Transcript Chapter 14

Chapter 14:
Periodic Trends
…and naming ions (chapter 6)
The Modern Periodic Table
Review: Energy Levels
• Principle quantum numbers (n) = energy
level
– Lower the number, lower the energy
Sublevels – s, p, d, f, g
Put it all together…
A new way to write electron
configurations
A new way to write electron
configurations
Use the Noble gas abbreviation, write the
electron configurations for:
• Na:
• Cl:
• W:
• Sn:
Organizing the Table
• Noble Gases –
– Filled outermost s and p orbitals (s2p6)
• Representative Elements- (s and p block)
– Partially filled outermost s and p orbitals
• Transition Metals (d block)
– Outermost s and nearby d orbitals contain
electrons
• Inner Transition Metals
– Outermost s and nearby f contain electrons
Representative Elements
•
•
•
•
•
Sometimes called “group A elements”
1A = alkali metals
2A = alkaline earth metals
7A = halogens
8A = noble gases
• The group number equals the number of
electrons in its outermost energy level
(this will be more important later on…)
Metals and Nonmetals
• From here on out you need to identify
whether an element is a metal or
nonmetal (FAST and accurate).
Other things you need to know… (Ch. 6)
• When atoms gain or lose electrons they
form IONS (positively or negatively charged
atoms).
• Charge determined by difference between p+ and e-
• Positive ions = cations (lose valence e-)
– Metal atoms lose valence e-
• Negative ions = anions (gain valence e-)
– Nonmetal atoms gain valence e-
Cations
• Formed when a metallic atom LOSES
electrons.
• Examples: Calcium and Magnesium
calcium ion
magnesium ion
• You can determine how many electrons
are lost based on location on the periodic
table. These must be memorized.
– Group 1A, 2A, Al, Ag, Zn
Naming Ions
• How to name: name of element + “ion”
• Examples:
• Cations
– Aluminum (Al)
– Sodium (Na)
:
:
Aluminum Ion (Al3+)
Sodium Ion (Na+)
Transition Metals…
• Can form different charges (you can’t
memorize)
• Here is how you know the charge:
– They are metals, so are cations.
– Roman Numerals (I, II, III, IV, V, VI, VII)
indicate how many e- lost.
– Copper (II) : two val e- : Cu2+
– Copper (I) : one val e- : Cu+
Anions
• Non-metal atoms that gain electrons
become anions.
• Examples: Bromine and Nitrogen
bromide ion
nitride ion
• You need to MEMORIZE the common
charges of the anions to be successful for
the rest of the school year…
Naming Anions
• Anions change the ending of the element
– unlike cations
• Stem-ide ion
• Examples:
• Chlorine = chloride ion
• Oxygen = oxide ion
• Anions
– Chlorine (Cl)
– Oxygen (O)
:
:
chloride ion (Cl-)
oxide ion (O2-)
Ion Size
• Cations are smaller in
size than the neutral
element.
• Anions are larger in size
than the neutral element.
Polyatomic ions
• Polyatomic = many atoms
• Ions = charged
• You will get a list of 10 polyatomic ions.
You must memorize the name and
formula and be able to recall them at any
time after the first test (i.e. I won’t feel
guilty if there is a pop quiz).
Now on to trends…
• There is no secret to success this chapter
other than to memorize the following
trends…
Types of Trends
• Periodic Trends: Trends across a period
(row) of the periodic table.
• Group Trends: Trends down a group
(family) of the periodic table
Nuclear Charge (+ in
nucleus)
• Periodic: Nuclear charge increases as
you go left to right across a period.
• Group: Nuclear charge increases as you
go down a group.
Atomic Radii and Size
• Periodic: Atomic Radii and size decrease as
you go L to R across the table.
– Same principle energy level (n)
– Add p+ and e-, increase nuclear charge, pulls in
orbitals closer to the nucleus
• Group: Atomic Radii and size increase as you
go down a group
– Electrons being added to outer orbital (increasing
principle energy level)
Ionization Energy
• Ionization Energy (IE): The energy
required to remove an electron from a
gaseous atom.
– Remove 1st electron = 1st IE
– Remove 2nd electron = 2nd IE
Trends in Ionization Energy
• Periodic: IE generally increases as you
move L to R across the period.
– Harder to remove an electron as you go L to R
because of greater attraction to nucleus
– Shielding effect -
• Group: 1st IE generally decreases as you
go down a group.
– Atom gets bigger, outermost e- farthest from
nucleus, easy to be removed.
Electronegativity
• Electronegativity: The tendency for
atoms to attract electrons when they are
chemically combined.
– Stronger attraction.
Electronegativity Trends
• Periodic: Electronegativity increases as
you go L to R across the period
– Elements want to be like noble gases!
• Group: Electronegativity decreases as
you go down a group.
The most electronegative element
is Fluorine.
Summary
Decreasing Atomic Radius
Increasing Electronegativty
Increasing Atomic Radius
Decreasing Electronegativity
Decreasing Ionization Energy
Increasing Ionization Energy