1. staff.tuhsd.k12.az.us/ldompier/ChmPhyPP%5CCh.%206.ppt

2. TEKS  7 (A) name ionic compounds containing main group or transition metals, covalent compounds, acids, and bases, using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules;  7 (B) write the chemical formulas of common polyatomic ions, ionic compounds containing main group or transition metals, covalent compounds, acids, and bases;  7 (C) construct electron dot formulas to illustrate ionic and covalent bonds;  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;

3. Objectives  Explain why atoms form bonds  Define chemical bond & name three types of chemical bonds  Compare and contrast the advantages and disadvantages of varying molecular models

4. Bonding Atoms  Why do atoms bond? - each atom wants a full outermost energy level - gain, lose, and share valence electrons to achieve the duet or octet rule aka: “being happy” - gives each atom an electron configuration similar to that of a noble gas ex. Group 18: He, Ne, Ar

5. Chemical Bonds  Chemical Bonds - attractive force that holds atoms or ions together - 3 types ionic, covalent, metallic - determines the structure of compound - structure affects properties - melting/boiling pts, conductivity etc.

6. Chemical Structure/Models  Chemical Structure/Molecular Models - arrangement of bonded atoms or ions - bond length: the average distance between the nuclei of two bonded atoms - bond angles: the angle formed by two bonds to the same atom

7. Molecular Models of Compounds  Ball and stick - atoms are represented by balls - bonds are represented by sticks * good for “seeing” angles  Structural - chemical symbols represents atoms - lines are used to represent bonds * good for “seeing” anglesH H O

8. Molecular Models Cont.  Space filling - colored circles represent atoms, and the space they take up - no bonds, no bond angles  Electron Dot/Lewis Structure - chemical symbol represent atom - dots represent valence electrons - 2 center dots represent a bond - no bond angles, no bond length

9. Objectives  Describe how an ionic, covalent and metallic bonds forms  Relate the properties of ionic compounds to the structure of crystal lattices  Compare polar and non polar bonds, and demonstrate how polar bonds affect polarity of a molecule  Describe the structure and strength of bonds in metals & relate their properties to their structure

10. Ionic Bonds / Ionic Compounds  Definition - bond formed by the attraction between oppositely charged ions cation: positive: lost e-’s anion: negative: gained e-’s - oppositely charged ions attract each other and form an ionic bond ex. Na + + Cl - = NaCl - electrons are transferred from one atom to another - negative ions attract more positive ions, and soon a network is formed

11. Ionic Bonds Cont.

12. Networks / Crystal Lattices  Networks - repeating pattern of multiple ions ex. NaCl - every Na ion is next to 6 Cl ions - strong attraction between ions creates a rigid framework, or lattice structure: aka: crystals ex, cubes, hexagons, tetragons

13. Properties of Ionic Compounds  Structure affects properties - strong attractions between ions: strong bonds - high melting/boiling pt - shatter when struck (think of it as one unit) - conductivity solid: ions are so close together, fixed positions, (can’t move) NO conductivity liquid: ions are freely moving due to a broken lattice structure Good conductivity

14. Covalent Bonds  Definition - chemical bond in which two atoms share a pair of valence electrons - can be a single, double, or triple bond single, 2e-’s (-); double, 4e-’s (=); triple, 6e-’s( ) - always formed between nonmetals - mostly low melting/boiling points  2 types of bonds - polar - non polar

15. Covalent Bond Cont.  Non Polar - bonded atoms that share e - ’s equally - same atoms bonded ex. Cl – Cl: Cl 2  Polar - bonded atoms that do not share e - ’s equally - different atoms bonded H ex. H – N – H: NH 3

16. Covalent Bonds Cont.

17. Metallic Bonds  Definition - a bond formed by the attraction between positively charged metal ion (cation) and the shared electrons that surround it (sea of electrons) ex. Cu  Properties - Conductivity: Good: electrons can move freely - Malleable: lattice structure is flexible

18. Metallic Bonds Cont.

19. Predicting Bond Type

20. Objectives  Recognize monoatomic ions, metals with multiple ions and polyatomic ions  Name and determine chemical formulas for monoatomic ions, metals with multiple ion and polyatomic ions

21. Naming Ions  Monoatomic Ions - cation -name of element with ion ex. (Na) Sodium (Na+) Sodium ion - anion - name of element with the suffix –ide ex. (Br) Bromine (Br-) Bromide  Ions with multiple cations - transition metals - most form 2 +, 3 + and 4 + ex. Cu +, Cu 2+

22. Naming Metals with Multiple Ions  Transition Metals - form multiple ions - in order to name the ion use a roman numeral to indicate the charge ex. Cu 2+ : Copper (II), Titanium (III): Ti 3+ Practice Problems: Fe 3+ : Iron (III) Mercury (III): Hg 3+ Pb 4+ : Lead (IV)Chromium (II): Cr 2+

23. Polyatomic Ions  Definition - an ion made of one or more atoms that are covalently bonded and that act as a unit (atoms that have lost or gained electrons) ex. CO 3 2-, NH 4 + - behave the same as other ions - polyatomic ions can combined like any other ion (as a unit) ex. NH 4 NO 3 1:1 ratio (NH 4 ) 2 SO 4 2:1 ratio

24. Polyatomic Ions  Naming polyatomic ions - not logical - rules for some compounds  -ite & -ate endings - indicates the presence of oxygen - called oxyanions - if (-) does not specify how many oxygen atoms are present ex. Sulfate:4, Nitrate:3, Acetate:2

25. Polyatomic Ions Cont. - often several oxyanions differ only in the number of oxygen atoms present ex. Sulfur - ion with more oxygen takes the –ate ending ex. SO 4 - ion with less takes the –ite ending ex. SO 3  Common Oxyanions * Make sure you know these: memorize

26. Polyatomic Ions Cont.  Common Polyatomic Ions

27. Objectives  Name ionic compounds from formulas  Determine the chemical formulas for ionic compounds from compound name

28. Naming Ionic Compounds  Naming ionic compounds (binary) Formula to Name - name of cation followed by the name of the anion ex. NaCl: Sodium Chloride ZnO: Zinc (II) Oxide CuCl 2 : Copper (II) Chloride - formulas must indicate the relative number of cations and ions if transitional

29. Naming Ionic Compounds  Practice Problems MgBr 2 Magnesium Bromide KI Potassium Iodide CuCl 2 Copper (II) Chloride Fe 2 S 3 Iron (III) Sulfide

30. Formulas of Ionic Compounds  Writing formulas for ionic compounds Name to Formula - balance the cation charge and anion charge, leaving NO net charge - use subscripts to denote the number of atoms in the formula ex. NaCl: Na + Cl - : NaCl CaCl: Ca 2+ Cl - : CaCl 2 **1 to 1 ratios do not designate charge** **Criss-Cross charges into subscripts**

31. Practice Problems  Write the formula for the following atoms a.lithium oxide Li 2 O b.beryllium chloride BeCl 2 c.titanium (III) nitride TiN d.cobalt (III) hydroxide Co(OH) 3

32. Objectives  Name Covalent compounds from formulas  Determine the chemical formulas for covalent compounds from compound name

33. Naming Covalent Compounds  Prefix System # of atomsprefix 1mono 2di 3tri 4tetra 5penta 6hexa 7hepta 8octa 9nona 10deca

34. Naming Covalent Compounds Cont.  Rules for the prefix system 1. less electronegative element is given first. It is given a prefix only if it contributes more than one atom to a molecule of the compound 2. The second element is named by combining (a) a prefix indicating the number of atoms contributed by the atom (b) the root of the name of the second element, and (c) the ending –ide 3. The o or a at the end of a prefix is usually dropped when the word following the prefix begins with another vowel ex. Monoxide or pentoxide

35. Naming Covalent Compounds Cont. Naming covalent compounds from formula 1. SiO 2 Silicon dioxide 2. PBr 3 Phosphorus tribromide 3. CI 4 Carbon tetraiodide 4. N 2 O 3 Dinitrogen trioxide

36. Writing Formulas for Covalent Compunds  Writing formulas from names 1.Carbon Dioxide CO 2 2.Dinitrogen Pentoxide N2O5N2O5 3.Triphosphorus monosulfide P3SP3S 4.Sulfur Monobromide SBr