1 PERIODIC TRENDS 2 Electron Filling Order Figure 8.5.
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Transcript of 1 PERIODIC TRENDS 2 Electron Filling Order Figure 8.5.
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PERIODIPERIODIC C
TRENDSTRENDS
PERIODIPERIODIC C
TRENDSTRENDS
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Electron Electron Filling Filling OrderOrder
Figure 8.5Figure 8.5
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Electron Configurations Electron Configurations and the Periodic Tableand the Periodic Table
Figure 8.7
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PhosphorusPhosphorusPhosphorusPhosphorus
All Group 5A All Group 5A elements have elements have [core ] ns[core ] ns2 2 npnp3 3
configurations configurations where n is the where n is the period number.period number.
Group 5AGroup 5A
Atomic number = 15Atomic number = 15
1s1s2 2 2s2s2 2 2p2p6 6 3s3s2 2 3p3p33
[Ne] 3s[Ne] 3s2 2 3p3p33
1s
2s
3s3p
2p
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LithiumLithiumLithiumLithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons
1s
2s
3s3p
2p
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Electron PropertiesElectron Properties
Diamagnetic: NOT attracted to a magnetic fieldParamagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.
Diamagnetic: NOT attracted to a magnetic fieldParamagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.
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NeonNeonNeonNeon
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
Diamagnetic
1s
2s
3s3p
2p
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BerylliumBerylliumBerylliumBeryllium
Group 2A
Atomic number = 4
1s22s2
Diamagnetic
1s
2s
3s3p
2p
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BoronBoronBoronBoron
Group 3A
Atomic number = 5
1s2 2s2 2p1
Paramagnetic
1s
2s
3s3p
2p
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CarbonCarbonCarbonCarbon
Group 4A
Atomic number = 6
1s2 2s2 2p2
Paramagnetic1s
2s
3s3p
2p
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FluorineFluorineFluorineFluorine
Group 7A
Atomic number = 9
1s2 2s2 2p5 --->
Paramagnetic
1s
2s
3s3p
2p
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Ion ConfigurationsIon ConfigurationsIon ConfigurationsIon Configurations
How do we know the configurations of ions?
Determine the magnetic properties of ions.
Ions with UNPAIRED ELECTRONS are
PARAMAGNETIC.
Without unpaired electrons DIAMAGNETIC.
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transition metal ionstransition metal ionstransition metal ionstransition metal ions
Fe [Ar] 4s2 3d6 loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
• loses 3 electrons ---> Fe3+ [Ar] 4s0 3d5
4s 3d 3d4s
Fe Fe2+
4s 3d 3d4s
Fe Fe2+
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Transition MetalsTransition Metals
How do they fill? How can we determine?
CopperCopperIronIron
ChromiumChromium
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Ion Configurations MnIon Configurations MnIon Configurations MnIon Configurations Mn
Mn [Ar] 4s2 3d5 ---> Mn5+ [Ar] 4s03d2
loses 5 electrons ---> Mn5+ [Ar] 4s2 3d0
4s 3d 3d4s
Fe Fe2+D P
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PERIODIPERIODIC C
TRENDSTRENDS
PERIODIPERIODIC C
TRENDSTRENDS
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PERIODICITYPERIODICITYPERIODICITYPERIODICITYPeriod Law-
-physical and chemical properties of elements are a
periodic function of their atomic numbers
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18General Periodic General Periodic TrendsTrends
• Atomic and ionic size
• Ionization energy
• Electron affinity, electronegativity
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19Effective Nuclear Effective Nuclear ChargeCharge
Z*Z*
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Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*
Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*
• Z* is the nuclear charge experienced by Z* is the nuclear charge experienced by the outermost electrons.the outermost electrons. See p. 295 and Screen 8.6.See p. 295 and Screen 8.6.
• Explains why E(2s) < E(2p)Explains why E(2s) < E(2p)
• Z* increases across a period owing to Z* increases across a period owing to incomplete shielding by inner electrons.incomplete shielding by inner electrons.
• Estimate Z* by --> [ Estimate Z* by --> [ Z - (no. inner electrons) Z - (no. inner electrons) ]]
• Charge felt by 2s e- in Li Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Z* = 3 - 2 = 1
• Be Be Z* = 4 - 2 = 2Z* = 4 - 2 = 2
• B B Z* = 5 - 2 = 3Z* = 5 - 2 = 3 and so on!and so on!
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EffectiveEffective Nuclear Charge, Z* Nuclear Charge, Z*
• Atom Z* Experienced by Electrons in Valence Orbitals
• Li +1.28
• Be -------
• B +2.58
• C +3.22
• N +3.85
• O +4.49
• F +5.13
Increase in Increase in Z* across a Z* across a periodperiod
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22General Periodic General Periodic TrendsTrends
Higher effective nuclear chargeElectrons held more tightly
Larger orbitals.
Electrons held less
tightly.
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Periodic Trend in the Reactivity of
Metals
Periodic Trend in the Reactivity of
Metals
LithiumLithium
SodiumSodium
PotassiumPotassium
MOST
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2. Reactivity for Metals
As you go down a group for metals the number of energy levels increase.
Because of this, reactivity increases because the atom is more willing to give away its electron (react).
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3.Nonmetalic Trends: Gain electrons Nonmetals on right side, form anions
Going right elements are more nonmetallic (better gainers of electrons)
Going UP elements become more nonmetallic (want to gain)
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8. Reactivity nonmetals: Gain e The reason Across = fill the energy level
Going UP a group, nonmetals have same valence but fewer total electrons
Flourine is the most reactive nonmetal.
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27Atomic Atomic RadiiRadiiAtomic Atomic RadiiRadii
Figure 8.9Figure 8.9
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Atomic Atomic SizeSize
Atomic Atomic SizeSize
• Size increasesSize increases, down a group. , down a group.
• Because electrons are added into Because electrons are added into additional energy levels, there is less additional energy levels, there is less attraction.attraction.
• Size decreases Size decreases across a period.across a period.
• Because, increased effective nuclear Because, increased effective nuclear charge.charge.
• Size increasesSize increases, down a group. , down a group.
• Because electrons are added into Because electrons are added into additional energy levels, there is less additional energy levels, there is less attraction.attraction.
• Size decreases Size decreases across a period.across a period.
• Because, increased effective nuclear Because, increased effective nuclear charge.charge.
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Atomic SizeAtomic SizeAtomic SizeAtomic Size
Size Size decreasesdecreases across a period owing across a period owing to increase in Z*. Each added electron to increase in Z*. Each added electron feels a greater and greater + charge.feels a greater and greater + charge.
LargeLarge SmallSmall
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Trends in Atomic SizeTrends in Atomic SizeSee Figures 8.9 & 8.10See Figures 8.9 & 8.10
0
50
100
150
200
250
0 5 10 15 20 25 30 35 40
Li
Na
K
Kr
He
NeAr
2nd period
3rd period 1st transitionseries
Radius (pm)
Atomic Number
0
50
100
150
200
250
0 5 10 15 20 25 30 35 40
Li
Na
K
Kr
He
NeAr
2nd period
3rd period 1st transitionseries
Radius (pm)
Atomic Number
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Ion SizesIon SizesIon SizesIon Sizes
• CATIONSCATIONS are are SMALLERSMALLER than the than the atoms from which they come.atoms from which they come.
• The electron/proton attraction has The electron/proton attraction has gone UP and so size gone UP and so size DECREASESDECREASES..
Li,152 pm3e and 3p
Li +, 78 pm2e and 3 p
+Forming Forming a cation.a cation.Forming Forming a cation.a cation.
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Ion SizesIon SizesIon SizesIon Sizes
• ANIONSANIONS are are LARGERLARGER than the atoms from than the atoms from which they come.which they come.
• The electron/proton attraction has gone DOWN The electron/proton attraction has gone DOWN and so size and so size INCREASESINCREASES..
• Trends in ion sizes are the same as atom sizes. Trends in ion sizes are the same as atom sizes.
Forming Forming an anion.an anion.Forming Forming an anion.an anion.F, 71 pm
9e and 9pF-, 133 pm10 e and 9 p
-
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Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes
Figure 8.13Figure 8.13
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Ionization EnergyIonization Energy
IE = energy required to remove an electron from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
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Mg (g) + 738 kJ ---> MgMg (g) + 738 kJ ---> Mg++ (g) + e- (g) + e-
MgMg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-
MgMg++ has 12 protons and only 11 has 12 protons and only 11 electrons. Therefore, IE for Mgelectrons. Therefore, IE for Mg++ > Mg. > Mg.
IE = energy required to remove an electron from an atom in the gas phase.
Ionization EnergyIonization Energy
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11stst IE: Mg (g) + 735 kJ ---> Mg IE: Mg (g) + 735 kJ ---> Mg++ (g) + e- (g) + e-
22ndnd IE: Mg IE: Mg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-
33rdrd IE: Mg IE: Mg2+2+ (g) + 7733 kJ ---> Mg (g) + 7733 kJ ---> Mg3+3+ (g) + e- (g) + e-
Energy cost is very high to dip into a Energy cost is very high to dip into a shell of lower n (core electrons). shell of lower n (core electrons). This is why ox. no. = Group no.This is why ox. no. = Group no.
Ionization EnergyIonization Energy
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Trends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization Energy
1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 350
500
1000
1500
2000
2500
1st Ionization energy (kJ/mol)
Atomic NumberH Li Na K
HeNe
ArKr
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Trends in Ionization Trends in Ionization EnergyEnergy
Trends in Ionization Trends in Ionization EnergyEnergy
• IE decreases down a group
• Because size increases.
• IE increases across a period
• Because effective nuclear
charge increases
• IE increases across a period
• Because effective nuclear
charge increases
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Electron AffinityElectron Affinity
A few elements GAIN electrons to form anions.
Electron affinity is the energy involved when an atom
gains an electron to form an anion.
X(g) + e- ---> X-(g) E.A. = ∆E
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Trends in Electron AffinityTrends in Electron Affinity
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• Affinity for electron increases across a period (EA becomes more negative).
• Affinity decreases down a group (EA becomes less negative).
Atom EAAtom EAFF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ
Atom EAAtom EAFF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ
Trends in Electron Trends in Electron AffinityAffinity
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Electron Affinity of OxygenElectron Affinity of Oxygen
∆E is EXOthermic
because O has an affinity for
an e-.
[He] O atom
EA = - 141 kJ
+ electron
O [He] - ion
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Electron Affinity of Electron Affinity of NitrogenNitrogen
∆∆E is E is zero zero for Nfor N- -
due to due to electron-electron-electron electron
repulsions.repulsions.EA = 0 kJ
[He] N atom
[He] N- ion
+ electron
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Electronegativity
• So how is this different from electron affinity?
• Electron Affinity – is rating of how well an atom wants to gain an electron
• Electronegativity – is rating of how well an atom keeps the electron once it is bonded to another atom
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Electronegativity
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46Electron Configurations and the Periodic Trends
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“Your Best Friend”
• Periodic table