Download Chapter 8: Periodic Relationships Among the Elements

Chapter 8: Periodic Relationships Among the Elements
Problems: 1-3, 5, 8, 15-16, 27-29, 31-41, 43-46, 49, 51-52, 55-56, 60-62, 66, 72, 74, 91, 128
8.1 Development of the Periodic Table
Dimitri Mendeleev proposed that elements should be arranged based on the
periodic recurrence of their properties, which generally matched ordering them in
terms of increasing atomic mass
→ Periodic Table originally arranged with elements in order of increasing atomic
mass
Henry G.J. Moseley’s high-energy X-ray radiation experiments of atomic nuclei
→ Repeating properties of elements more clearly reflected by the arrangement
of elements according to increasing atomic number
→ Periodic Table’s arrangement today
Trends for increasing atomic mass are identical with those forincreasing
atomic number, except for Ni & Co, Ar & K, Te & I.
Neils Bohr’s introduction of electron energy levels
ٛ → Periodic Table’s shape
– Indicates filling of electron orbitals and element’s electron configuration
8.2
PERIODIC CLASSIFICATION OF THE ELEMENTS
Main-Group (Representative or A Group) Elements
Those elements in groups 1, 2, 13, 14, 15, 16, 17, 18 (or IA to VIIIA)
– Group 1 or IA: alkali metals
– Group 2 or IIA: alkaline earth metals
– Group 17 or VIIA: halogens
– Group 18 or VIIIA: noble gases (because they are all unreactive gases)
Transition Metals (or B Group Elements)
– Elements in Groups 3 to 12 (middle of the Periodic Table)
Inner Transition Elements (beneath the main body of Periodic Table)
– lanthanide series: Ce-Lu, also called rare earth metals because they make up
<0.005% of the earth's crust
– actinide series: Th-Lr, also called transuranium elements, generally all manmade, exist for very short periods of time before decaying to other elements
CHEM 150: Chapter 8 Notes
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valence electrons: Electrons in the outermost shell
– Since atoms want filled electron shells to be most stable, they’ll combine with
other atoms with unfilled shells (gaining or losing e–s) to get stability.
→ Valence electrons lead to chemical bonds and reactions between atoms
For Main Group (A) elements, Group # → # of valence electrons
– Elements in Group IA: Each has 1 valence electron
– Elements in Group IIA: Each has 2 valence electrons
– Elements in Group IIIA: Each has 3 valence electrons
– Elements in Group IVA: Each has 4 valence electrons
– Elements in Group VA: Each has 5 valence electrons
– Elements in Group VIA: Each has 6 valence electrons
– Elements in Group VIIA: Each has 7 valence electrons
– Elements in Group VIIIA: Each has 8 valence electrons
Example: Indicate the number of valence electrons for each element below:
Mg: ____ valence electrons
N: ____ valence electrons
Br: ____ valence electrons
Al: ____ valence electrons
Rb: ____ valence electrons
Si: ____ valence electrons
Se:
Xe: ____ valence electrons
____ valence electrons
Electron Configurations of Cations and Anions
Ions of the Main-Group Elements
– Representative elements generally form ions—ie. gain or lose electrons—
to achieve a noble gas electron configuration
→ Ions from representative metals are usually isoelectronic with—
i.e. have the same electron configuration as—one of the noble gases!
For IONS, one must account for the loss or gain of electrons:
# electrons = atomic # – (charge = change in # of valence electrons)
CHEM 150: Chapter 8 Notes
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Or you can simply use the Periodic Table
– Find out with which element the ion is isoelectronic
– Move to the left for electrons lost
– Move to the right for electrons gained
→ write the electron configuration for that element
Example 1: Fill in the blanks for the following ions:
Ion
Isoelectronic
with what
atom?
Electron
Configuration
using core
notation
Ion
Na+
Ba+2
P–3
O–2
Al+3
I–
Se–2
Ti+4
Electron
Isoelectronic
Configuration
with what
using core
atom?
notation
Cations from Transition Metals, Sn, Pb
– Transition metals lose s electrons before the d electrons when forming cations
Ion
Isoelectronic
with what
atom?
Electron
Configuration
using core
notation
Ion
Ag+
Cu+
Zn+2
Pb+4
Co+3
Fe+2
Sn+2
Fe+3
Electron
Isoelectronic
Configuration
with what
using core
atom?
notation
Example: Given the electron configurations of Fe+2 and Fe+3, which might be
more stable? Explain.
CHEM 150: Chapter 8 Notes
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8.3 PERIODIC VARIATIONS IN PHYSICAL PROPERTIES
Atomic radius: distance from nucleus to outermost electrons
– Increases down a group: More p+, n, and e– ٛ → bigger radius
– Decreases from left to right along a period.
Effective Nuclear Charge:
– positive charge an electron experiences because of the positive charge
from protons in the nucleus minus any shielding due to electrons closer to
the nucleus
Consider the atoms of Ti and Ni:
Thus, the higher the effective nuclear charge, the smaller the atomic radius!
Atomic Radius trend:
– Trend from top to bottom → like a snowman
– Trend from left to right → like a snowman that fell to the right
CHEM 150: Chapter 8 Notes
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IONIC RADIUS: distance from nucleus to outermost electrons in an ion
cation radius < neutral atom radius < anion radius
(loses e-s)
(gains e-s)
Cl atom
Na atom
8.4
Na+ ion
Cl – ion
IONIZATION ENERGY
First Ionization Energy:
Energy necessary to remove first electron from a
neutral atom in gaseous state to form negatively
charged ion.
First Ionization Energy TRENDS
– Decreases down a group:
Bigger the atom, the further away e–s are from +vely charged nucleus
→ e–s held less tightly and are more easily removed
– Increases from left to right along a period:
– Elements with fewer (1–3) valence e–s can more easily give up e–s to
gain noble gas configuration (stability)
– Elements with more (4–7) valence e–s can more easily gain e–s togain
noble gas configuration (stability)
Trend from top to bottom → like an upside-down snowman
Trend from left to right →ٛ like a upside-down snowman that fell to the right
CHEM 150: Chapter 8 Notes
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Variations in Successive Ionization Energies
– Recognize that it becomes more difficult to remove electrons from stable ions.
– See Table 8.2 on p. 320
– For example, the ionization energy to remove the first electron from Li is
much smaller than the any successive ionization energy because electrons
are now removed from a stable ion (with a positive charge AND at times a
noble gas electron configuration).
We can indicate first and successive ionization energies in the following way:
First ionization energy = IE1
Second ionization energy = IE2
Third ionization energy = IE3
Ex. 1:Between which two ionization energies (IE1 and IE2, IE2 and IE3, etc.)
elements? Explain.
a. Mg: Between _____ and _____
Explain why:
b. K: Between _____ and _____
Explain why:
c. Al: Between _____ and _____
Explain why:
Ex. 2:This 3rd period element has a large jump between IE5 and IE6: _____
Explain how you can identify the element.
CHEM 150: Chapter 8 Notes
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8.5
ELECTRON AFFINITY
– A quantitative measure of the an atoms ability to accept an electron
– The higher the electron affinity, the more likely an atom can gain an
electron
– F has the highest electron affinity, so the closer an atom is to F, the
higher its electron affinity
– Exceptions: H’s electron affinity is between B & C, and because
Noble gases are stable, they have the lowest electron affinities
Electron Affinity follows the same snowman trend as atomic radius:
8.6
VARIATIONS IN CHEMICAL PROPERTIES OF REP. ELEMENTS
Metallic character:
– Decreases from left to right along a period:
Metals concentrated on left-hand side of P.T., nonmetals on right-hand side
– Increases down a group: Looking at groups IVA and VA, go from nonmetals
(C & N) to semimetals (Si & As) to metals (Sn & Bi)
→ Same snowman trends as for atomic radius and electron affinity!
Properties of Oxides Across a Period
Nonmetal Oxides
– Nonmetal oxides react with water to form metal hydroxides
P4O10 (s) +
6 H2O (l) → 4 H3PO4 (aq)
→ Nonmetal oxides are often called acidic oxides
Note:
You generally must be told what the ternary oxyacid (acid containing
oxygen as well as H) product is for reactions of nonmetals oxides
with water.
CHEM 150: Chapter 8 Notes
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Metal Oxides
– Some metal oxides react with water to form metal hydroxides:
Na2O (s) + H2O (l) → 2 NaOH (aq)
– Insoluble metal oxides react with acids in the same way as metal
hydroxides:
MgO (s) + HCl (aq) → H2O (l) + MgCl2 (aq)
Note:
The products are water and a salt like in other acid-base
neutralization reactions; thus, MgO acts like Mg(OH)2 with an acid.
→ Metal oxides are often called basic oxides
Ex. 1:Write a balanced chemical equation for each of the following sets of
reactants, then identify the oxide as acidic or basic:
(Note: The most common oxyacids are HNO3(aq) and H2SO4(aq).)
a.
K2O (s)
+
H2O (l) →
Thus, K2O is _______ oxide. (Circle one)
b.
N2O5 (s)
+
SO3 (s)
+
BaO (s)
+
a basic
an acidic
a basic
an acidic
a basic
H2O (l) →
Thus, BaO is _______ oxide. (Circle one)
CHEM 150: Chapter 8 Notes
an acidic
H2O (l) →
Thus, SO3 is _______ oxide. (Circle one)
d.
a basic
H2O (l) →
Thus, N2O5 is _______ oxide. (Circle one)
c.
an acidic
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