10 FE Chemistry Atoms F12
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An element - is a pure substance that cannot be decomposed into other pure substances. - each element is homogeneous. - each element is composed of only one kind of atom. - each element is represented by a chemical symbol; e.g., gold is an element represented by the symbol, Au. 4.1 Atom 4.1.1 Atomic Structure and the Properties of the Elements An atom - is the smallest subdivision or particle of an element that enters into the chemical reactions of that element. - consists of an extremely dense nucleus of protons and neutrons, and very diffuse surrounding electron cloud containing enough electrons to exactly balance the nuclear charge. Nucleus - consists of protons and neutrons (electrons exist outside of nucleus) Protons - particles with a unit positive charge and a mass very close to 1 amu (amu = atomic mass unit); it has +1 charge Neutrons - neutral particles of essentially the same mass as the proton (1 amu); it has 0 charge Electrons - entities with a unit negative charge and a comparatively negligible mass (assumed to be 0 amu) - actually ~ 5.5 x 10-4 amu); it -1 charge - appear to have properties of both waves and particles Note: - For any atoms, total number of protons = total number of electrons. - Each element has unique number of protons; e.g., H has 1 proton. - A negative ion of a certain element can be formed by addition of an electron to an atom of that element. Atomic weight of an element - is the weight of an atom of that element in atomic mass units. - One atomic mass unit (amu) = 1/12 the mass of one normal C atom ~ mass of one H atom - the sum of the masses of the protons, neutrons, and electrons in the atom (Atomic weight = # of protons + # of neutrons ) -
is approximately equal to the number of protons and neutrons, with a little extra added by the electrons. An atomic weight (relative atomic mass) of an element from a specified source is the ratio of the average mass per atom of the element to 1/12 of the mass of 12C" in its nuclear and electronic ground state
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Atomic number (Z) - also known as Proton number - corresponds to the number of protons in the atomic nucleus (+ charge on nucleus). - also, the number of electrons in the electron cloud. Mass number Mass number = # of protons + # of neutrons
(In Periodic Table) Atomic number Atomic weight
1
H Chemical symbol of
the element
1.0080
(In other periodic table) Mass number Atomic number
4 2 He Chemical
http://en.wikipedia.org/wiki/File:Atomic_number_depiction.jpg
2
symbol of the element
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Periodic Tables of the Elements - Elements are ordered according to their atomic weights. - results in ordering according to atomic numbers. - Atomic number is actually the parameter which determines the position of an element in the periodic table, rather than the atomic weight.
Alkali metals
Atomic number Atomic weight
1
H Chemical symbol of
the element
1.0080
Noble gas Alkaline earth metals
Halogens
Periodic Law - the properties of the elements are a periodic function of their atomic numbers. Alkali metals Alkaline earth metals Halogens Lanthanides or rare earths Actinides
Group I (IA) Group II (IIA) Group VIII (VII A) Elements 58-71 Elements 90-103
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Metals - are conducting and have low electron affinities (tendency to accept an electron) - tend to give up electrons easily to form positive ions - “reducing agents” Non-metals - are non-conducting and have high electron affinities - tend to accept electrons to form negative ions - “oxidizing agents” Note: - Conduction is a flow of electrons. - Metallic conduction involves passage of electrons from one atom of a metal to another. Metallic properties - tend to increase for the elements of higher atomic number. e.g., polonium (po, element number 84) is more metallic than tellurium (Te, element number 52); Po is classified as a metal, Te as non-metal. - metallic character increases down the table. Atomic radius - increases down the periodic table, and decreases across (left to right).
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Electron Orbitals Electrons occupy regions in space of specific shape and size called orbitals. - Orbitals are energy states.
Fig 4.1 Boundary surface representation of an electron in the n=1 state of the hydrogen atom. 1s orbital. (Volume encloses 90% of the electron density. Nucleus is at the origin.) n = Principal quantum number
Fig 4.2 Boundary surface diagrams for the 2p orbitals (n=2).
Fig 4.3 Boundary surface diagrams for the 3d orbitals (n=3).
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An orbital can be defined by three quantum numbers (n, ℓ, m).
Principal quantum number, n
defines the size of the orbital shell designates the main energy level in which an electron is found. the larger the principal quantum number (n), - the larger the orbital - the further the average negative charge of the electron from the positive nucleus - the higher the energy.
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Subsidiary quantum number (or secondary quantum number), ℓ defines the shape of the orbital subshell - For a principal quantum mnumber n = 4, there are different orbital shapes corresponding to the subsidiary quantum numbers ℓ = 0, 1, 2, 3. - These shapes are often identified by letter symbols (s, p, d, f ). An s A p A d An f
orbital has ℓ = 0 and is spherical in shape (Fig. 4.1) orbital has ℓ = 1 and is dumbbell shaped (Fig, 4.2) orbital has ℓ = 2 and is still more complex in shape (Fig. 4.3) orbital has ℓ = 3 and so on.
- There is only one orientation in space for a spherical shape, and therefore only one value (0) for m, the magnetic orbital quantum number, for an s orbital. • The dumbbel-shaped p orbital has three perpendicular orientations in space (Fig. 4.2), and has three values for m, (m = +1, 0, -1). • That is, for a given principal quantum number n, there is only one s orbital, but there are three p orbitals. • For the d orbital, there are five values for m (m = +2, +1, 0, -1, -2), hence five different d orbitals. This superscript (2) says that there are two electrons in this orbital. This number (1) designates the principal level
—> 1s2 This letter (s) specifies the kind of orbital
Quantum number Orbital
Defines
Comment
Principal
n = 1, 2, 3, ......
the size of orbital shell in space
the larger the n, the larger the orbital
Subsidiary
ℓ = 0, 1, 2 ... n-1
Magnetic orbital
m = 0, ±1, ±2 …. ± ℓ
the shape of the orbital often defined by letter subshell in space symbols (s, p, d, f). ex. If n = 3, ℓ = 0, 1, 2 If n = 4, ℓ = 0, 1, 2, 3 the orientation of the If ℓ = 0, m = 0 orbital subshell in If ℓ = 1, m = 0, ±1 space If ℓ = 2, m = 0, ±1, ±2 If ℓ = 3, m = 0, ±1, ±2, ±3 8
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Table 4.2. The Orbitals of the First Four Shells Principal quantum number (Shell) n 1 2 3
4
Subsidiary quantum number (Subshell) ℓ 0 (s) 0 1 0 1 2 0 1 2 3
(s) (p) (s) (p) (d) (s) (p) (d) (f)
Magnetic orbital quantum number
Subshell Notation
Orbital per Subshell
max. # of eper subshell
Electron Configuration
0
1s
1
2
1s2 or 1s (2)
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
1 3 1 3 5 1 3 5 7
2 6 2 6 10 2 6 10 14
2s2 or 2s(2) 2p6 or 2p(6) 3s2 3p6 3d10 * 4s2 4p6 4d10 4d14
(Orbital) m
* 3d is in higher energy state than 4s, thus is filled after 4s .
Electron Configurations of the Elements The aufbau rules: 1) The atomic number tells how many electrons to distribute. 2) Electrons are placed into the orbitals of lowest energy that are available, provided that: o No more than two electrons go into the same orbital, and then only if they have opposite spin (Pauli exclusion principle) o Electrons are spread out as much as possible -retaining the same spin- when orbitals of the same sublevel are open (Hund’s rule).
Example 4.1. What is the electron configuration for an aluminum atom? Problem 4.11 (same as Example 4.1 above) (Solution) see below.
Example 4.2 What is the electron configuration of a phosphorus atom? (Solution) see below. 9
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Examples 4.2 and 4.2
Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
Example 4.1
Example 4.2
# of e-
Z=13 Al
Z=15 P
(Orbital) m 0
1s
1s2
2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2 6 2 6 10 2 6 10 14
2s2 2p6 3 s2 3p1
2s2 2p6 3 s2 3p3
1s2 2s22p6 3s2 3p1 1s2 2s22p6 3s23p3
Note: Z = atomic number = number of protons Hund’s rule of maximum multiplicity o Whenever possible, the electrons in a subshell will have the same value for the magnetic spin quantum number, ms. o Electrons are distributed among the orbitals of a subshell in a way that gives the maximum number of unpaired electrons with parallel spins. o Electrons at the same sublevel spread out among the sublevels’s oebitals as much as possible.
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Example from Example 4.2 Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
# of e-
Z=15 P
(Orbital) m 0
1s
1s2
2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2 6 2 6 10 2 6 10 14
2s2 2p6 3 s2 3Py1 3Pz1
3Px1
1s2 2s22p6 3s2 3Px1 3Py1 3Pz1
Example 4.3 Determine the electron configuration of the valence electrons of an atom of N, and compare it with that of As. (Solution) Example 4.3 Principal Magnetic orbital quantum quantum number number (Shell) (Orbital) n m 1 0 2 3
4
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
Subshell Notation
# of e-
Z=7 N
Z=33 As
1s2
1s2
1s
1s2
2
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2 6 2 6 10 2 6 10 14
2px1
2s2 2py1 2pz1
1s2 2s22px1 2py1 2pz1
2s2 2p6 3 s2 3p6 3d10 4s2 1 4px 4py1 4pz1 1s2 2s2 2p63s23p63p10 4s2 4px1 4py1 4pz1
Note: the similarity is both have filled s orbitals and three singly occupied p orbitals.
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4.1.4 The Properties of the Families of Elements Group I A - Alkali Metals -
have one s electron beyond a closed-shell configuration. tend to lose an electron very readily to give a positively charged species (cation) with a closed-shell configuration. e.g., Na -----> Na+ + ehave a low electron affinity have a low ionization energy; i.e., it doesn’t take much energy to make them lose an electron to form a cation. “Reducing agent” - give up electrons easily to other substances and are themselves oxidized.
-
e.g.,
Na + Cl -----> Na+ Cl-
An electron is transferred from sodium to chloride. The sodium atom is oxidized (REO) and the chlorine atom is reduced (GER). Since the higher atomic number elements in a family have larger electron clouds which screen the nuclear positive charge, electron affinity decreases as the atomic number increases within a family. e.g., potassium (K, atomic number Z=19) is a more powerful reducing agent than lithium (Li, Z=3). Example Principal quantum number (Shell) n 1 2 3
4
Na Magnetic orbital quantum number
Alkali Metal
Halogen
Z=11 Na
Z=17 Cl
-
Subshell Notation
# of e
0
1s
2
1s2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2 6 2 6 10(*) 2 6 10 14
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2s2 2p6 3 s1 (**)
2s2
(Orbital) m
1s2 2s22p6 3s1 Note: * 3d is in higher energy state than 4s, thus is filled after 4s . 12
2p6 3 s2
3p5 (***)
1s2 2s22p6 3s2p5
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2 6
1
0
** 1s 2s p 3s is sodium (Na , elemental sodium, very unstable, oxidation state is 0); if e- is released from 3s, Na0 Na+ 1s2 2s2p6 (sodium ion, stable), oxidation state is +1. *** 1s2 2s2p6 3s23p5 is chlorine; if 3p is filled (3p6 ), Cl Cl-1 1s2 2s2p6 3s23p6 (chloride), oxidation state is -1.
Example 4.4 Complete the following chemical reaction: Rb + F –> (Solution) Since rubidium (Rb) is in the same family (Alkali metals) as sodium, it is a reducing agent. Since fluorine (F) is the same family as chlorine, it is an oxidizing agent. Hence, an electron will trabsfer from rubidium to fluorine: Rb + F -----> Rb+ + F The rubidium atom is oxidized (REO) and the fluorine atom is reduced (GER).
Group II A - Alkaline Earth Metals Example
Berylliun (Be, atomic number 4) Alkaline Earth
Principal quantum number (Shell) n 1 2
3
4
Subshell Notation
# of e-
0
1s
2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2 6 2 6 10 * 2 6 10 14
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2s2
Magnetic orbital quantum number
Z=4 Be
(Orbital) m
1s2 2s2
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The s electrons in the second shell are easily lost to give Be . Alkaline earth metals are typically reducing agents that give up two electrons (divalent). - less powerful reducing agents than the alkali metals since the higher nuclear charge tend to hold the valence electrons more tightly. - electron affinity increases across the periodic table, and decreases down the periodic table. - reducing power decreases across the periodic table and increases down the table. - have a metallic luster and are harder than the alkali metals. - are good conductors of electricity.
Example 5.4 Complete the following chemical reaction: Ca + 2 Br —> (Solution) Calcium (Ca, atomic number 20; electron configuration, 1s2 2s22p63s2 3p6 4s2) Alkaline Earth Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
# of e-
Z=20 Ca
(Orbital) m 0
1s
1s2
2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6
2 6 2 6 (*)
2s2 2p6 3s2 3p6 4s2
1s2 2s22p63s2 3p6 4s2 Note: * 3d is in higher energy state than 4s, thus is filled after 4s .
Ca
- is an alkaline earth metal with two valence electrons in the fourth shell. - It is a reducing agent with a tendency to give up two electrons.
Br
- is like chlorine, and can accept one electron to give a closed-shell anion. - two bromine atoms can accept one electron each from one calcium atom:
Thus, Ca + 2 Br
Ca 2+ + 2 Br14
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Example 4.6 (p. 187) Assign the ionization energies in the column on the right to the elements in the column on the left: (Solution)
Problem Element
Ionization Energy
Cs
738
Z= atomic number 55 Alkali metal
Rank
Ionization Energy
4
376
Mg
503
12
Alkaline earth
1
738
Ca
590
20
Alkaline earth
2
590
Ba
376
56
Alkaline earth
3
503
lower ionization energy than alkali earth
Group VII A - The Halogens Example
Principal quantum number (Shell) n 1 2 3
4
Cl
Magnetic orbital quantum number
Subshell Notation
Alkali Metal
Halogen
# of e-
Z=11 Na
Z=17 Cl
(Orbital) m 0
1s
1s2
2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6
2 6 2 6 (*)
2s2 2p6 3 s1 (**)
2s2 2p6 3 s2 3p5 (***)
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2 p5
Note: * 3d is in higher energy state than 4s, thus is filled after 4s . ** 1s2 2s2 p6 3s1 is sodium (Na0, elemental sodium, very unstable, oxidation state is 0); if e- is released from 3s, Na0 Na+ 1s2 2s2p6 (sodium ion, stable), oxidation state is +1. 15
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*** 1s2 2s2p6 3s23p5 is chlorine; if 3p is filled (3p6 ), Cl -> Cl-1 1s2 2s2p6 3s23p6 (chloride), oxidation state is -1. Cl + e- –> Cl Halogen atoms -
are one electron short of a closed-shell configuration.
-
have high electron affinity because they have high nuclear charges which hold their valence electrons very tightly.
-
tend to accept electrons readily to give closed-shell configuration anions.
-
are oxidizing agents.
Example,
I + K --> I -
+ K+
- oxidizing power of halogens is greatest for fluorine (F, atomic number 9), and decreases down the table.
- tend to be colored. - chlorine (Cl) is a green gas, - bromine (Br) is a red volatile liquid, - iodine (I) is a deep purple, volatile solid.
Example 4.7
Which of the following chemical reactions has the greatest tendency to take place?
Cl + K —>
Cl -
+
K+ K+
Br + K
—>
Br - +
Cl + Li
—>
Cl -
Br + Li
—>
Br - +
+ Li + Li +
(Solution) Since oxidizing power decreases down the table, Cl is more powerful oxidizing agent than Br.
-
Since reducing power increases down the table, K is more powerful reducing agent than Li.
The reaction with the greatest tendency to take place is the one between the best oxidizing and best reducing agents, i.e., Cl -
+ K+
Group 0 - The Novel Gases
- have closed-shell configurations. - have low electron affinity and have high ionization potentials. - enter into chemical reactions with extreme reluctance. - only non-metallic elements that exist as single atoms in the elemental state.
Problem 4.6
Group VI is 2 electrons short of a closed shell configuration.
(Solution)
16
Cl + K
—>
10 FE_Chemistry-Atoms_F12 Example 4.17 Principal
Magnetic orbital
Subshell
quantum
quantum number
Notation
# of e-
Z=35 Br
number (Shell)
(Orbital)
n
m
1
0
1s
1s2
2
1s2
2
0
2s
2s2
2
2s2
+1, 0, -1
2p
2p6
6
2p6
0
3s
3 s2
2
3 s2
+1, 0, -1
3p
3p6
6
3p6
+2, +1, 0, -1, -2
3d
3d10
10
3d10 4s2 5 4p
3
4
0
4s
4s2
2
+1, 0, -1
4p
4p6
6
+2, +1, 0, -1, -2
4d
4d10
10
+3, +2, +1, 0, -1,-2,-3
4f
4d14
14
1s2 2s2 2p63s23p63p10 4s2 4p5
answer is d) 30 electrons.
If the 5 electrons in 4p orbitals are removed, that leaves 30 electrons.
Problem 4.18 (Solution) similar to Example 4.2 Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
# of e-
Z=15 P
(Orbital) m 0
1s
1s2
2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 10 4d 4d14
2 6 2 6 10 2 6 10 14
2s2 2p6 3 s2 3p3
1s2 2s22p6 3s2 3p3 Answer is d) 3.
Phosphorus (z = 15) has electron configuration of 1s2 2s22p6 3s2 3p3 .
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An element - is a pure substance that cannot be decomposed into other pure substances. - each element is homogeneous. - each element is composed of only one kind of atom. - each element is represented by a chemical symbol; e.g., gold is an element represented by the symbol, Au. 4.1 Atom 4.1.1 Atomic Structure and the Properties of the Elements An atom - is the smallest subdivision or particle of an element that enters into the chemical reactions of that element. - consists of an extremely dense nucleus of protons and neutrons, and very diffuse surrounding electron cloud containing enough electrons to exactly balance the nuclear charge. Nucleus - consists of protons and neutrons (electrons exist outside of nucleus) Protons - particles with a unit positive charge and a mass very close to 1 amu (amu = atomic mass unit); it has +1 charge Neutrons - neutral particles of essentially the same mass as the proton (1 amu); it has 0 charge Electrons - entities with a unit negative charge and a comparatively negligible mass (assumed to be 0 amu) - actually ~ 5.5 x 10-4 amu); it -1 charge - appear to have properties of both waves and particles Note: - For any atoms, total number of protons = total number of electrons. - Each element has unique number of protons; e.g., H has 1 proton. - A negative ion of a certain element can be formed by addition of an electron to an atom of that element. Atomic weight of an element - is the weight of an atom of that element in atomic mass units. - One atomic mass unit (amu) = 1/12 the mass of one normal C atom ~ mass of one H atom - the sum of the masses of the protons, neutrons, and electrons in the atom (Atomic weight = # of protons + # of neutrons ) -
is approximately equal to the number of protons and neutrons, with a little extra added by the electrons. An atomic weight (relative atomic mass) of an element from a specified source is the ratio of the average mass per atom of the element to 1/12 of the mass of 12C" in its nuclear and electronic ground state
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Atomic number (Z) - also known as Proton number - corresponds to the number of protons in the atomic nucleus (+ charge on nucleus). - also, the number of electrons in the electron cloud. Mass number Mass number = # of protons + # of neutrons
(In Periodic Table) Atomic number Atomic weight
1
H Chemical symbol of
the element
1.0080
(In other periodic table) Mass number Atomic number
4 2 He Chemical
http://en.wikipedia.org/wiki/File:Atomic_number_depiction.jpg
2
symbol of the element
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Periodic Tables of the Elements - Elements are ordered according to their atomic weights. - results in ordering according to atomic numbers. - Atomic number is actually the parameter which determines the position of an element in the periodic table, rather than the atomic weight.
Alkali metals
Atomic number Atomic weight
1
H Chemical symbol of
the element
1.0080
Noble gas Alkaline earth metals
Halogens
Periodic Law - the properties of the elements are a periodic function of their atomic numbers. Alkali metals Alkaline earth metals Halogens Lanthanides or rare earths Actinides
Group I (IA) Group II (IIA) Group VIII (VII A) Elements 58-71 Elements 90-103
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Metals - are conducting and have low electron affinities (tendency to accept an electron) - tend to give up electrons easily to form positive ions - “reducing agents” Non-metals - are non-conducting and have high electron affinities - tend to accept electrons to form negative ions - “oxidizing agents” Note: - Conduction is a flow of electrons. - Metallic conduction involves passage of electrons from one atom of a metal to another. Metallic properties - tend to increase for the elements of higher atomic number. e.g., polonium (po, element number 84) is more metallic than tellurium (Te, element number 52); Po is classified as a metal, Te as non-metal. - metallic character increases down the table. Atomic radius - increases down the periodic table, and decreases across (left to right).
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Electron Orbitals Electrons occupy regions in space of specific shape and size called orbitals. - Orbitals are energy states.
Fig 4.1 Boundary surface representation of an electron in the n=1 state of the hydrogen atom. 1s orbital. (Volume encloses 90% of the electron density. Nucleus is at the origin.) n = Principal quantum number
Fig 4.2 Boundary surface diagrams for the 2p orbitals (n=2).
Fig 4.3 Boundary surface diagrams for the 3d orbitals (n=3).
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An orbital can be defined by three quantum numbers (n, ℓ, m).
Principal quantum number, n
defines the size of the orbital shell designates the main energy level in which an electron is found. the larger the principal quantum number (n), - the larger the orbital - the further the average negative charge of the electron from the positive nucleus - the higher the energy.
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Subsidiary quantum number (or secondary quantum number), ℓ defines the shape of the orbital subshell - For a principal quantum mnumber n = 4, there are different orbital shapes corresponding to the subsidiary quantum numbers ℓ = 0, 1, 2, 3. - These shapes are often identified by letter symbols (s, p, d, f ). An s A p A d An f
orbital has ℓ = 0 and is spherical in shape (Fig. 4.1) orbital has ℓ = 1 and is dumbbell shaped (Fig, 4.2) orbital has ℓ = 2 and is still more complex in shape (Fig. 4.3) orbital has ℓ = 3 and so on.
- There is only one orientation in space for a spherical shape, and therefore only one value (0) for m, the magnetic orbital quantum number, for an s orbital. • The dumbbel-shaped p orbital has three perpendicular orientations in space (Fig. 4.2), and has three values for m, (m = +1, 0, -1). • That is, for a given principal quantum number n, there is only one s orbital, but there are three p orbitals. • For the d orbital, there are five values for m (m = +2, +1, 0, -1, -2), hence five different d orbitals. This superscript (2) says that there are two electrons in this orbital. This number (1) designates the principal level
—> 1s2 This letter (s) specifies the kind of orbital
Quantum number Orbital
Defines
Comment
Principal
n = 1, 2, 3, ......
the size of orbital shell in space
the larger the n, the larger the orbital
Subsidiary
ℓ = 0, 1, 2 ... n-1
Magnetic orbital
m = 0, ±1, ±2 …. ± ℓ
the shape of the orbital often defined by letter subshell in space symbols (s, p, d, f). ex. If n = 3, ℓ = 0, 1, 2 If n = 4, ℓ = 0, 1, 2, 3 the orientation of the If ℓ = 0, m = 0 orbital subshell in If ℓ = 1, m = 0, ±1 space If ℓ = 2, m = 0, ±1, ±2 If ℓ = 3, m = 0, ±1, ±2, ±3 8
10 FE_Chemistry-Atoms_F12
Table 4.2. The Orbitals of the First Four Shells Principal quantum number (Shell) n 1 2 3
4
Subsidiary quantum number (Subshell) ℓ 0 (s) 0 1 0 1 2 0 1 2 3
(s) (p) (s) (p) (d) (s) (p) (d) (f)
Magnetic orbital quantum number
Subshell Notation
Orbital per Subshell
max. # of eper subshell
Electron Configuration
0
1s
1
2
1s2 or 1s (2)
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
1 3 1 3 5 1 3 5 7
2 6 2 6 10 2 6 10 14
2s2 or 2s(2) 2p6 or 2p(6) 3s2 3p6 3d10 * 4s2 4p6 4d10 4d14
(Orbital) m
* 3d is in higher energy state than 4s, thus is filled after 4s .
Electron Configurations of the Elements The aufbau rules: 1) The atomic number tells how many electrons to distribute. 2) Electrons are placed into the orbitals of lowest energy that are available, provided that: o No more than two electrons go into the same orbital, and then only if they have opposite spin (Pauli exclusion principle) o Electrons are spread out as much as possible -retaining the same spin- when orbitals of the same sublevel are open (Hund’s rule).
Example 4.1. What is the electron configuration for an aluminum atom? Problem 4.11 (same as Example 4.1 above) (Solution) see below.
Example 4.2 What is the electron configuration of a phosphorus atom? (Solution) see below. 9
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Examples 4.2 and 4.2
Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
Example 4.1
Example 4.2
# of e-
Z=13 Al
Z=15 P
(Orbital) m 0
1s
1s2
2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2 6 2 6 10 2 6 10 14
2s2 2p6 3 s2 3p1
2s2 2p6 3 s2 3p3
1s2 2s22p6 3s2 3p1 1s2 2s22p6 3s23p3
Note: Z = atomic number = number of protons Hund’s rule of maximum multiplicity o Whenever possible, the electrons in a subshell will have the same value for the magnetic spin quantum number, ms. o Electrons are distributed among the orbitals of a subshell in a way that gives the maximum number of unpaired electrons with parallel spins. o Electrons at the same sublevel spread out among the sublevels’s oebitals as much as possible.
10
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Example from Example 4.2 Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
# of e-
Z=15 P
(Orbital) m 0
1s
1s2
2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2 6 2 6 10 2 6 10 14
2s2 2p6 3 s2 3Py1 3Pz1
3Px1
1s2 2s22p6 3s2 3Px1 3Py1 3Pz1
Example 4.3 Determine the electron configuration of the valence electrons of an atom of N, and compare it with that of As. (Solution) Example 4.3 Principal Magnetic orbital quantum quantum number number (Shell) (Orbital) n m 1 0 2 3
4
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
Subshell Notation
# of e-
Z=7 N
Z=33 As
1s2
1s2
1s
1s2
2
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2 6 2 6 10 2 6 10 14
2px1
2s2 2py1 2pz1
1s2 2s22px1 2py1 2pz1
2s2 2p6 3 s2 3p6 3d10 4s2 1 4px 4py1 4pz1 1s2 2s2 2p63s23p63p10 4s2 4px1 4py1 4pz1
Note: the similarity is both have filled s orbitals and three singly occupied p orbitals.
11
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4.1.4 The Properties of the Families of Elements Group I A - Alkali Metals -
have one s electron beyond a closed-shell configuration. tend to lose an electron very readily to give a positively charged species (cation) with a closed-shell configuration. e.g., Na -----> Na+ + ehave a low electron affinity have a low ionization energy; i.e., it doesn’t take much energy to make them lose an electron to form a cation. “Reducing agent” - give up electrons easily to other substances and are themselves oxidized.
-
e.g.,
Na + Cl -----> Na+ Cl-
An electron is transferred from sodium to chloride. The sodium atom is oxidized (REO) and the chlorine atom is reduced (GER). Since the higher atomic number elements in a family have larger electron clouds which screen the nuclear positive charge, electron affinity decreases as the atomic number increases within a family. e.g., potassium (K, atomic number Z=19) is a more powerful reducing agent than lithium (Li, Z=3). Example Principal quantum number (Shell) n 1 2 3
4
Na Magnetic orbital quantum number
Alkali Metal
Halogen
Z=11 Na
Z=17 Cl
-
Subshell Notation
# of e
0
1s
2
1s2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2 6 2 6 10(*) 2 6 10 14
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2s2 2p6 3 s1 (**)
2s2
(Orbital) m
1s2 2s22p6 3s1 Note: * 3d is in higher energy state than 4s, thus is filled after 4s . 12
2p6 3 s2
3p5 (***)
1s2 2s22p6 3s2p5
10 FE_Chemistry-Atoms_F12 2
2 6
1
0
** 1s 2s p 3s is sodium (Na , elemental sodium, very unstable, oxidation state is 0); if e- is released from 3s, Na0 Na+ 1s2 2s2p6 (sodium ion, stable), oxidation state is +1. *** 1s2 2s2p6 3s23p5 is chlorine; if 3p is filled (3p6 ), Cl Cl-1 1s2 2s2p6 3s23p6 (chloride), oxidation state is -1.
Example 4.4 Complete the following chemical reaction: Rb + F –> (Solution) Since rubidium (Rb) is in the same family (Alkali metals) as sodium, it is a reducing agent. Since fluorine (F) is the same family as chlorine, it is an oxidizing agent. Hence, an electron will trabsfer from rubidium to fluorine: Rb + F -----> Rb+ + F The rubidium atom is oxidized (REO) and the fluorine atom is reduced (GER).
Group II A - Alkaline Earth Metals Example
Berylliun (Be, atomic number 4) Alkaline Earth
Principal quantum number (Shell) n 1 2
3
4
Subshell Notation
# of e-
0
1s
2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2 6 2 6 10 * 2 6 10 14
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 4d10 4d14
2s2
Magnetic orbital quantum number
Z=4 Be
(Orbital) m
1s2 2s2
13
10 FE_Chemistry-Atoms_F12 2+
The s electrons in the second shell are easily lost to give Be . Alkaline earth metals are typically reducing agents that give up two electrons (divalent). - less powerful reducing agents than the alkali metals since the higher nuclear charge tend to hold the valence electrons more tightly. - electron affinity increases across the periodic table, and decreases down the periodic table. - reducing power decreases across the periodic table and increases down the table. - have a metallic luster and are harder than the alkali metals. - are good conductors of electricity.
Example 5.4 Complete the following chemical reaction: Ca + 2 Br —> (Solution) Calcium (Ca, atomic number 20; electron configuration, 1s2 2s22p63s2 3p6 4s2) Alkaline Earth Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
# of e-
Z=20 Ca
(Orbital) m 0
1s
1s2
2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6
2 6 2 6 (*)
2s2 2p6 3s2 3p6 4s2
1s2 2s22p63s2 3p6 4s2 Note: * 3d is in higher energy state than 4s, thus is filled after 4s .
Ca
- is an alkaline earth metal with two valence electrons in the fourth shell. - It is a reducing agent with a tendency to give up two electrons.
Br
- is like chlorine, and can accept one electron to give a closed-shell anion. - two bromine atoms can accept one electron each from one calcium atom:
Thus, Ca + 2 Br
Ca 2+ + 2 Br14
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Example 4.6 (p. 187) Assign the ionization energies in the column on the right to the elements in the column on the left: (Solution)
Problem Element
Ionization Energy
Cs
738
Z= atomic number 55 Alkali metal
Rank
Ionization Energy
4
376
Mg
503
12
Alkaline earth
1
738
Ca
590
20
Alkaline earth
2
590
Ba
376
56
Alkaline earth
3
503
lower ionization energy than alkali earth
Group VII A - The Halogens Example
Principal quantum number (Shell) n 1 2 3
4
Cl
Magnetic orbital quantum number
Subshell Notation
Alkali Metal
Halogen
# of e-
Z=11 Na
Z=17 Cl
(Orbital) m 0
1s
1s2
2
1s2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6
2 6 2 6 (*)
2s2 2p6 3 s1 (**)
2s2 2p6 3 s2 3p5 (***)
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2 p5
Note: * 3d is in higher energy state than 4s, thus is filled after 4s . ** 1s2 2s2 p6 3s1 is sodium (Na0, elemental sodium, very unstable, oxidation state is 0); if e- is released from 3s, Na0 Na+ 1s2 2s2p6 (sodium ion, stable), oxidation state is +1. 15
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*** 1s2 2s2p6 3s23p5 is chlorine; if 3p is filled (3p6 ), Cl -> Cl-1 1s2 2s2p6 3s23p6 (chloride), oxidation state is -1. Cl + e- –> Cl Halogen atoms -
are one electron short of a closed-shell configuration.
-
have high electron affinity because they have high nuclear charges which hold their valence electrons very tightly.
-
tend to accept electrons readily to give closed-shell configuration anions.
-
are oxidizing agents.
Example,
I + K --> I -
+ K+
- oxidizing power of halogens is greatest for fluorine (F, atomic number 9), and decreases down the table.
- tend to be colored. - chlorine (Cl) is a green gas, - bromine (Br) is a red volatile liquid, - iodine (I) is a deep purple, volatile solid.
Example 4.7
Which of the following chemical reactions has the greatest tendency to take place?
Cl + K —>
Cl -
+
K+ K+
Br + K
—>
Br - +
Cl + Li
—>
Cl -
Br + Li
—>
Br - +
+ Li + Li +
(Solution) Since oxidizing power decreases down the table, Cl is more powerful oxidizing agent than Br.
-
Since reducing power increases down the table, K is more powerful reducing agent than Li.
The reaction with the greatest tendency to take place is the one between the best oxidizing and best reducing agents, i.e., Cl -
+ K+
Group 0 - The Novel Gases
- have closed-shell configurations. - have low electron affinity and have high ionization potentials. - enter into chemical reactions with extreme reluctance. - only non-metallic elements that exist as single atoms in the elemental state.
Problem 4.6
Group VI is 2 electrons short of a closed shell configuration.
(Solution)
16
Cl + K
—>
10 FE_Chemistry-Atoms_F12 Example 4.17 Principal
Magnetic orbital
Subshell
quantum
quantum number
Notation
# of e-
Z=35 Br
number (Shell)
(Orbital)
n
m
1
0
1s
1s2
2
1s2
2
0
2s
2s2
2
2s2
+1, 0, -1
2p
2p6
6
2p6
0
3s
3 s2
2
3 s2
+1, 0, -1
3p
3p6
6
3p6
+2, +1, 0, -1, -2
3d
3d10
10
3d10 4s2 5 4p
3
4
0
4s
4s2
2
+1, 0, -1
4p
4p6
6
+2, +1, 0, -1, -2
4d
4d10
10
+3, +2, +1, 0, -1,-2,-3
4f
4d14
14
1s2 2s2 2p63s23p63p10 4s2 4p5
answer is d) 30 electrons.
If the 5 electrons in 4p orbitals are removed, that leaves 30 electrons.
Problem 4.18 (Solution) similar to Example 4.2 Principal quantum number (Shell) n 1 2 3
4
Magnetic orbital quantum number
Subshell Notation
# of e-
Z=15 P
(Orbital) m 0
1s
1s2
2
1s2
0 +1, 0, -1 0 +1, 0, -1 +2, +1, 0, -1, -2 0 +1, 0, -1 +2, +1, 0, -1, -2 +3, +2, +1, 0, -1,-2,-3
2s 2p 3s 3p 3d 4s 4p 4d 4f
2s2 2p6 3 s2 3p6 3d10 4s2 4p6 10 4d 4d14
2 6 2 6 10 2 6 10 14
2s2 2p6 3 s2 3p3
1s2 2s22p6 3s2 3p3 Answer is d) 3.
Phosphorus (z = 15) has electron configuration of 1s2 2s22p6 3s2 3p3 .
17
