalent Bond. Ternary Compounds – Contain Polyatomic ions. The formula will have three or more elements in it.

3. Metal Binary Compounds Name the first element- a metal. Replace the ending on second element (non-metal) with an “-ide” ending. Examples: NaCl Sodium + Chlorine = Sodium Chloride MgS Magnesium + Sulfur = Magnesium Sulfide

4. Naming Compounds with a Transition metal When some atoms can have more than one possible charge, you name the charge on the atom. Copper is +1 or +2Iron is +2 or +3 Cu +1 is Copper I Fe +2 is Iron II Cu +2 is Copper II Fe +3 is Iron III CuCl = Copper (I) Chloride FeCl 2 = Iron (II) Chloride CuCl 2 = Copper (II) Chloride FeCl 3 = Iron (III) Chloride

5. 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. CO 2 is Carbon + Oxygen = Monocarbon Dioxide N 2 O is Nitrogen + Oxygen = Dinitrogen Monoxide H 2 O is Hydrogen + Oxygen = Dihydrogen Monoxide

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

7. Ternary Compounds Compounds with Polyatomic Ions Name the first part of the compound- an element or polyatomic ion. Name the second part of the compound- an element or polyatomic ion. Examples: MgSO 4 NH 4 OH Magnesium SulfateAmmonium Hydroxide K 3 PO 4 Potassium Phosphate

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

9. 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

10. 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 1 (+2) + 2 (-1) = 0Yes MgCl 2

11. 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 Fe 1 Cl 3 1 (+3) + 3 (-1) = 0 YesFeCl 3

12. 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

13. 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 1 (+2) + 2 (-1) = 0 Yes Ca(OH) 2 Keep the parenthesis because there is a number greater than one outside the parenthesis

14. C.7.C construct electron dot formulas to illustrate ionic and covalent bonds

15. 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.

16. Electronegativities 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.

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

19. NaCl Sodium Chloride Sodium: (1.01) Chlorine: (2.83) Na: 1s 2 2s 2 2p 6 3s 1 Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 Sodium transfers the 3s 1 to Chlorine to complete the 3p energy level. The electronegativity difference is 1.72 An ionic bond is formed.

20. 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.

21. AsI 3 Arsenic Triiodide Arsenic (2.20) Iodine (2.21) As: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 I: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 The electronegativity difference is.01 A covalent bond is formed. The atoms share the electrons.

22. 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.

23. 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

24. 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.

25. Linear molecule Total number of electron pairs. Number of shared pairs Number of unshared pairs ShapeBond Angle 220 Linear 180 0

26. Trigonal planar molecule Total number of electron pairs. Number of shared pairs Number of unshared pairs ShapeBond Angle 330 Trigonal Planar 120 0

27. Tetrahedral molecule Total number of electron pairs. Number of shared pairs Number of unshared pairs Shape Bond Angle 440Tetrahedral109.5 0

28. Trigonal Pyramidal Total number of electron pairs. Number of shared pairs Number of unshared pairs ShapeBond Angle 43 1 Trigonal Pyramidal107.3 0

29. Bent molecule Total number of electron pairs. Number of shared pairs Number of unshared pairs ShapeBond Angle 422Bent104.5 0

30. Linear Trigonal Planar TetrahedralTrigonal Pyramidal Bent

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

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

33. Naming Acids with Oxygen Some acids with oxygen have several forms and use suffixes with “-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

34. Once you know the “-ic” ending, count the number of oxygens 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

35. 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