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-Periodic table with color-coding of main group elements, noble gases, transition metals, group
2B metals, lanthanides, and actinides

Electron Configurations of Group 1A and 2A Elements
Group 1AGroup 2A
Li
[He]2s1 Be [He]2s2
Na
[Ne]3s1 Mg[Ne]3s2
K
[Ar]4s1 Ca [Ar]4s2
Rb
[Kr]5s1 Sr [Kr]5s2
Cs
[Xe]6s1 Ba [Xe]6s2
Fr
[Rn]7s1 Ra [Rn]7s2

Fr

[Rn]7s1

Ra [Rn]7s2

*Although hydrogen’s electron configuration is 1s1, it is a nonmetal and is not really a member of
Group 1A
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e
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Valence Electron Configurations

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Shielding (Screening): the partial obstruction of nuclear charge by core electrons
*Although al the electrons in an atom shield one another to some extent, those that are most
effective at shielding are the core electrons
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Core Electrons: those in the completed inner shells
* As we move to the right across period 2, the nuclear charge increases by 1 with each new
element, but the effective nuclear charge increases only by an average of 0
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(If the valence
electrons did not shield one another, the effective nuclear charge would also increase by 1 each
time a proton was added to the nucleus
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4: Periodic Trends in Properties of Elements
*The value of the principal quantum number (n) increases as the distance from the nucleus
increases
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The atomic radius decreases as we move from left to right across a period
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2
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• As we step down a column, the outermost occupied shell has an everincreasing value of n, so it lies farther from the nucleus making the radius
bigger
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The result is that
as we step across a period the valence shell is drawn closer to the nucleus making the atomic
radius smaller
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Ionization Energy (IE): the minimum energy required to remove an electron from an atom in
the gas phase, units: kJ/mol
*First ionization energy corresponds to the removal to the most loosely held electron
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(a): It’s harder to
remove an electron
from an s orbital than it
is to remove an
electron from a p
orbital with the same
principal quantum
number n
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orbital
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move from top to bottom within a group due to the *In general, as effective nuclear charge
increasing atomic radius
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*According to Coulomb’s Law, the attractive force *IE increases from left to right across a
between a valence electron and the effective nuclear period
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This makes it easier to remove an electron, a higher energy (are less tightly held by
and so IE, decreases
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possible, requires ever-increasing amounts of energy,
because it is harder to remove an electron from a
cation than from an atom
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*It takes much more energy to remove core electrons than to remove valence electrons because
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Core electrons are closer to the nucleus
​2
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Electron Affinity (EA): the energy released (the negative of the enthalpy change ΔH) when an
atom in the gas phase accepts an electron
*In general, the larger and more positive (a): Electron affinities (kJ/mol) of the main group
the EA value, the more favorable the
elements
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(b): Electron affinity as a function of atomic number
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• It becomes progressively easier to

*Electron affinity increases from left to right across a
period
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*Although IE and EA both increase from left to right
across a period, an increase in IE means that it is less
likely that an electron will be removed from an atom
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(a): It’s easier to add an electron to an s orbital than to add one to a p orbital with the same
principal quantum number n
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*While many first electron affinities are positive, subsequent electron affinities are always
negative
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Metalloid: elements with properties intermediate between metals and nonmetals
Coulomb’s Law: the force (F) between two charged objects (Q1 and Q2) is
directly proportional to the product of the two charges and inversely
proportional to the distance (d) between the objects squared
*When the charges have opposite signs, F is negative – indicating an attractive force between
the objects
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Distance (d = 1) from Each Other
Q1 Q2 Attractive force is proportional to
+1−11

1

2

+1−11
+2−24
+3−39

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Isoelectric: describes two or more species with identical electron configurations
*To write the electron configurations of an ion formed by a main group element,
​1
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Either add or remove the appropriate number of electrons
*d-block elements/transition metals always lose electrons first from the s shell with the highest
value of n
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Ionic Radius: the radius of a cation or anion
• Affects the physical and chemical properties of an ionic compound
*When an atom loses an electron and becomes a cation, its radius decreases due in part to a
reduction in electron-electron repulsions (and consequently a reduction in shielding) in the
valence shell
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Adding an electron causes the rest of the electrons in the
valence shell to spread out and take up more space to maximize the distance between them
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*Electron affinity is a measure of how powerfully an atom can attract electrons from another
source
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Elements in the same group resemble one another in chemical behavior because
they have similar valence electron configurations
a
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b
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i
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2
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Hydrogen (1s1)
• No completely suitable position for hydrogen in the periodic table
• Can be a cation (H+) or an anion (H-…called hydride)
Group 1A Elements (ns1, n ≥ 2)
• Low ionization energies
• Tend to form M+ cations
• Very reactive
• React with water to produce H2(g) and the corresponding metal hydroxide

• React with water to produce H2(g) and the corresponding metal hydroxide
▪ M denotes an alkali metal
• Lithium forms lithium oxide (containing the oxide ion, O2-)
• Other alkali metals form oxides or peroxides (containing the peroxide ion,
O22-)
• K, Rb, Cs also form superoxides (containing the superoxide ion O2-)
Group 2A Elements (ns2, n ≥ 2)
• Somewhat less reactive than alkali metals
• Both the 1st and 2nd IE decreases (and metallic character increases from Be à
Ba)
• Tend to form M2+ ions, where M denotes an alkaline earth metals atom
• Reactions of alkaline earth metals with water vary considerably
o Be does not react with water
o Mg reacts slowly with steam
o Ca, Sr, and Ba react vigorously with cold water

• Reactivity toward oxygen increases from Be à Ba
o BeO and MgO form only at elevated temperatures
o CaO, SrO, and BaO form at room temperature
Group 3A Elements (ns2np2, n ≥ 2)
• B is a metalloid while the others (Al, Ga, In, and Tl) are metals
• B does not form binary ionic compounds and is unreactive toward both
oxygen and water
• Al readily forms aluminum oxide when exposed to air
• Al forms the A3+ ion
• Ga, In, and Tl can form both M+ and M3+ ions
• As we move down the group, the M+ ion becomes the more stable of the two
• The metallic elements also form many molecular compounds
Group 4A Elements (ns2np2, n ≥ 2)
• C is a nonmetal
• Si and Ge are metalloids
• Sn and Pb are metals and do not react with water but do react with aqueous
acid to produce H2(g)

• Group 4A elements form compounds in both the +2 and +4 oxidation states
o For C and Si, the 4+ oxidation state is the more stable one
o Sn compounds à the +4 oxidation state is only slightly more stable
than the +2 oxidation state
o Pb comounds à the +2 oxidation state is more stable
o Outer electron configuration of lead is 6s26p2 and it tends to lose
only the 6p electrons to form Pb2+ rather than both the 6p and 6s
electrons to form Pb4+
2
3
Group 5A Elements (ns np , n ≥ 2)
• N and P are nonmetals

Group 5A Elements (ns2np3, n ≥ 2)
• N and P are nonmetals
• As and Sb (Antimony) are metalloids
• Bi is a metal
• Elemental nitrogen is a diatomic gas (N2)
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All of which are gases except N2O5 which is a
solid at room temperature
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Most metal nitrides, such as Li3N and Mg3N2, are ionic compounds
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o It forms two solid oxides with the formulas P4O6 and P4O10
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• Oxygen is a colorless, odorless, diatomic gas; elemental sulfur and selenium
exist as the molecules S8 and Se8, respectively; and tellurium and
polonium have more extensive three-dimensional structures
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• Sulfur, selenium, and tellurium also form ions by accepting two electrons:
S2–, Se2–, and Te2–
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Group 7A Elements (ns2np5, n ≥ 2)
• All the halogens are nonmetals with the general formula X2, where X
denotes a halogen element
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• The halogens have high ionization energies and large, energetically
favorable electron affinities
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• The vast majority of alkali metal halides are ionic compounds
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• The halogens react with hydrogen to form hydrogen halides:
Group 8A Elements (ns2np6, n ≥ 2)
• All the noble gases exist as monatomic species
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• Their electron configurations give the noble gases their great stability
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• Their electron affinities are all less than zero, so they have no tendency to
accept extra electrons
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• Because the 1B IE values are considerably larger than those of the alkali
metals, the Group 1B elements are much less reactive
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• The higher ionization energies of the Group 1B elements result from
incomplete shielding of the nucleus by the inner d electrons (compared
with the more effective shielding by the completely filled noble gas cores)
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Variation in Properties of Oxides Within a Period

• Most oxides can be classified as acidic or basic depending on whether they
produce acidic or basic solutions when dissolved in water (or whether they
react as acids or bases)
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