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PROPERTIES OF THE ELEMENTS AND PERIODICITY
- Most chemical properties are dependent of structure of valence electrons.
- Elements in a single group (column) have similar electronic structures.
- Therefore elements in a single group have similar chemical properties.
ELECTROSTATIC EFFECTS ON ATOMIC STRUCTURE
1. Increasing nuclear charge, increases attraction of electron to the nucleus.
Example: Which has a lower energy: a 2s electron in a hydrogen atom or a 2s
electron in a helium ion?
H:
He+:
1s1
1s1
Helium has a higher nuclear charge.
2. Electrons that have other electrons closer to the nucleus are not as attracted to
the nucleus.
Example: Which is more tightly bound to the nucleus: a 3s electron in the
fluorine atom or a 3s electron in a fluorine anion?
F*:
F-*:
1s22s22p43s1
1s22s22p53s1
- The 3s electron in the atom has a lower energy (more negative) energy
than the 3s electron in the ion.
- The F- ion has a lower effective nuclear charge than the F atom due to
shielding.
Effective nuclear charge
Electrons that are in between an outer electron and the nucleus reduce the attraction
between the outer electron and the nucleus. The reduction in attraction can be
described using the effective nuclear charge, Zeff.
The number of electrons in between are described by a screening constant, S, so
that the connection between the nuclear charge and effective nuclear charge is
Zeff = Z – S
Examples:
For a 3s electron in sodium, S would be 10 (that is 2 “1s” electrons, 2 “2s”
electrons plus 6 “2p” electrons).
Zeff = 11 – 10 = +1
For a 3p electron in phosphorus, S would be 2 + 2 + 6 + 2 = 12;
therefore,
Zeff = 15 – 12 = +3
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3. Electrons with the same subshell repel each other.
Example: Which has a lower energy level: a 3p electron in a chlorine atom or a
chlorine anion?
Cl:
Cl-:
1s22s22p63s23p5
1s22s22p63s23p6
- The 3p electrons in the chlorine atom will have a more negative energy
since the 5 3p electrons in the atom will repel each other less than the 6
3p electrons in the chlorine anion.
4. Electrons with the same orbital repel each other.
Example: Which electron is easiest to remove: a 2p electron in oxygen or a 2p
electron in nitrogen?
N:
O:
1s22s22p3
1s22s22p4
- The 2p electron in the oxygen atom will be easier to remove since it is
being repulsed by the other electron in its orbital.
- The 2p electron in the oxygen atom has a higher energy (less negative)
than the 2p electron in nitrogen.
Order of strength of electrostatic effects.
1. Effective charge of nucleus
- electron-nuclear attraction
- nuclear charge including shielding
2. Subshell electron repulsion
3. Orbital electron repulsion
PRINCIPLES OF ELECTRONIC CONFIGURATION STABILITY
- Electronic configurations with a filled s subshell and a filled p subshell are
most stable.
- Electronic configurations with other filled subshells are somewhat stable.
- Electronic configurations with half-filled subshells are slightly stable.
- Filled shells and subshells “shield” or “screen” nuclear charge from valence
electrons.
- Principles of stable electronic configuration are very important in explaining
chemical properties.
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MAGNETIC PROPERTIES OF ATOMS
Diamagnetic
- atom with no unpaired e- atom not attracted to magnet
- in fact, atom is slightly repelled by magnet
Examples: Be, Ar, Zn
Paramagnetic
- atom with unpaired e- spinning e- acts like bar magnet
- e- has magnetic moment
- atom behaves as if it is magnet, it is attracted to other magnets
Examples: K, N, Ag
IONIC CHARGE
- Atoms in same group tend to have same ionic charge.
- alkali metals
1+
- alkaline earth metals
2+
- chalcogens
2- halogens
1- Why does this pattern exist??
Alkali Metals
- valence electron is ‘s’ electron
- atom becomes ion by losing ‘s’ electron to have filled s and p subshells
Example
Na:
1s22s22p63s1
+
Na : 1s22s22p6 (filled shell)
Alkaline Earth Metals
- valence electrons are ‘s’ electrons
- atom becomes ion by losing two ‘s’ electrons to have filled s and p subshells
Example
Ca:
1s22s22p63s23p64s2
Ca2+: 1s22s22p63s23p6
Halogens
- valence electrons are ‘p’ electrons
- atom becomes ion by gaining ‘p’ electron to have filled s and p subshells
Example
F:
1s22s22p5
F:
1s22s22p6
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Chalcogens
- valence electrons are p electrons
- atom becomes ion by gaining two ‘p’ electrons to have filled s and p
subshells
Example
Se:
1s22s22p63s23p64s23d104p4
2Se : 1s22s22p63s23p64s23d104p6
**Note: Electron configurations of the above ions have the same electron
configurations as noble gases.**
*Caution: Charges of transition metals are found with much subtler rules.*
E.g., transition metal atoms lose their s electrons before they lose their d electrons.
Example
Fe:
1s22s22p63s23p64s23d6
2+
Fe : 1s22s22p63s23p63d6
Fe3+: 1s22s22p63s23p63d5
The s electrons in a transition metal are “first in, first out”.
ATOMIC RADIUS
- atomic radius is distance of furthest electron from nucleus
- since electrons are in orbitals, measuring distance is subjective
- 3 common ways to consider atomic radii
1. ½ distance between atoms in a solid
r = d/2
d
2. ½ distance between atoms in a molecule (covalent radius).
- cut the homonuclear bond length, 1, in half.
l
r = l/2
l(Br2) = 228 pm  r(Br) = 114 pm
5
3. surface that contains 95% of electron density
50%
75%
95%
4. furthest peak in radial distribution function of valence electrons
RADIAL DISTRIBUTION FUNCTIONS
- Electron density within a spherical shell.
0.045
0.043
e- density
 r2
0.04
1s
0.035
0.03
0.025
a ( r)
b ( r)
d ( r)
0.02
2s
0.015
3s
0.01
0.005
0
0
0
0
5
10
15
20
25
25
r
radius
Note: As n quantum number increases, electron density spreads further from
nucleus.
As an approximation, we can say that the 2s electrons are “outside” of the 1s
electrons.
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Radial distribution functions help to find screening constants and effective nuclear
charges.
0.05
0.047
2p
0.045
e- density
 r2
0.04
0.035
0.03
3d
c ( r)
e ( r)
3p
0.025
f ( r)
0.02
0.015
0.01
0.005
0
0
0
5
0
10
15
20
r
25
25
Note some of the 3d electrons are “inside” of the 3p electrons.
The 2p electrons shield the 3p and 3d electrons (as does the 2s and 1s electrons).
The concept of effective nuclear charge is very useful in explaining periodic
trends of atomic properties.
ATOMIC RADIUS AND PERIODICITY
increasing
1. Atomic radius increases when row increases
- when valence electrons increase shell, they are further from the nucleus.
r(Ca) = 1.97 Å
r(Sr) = 2.15 Å
2. Atomic radius decreases going from left to right on table.
- valence electrons are in same subshell, but feel greater attraction due to
increased charge of the nucleus.
r(C) = 0.77 Å
r(N) = 0.75 Å
r(O) = 0.73 Å
Zeff = 3.136
Zeff = 3.834
Zeff = 4.453
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IONIC RADII
- size of ions depends on three things
1. Outermost shell of electrons
r(Rb+) > r(K+)
4p e- > 3p e2. Since many ions have noble gas e- config., charge of nucleus is very important.
Consider e- config. of following ions
1s22s22p6
O2-:
F:
1s22s22p6
Ne:
1s22s22p6
Na+:
1s22s22p6
2+
Mg :
1s22s22p6
What is relative order of the ionic radii?
r(O2-) > r(F-) > r(Ne) > r(Na+) > r(Mg2+)
- increased nuclear charge pulls e- closer to nucleus
O2-, F-, Ne, Na+, Mg2+ is an example of an isoelectronic series.
isoelectronic – same number of electrons
3. Repulsion of electrons in the valence subshells
Radius of anion is greater than neutral atom
r(O) = 73 pm
r(O2-) = 126 pm
Radius of cation is smaller than neutral atom
r(Fe) = 125 pm
r(Fe2+) = 75 pm
r(Fe3+) = 69 pm
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IONIZATION ENERGY AND PERIODICITY
increasing
1. First ionization energy decreases when row increases.
- Electron further away from nucleus is less attracted, therefore less energy is
needed to remove it.
EI (F) = 1681 kJ/mol
1
EI (Cl) = 1251 kJ/mol
1
2. First ionization energy increases going from left to right.
- effective nuclear charge increases creating greater attraction to nucleus
EI (Al) = 578 kJ/mol
1
EI (Si) = 786 kJ/mol
1
EI (P) = 1012 kJ/mol
1
3. Removing electrons from lower energy subshells takes larger amounts of energy.
EI (Na) = 496 kJ/mol (3s)
EI (Na) = 4560 kJ/mol (2p)
EI (Mg) = 738 kJ/mol (3s)
EI (Mg) = 1450 kJ/mol (3s)
1
1
2
2
- For the second electron removed, the 2p of Na is closer than the 3s of Mg.
4. Removing the first paired electron in a subshell takes a slightly smaller
amount of energy.
EI (N) = 1402 kJ/mol
1
EI (O) = 1314 kJ/mol
1
- Paired electron in oxygen is easier to remove than unpaired electron of
nitrogen.
5. A break in the horizontal trend occurs between elements straddling the
boundary between “blocks” on the periodic table especially between the “s”
and “p” blocks and the “d” and “p” blocks.
EI (Mg) = 738 kJ/mol (3s)
EI (Al) = 578 kJ/mol
(3p)
EI (Cd) = 868 kJ/mol (5s)
EI (In) = 558 kJ/mol
(5p)
1
1
1
1
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ELECTRON AFFINITY AND PERIODICITY
- energy change (usually loss) when e- is added to atom in gas phase
- each electron added has an electron affinity
- positive electron affinity indicates atom has no desire for additional electron
Source of Confusion: The more negative the number, the greater the electron
affinity.
increasing
1. Electron affinity tends to decrease when row increases
- two reasons
a) attraction of e- to nucleus is weaker
b) repulsion of e- with other e- is less
- distance between e- is greater
- two effects work together, sometimes one is more important than other
- for example, note exception to rule
EA(F) = – 328 kJ/mol
EA(Cl) = – 349 kJ/mol
- 2p e- of fluorine are more strongly attracted to the nucleus; however, they are
also more repelled by other electrons in the open subshell.
2. Electron affinity tends to increase going from left to right.
- greater effective nuclear charge causes attraction of e- to nucleus to be
stronger
3. Stability of electron configuration is also very important.
EA(Li) = – 60 kJ/mol Be has filled subshell.
EA(Be)  0 kJ/mol
EA(C) = – 60 kJ/mol
EA(N)  +0 kJ/mol
EA(O) = –141 kJ/mol
N has half-filled subshell. Takes energy to pair e-.
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METALS AND NONMETALS
- Physical characteristics of metals
- shiny luster
- malleability – ability of material to be pounded into thin sheet without
breaking
- ductility – ability of material to stretched in thin wire without breaking
- high electrical conductivity
- high heat conductivity
- Chemical characteristics of metals
- ions of metals are cationic in solution
- ions of nonmetals are anionic in solution
- oxides
- metal oxides react with water to form bases
- ionic bonding with metal oxides
- O2- + H2O  2 OH- nonmetal oxides react with water to form acids
- covalent bonding within nonmetal oxides
- Cl2O7 + H2O  2 HClO4
- N2O5 + H2O  2 HNO3
- metalloid oxides are amphoteric
- oxide may behave as acid or base
METALLIC CHARACTER AND PERIODICITY
increasing
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CHEMICAL CHARACTERISTICS OF MAIN-GROUP ELEMENTS
Alkali metals
- ionic charge +1
- oxides have M2O formula.
- chlorides, etc… have MX formula.
Alkaline earth metals
- ionic charge +2
- oxides have MO formula.
- chlorides, etc… have MX2 formula.
Group 13
- ionic charge +3
- oxides have M2O3 formula.
- B is a metalloid.
- chlorides, etc… have MX3 formula.
Group 14
- Sn, Pb have ionic charges +2 and +4.
- C, Si covalently bond.
- Si, Ge are metalloids.
- oxides have MO2 formula.
- chlorides, etc… have MX4 formula.
Pnictogens
- Bi has an ionic charge of +3
- N, P have an ionic charge of -3.
- As and Sb are metalloids.
- oxides have M2O5 formula.
- chlorides, etc… have MX3 formula.
Chalcogens
- ionic charge -2
- oxides have MO3 formula.
- Se, Te and Po(?) are metalloids.
- chlorides, etc… have MX2 formula.
Halogens
- ionic charge -1
- oxides have M2O7 formula.
- At is metalloidish! 
- chlorides, etc… have MX formula.
Noble gases
- Kr, Xe and Rn are reactive toward F and O.
- He, Ne and Ar are unreactive.