Presentation on theme: "Chapter 6 “The Periodic Table”"— Presentation transcript:
1 Chapter 6 “The Periodic Table” The Elements by Tom Lehrer
2 Organizing the Elements used properties of elements to sort into groups.1829 J. W. Dobereiner arranged elements into triads – groups of 3 w/ similar propertiesOne element in each triadhad properties intermediateof the other two elementsCl, Br, and I look different,but similar chemically
3 Mendeleev’s Periodic Table mid-1800s, about 70 elements knownDmitri Mendeleev – Russian chemist & teacherArranged elements byincreasing atomic mass
4 blanks for undiscovered elements Mendeleevblanks for undiscovered elementsWhen discovered, his predictions accurateProblems w/ orderCo to NiAr to KTe to I
5 A better arrangement1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number
7 Periodic LawWhen elements arranged in order of increasing atomic #, periodic repetition of phys & chem propsHorizontal rows = periods7 periodsVertical column = group (or family)Similar phys & chem prop.ID’ed by # & letter (IA, IIA)
8 Areas of periodic table 3 classes of elements:1) Metals: electrical conductors, have luster, ductile, malleable2) Nonmetals: generally brittle and non-lustrous, poor conductors of heat and electricitySome gases (O, N, Cl)some brittle solids (B, S)fuming red liquid (Br)
9 3) Metalloids: border the line-2 sides Properties are intermediate between metals and nonmetals
10 Section 6.2 Classifying the Elements OBJECTIVES:Describe the information in a periodic table.
11 Section 6.2 Classifying the Elements OBJECTIVES:Classify elements based on electron configuration.
12 Section 6.2 Classifying the Elements OBJECTIVES:Distinguish representative elements and transition metals.
14 Groups of elements - family names Group IA – alkali metalsForms “base” (or alkali) when reacting w/ H2O (not just dissolved!)Group 2A – alkaline earth metalsAlso form bases with H2O; don’t dissolve well, hence “earth metals”Group 7A – halogens“salt-forming”
15 Electron Configurations in Groups Elements sorted based on e- configurations:Noble gasesRepresentative elementsTransition metalsInner transition metalsLet’s now take a closer look at these.
16 Electron Configurations in Groups Noble gases in Group 8A (also called Group 18)very stable = don’t reacte- configuration w/ outer s & p sublevels full
17 Electron Configurations in Groups Representative Elements Groups 1A - 7Awide range of properties“Representative” of all elementss & p sublevels of highest energy level NOT filledGroup # equals # of e- in highest energy level
18 Electron Configurations in Groups Transition metals in “B” columnsouter s sublevel fullStart filling “d” sublevel“Transition” btwn metals & nonmetals
19 Electron Configurations in Groups Inner Transition Metals below main body of PT, in 2 horizontal rowsouter s sublevel fullStart filling “f” sublevelOnce called “rare-earth” elementsnot true b/c some abundant
20 Elements 1A-7A groups called representative elements outer s or p filling1A8A2A3A4A5A6A7A
21 The group B called transition elements These are called the inner transition elements, and they belong here
22 Group 1A called alkali metals (but NOT H) Group 2A called alkaline earth metalsH
23 Group 8A are noble gasesGroup 7A called halogens
25 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1H1Li3Na11K19Rb37Cs55Fr87Do you notice any similarity in these configurations of the alkali metals?
26 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6HeDo you notice any similarity in the configurations of the noble gases?2Ne10Ar18Kr36Xe54Rn86
27 Elements in the s - blocks Alkali metals end in s1Alkaline earth metals end in s2should include He, but…He has properties of noble gaseshas a full outer level of e-’sgroup 8A.
28 Transition Metals - d block Note the change in configuration.s1 d5s1 d10d1d2d3d5d6d7d8d10
35 Section 6.3 Periodic Trends OBJECTIVES:Describe trends among the elements for atomic size.
36 Section 6.3 Periodic Trends OBJECTIVES:Explain how ions form.
37 Section 6.3 Periodic Trends OBJECTIVES:Describe periodic trends for first ionization energy, ionic size, and electronegativity.
38 } Trends in Atomic Size Radius Measure Atomic Radius - half distance btwn 2 nuclei of diatomic molecule (i.e. O2)Units of picometers (10-12 m… 1 trillionth)
39 ALL Periodic Table Trends Influenced by 3 factors:1. Energy LevelHigher energy levels further away from nucleus.2. Charge on nucleus (# protons)More charge pulls electrons in closer. (+ and – attract each other)3. Shielding effect
40 What do they influence?Energy levels & Shielding have effect on GROUP ( )Nuclear charge has effect on PERIOD ( )
41 #1. Atomic Size - Group trends HGoing down a group, each atom has another energy level (floor)atoms get biggerLiNaKRb
42 #1. Atomic Size - Period Trends left to right across period:size gets smallere-’s occupy same energy levelmore nuclear chargeOuter e-’s pulled closerHere is an animation to explain the trend.NaMgAlSiPSClAr
43 RbKPeriod 2NaLiAtomic Radius (pm)KrArNeH310Atomic Number
45 ion is atom (or group of atoms) w/ + or - charge IonsSome compounds composed of “ions”ion is atom (or group of atoms) w/ + or - chargeAtoms are neutral because the number of protons = electrons+ & - ions formed when e- transferred (lost or gained) btwn atoms
46 Metals LOSE electrons, from outer energy level IonsMetals LOSE electrons, from outer energy levelSodium loses 1 e-more p+ (11) than e- (10)+ charge particle formed…“cation”Na+ called “sodium ion”
47 Nonmetals GAIN one or more electrons IonsNonmetals GAIN one or more electronsCl gains 1 e-p+ (17) & e- (18), so charge of -1Cl1- called “chloride ion”anions
48 #2. Trends in Ionization Energy Ionization energy - energy required to completely remove e- (from gaseous atom)energy required to remove only 1st e-called first ionization energy.
49 Ionization Energysecond ionization energy is E required to remove 2nd e-Always greater than first IE.third greater than 1st or 2nd IE.
50 Symbol First Second Third Table 6.1, p. 173Symbol First Second ThirdHHeLiBeBCNO F Ne
51 Symbol First Second Third HHeLiBeBCNO F NeWhy did these values increase so much?
52 What factors determine IE greater nuclear charge = greater IEGreater distance from nucleus decreases IEFilled & half-filled orbitals have lower energyEasier to achieve (lower IE)Shielding effect
53 Shieldinge-’s in outer energy level “looks through” all other energy levels to see nucleus
54 Ionization Energy - Group trends going down groupfirst IE decreases b/c...e- further away from nucleus attractionmore shielding
55 Ionization Energy - Period trends Atoms in same period:same energy levelSame shieldingIncreasing nuclear chargeSo IE generally increases left - rightExceptions…full & 1/2 full orbitals
56 Both have same shielding (e- in 1st level) HeHe has greater IE than H.Both have same shielding (e- in 1st level)He = greater nuclear chargeHFirst Ionization energyAtomic number
57 These outweigh greater nuclear charge Li lower IE than Hmore shieldingfurther awayThese outweigh greater nuclear chargeHFirst Ionization energyLiAtomic number
58 greater nuclear charge HeBe higher IE than Lisame shieldinggreater nuclear chargeFirst Ionization energyHBeLiAtomic number
59 greater nuclear charge HeB has lower IE than Besame shieldinggreater nuclear chargeBy removing an electron we make s orbital half-filledFirst Ionization energyHBeBLiAtomic number
60 First Ionization energy HeFirst Ionization energyHCBeBLiAtomic number
61 First Ionization energy HeNFirst Ionization energyHCBeBLiAtomic number
62 HeOxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbitalNFirst Ionization energyHCOBeBLiAtomic number
63 First Ionization energy HeFNFirst Ionization energyHCOBeBLiAtomic number
64 Ne has a lower IE than He Both are full, Ne has more shielding Greater distanceFNFirst Ionization energyHCOBeBLiAtomic number
65 Na has a lower IE than Li Both are s1 Na has more shielding HeNeNa has a lower IE than LiBoth are s1Na has more shieldingGreater distanceFNFirst Ionization energyHCOBeBLiNaAtomic number
68 Driving Forces Full Energy Levels require high E to remove e- Noble Gases = full orbitalsAtoms want noble gas configuration
69 2nd Ionization EnergyFor elements w/ filled or ½ filled orbital by removing 2 e-, 2nd IE lower than expected.True for s2Alkaline earth metals form 2+ ions.
70 3rd IEUsing the same logic s2p1 atoms have an low 3rd IE.Atoms in the aluminum family form 3+ ions.2nd IE and 3rd IE are always higher than 1st IE!!!
71 Trends in Ionic Size: Cations Cations form by losing electrons.metalsCations are smaller than the atom they came from –they lose electronsthey lose an entire energy level.Cations of representative elements have noble gas configuration before them.
72 Trends in Ionic size: Anions Anions gain electronsAnions bigger than the atom they came from –same energy levelgreater area the nuclear charge needs to coverNonmetals
73 Configuration of IonsIons always have noble gas configurations (full outer level)Na atom is: 1s22s22p63s1Forms a 1+ sodium ion: 1s22s22p6Same as Ne
74 Configuration of IonsNon-metals form ions by gaining electrons to achieve noble gas configuration.They end up with the configuration of the noble gas after them.
75 Ion Group trends Each step down a group is adding an energy level Li1+Na1+Each step down a group is adding an energy levelIons get bigger going down, b/c of extra energy levelK1+Rb1+Cs1+
76 Ion Period Trends Across period nuclear charge increasesIons get smaller.energy level changes between anions and cations.N3-O2-F1-B3+Li1+Be2+C4+
77 Size of Isoelectronic ions Iso- means “the same”Isoelectronic ions have the same # of electronsAl3+ Mg2+ Na1+ Ne F1- O2- and N3-all have 10 electronsall have the same configuration: 1s22s22p6 (which is the noble gas: neon)
78 Size of Isoelectronic ions? Positive ions that have more protons would be smaller (more protons would pull the same # of electrons in closer)N3-O2-F1-NeNa1+Al3+10987131211Mg2+
79 #3. Trends in Electronegativity Electronegativity is tendency for atom to attract e-’s when atom in a compoundSharing e-, but how equally do they share it?Element with big electronegativity means it pulls e- towards itself strongly!
80 Electronegativity Group Trend Further down a group, farther e- is away from nucleus, plus the more e-’s an atom hasmore willing to shareLow electronegativity
81 Electronegativity Period Trend Metals let e-’s go easilylow electronegativityNonmetals want more electronstake them away from othersHigh electronegativity.