Transcript Document


Periodic table: trends in chemical and
physical properties that occur within the
same groups and periods
 First attempt (Mendeleev and Meyer) elements
arranged in order of increasing atomic weight

Modern periodic table: elements arranged
in order of increasing atomic number

Effective nuclear charge (Zeff): positive
charge felt by a valence electron in a
many-electron atom
 Depends on its distance from the nucleus and
the number of core electrons
 Core electrons block, or screen, valence e- from the full
attraction of the nucleus
 As average number of screening/core electrons (S)
increases, effective nuclear charge decreases
 As distance from the nucleus increases, S increases and Zeff
decreases

Mg: [Ne]3s2
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Z = 12
S = core e- = 10
Z – S = 12 – 10 = 2

Effective nuclear charge…
› Increases as you move across a period
 Number of core e-, S, remains the same
 Atomic number increases; higher nuclear
charge, Z
 Valence e- added to counter charge of p+, but
valence e- don’t shield one another
› Decreases as you go down a group
 Large e- cores (think atomic radius) aren’t able
to screen valence e- from the nuclear charge
Bounce two of the same atoms off each
other (B)
 The distance between two nuclei is the
nonbonding atomic radius (aka: van der
Waals radius)
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-
-
-
-
Brad
-
-
+
-
+
- Brad
-
-
Homonuclear
diatomic molecule
 Distance between 2
nuclei is bond length
 Half the bond length
is the bonding atomic
radius
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Decrease in atomic radius across a period
› Effective nuclear charge increases, drawing the electrons
in towards the nucleus very tightly
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Increase in atomic radius down a group
› As the principal quantum number increases, the
probability of finding the electrons further from the nucleus
increases

Arrange the following atoms in order of
increasing atomic size:
› Mg, Ca, Sr
› B, F, Ge, Pb
Ionic radius: distance between ions in an
ionic compound
 When atoms form ions, they tend to
achieve a full valence shell by either losing
e- or gaining e Cations are smaller than parent atom

› Outermost orbital is emptied - decreasing radius
of the ion

Anions are larger than parent atom
› Outermost orbital fills up, taking up more space
 For
ions of the same charge, ion size
increases down a group
 Ion size decreases across a period
 ASIDE:
Transition metals lose electrons
in outermost orbital first!
› Ex.: Fe(3d64s2)  Fe3+(3d5) + 3e-
Isoelectric series: group of ions containing
the same number of electrons
 Ex: Se2-, Br-, Rb+, Sr2+, Y3+
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Se2- = atomic # 34 (36 e-)
Br- = atomic # 35 (36 e-)
Rb+ = atomic # 37 (36 e-)
Sr2+ = atomic # 38 (36 e-)
Y3+ = atomic # 39 (36 e-)
Ionic size in isoelectric series decreases with
increasing atomic number due to
increasing effective nuclear charge

Ionization energy (I): minimum energy
required to remove an e- from a gaseous
atom or ion
› The larger the ionization energy, the harder it is
to remove the e-


First ionization energy (I1): energy needed to
remove 1st e- from neutral atom
› Na(g)  Na+(g) + e-(g)
I1 = +496 kJ/mol
› Na+(g)  Na2+(g) + e-(g)
I2 = +4562 kJ/mol
Second ionization energy (I2): energy
needed to remove 2nd e-
Each successive ionization energy is
higher than the previous
 Sharp increase in ionization energy when
a core e- is removed

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Decreases as you move down a group
and atomic number gets larger
› Outermost e- is easier to remove as you go
down a group and atom gets bigger

Increases as you move across a period
› Outermost e- is harder to remove as you go
across a period and atom gets smaller

Electron affinity (E): energy change
associated with adding an e- to a
gaseous atom or ion
› I(g) + e-(g)  I-(g)

ΔE = -295 kJ/mol
For most atoms, energy is usually
released when an e- is added
› Negative value

Puts an e- in a new (higher energy)
subshell (unfavorable)


Ne(g) = e-(g)  Ne-(g)
ΔE > 0
To determine whether e- affinity is (+) or (-),
look at e- configuration:
› Do you have to add the e- to a higher energy
orbital or subshell?
› Ne = 1s22s22p6
› Ne- = 1s22s22p63s1
Ionization Energy
Electron Affinity
How easily an atom or ion loses an e Forms consecutively more (+) species
 Measure of how tightly an e- is held on to
 Usually requires energy [(+) value]
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How easily an atom or ion gains an eForms consecutively more (-) species
Measure of attraction for an outside eUsually releases energy [(-) value]
What is the ionization energy for F-(g) if the
electron affinity for F(g) is -328 kJ/mol?
 Which member of each ion pair has the
higher ionization energy? Why?

› Sodium or rubidium
› Silicon or phosphorous

Which element has the higher electron
affinity, sulfur or chlorine?
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Metals
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Nonmetals
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Shiny luster
Various colors, but most are silvery
Malleable and ductile
Good conductors
Metal oxides form bases in H2O
Form cations in solution
All are solids at RT (except Hg)
No luster
Various colors
Brittle
Poor conductors
Nonmetal oxides form acids in H2O
Form anions in solution
Solids, gases, and liquid at RT
Metals form cations
 Low ionization energies
 Metallic character:

› Increases down a group
› Decreases across a period

When metals are oxidized they form
characteristic cations:
› Group 1A metals:
 1+ ions
› Group 2A metals:
 2+ ions
› Most transition metals have variable charges

Form anions
› When nonmetals react with metals, nonmetals gain e- and
metals lose e-
 Metal + Nonmetal  Salt
 2Al(s) + 3Br2(l)  2AlBr3(s)

When nonmetals are reduced they form
characteristic anions
› Group 7A nonmetals
 1- ions
› Group 6A nonmetals
 2- ions
› Group 5A nonmetals
 3- ions
Soft
 Chemistry dominated by loss of single e
› M  M+ + e-
Reactivity increases down the group b/c
easier to lose e Alkali metals react with water to form
MOH and H2 gas

› 2M(s) + 2H2O(l)  2MOH(aq) + H2(g)


Alkali metal ions are reduced to metal gas atoms that emit
characteristic colors when placed in a high temp flame
“s” e- is excited by flame and jumps to “p” sublevel and then
emits light energy when it returns to ground state
Li line: 2p  2s
transition
Na line (589 nm): 3p
 3s transition
K line: 4p  4s
transition
Harder and more dense
 Chemistry dominated by loss of 2 e
› M  M2+ + 2e
Reactivity increases down the group b/c
easier to lose e-

Very unique element that forms
› Colorless diatomic gas, H2
› Metallic solid at high pressures
› Can behave as a cation or anion:
2Na(s) + H2(g)  2NaH(s)
less common (H-; hydride)
H2(g) + Cl2(g)  2HCl(g)
most common (H+; proton)

Metallic character increases down the
group
› O2 colorless gas; others are solids
› Te is a metalloid, Po is a metal, etc.
Gain e- to form 2- anions
2 important forms of oxygen: O2 & O3
3O2(g)  2O3(g)
H = +284.6 kJ
 O2: potent oxidizing agent
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› Two oxidation states for oxygen:
› 2- (Ex. H2O)
› 1- (Ex. H2O2)

All form diatomic molecules
› F2 and Cl2  gases; Br2  liquid; I2  solid

Gain an e- to form an anion
X2 + 2e-  2X-

Fluorine is one of the most reactive substances known
2F2(g) + 2H2O(l)  4HF(aq) + O2(g)

H = -758.7 kJ
All react with metals to produce ionic halide salts
Mg(s) + Cl2(g)  MgCl2(s)

All react with hydrogen gas to form gaseous
hydrogen halide compounds
Cl2(g) + H2(g)  2HCl(g)
 All
monatomic
 Notoriously unreactive
 Have completely filled s and p
subshells
 In 1962 the first compound of the
noble gases was prepared: XeF2, XeF4,
and XeF6
› Why these particular compounds?