Download Unit 5 – The Periodic Table

Unit 5 – The Periodic Table
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How is the periodic table arranged?
How does periodicity explain the chemical and physical properties of the
elements?
How can chemical and physical properties of the elements be predicted
based on their positions in the periodic table?
Origins of the Periodic Table
• By 1700, only 14 elements had been isolated
• Scientific discovery led to a higher rate of
element discovery (20 in the 1700s)
• A logical organization of elements was needed
for all the new elements
Early Organization
• Triads
– J.W. Döbereiner (1829)
– 3 elements with similar
properties
– Based on average
properties
– Ex: Cl, Br, I
• Average of I and Cl
atomic mass = Br atomic
mass
The Periodic Table
• Periodicity
– Repeating pattern to properties
• Periodic Table
– Dmitri Mendeleev (1869)
– Arranged elements according to
their properties
– Mendeleev left some spaces blank
• Accurately predicted properties
of unknown elements
Early Organization
• Octaves
– John Newlands
(1864)
– Repeating group of 8
elements
The Periodic Table
• Modern Periodic
Table
– Henry Moseley (1913)
– Arranged elements
according to atomic
number
• Periodic Law
– Repetition of
properties when
arranged by atomic
number
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Modern Periodic Table
• The modern periodic table consists of Rows
and Columns
• Rows - Horizontal
– Also known as Periods
– Numbered 1-7
• Columns - Vertical
Classifying Elements
• Representative Elements
– Groups 1,2,13-18 (1A-8A)
– S-block and p-block
– All elements in same group share:
• Oxidation state
• Valence electrons
– Also known as Groups and Families
– Numbered 1-18
Classifying Elements
• Transition Metals
– Groups 3-12 (3B-2B)
– d-block
– Usually have multiple oxidation states
• Inner Transition Metals
– Lanthanide and Actinide series
– f-block
Oxidation State
• Octet
– 8 valence e– Full valence shell
• Stability
– Atoms behave in ways to achieve an octet
• Noble Gases have octets – don’t bond
• Exception – 1st energy level
– Duet (2 e-)
Oxidation State
What is the trend for Oxidation State?
• Oxidation State
– Number of e- gained or lost by an atom in forming a
bond
– Ion whose configuration matches nearest noble gas
– Lost e- = Positive Oxidation State
– Gained e- = Negative Oxidation State
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Classifying Elements
• Metals
–
–
–
–
Classifying Elements
• Metalloids
Left side of table
Most common class of elements
Lose electrons to become cations
Generally:
• Solid (except Hg)
• Conductive of heat and electricity
• Malleable
• Ductile
• Lustrous (Shiny)
Classifying Elements
• Nonmetals
– Right side of staircase
– Gain electrons to become anions
– Generally:
• Gases or dull, brittle solids
• Poor conductivity
• Poor ductility
• Non-malleable
• Non-lustrous
Classifying Elements
• Alkaline Earth Metals
– Group 2 elements
– Be, Mg, Ca, Sr, Ba, and Ra
– 2 valence e• +2 ox. state
– Harder, denser, stronger, less reactive
than alkali metals
• Can occur in nature, but are usually
in compounds
– Border the staircase
– B, Si, Ge, As, Sb, Te, Po,
At
– Properties in between
metals and nonmetals:
• Brittle
• Lustrous
• Semiconductors
Classifying Elements
• Alkali Metals
– Group 1 elements
• Li, Na, K, Rb, Cs, Fr
– 1 valence e• +1 ox. state
– Generally dull, soft, and reactive
• Never free elements in nature –
react with air/water
Classifying Elements
• Transition Metals
– Mn, Fe, Ag, Au, Mo, etc.
• Inner Transition Metals
– Lanthanide Series (Rare Earth Metals)
• Top row of f-block
– Actinide Series
• Bottom row of f-block
• Post-Transition Metals
– Al, Ga, In, Sn, Tl, Pb, Bi
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Classifying Elements
• Halogens
Classifying Elements
• Noble Gases
– Group 17 elements
– F, Cl, Br, and I
– 7 valence e• -1 ox state
– Most reactive nonmetals
• As reactive as alkali metals
• Rarely free elements
Classifying Elements
• Other groups can be named by the top-most
element
– Ex: Group 15
• Nitrogen Group
• Oxidation State: -3
– Q: What is another name for Group 16?
• A: Oxygen group
– Q: Oxidation State
• A: -2
Periodic Trends
• Ionization Energy (IE)
– Energy required to remove an electron
– Increases from left to right across a period
• Smaller atom
– Decreases from top to bottom in a group
• e- is further away
– Larger change across period than down group
– Elements in group 18
– He, Ne, Ar, Kr, Xe, Rn
– 8 valence e• No ox state
– Extremely unreactive and stable
• Almost never bond
Periodic Trends
• Atomic Radius
– Overall size of the atom
– Decreases from left to right across a period
• Stronger nucleus, same energy level
– Increases from top to bottom in a group
• More energy levels
– Larger change across period than down group
Periodic Trends
• Electronegativity (EN)
– Ability for an atom to attract electrons in a
compound
• Noble gases aren’t considered
– Increases from left to right in a period
– Decreases from top to bottom in a group
– Larger change across period than down group
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Periodic Trends
• Ionic Radius
– Atomic radius of the ion
– Cations
• Smaller than atoms from which they form
–Less electrons
– Anions
• Larger than atoms from which they form
–More electrons
Periodic Trends
• Ionic Radius
– Energy level changes between cations and
anions
– Cations – previous noble gas
– Anions – next noble gas
Li1+
B3+
N3-
O2-
F1-
Be2+ C4+
Periodic Trends
• Ionic Radius
– Decreases from left to right across a period
– Increases from top to bottom in a group
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