PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 7

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PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 7
DR. AUGUSTINE OFORI AGYEMAN Assistant professor
of chemistry Department of natural
sciences Clayton state university
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CHAPTER 7 PERIODIC PROPERTIES OF THE ELEMENTS
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EFFECTIVE NECLEAR CHARGE

• Negatively charged electrons are attracted to the
  positively
• charged nucleus
• The force of attraction between an electron and
  the nucleus
• - Depends on the magnitude of the net nuclear
  charge acting
• on the electrons
• - Depends on the average distance between the
  nucleus and
• the electron
• (Coulombs law)

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EFFECTIVE NECLEAR CHARGE
The force of attraction - Increases with
increasing nuclear charge - Decreases with
increasing average distance between electrons
and the nucleus - Electrons also experience
repulsion by other electrons in the atom Zeff
Z S Zeff effective nuclear charge Z
actual nuclear charge (number of protons in the
nucleus)( Zeff) S screening constant
(represents number of core electrons)
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EFFECTIVE NECLEAR CHARGE
Atomic number of Na 11 Number of valence
electrons 1 Number of core electrons 10 As
simplified from this model Z 11 S 10 Zeff
11-10 1 In actual fact Zeff in Na is about
2.5
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EFFECTIVE NECLEAR CHARGE
In general Zeff in s orbital Zeff in p
orbital Zeff in d orbital Zeff in f
orbital - This is the result of the trend in
energy levels ns increases across the periods (from left to right)
in the periodic table (Z increases but S
remains the same) - Zeff increases slightly down
the groups in the periodic table
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SIZES OF ATOMS
- Atomic radius tends to decrease across the
periods (from left to right) in the periodic
table - Due to increase in effective nuclear
charge which draws valence electrons closer to
the nucleus - Atomic radius tends to increase
down the groups (from top to bottom) of the
periodic table - Due to increase in principal
quantum number of the outer electrons (number
of shells)
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SIZES OF ATOMS
- Cations are smaller than their parent atoms -
Decrease in the number of electrons decreases
electron-electron repulsions - Anions are
larger than their parent atoms - Increase in the
number of electrons increases electron-electron
repulsions - Ionic size increases down the group
of the periodic table for ions carrying the same
charge
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ISOELECTRONIC SPECIES
- A group of ions containing the same number of
electrons Due to the same number of electrons -
Ionic radius decreases with increasing nuclear
charge (electrons are more strongly attracted
to the nucleus) increasing nuclear charge O2-
, F- , Na , Mg2 , Al3 S2- , Cl- , K ,
Ca2 , Ga3 Se2- , Br- , Rb , Sr2 ,
In3 decreasing ionic radius
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ELECTRON CONFIGURATION OF IONS
Cl Ne3s23p5 Cl- Ne3s23p6 Na
Ne3s1 Na Ne
F 1s22s22p5 F- 1s22s22p6 Co
Ar3d74s2 Co3 Ar3d6
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IONIZATION ENERGY

• The energy required to remove an electron from
• a gaseous atom or ion
• X(g) ? X(g) e-
• ?E is positive
• - The atom or ion is assumed to be in its ground
  state
• - The highest energy electron is always removed
  first
• Units kJ/mol
• 96.485 kJ/mol 1 eV

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IONIZATION ENERGY
Ionization Energies of Magnesium (Mg) Mg(g) ?
Mg(g) e- I1 735 kJ/mol Mg(g) ? Mg2(g)
e- I2 1450 kJ/mol Mg2(g) ? Mg3(g) e- I3
7730 kJ/mol I1
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IONIZATION ENERGY
First Ionization Energy (I1) - Energy required to
remove the highest energy electron of an atom -
The first electron is removed from a neutral
atom - The second electron is removed from a
positive ion (more difficult) - Increase in
positive charge binds electrons more tightly -
Large jump is observed in going from removal of
valence electrons to removal of core electrons
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IONIZATION ENERGY
- First ionization energy increases across the
period of the periodic table (from left to
right) - Electrons added in the same principal
quantum number do not completely shield
increasing nuclear charge - First ionization
energy decreases down the group of the periodic
table (from top to down) - As n increases, the
size of orbital increases (distance from
nucleus increases) and electrons are easier to
remove
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IONIZATION ENERGY
- Discontinuities are due to electron repulsions
and shielding (Be to B, N to O) -
Representative elements show a larger range of
values of I1 than the transition-metal
elements - Smaller atoms have higher ionization
energies
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ELECTRON AFFINITY
- The energy change that occurs when an electron
is added to a gaseous atom - A measure of the
attraction of the atom for the added electron -
Energy is released when an electron is added to
most atoms X(g) e- ? X-(g) ?E is
negative Units kJ/mol
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ELECTRON AFFINITY
Cl(g) e- ? Cl-(g) ?E -349 kJ/mol - The
greater the attraction between an atom and an
added electron the more negative the atoms
electron affinity - Electron affinities for
noble gases are positive values (?E 0) -
Halogens have the most negative electron
affinities - Electron affinity changes slightly
down the group of the periodic table
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METALS

• - Refer to chapter 2 for properties of metals
• - Metals tend to have low ionization energies
• - Metals form positive ions relatively easily
• - Metals lose electrons (oxidize) when they
• undergo chemical reactions
• Charge on metals
• Alkali metals 1
• Alkaline earth metals 2

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METALS
- The charge on transition metals do not follow
any obvious pattern - Transition metals are
able to form more than one positive ion -
Compounds of metals and nonmetals are ionic
2Na(s) Cl2(g) ? 2NaCl(s) (contains Na and
Cl- ions)
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METALS
Most Metal Oxides are Basic - Dissolve in water
to form metal hydroxides Na2O(s) H2O(l) ?
2NaOH(aq) O2-(aq) H2O(l) ? 2OH-(aq) (net
ionic equation) - React with acid to form salt
and water NiO(s) 2HNO3(aq) ? Ni(NO3)2(aq)
H2O(l)
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NONMETALS
- Refer to chapter 2 for properties of
nonmetals - Nonmetals tend to have high
electron affinities - Nonmetals tend to gain
electrons when they react with metals -
Compounds composed of only nonmetals are
generally molecular substances
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NONMETALS
Most Nonmetal Oxides are Acidic CO2(g)
H2O(l) ? H2CO3(aq) (acidity of rainwater) -
Dissolve in basic solutions to form salt and
water CO2(g) 2NaOH(aq) ? Na2CO3(aq)
H2O(l)
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ALKALI METALS (GROUP IA)
- Alkali means ashes Relatively abundant in
the - earths crust (Na, K) - sea water -
human bodies - Have low densities and melting
points - Very reactive and readily lose an
electron to form 1 ions
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ALKALI METALS (GROUP 1A)
- Form hydrides with hydrogen and sulfides with
sulfur 2M(s) H2(g) ? 2MH(s) 2M(s) S(s)
? M2S(s) - React vigorously with water to
produce hydrogen gas and alkali metal
hydroxide (very exothermic and may explode)
2M(s) 2H2O(l) ? 2MOH(aq) H2(g) - Can react
with oxygen to form oxides and peroxides 4Li(s)
O2(g) ? 2Li2O(s) (lithium oxide) 2Na(s)
O2(g) ? Na2O2(s) (sodium peroxide)
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ALKALINE EARTH METALS (GROUP 2A)
Compared to alkali metals - harder - more dense
- higher melting points - less reactive than the
respective adjacent alkali metal - Tend to lose
two outer s electrons to from 2 ions - Give
off characteristic colors when heated in a flame
(salts used in fireworks)
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ALKALINE EARTH METALS (GROUP 2A)
- Reactivity increases down the group -
Beryllium does not react with water - Magnesium
reacts slowly with water - Calcium and elements
below react readily with water Ca(s) 2H2O(l)
? Ca(OH)2 H2(g)
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HYDROGEN
- First element in the periodic table (1s1
electron configuration) - Nonmetal - Can be
metallic under extreme pressures - Colorless
diatomic gas - Has very high ionizaton energy -
More than double those of alkali metals - Due to
absence of nuclear shielding of the 1s electron
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HYDROGEN
- Does not easily lose its valence electron -
Share with nonmetals to form molecular
compounds - Can lose its electron to form a
cation (H) - Can gain electron to form the
hydride ion (H-)
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CHALCOGENS (GROUP 6A) THE OXYGEN GROUP
- Properties change from nonmetallic to metallic
down the group Nonmetallic properties oxygen,
sulfur, selenium Metallic properties tellurium
and below - Oxygen is a colorless gas at room
temperature - The other group members are
solids - Oxygen exists as O2 (oxygen gas) and O3
(ozone) - Allotropes (different forms of the
same element in the same state)
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CHALCOGENS (GROUP 6A) THE OXYGEN GROUP
- O2 can produce O3 in lightning storms 3O2(g)
? 2O3(g) ?Ho 284.6 kJ - Sulfur also has
several allotropic forms - The most common is S8
(yellow solid)
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THE HALOGENS (GROUP 7A)
- All halogens are nonmetals - Melting and
boiling points increase with increasing atomic
number - Consist of diatomic molecules (F2,
Cl2, Br2, and I2) - Form colored gases
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THE HALOGENS (GROUP 7A)
At room temperature - Fluorine and chlorine are
gases - Bromine is a liquid - Iodine is a
solid - Have highly negative electron
affinities - Tend to gain electrons to form 1-
ions - Reactivity decreases down the group
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THE HALOGENS (GROUP 7A)
- React readily with most metals to form ionic
halides - React with hydrogen to form gaseous
hydrogen halides H2(g) X2 ? 2HX(g) -
Hydrogen halides dissolve in water to form acids
HCl(aq) - Fluorine is very reactive
(dangerous) - Chlorine is the most industrially
useful
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THE NOBLE GASES (GROUP 8A)
- Nonmetals - Monatomic - Gases at room
temperature - Have completely filled s and p
subshells - Have high first ionization
energies - Have stable electron configuration
- Unreactive