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 Valence Electrons: › Electrons in the highest occupied energy level › Determines the chemical properties › = to the group number (groups 1,2, 13-18.

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Presentation on theme: " Valence Electrons: › Electrons in the highest occupied energy level › Determines the chemical properties › = to the group number (groups 1,2, 13-18."— Presentation transcript:

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2  Valence Electrons: › Electrons in the highest occupied energy level › Determines the chemical properties › = to the group number (groups 1,2, minus 10)

3  Electron dot structures (Lewis structures): › Shows valence electrons as dots Ex. Ca

4  Octet Rule: › In compounds – atoms want noble gas (octet) configuration › Metallic atoms – lose electron(s) = noble gas (octet) configuration from next lowest energy level  ex. Na + = Ne configuration › Non-metallic atoms – gain or share electron(s) = noble gas (octet) configuration for that energy level  ex. F - = Ne configuration

5  Exceptions – transition metals achieve pseudo noble gas configurations (18 electrons in outer energy level) › ex. Ag = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 4d 10 Ag + = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 4d 10

6  Formation of Ions: › Neutral atom loses or gains electrons › Produces cations (+) charge or anions (-) charge › Monotomic Ions- composed of 1 atom with a +/- charge › Polyatomic ions- composed of 2 or atoms with a +/- charge

7  Formation of Cations: › + charge › Lose electron(s) › Mostly metals

8  Examples: Ionization of Na: Na loss of valence electron Na + + e - Ionization Na 1s 2 2s 2 2p 6 3s 1 Na + 1s 2 2s 2 2p 6 Ionization of Mg: Mg loss of 2 valence electron Mg e - Ionization Mg 1s 2 2s 2 2p 6 3s 2 Mg +2 1s 2 2s 2 2p 6

9  Formation of Anions: › - charge › Gain electron(s) › Mostly non-metals

10  Examples: Cl + e - gain one valence electron Cl -1 chlorine atom chloride ion Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Cl -1 1s 2 2s 2 2p 6 3s 2 3p 6 O t 2e - gain 2 valence electron O -2 oxygen atom oxide ion O 1s 2 2s 2 2p 4 O -2 1s 2 2s 2 2p 6

11  Naming of Cations: › Name = element’s name + ion (ex. Na + = sodium ion) › Transition Metals -Stock system – use Roman numerals to ID the charge (ex. Fe +2 = iron(II) ion, Fe +3 = iron(III) ion)

12  Naming of Anions: › Name = ending of the element’s name dropped and ide added (ex. fluorine (F - ) = fluoride)

13  Oxidation number : › + or – # that indicates how many electrons an atom has gained, lost or shared to become stable › Some elements have more than one oxidation #

14  Rules for Assigning Oxidation #s: 1. Monatomic ions = + or – what it takes to gain the octet (1-2 group #, group # -10) = ion’s charge 2. Hydrogen = + 1. Exception = Metal hydrides (ex. LiH), = Oxygen = - 2 (exception- when combined with F then = + 2) 4. Neutral compounds = sum oxidation #s = 0 5. Polyatomic ions-the sum of the oxidation # = the charge of the ion.

15  Examples: +2-2=0 NO O = - 2, so N = + 2 (+2-2=0 rules: 1,3,4) +4 –2(2) = 0 NO 2 O = - 2(2), so N = + 4 (+4 – 2(2) = 0 rules: 1,3,4) +6-2(4)= - 2 SO 4 -2 O = - 2(4) S -2(4) = - 2  S - 8 = - 2 S = + 6 (+6-2(4)= - 2 rules: 1,3,5)

16  Assign oxidation numbers to each element in the following: H 2 S MgF 2 PO 4 -3

17  Polyatomic Ions: › Tightly bound groups of atoms behaving as a unit and carry a charge › Most – anions (except ammonium)

18  Naming Polyatomic Ions: › Those containing oxygen (oxyanions) in different ratios have different endings (ite or ate) depending on the # of oxygens presents:

19 › –ite = one less oxygen then ate  ex, chlorite- ClO 2 -1 ; chlorate- ClO 3 -1 › Hypo = one less oxygen then ite  Ex. Hypochlorite- ClO -1 › Per = one more oxygen then ate  Ex. Perchlorate- ClO 4 -1

20  Formation of Ionic Compounds: › Composed of cations and anions › Electrically neutral

21  Ionic Compounds Properties: › Crystalline solids › Metals/Nonmetals › High melting points › Good conductors of electricity when in water or melted › Soluble in water

22  Ionic Bonds: › Electrostatic force that holds ions together Ex. Na gives away its lone electron to chlorine (7 valence electrons) = both have noble (octet) configs in a 1-1 ratio

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25  Chemical formula: › Shows kinds and # s of atoms in the smallest representative unit › Subscript – indicates the # of atoms of each element in the compound = Oxidation #

26  Formula Units: › Lowest whole number ratio of ions in an ionic compound Ex. NaCl lowest whole number ratio = 1:1 Ex. MgCl 2 lowest whole number ratio = 1:2 (1 Mg +2 ion to 2 Cl -1 )

27  Naming Ionic Compounds: › Cation name first followed by anion name Ex. Cs 2 O = cesium oxide › Cations that form more than 1 oxidation # - include roman numeral in name Ex. CuO = copper(II) oxide

28  Writing Formulas for Ionic Compounds: › Cation symbol followed by anion symbol  Add subscripts  Assign Oxidation #s  Crisscross method- oxidation # (charge) crosses over to become the subscript

29 ex. Fe +3 O -2 Fe 2 O 3

30  Naming Polyatomic Compounds: › Cation name first followed by the polyatomic anion name (exception – ammonium ion = + 1 charge) Ex. NaClO = sodium hypochlorite › Cations that form more than 1 ion – include roman numeral Ex. Cr(NO 2 ) 3 = Chromium(III) nitrite

31  Writing Polyatomic Compounds: › Cation symbol followed by the polyatomic ion symbol › Add subscripts › ( )needed if more than 1 unit of the same polyatomic ion is needed Ex. Ca + 2 (NO 3 ) 2 -1 Ca (NO 3 ) 2

32  Bases: › Ionic compound that produces hydroxide ions (OH -1 ) when dissolved in water

33  Naming Bases: › Named same way ionic compounds › Anion =hydroxide ion (OH -1 ) Ex. NaOH = sodium hydroxide

34  Writing Formulas for Bases: › Same way as ionic compounds Ex. Iron(III) Hydroxide = Fe(OH) 3

35  Acids: › Ionic compound that contains 1 or more hydrogen atoms › Produces hydrogen ions (H + ) when dissolved in water

36  Naming Acids: › Consists of a cation (hydrogen ions) and an anion › H n X formula – H = hydrogen; n = # of Hydrogens; X= monatomic or polyatomic anion

37 Naming Acids Rules Anion endingExampleAcid nameExample -ideChloride, Cl -1 Hydro-(stem)-ic acidHydrochloric acid -iteSulfite, SO 3 -2 (stem)-ous acidSulfurous acid -ateNitrate, NO 3 -1 (stem)-ic acidNitric acid

38  Writing Formulas for Acids: › Use the naming rules in reverse: Ex. 1 Hydrochloric acid = HCl Ex. 2 Sulfurous acid = H 2 SO 3 Ex. 3 Nitric acid = HNO 3

39  Metallic bonds : › Metals made up of closely packed cations rather than neutral atoms › Valence electrons – “sea of electrons” › Metallic bonds:  Consist of the attraction of the free-floating valence electrons  Forces of attraction that hold metals together

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41  Metallic Properties: › Good conductors b/c electrons can flow freely in them – as electrons enter one end = # leave the other end › Ductile and malleable b/c of mobility of valance electrons- metal cations easily slide past one another like ball bearings immersed in oil when subjected to pressure.


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