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Periodicity. Metals, nonmetals, & metalloids Transition metals vs. representative elements.

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Presentation on theme: "Periodicity. Metals, nonmetals, & metalloids Transition metals vs. representative elements."— Presentation transcript:

1 Periodicity

2 Metals, nonmetals, & metalloids

3 Transition metals vs. representative elements

4 Groups of Elements  Alkali Metals  Alkaline Earth Metals  Transition Metals  Halogens  Noble Gases

5 Trends in the periodic table Groups (families)  They all have the same # of outer electrons= VALENCE ELECTRONS Periods  They have valence electrons in the same highest energy level

6 Periodic trends  Periodic = happens according to a repeating pattern  Periodic table has repeating pattern of valence e- configurations Leads to periodic trends:  Atomic Radius  Ion Formation and Ion size  Ionization Energy  Electronegativity

7 Periodic trend: Atomic Radius  Size of atoms depends on: Number of energy levels Number of protons  Down a column:  Across a row:

8 What is the trend across a period?

9 A closer look at the trend across a period.

10 What is the trend down a group?

11 Atomic radius increases down a group.  The number of energy levels increases down a group  Each subsequent energy level is further from the nucleus. n=1 n=2 n=3 n=4 n=5

12 Atomic radius decreases across a row  Across a row  valence electrons are in the same energy level (n=3)  Positive charge of nucleus is partially cancelled out by negative charge of core electrons =electron shielding  With each additional proton, there is a stronger force pulling the electrons closer to the nucleus. This results in a smaller atomic radius. Electron shielding in Na vs. Cl

13 Shielding

14 Which has the greater atomic size, C or O?

15 Which has the greater atomic size, Li or K?

16 Ionization energy (kj/mol) Recall emission lab….. Ionization energy – the energy required to remove an electron from an atom. (pg. 358)

17 Ionization Energy- what is the trend?

18 Ionization Energy Can you explain this trend? Click HERE for a tutorial.HERE

19 Ionic Radii  What happens to the size of the atomic radius when an electron is removed?  Is an anion bigger or smaller than its neutral atom? Why?

20 Radii of neutral and charged atoms in pm (10 -12 m) Purple models represent neutral atoms; Red=cations; Blue=anions

21 Check your understanding  Which atom would be larger, and why? N or O K or Rb Ne or Na  Which atom would be more difficult to turn into a cation (higher IE), and why? Li or Be Mg or Ca F or Ne

22 Electronegativity  Electronegativity – is the ability of an atom to ATTRACT electrons when the atom is in a compound. (pg 363)  Credit to Linus Pauling

23 Ionization & Ionic Compounds Unit 3 Chemistry 1 Spring 2012

24 What’s an ion?  Ion = Cation Anion  Atoms form ions when they are able to give away or accept electrons.  The most likely type of ion an atom will form is based on the number of valence electrons it starts with.

25 Comparing cations and anions Cations  Positive charge  Valence e- removed  Size vs. atom: Na Na + Cu Cu 2+ Anions  Negative charge  Valence e- added  Size vs. atom: O O - O 2-

26 Ionization-how does it happen?  How do ions form?  Will it be endothermic or exothermic?  What happens next?

27 Specific ions form for each element due to its electron configuration  Elements with electrons that completely fill the s and p orbitals of the highest energy level do not react.  Elements without “full” outer energy levels take, give up, or share electrons (chemical bonding).  Vocabulary: OCTET RULE ISOELECTRONIC

28 Cation formation  How many valence electrons does Na have?  Would Na give up its valence electron spontaneously?  Why does Na tend to form an Na + cation, rather than an Na 2+ (or any other) cation?  What about Mg? What about F?

29 Anion formation  How many valence electrons does F have?  Is it possible to add an extra electron into the valence energy level of F?  Why does F tend to form an F - anion, rather than an F 2- (or any other) anion?  What about O? What about Be?

30 Electron configurations  Write the electron configuration for each neutral atom. What ion will form for each? F O Li He Ne Cl

31 Symbols for ions:  Use the chemical symbol and the charge for each ion. Notice PT trends! Some common ions

32 Periodic relationship of simple ions Transition Metals?

33 Ionic bonding  Oppositely charged ions attract!  This forms a strong electrostatic attraction  Called an Ionic Bond.

34 Atoms will become charged before making ionic bonds  Ion charged atom loss or gain of electrons  Cations:  Anions: (fulfills octet or duet rule)  Use periodic table to predict the ion formed for O, Li, He, Ne, and Cl


36 Ionic bonds-form a crystal lattice  Ions form ionic crystals.  pattern of alternating + and – ions  An exothermic process.  Crystal structure extremely stable.  Click HERE for an animated tutorialHERE

37 Ionic Solids Salts are held together by ionic bonds

38 Ionic compounds  When cations and anions come together to form ionic bonds, an ionic compound is the result.  All compounds are neutral.  Ex: Na reacts with S

39 Ionic Compounds  Most ionic compounds are composed of a metal and a nonmetal. Example: Sodium Chloride Example: Magnesium Oxide  Are ionic compounds charged?  Ionic compounds are NEUTRAL (no charge), so the two ions together must cancel each other’s charge. Example: Mg +2 and O -2 or Na + and Cl - One of each ion, makes a neutral ratio.

40 Ionic Compounds  What if the ions are of different charges? Example: Sodium and Oxygen Na +1 and O -2  The ions exist in ratios, producing an electrically neutral compound. 2 sodium ions for every oxygen ion  How can we write this as a formula? Na 2 O This formula shows NO charges, it is neutral There are two sodium ions for every oxygen It represents a RATIO of ions.

41 Practice writing formulas for the following ionic compounds:  potassium fluoride  calcium bromide  aluminum chloride  sodium sulfide  beryllium oxide

42 Formulas of ionic compounds  potassium fluoride KF  calcium bromide CaBr 2  aluminum chloride AlCl 3  sodium sulfide Na 2 S  beryllium oxide BeO

43 Rules for naming ionic compounds  Salts (ionic compounds) are named: Cation first, anion second.  Cations are named the same as the neutral atom, i.e. sodium, aluminum… they almost always end in “-ium”.  Anions are named with the ending “-ide”. i.e. fluorine  fluoride; oxygen  oxide  Examples: sodium chloride, magnesium oxide

44 Ions of transition metals  Some transition metals can form more than one ion  Roman numerals indicate which it is.  Example: Name: Nickel (I) oxide: Ions: Ni +1 and O -2 Formula:Ni 2 O  Example: Name: Cobalt (II) Chloride Ions:Co +2 and Cl -1 Formula:CoCl 2

45 Roman numeral exceptions  Some transition metals only form one stable ion and do not require roman numerals: Ag + Zn +2  A few other metals have more than one stable ion and require roman numerals: tin(II) and tin(IV) lead(II) and lead(IV) Memorize these...

46 Practice writing names for ionic compounds of transition metals  ZnCl 2  PbBr 2  AuO

47 Names of ionic compounds  ZnCl 2 zinc chloride  PbBr 2 lead (II) bromide  AuO gold (II) oxide

48 Polyatomic ions: to memorize Polyatomic ion Name:  SO 4 -2 sulfate  NO 3 -1 nitrate  CO 3 -2 carbonate  HCO 3 -1 bicarbonate  NH 4 +1 ammonium  OH -1 hydroxide  PO 4 -3 phosphate  CH3COO -1 (C 2 H 3 O 2 -1 ) acetate

49 Formulas for ionic compounds:  Iron(II) phosphate  Zinc carbonate  Aluminum nitride

50 Formulas for ionic compounds:  Iron(II) phosphateFe 3 (PO 4 ) 2  Zinc carbonateZnCO 3  Aluminum nitrideAlN

51 Naming ionic compounds:  MgSO 4  Cu(OH) 2  NH 4 Cl

52 Naming ionic compounds:  MgSO 4 magnesium sulfate  Cu(OH) 2 copper(II) hydroxide  NH 4 Clammonium chloride

53 Introducing dimensional analysis.  Suppose you work at a hardware store and a customer comes in asking to buy 150 bolts from a bulk bin. What would be the most efficient way to get the customer his 150 bolts.

54 Counting large numbers  Atoms are too small to count one at a time!  The “mole” is a useful unit for counting atoms or other small particles.  1 mole = 6.022 x 10 23 objects

55 Compare and contrast  Compare 5 dozen bagels to 5 dozen elephants How are they the same? How are they different?  Compare 2.5 moles of hydrogen atoms to 2.5 moles of gold atoms How are they the same? How are they different?

56 Avogadro’s Number  Ratio between grams and a.m.u.  Therefore, atomic masses on PT are also the molar masses of elements (in g/mol)

57 Counting by weighing  We know the mass of 1 mole of any element AND  We know how many atoms are in 1 mole of any element  THEREFORE we can “count” a certain number of atoms using mass.

58 Molar mass of compounds  To calculate the mass of one mole of a compound, add the molar masses of the atoms  EX: 1 mole of Al 2 O 3 has 2 mol Al + 3 mol O  What is the molar mass of: Li 3 P NaOH Mg(NO 3 ) 2

59 Mole ratios for compounds  One mole of CaCl 2 has mole(s) of Ca 2+ ions mole(s) of Cl - ions  One mole of Al 2 (CO 3 ) 3 has mole(s) of Al 3+ ions mole(s) of CO 3 2- ions

60 Examples:  What is the mass of 1 mole of calcium atoms? What is the mass of 2 moles of calcium atoms? What about 5.3 moles?  How many atoms are there in 2.25 moles of sulfur? What is the mass of this sample?

61 Conversion blueprint Atoms (particles) MolesMass (g)

62 Calculations with ionic compounds  How many moles in 5000 kg of iron(III) oxide?  What is the mass of 2.50 mol of calcium chloride?  How many potassium ions are there in 4.5 moles of potassium oxide?  How many formula units are there in 285 g of copper(II) oxide?

63 Test Preparation:  Start the review sheet…but don’t end there.  Use practice problems from the book at the end of the chapter.  Power point  DO PRACTICE PROBLEMS! Do the worksheets or look over labs again.  Be able to explain ideas completely. Cause and effect.

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