1. Physical Science 11/27/12

2. Bonding Atoms  A compound is made of two or more elements that are chemically combined. Held together by chemical bonds  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

3. Chemical Bonds  3 types ionic, covalent, metallic  determines the structure of the compound  structure affects properties of the compound melting/boiling pts, conductivity etc.

4.  Chemical Structure The arrangement of atoms in a substance  Molecular Models Models used to show different aspects of chemical structures Some models represent bond length and angles. bond length: the average distance between the nuclei of two bonded atoms bond angles: the angle formed by two bonds to the same atom

5. Molecular Models of Compounds  Ball and stick atoms are represented by balls bonds are represented by sticks good for “seeing” angles  Structural formula chemical symbols represents atoms lines are used to represent bonds good for “seeing” angles

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

8.  Ionic bonds are formed by the attraction between oppositely charged ions (a metal and a nonmetal) cation: positive ion: lost electrons ○ Formed by metals anion: negative ion: gained electrons ○ Formed by nonmetals

9. ex. Na + + Cl - = NaCl Each positive sodium ion attracts many negative chloride ions. These negative chloride ions attract many more positive sodium ions, and so on.

10.  Ionic bonds form networks instead of molecules repeating pattern of multiple ions Chemists refer to the smallest ratio of ions in ionic compounds. The ratio of sodium to chloride ions in sodium chloride is 1:1. There is one sodium ion for every one chloride ion. One Na + and one Cl - form what is known as a formula unit.

11. Properties of Ionic Compounds  Structure affects properties strong attraction between ions creates a rigid framework, or lattice structure: aka: crystals ex, cubes, hexagons, tetragons high melting/boiling pt shatter when struck (think of it as one unit) Solids at room temperature conductivity ○ solid: ions are close together and in fixed positions (can’t move) - NO conductivity ○ Dissolved in liquid: ions are freely moving due to a broken lattice structure - Good conductivity

12.  Covalent bonds chemical bond in which two atoms share a pair of valence electrons Occurs between two atoms with similar electronegativities (neither atom is “strong” enough to take the electron away from the other) ○ Electronegativity = the power of an atom in a molecule to attract electrons to itself can be a single, double, or triple bond single: share 2 e-’s double: share 4 e-’s triple: share 6 e-’s usually formed between nonmetals mostly low melting/boiling points most do not conduct electricity because they are not charged Liquid or gas at room temperature

13.  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 Covalent Bonds IONIC AND COVALENT BONDS ANIMATION

14. Example: H 2 O Example: H 2

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

16. Polyatomic Ions  An ion made of two or more atoms  Charged molecules (have an overall negative or positive charge as a group)  Act as a single unit in a compound  Basic building blocks of many ionic compounds

17. Polyatomic ion examples CO 3 2- carbonate HCO 3 - hydrogen carbonate (bicarbonate) NO 3 - nitrateitrateitrate NO 2 - nitrite SO 4 2- sulfate SO 3 2- sulfite PO 4 3- phosphate HPO 4 2- hydrogen phosphate OH - hydroxide O 2 2- peroxide

18. Parenthesis are used to group atoms of a polyatomic ion. Copper (II) chlorate: Cu(ClO 3 ) 2 Iron (II) nitrate: Fe(NO 3 ) 2 No parenthesis are needed if there is not more than one of the polyatomic ion present, such as in NaOH.

19. 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

20. 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+

21. 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+

22. 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

23. 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

24. 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

25. Polyatomic Ions Cont.  Common Polyatomic Ions

26. Objectives  Name ionic compounds from formulas  Determine the chemical formulas for ionic compounds from compound name

27. 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

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

29. 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**

30. 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

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

32. 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

33. 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

34. Writing Formulas for Covalent Compunds  Writing formulas from names 1. Carbon Dioxide CO 2 2. Dinitrogen Pentoxide N 2 O 5 3.Triphosphorus monosulfide P 3 S 4.Sulfur Monobromide SBr