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Chapter 4
The Periodic Table
4.1 Element Organization
Patterns in Element Properties


groups of elements have much in common
discovery (invention) of periodic table

1861
4.1 Element Organization
Patterns in Element Properties

discovery (invention) of periodic table

1865: John Newlands (English chemist)
noticed repeating chemical and physical
properties every eight elements when arranged
by increasing atomic mass
 thought to be absurd
4.1 Element Organization
Patterns in Element Properties

1869: Dmitri Mendeleev made first periodic
table
 arranged elements by increasing atomic
mass
 switched problem elements
 made predictions for gaps
 he was right
 nominated for Nobel prize
4.1 Element Organization
Patterns in Element Properties

Henry Moseley (English chemist) found that
pattern of table was due to atomic number,
not atomic mass
4.1 Element Organization
The Periodic Law: when elements are
arranged by atomic #, elements with
similar properties appear at regular
intervals

how convenient!
Periodic Table
Symbols in Table
4.1 Element Organization
The Periodic Law

similar properties are due to similar
electron configuration

each column (group) contains elements with
the same number of electrons in the outer
energy level (valence electrons)
 valence electrons participate in reactions
(bonding)
4.1 Element Organization
The Periodic Law

similar properties are due to similar
electron configuration

each row (period) contains elements with the
same number of occupied energy levels
Main-Group Elements
4.2 Table Tour
Main-Group Elements



elements in groups 1,2,13-18
s- and p- blocks
electron configuration varies consistently

examples
4.2 Table Tour
Main-Group Elements

Group 1: alkali metals
react with water to make alkaline (basic)
solutions
 examples:
 very reactive; one valence e
 not found as pure elements
 soft, shiny (on a fresh surface), good
conductors of electricity

4.2 Table Tour
Main-Group Elements

Group 2: alkaline-earth metals
also very reactive, but less than Group 1
 usually found as compounds
 harder, higher melting points than alkali metals

4.2 Table Tour
Main-Group Elements

Group 17: halogens
halogen means “salt maker”
 very reactive; have seven valence electrons
 react readily with alkali metals
 wide range of properties (e.g., gases, liquids,
solids)

4.2 Table Tour
Main-Group Elements

Group 18: noble gases
unreactive; full set of electrons in outer energy
level
 used to be called inert, but xenon was found to
react
 good for electric lights

4.2 Table Tour
Main-Group Elements

Hydrogen
about 75% of the universe
 behaves unlike any other element
 reactive
 major component in organic compounds
(carbs, proteins, lipids, nucleic acids)

Metals
4.2 Table Tour
Metals

excellent conductors of electricity


100 000 x better than most conductive
nonmetal
excellent conductors of heat
4.2 Table Tour
Metals



often ductile (can be squeezed into a wire)
often malleable (can be hammered into a
sheet)
often shiny
4.2 Table Tour
Metals

Transition Metals (Groups 3-12)
d- block
 electron configuration is not identical in each
group
 do not always lose the same numbers of
valence electrons in each reaction (depends
on other element)

4.2 Table Tour
Metals

Transition Metals (Groups 3-12)
can be found as pure elements (e.g., gold,
platinum)
 good conductors (both heat and electricity)
 ductile and malleable

4.2 Table Tour
Metals

Lanthanides and Actinides
f- block
 lanthanides: atomic numbers follow lanthanum
 actinides: atomic numbers follow actinium
 shiny metals

4.2 Table Tour
Metals

Lanthanides and Actinides
reactive like alkaline-earth metals
 all actinides are radioactive (e.g., uranium)

4.2 Table Tour
Metals


melting points vary widely
metals can be mixed with other metals to
form alloys

examples: brass (copper and zinc), sterling
silver (silver and copper), steel (iron, carbon,
manganese, nickel, chromium)
4.3 Trends in Periodic Table
Trend: predictable change in a
particular direction
Ionization Energy


Ionization: removing an electron from an
atom or ion
Ionization energy: energy required to
remove an electron from an atom or ion

must overcome attraction
Ionization
4.3 Trends in Periodic Table
Ionization Energy

decreases as you move down a group


due to electron shielding (inner electrons
shield outer electrons from the full attractive
force of the nucleus)
increases as you move across a period
Ionization Energy Trends
Ionization Energy Graph
4.3 Trends in Periodic Table
Atomic Radius

bond radius: half the distance from center
to center of two like atoms that are bonded
Bond Radius
4.3 Trends in Periodic Table
Atomic Radius

increases as you move down a group


due partly to electron shielding
decreases as you move across a period

levels as electrons get close to one another
Atomic Radii Trends
Atomic Radii Graph
4.3 Trends in Periodic Table
Electronegativity

measure of the ability of an atom in a
compound to attract electrons is called
electronegativity

higher electronegativity means a stronger pull
on electrons
4.3 Trends in Periodic Table
Electronegativity

decreases as you move down a group


due to electron shielding
increases as you move across a period

not adding electrons to inner levels, so
stronger effective nuclear charge
Electronegativity Trends
Electronegativity Graph
4.3 Trends in Periodic Table
Ionic Size

increases as you move down a group


due to electron shielding
decreases as you move across a period

due to increasing nuclear charge, whether the
ion is positive or negative
Ionic Radii Trends
4.3 Trends in Periodic Table
Electron Affinity



the energy change that occurs when a
neutral atom gains an electron is called the
atom’s electron affinity
decreases as you move down a group due
to…take a guess
increases as you move across a period
Electron Affinity Trends
4.3 Trends in Periodic Table
Melting and Boiling Points


peak as d and p orbitals fill
bonds are stronger with half-filled orbitals
Melting and Boiling Points