Chapter 14 – Chemical Periodicity
Reading Assignment C14!
1. Read C14 pp. 390-408 and while reading, continue building your science vocabulary
table that includes all terms in bold face type and all terms you are unfamiliar with or
unsure of. Your vocabulary table needs to include three things:
1.
The term
2.
The definition
3.
A picture of what the term means to you.
2. On page 408, organize the Concept Map into a form that correctly connects the major
ideas of the chapter.
I.
Chemical Periodicity.
a. History of the Periodic Table
i. Sb, C, Cu, Au, Fe, Pb, Hg, Ag, S, Sn were only elements known
before 1250- As comes slightly later.
ii. Prior to late 1700’s only 24 elements were known.
iii. Johann Wolfgang Döbereiner- 1817- Germany
1. Triads - groups of 3 elements that share similar properties:
a. Li, Na, K
b. Cl, Br, I
2. John Newlands - 1867- England
a. Law of Octaves - when the elements are placed in
order of atomic weights, a cycle of properties is
repeated with every eight elements.
3. Dmitri Mendeleev- 1869 Russia
a. Devised first periodic table - coupled properties of
elements and organization of increasing atomic
mass.
b. Law - “Elements arranged according to their atomic
masses present a clear periodicity of properties.
c. Groupings- organized elements into an eightcolumn format.
d. Some spaces had to be skipped for not-yet
discovered elements.
i. Eka-Silicon – Germaninum – similar
properties to that of Silicon & Tin
ii. Gallium and Scandium
Predicted properties for Mendeleev’s Eka-Silicon and properties of
Germanium:
Atomic
Oxide
Chloride
Element
Density
Weight
formula
formula
Eka-Silicon
72
5.5 g/cm3
EsO2
EsCl4
(predicted
1871)
Germanium
72.59
5.32 g/cm3
GeO2
GeCl4
(discovered
1886)
II.
e. Problems with Mendeleev’s Periodic Table
i. Following increasing atomic mass,
sometimes elements had to be reversed.
1. Ex: Ni & Co, I & Te, Ar & K
ii. Newly discovered elements had no spaces
available.
1. Ex: Holmium and Samarium
iii. Elements in the same group were sometimes
quite different in their reactivity
1. Ex: Group I- alkali metals &
“Coinage” metals.
4. Lother Meyer - c. 1870 Germany
a. Worked on periodic relationship similar to that of
Mendeleev. Acknowledged that Mendeleev had
original idea.
5. Henry Moseley - 1913 England
a. Atomic number- the number of protons found in the
nucleus of a specific element.
b. Current periodic table utilizes increasing atomic
numbers instead of atomic masses.
c. Moseley worked on High Frequency Spectra of the
Elements.
Modern Periodic Table
a. Periodic Law - When the elements are arranged in order of increasing
atomic number, there is a periodic pattern in their physical and chemical
properties.
i. Elements
1. Identified by:
a. Name
b. Symbol
c. Atomic number
d. Atomic mass
2.
Characterized by:
III.
a. Physical Properties- Boiling/Melting Points,
Density, Color, Crystalline Structure, State of
Matter, etc…
b. Chemical Properties- Oxidation States, Acid/Basic
Properties, Ionization Energy, Electron Affinity,
etc…
ii. Periods- (Rows) - represents the number of energy levels.
iii. Columns- (Families/Groups)- representative of elements with
similar properties.
1. Analagous to the number of electrons in outer energy
level(s)
2. Representative Elements (A groups) IA- VIIIA- fills s and
p orbitals:
a. IA. Alkali Metals
b. IIA. Alkaline Earth Metals
c. IIIA. Aluminum Family
d. IVA. Carbon Family
e. VA. Pnicogens (Nitrogen Family)
f. VIA. Chalcogens
g. VIIA. Halogens
h. VIIIA. Noble Gases
i. Completely filled outer s and p orbitals.
These demonstrate chemical stability.
3. Transition Metals - (B groups) - fills outermost s (ns) and
prior d orbitals (n-1d).
4. Inner Transition Metals - fills outermost s orbital (ns) and
twice prior f orbitals (n-2f). (a.k.a. Rare Earth Metals).
a. Lanthanide Series. The elements in the 4f orbitals
(top row of the inner transition metals).
b. Actinide Series. The elements in the 5f orbitals
(botton row of the inner transition metals).
5. Line of Demarkation: Staggered line used to separate
metals from nonmetals.
a. Metals - elements to the left of the line.
b. Nonmetals - elements to the right of the line.
c. Metalloids/weak-metals - elements along the line.
Periodic Properties of the Elements - The properties are based on the electrons
and their positions.
a. Factors that affect the properties:
i. The number of valence electrons.
ii. The magnitude of the nuclear charge (Z) and the total number of
electrons surrounding the nucleus.
iii. The number of filled shells lying between the nucleus and the
valence shell.
iv. The distances of the electrons in the various shells from each other
and from the nucleus.
1. Atomic Radius
a. Ways to measure the radius of an atom:
i. Covalent radius - ½ distance from nuclei of
2 identical atoms joined by a single covalent
bond.
ii. Van der Waals radius - ½ distance from
nuclei of 2 atoms of neighboring molecules.
iii. Metallic radius - ½ distance from nuclei of 2
atoms in a solid metal
iv. Atomic Radius - based on the quantum
model. Theoretical/Mathematical approach.
b. Trends in the periodic table:
i. Period- radius decreases from left to rightincrease in (Z) with same number of energy
levels.
ii. Group- radius increases from top to bottom increase in the number of energy levels
(Principal Q.N. increases).
The atomic radii for the elements in the first 3 energy levels
b. Effective Nuclear Charge (Zeff ) - dependent upon (Z) and the shielding
effect of other electrons.
i. Shielding - interior orbitals that contain electrons shield the
attraction of the nucleus on the valence electrons.
ii. Slater’s Rule (Zeff = Z –σ) where σ is the shielding factor for
valence electrons.
1. Calculating σ for:
a. Valence electrons in s and p type orbitals.
i. (ns & np) electrons shield at 35%
ii. (n-1) orbitals, these shield at 85%
iii. (n-2) orbitals, these shield at 100%
b. Valence electrons in d and f type orbitals.
i. (nd, nf) shield at 35%
ii. Higher orbitals (n+1) shield at 0% (ns, np).
iii. s and p electrons in the same energy level
(ns & np)and lower energy levels (n-1&
et.al.) shield at 100
2. Ionization Energy - The amount of energy needed to
remove an electron from a gaseous atom. Production of a
positive ion.
a. Ion- an atom which has gained or lost electronsdependent upon the Effective Nuclear Charge.
i. cation- positive ion- due to a loss of
electrons.
ii. anion- negative ion- due to a gain of
electrons.
b. Trends:
i. Groups - First Ionization Energies decrease
from top to bottom.
ii. Periods- General- First Ionization energies
increase from left to right - some exceptions
occur:
1. Due to shielding of ns on 1st electron
in a np orbital (B, Al, Ga).
2. Losing 1 paired electron is less than
losing a parallel electron (O, S, Se)
3. These exceptions fail at higher
energy levels.
1st Ionization energy for the elements in the first 3 energy levels
c. Successive Ionization Energies
i. First Ionization- removing the first valence
electron.
ii. Second Ionization- removing the second
valence electron.
iii. Third Ionization- removing the third valence
electron.
1. Removing successive electrons
reduces the shielding factor but
maintains Z. This increases the Zeff
on the remaining electrons.
3. Electron Affinity - The energy change that accompanies the
addition of an electron to a gaseous atom.
a. For most atoms, energy is released when an electron
is gained. This is seen as a negative energy change.
b. General Trend:
i. Atoms that have a high ionization energy,
will typically have a larger negative electron
affinity.
ii. Noble Gases have a positive electron affinity
because it would take energy to add
electrons to these atoms.
4. Ion Size.
a. Cations are always smaller than their neutral atom
counterpart.
i. Losing electrons may: lose valence shells
and/or increase zeff values.
b. Anions are always bigger than their neutral atom
counterpart.
i.
ii.
Gaining electrons decrease zeff values.
Trends:
1. Period- Ion size decreases from left
to right, (cation/anion specific).
2. Groups- Ion size increases from top
to bottom.
3. Isoelectric species- atoms and ions
that have the same electron
configurations
a. (ex. N-3, O-2, F-1, Ne, Na+1,
Mg +2, Al+3).
4. Octet Rule- All atoms strive to have
full valence shells (typically 8
electrons in the ns & np orbitals)
exception is 1st energy level (no p
orbital: 2 electrons)
5. Oxidation Number- the charge of the
stable ion after gaining/losing
electrons. Net difference between
protons & electrons.
5. Electronegativity - The attraction that an atom has on an
electron when it is chemically combined with another atom.
a. Trends:
i. Periods- electronegativities increase from
left to right
ii. Groups- decrease from top to bottom.