1 Chapter Eight The Periodic Table
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Chapter Eight
The Periodic Table
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Quantum Numbers and the Periodic Table Principle quantum number, n =1 to 7 Row number of periodic table Angular momentum quantum number, l= 0 to (n-1) Specific area of periodic table, spdf Magnetic quantum number, ml = –l to +l Number of orbitals = = 2l+1 Spin quantum number, ms=+1/2 or -1/2 Number of electrons, look at atomic number
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Electronic Configurations and the Periodic Table Add 1 electron for each block in the periodic table
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Mendeleev’s Periodic Table 1869 Arranged the known elements in order of increasing atomic mass from left to right and from top to bottom in groups. Elements with similar properties are placed in same column. Used table to predict properties of undiscovered elements Eka-Silicon Germanium MM: 72 MM: 72.6 Density: 5.5g/mL Density: 5.47g/mL Color: dirty gray Color: grayish white
O I II III IV V VI VII
H1
He Li Be B C N O F4 7 9 11 12 14 16 19
Ne Na Mg Al Si P S Cl20 23 24 27 28 31 32 35Ar K Ca Ge As Se Br40 39 40 75 79 80Kr Rb Sr In Sn Sb Te I84 85 88 115 119 122 128 127Xe Cs Ba Tl Pb Bi131 133 137 204 207 209Rn
(222)
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The Modern Periodic Table
Representative Elements: (main group elements) Incomplete s or p shell determine elemental properties
6 Valence and Core Electrons Valence electrons:
Highest shell farthest from nucleus Largest principal quantum number (n) Located on the outside of the atom Determine the behavior of the atom
Core electrons Located on the inside in inner shells. Principal quantum number is lower
Example Oxygen, O Z = 8 valence electrons e- = 6 Core electrons e- = 2
422 221 pss42 22 ps
21s
7 Effect of Valence electrons on Elements Elements most stable with noble gas configuration
Last column in periodic table No electrons want to be added or removed Octet rule satisfied (8 electrons= 2s + 6p)
Isoelectronic: ion has same spdf as noble gas Helium (and Hydrogen) follow duet rule (2 electrons)
Main group elements Valence electrons: s2p6 Electrons are added: form anions Electrons removed: form cations Transition metals
All form cations Remove electrons from shell 4s before 3d, 5s before 4d, etc.
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Periodic Properties in Main Group Elements
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Atomic Radius Atomic radius increases from top to bottom in a group Electrons are shielded from nucleus Previous shells blocks attraction Effective nuclear charge decreases Large size difference between shells
Atomic radius decreases from left to right across a row Little shielding as all electrons in same shell Effective nuclear charge higher as protons added to nucleus Electrons draw closer to nucleus
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Ionic Radius Anions larger than atoms Low effective nuclear charge More electrons More repulsion
Cations smaller than atoms High effective nuclear charge Fewer electrons Less repulsion
11 Ionization Energy Energy needed to remove an e- from a gaseous atom (or ion)
X(g)→ X++ (g) + 1e- Endothermic Decreases top to bottom: Electrons shielded from nucleus Increases from left to right: Atoms want to now gain electrons 3rd ionization energy > 2nd >1st: Higher effective nuclear charge
12 Elemental Ionization Energies
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Electron Affinity Energy released when an electron is added to a gaseous atom
X(g) + 1e- → X- (g) Exothermic Increases bottom to top
Small atom (F) High nuclear:e- attraction
Increases left to right Small atom (F) High nuclear:e- attraction
2nd electron affinities lower: Electrons add to already negative ion
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Physical and Chemical Properties
in Main Group Elements
15 Valence Properties of Hydrogen
Valence Characteristics: 1s1
Single electron in valence shell Wants to lose electron to form H+ (H3O+ in water) Can gain a second electron to form hydrides, NaH, CaH2. Diatomic: exists as H2(g)
Location in Periodic Table
Normally placed with group 1 elements Reactions:
Hydride reactions: 2NaH(s) + 2H2O (l)→ 2NaOH(aq) + H2(g) Combustion: 2H2(g) + O2 (g)→ 2H2O (l) Acid-Base: NaOH(aq) + HCl (aq)→ NaCl(aq) + H2O (l)
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Valence Properties of Alkali Metals
Valence Characteristics: ns1, n ≥ 2 Single electron in valence shell is lost Forms +1 cation Low ionization energy Extremely reactive
Location in Table 1st column, Group 1A Li, Na, K, Rb, Cs Reactions:
With water: 2Na(s) + 2H2O (l)→ 2NaOH(aq) + H2(g) + heat Oxide formation: 4Li(s) + O2 (g)→ 2Li2O (s) Other oxides can form: peroxide: Na2O2, superoxide: KO2
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Valence Properties of Alkali Earth Metals
Valence Characteristics: ns2, n ≥ 2 Lose both s electrons Forms +2 cation 2 ionization energies Less reactive than 1A
Location in Table 2nd column, Group 2A Be, Mg, Ca, Sr, Ba, Ra Reactions: With oxygen: 2Be(s) + O2 (g)→ 2BeO (s)
With water/steam: Ba(s) + 2H2O(l)→ Ba(OH)2(s) + H2(g) With acid: Mg(s) + H+(aq) → Mg2+(aq) + H2(g) Reactivity increases going down, Be only reacts with O2
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Valence Properties of Group 3A Elements
Valence Characteristics: ns2np1, n ≥ 2 Lose s electrons & p electron Form +3 or +1 cations Weak metal characteristics Form molecular compounds
Location in Table 3rd column, Group 3A B, Al, Ga, I Reactions:
With oxygen: Al(s) + 3O2 (g)→ 2Al2O3 (s) With acid: 2Al(s) + H+(aq) → Al3+(aq) + 3H2(g) Hydride formation: 2Al(s) + 3H2 (g)→ 2AlH3(s) B is a metalloid and does not fully ionize Forms molecular compounds only
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Valence Properties of Group 4A Elements
Valence Characteristics: ns2np2, n ≥ 2 Lose s electrons & p electrons +2 & +4 oxidation states Primarily molecular compounds
Location in Table 4th column, Group 4A C, Si, Ge, Sn, Pb Reactions:
With oxygen: C(s) + O2 (g)→ CO2 (g) also CO With acid: Pb(s) + 2H+(aq) → Pb2+(aq) + H2(g) C is a nonmetal, Si and Ge are metalloids Pb and Sn are metals and can ionize C, Si and Ge form molecular compounds only
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Valence Properties of Group 5A Elements
Valence Characteristics: ns2np3, n ≥ 2 May gain or lose electrons Variable oxidation states Can act as anions -3 charge in salt
Location in Table 5th column, Group 5A N, P are nonmetals As, Sb, Bi more metallic Reactions:
With oxygen: N(s) + O2 (g)→ NO2 (g) also NO, N2O, N2O4 Acidic Oxides in water: N2O5(s) + H2O(l) → 2HNO3(aq) P4O10(s) + H2O(l) → 4H3PO4(aq)
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Valence Properties of Group 6A Elements
Valence Characteristics: ns2np4, n ≥ 2 Tend to gain electrons -2 most common charge Many molecular compounds
Location in Table 6th column, Group 6A O, S, Se, Te, Po Reactions:
With water: SO3 (g) + H2O(g)→ H2SO4(aq) Will form many nonmetal molecular compounds
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Valence Properties of the Halogens
Valence Characteristics: ns2np5, n ≥ 2 Gain 1 electron -1 charge as an anion Positive oxidation state if bonded to each other, BrF3
Location in Table 7th column, Group 7A F, Cl, Br, I, At Reactions: Ref only
With water: 2F2 (g) + 2H2O(g)→ 4HF (aq) + O2 (g) With hydrogen: H2 (g) + X2 (g) → 4HX (g) (F2 most reactive) All nonmetals, often designated with an X Larger halogens can have positive oxidation states All halogens are diatomic
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Valence Properties of the Noble Gases
Valence Characteristics: ns2np6, n ≥ 2 Full octet: No desire to gain or lose electrons + charge if larger gases are forced to bond, XeF4
Location in Table 8th column, Group 8A: He, Ne, Ar, Kr, Xe Reactions: Ref only
Xe can react with oxygen and fluorine, but not easily All gases due to no desire to associate with other atoms
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