Download The Periodic Table

The Periodic Table
Chapter 4
What information can be determined
from the periodic table?
How have elements been organized
into the periodic table used today?
Periodic Table with the f block in
its proper location:
Section 1: How Are Elements Organized?
• In 1869 Dmitri Mendeleev found that by placing
elements in order of increasing atomic mass
properties of elements were repeated.
▫ Each new row = properties repeated.
▫ This resulted in each column having elements
with similar properties.
• Made the first periodic table!
▫ Able to predict missing elements using this
repetition.
▫ Problem with ordering elements by atomic mass: some
did not match properties of other elements in the same
column. Needed to be switched around.
Adjusting the Periodic Table
• About 40 years after Mendeleev’s table, Henry
Moseley made an important change to the periodic
table:
▫ Organized elements by atomic number instead of
atomic mass.
▫ Elements that had not previously fit into the correct
column when ordered by atomic mass were fixed.
• This is the periodic table we still use today.
• Properties of elements repeat as a result of being
ordered by atomic number. In other words, they
exhibit periodicity.
• This is called the periodic law.
Additional Information
• Columns on the periodic table are called groups or
families.
▫ Recall that these elements all have similar properties!
▫ This is because they have the same number of valence
electrons, which means they will react in similar ways.
• Valence electrons: outermost electrons in an atom.
• We can easily determine the number of valence electrons
by looking at group numbers in the s & p blocks.
• Rows are called periods (indicates the energy level).
▫ Recall that each row begins when properties begin
repeating again.
Group Numbers & Valence e1
1
IA
1A
8
2
Numbers in bold are
the number of valence
e- for each group.
3
4
5
6
7
18
VIIIA
8A
How are elements grouped on the
periodic table?
Yellow = nonmetals
Blue = metalloids
Green = metals
Metals
• Most elements on the periodic table are metals.
• Conduct electricity and heat.
• Ductile
▫ Can be drawn into a wire.
• Malleable
▫ Can be hammered or rolled into sheets.
• Usually lustrous
▫ Look shiny.
▫ Dull in air or oxygen.
• Solids at room temperature (except Hg).
Nonmetals
• Opposite characteristics from metals:
▫ Do not conduct electricity and heat well.
▫ Not very ductile.
▫ Are not lustrous.
▫ Can be solids, liquids, or gases at room
temperature.
Transition Metals
• Groups 3-12.
▫ d block elements.
• Can lose a different number of valence electrons.
▫ Less reactive than other metals we will look at (alkali
and alkaline earth metals).
 Some like Pd, Pt, and Au are very unreactive.
Rare Earth Metals
• f block- 2 rows at the bottom of the table.
▫ Fit into rows 6 & 7 (look for * or other symbol).
• Lanthanide & Actinide series
▫ Lanthanides = 4f
 Reactive (like alkaline earth metals we will look
at).
▫ Actinides = 5f
 All of them are radioactive.
 Nuclei are unstable and break down.
Other Properties of Metals
• Varying melting points.
▫ Example: W = 4322oC and Hg = -39oC
• Used to make alloys.
▫ Alloys: Homogeneous mixtures of metals.
 New properties result from mixing metals.
 Example: Brass = copper and zinc.
▫ Harder than copper alone.
▫ More resistant to corrosion.
 Others include steel, stainless steel, sterling
silver.
Groups
• Main group elements – s & p block elements
▫ Groups 1,2 and 3-8 (or 13-18).
• Group 1(A) = Alkali Metals
▫ H is NOT included!
• Group 2(A) = Alkaline Earth Metals
• Group 7(A) (or 17) = Halogens.
• Group 8(A) (or 18) = Noble Gases.
• Remember: the group/column number tells you
how many valence electrons those elements have!
Alkali Metals
• VERY REACTIVE !
▫ React with water to make alkaline/basic solutions.
▫ Stored in oil to keep them from reacting with air and
water.
▫ Only 1 valence electron to lose- a filled valence shell is
very stable.
• Not found pure in nature, but combined with other
elements (as compounds).
• Soft – can be cut with a knife.
• Usually lustrous but will dull in contact with air.
▫ Form an oxide layer.
Alkaline Earth Metals
• Also highly reactive.
▫ Less reactive than alkali metals.
▫ Have 2 valence electrons to lose.
• Also found as compounds, rather than pure
substances.
• Harder and higher melting points than group 1.
• Often found as minerals and ores in the Earth’s
crust.
Halogens
• Most reactive nonmetals.
• 7 valence electrons.
▫ Only need to gain one more electron to have a full
valence shell and be stable.
• Frequently react with alkali metals.
▫ Recall that alkali metals have 1 valence electron to
lose.
▫ Ex: NaCl, KF, LiBr
• Compounds formed from halogens typically are
called salts.
Noble Gases
• Outermost energy level is completely filled with e-.
▫ s2p6 = 8 valence electrons
 Exception: He, which is 1s2. But the 1st energy level
does not have a p sublevel, so it is filled.
• Low chemical reactivity – very stable. They have no
desire to gain or lose electrons!
▫ Example – He used for blimps.
▫ Typically inert – thought to be completely unreactive.
 Exception: 1962, chemists were able to make some
compounds with Xe.
• Recall the Hindenberg
Hydrogen
• Most common element in the universe.
• Group by itself – very unique.
▫ Doesn’t fit perfectly into the alkali metals or halogens.
▫ Only 1 proton and 1 electron.
▫ Can gain or lose an electron- very reactive and forms
compounds with many other elements.
 If it loses its 1 electron, only a proton remains! This
is unlike any other element!
What trends can be found on the
periodic table?
Section 3: Trends in the Periodic Table
• Periodic trends exist since properties of
elements repeat in the table.
• We will look at the following trends:
▫
▫
▫
▫
▫
▫
ionization energy (IE)
atomic radius
electronegativity (e- neg)
ionic size
electron affinity
MP/BP, density
Ionization Energy (IE)
• Ionization energy: energy needed to remove an
electron from an atom (forms an ion- atom with a
charge).
IONIZATION ENERGY
Atomic Radius (size)
• Atomic Radius: Half the distance between two
bonded atoms’ nuclei.
• Hard to measure with only one atom due to e- cloud.
• How do we determine where it ends?
• Bond distance is easier to measure- then cut in half.
Atomic Radius Diagram
Where should
we consider
the outside of
the atom to be?
Measured in
picometers (pm)
or Angstroms (Å).
distance between two
bonded atoms’ nuclei
2
Electronegativity
• Ability of an atom to attract electrons in a bond.
• Electrons from each atom are involved when
atoms bond.
• Each atom’s ability to attract e- is different.
▫ Linus Pauling invented a scale to indicate how well
an atom can attract an e- in a bond.
▫ No units, just numbers.
▫ Ranges from 0 – 4.0.
 F assigned 4.0 (highest value- has the greatest
ability to attract e- when bonded).
 Noble gases don’t have a value (don’t need to
form bonds- they are stable).
ELECTRONEGATIVITY
Preview: Shielding
The Basics
• Shielding: inner electrons shield/block the valence
electrons from the positive nucleus.
Going
j
Li
down a
group.
Li
Going across
a period.
Q1: What happens to shielding as you move
down a group? Why?
Q2: What happens to shielding as you move
across a period? Why?
A1: Increases going down a group because
more shells/energy levels are being added.
A2: Stays the same going across a period
because you’re in the same shell/energy level.
The Basics
• Nuclear charge: positive charge in the nucleus.
Increases as atomic number/protons increase.
Li
Be
B
j
Going
+3
Li
down a
group.
+3
+4
+5
Going across
a period.
+11
Na
+19
K
Q1: What happens to nuclear charge as you
move down a group? Why?
Q2: What happens to nuclear charge as you
move across a period? Why?
A1: Increases going down a group because
atomic number increases.
A2: Increases across a period because atomic
number increases.
The Basics
• Effective nuclear charge: how well valence e- can feel
the positive nucleus based on shielding & nuclear charge.
Greatest effective nuclear charge.
+3
Li
Li
Be
+3
B
+4
+5
Going across
a period.
+11
Na
+19
K
Going
down a
group.
Q1: What happens to effective nuclear
charge as you move down a group? Why?
Q2: What happens to effective nuclear
charge as you move across a period? Why?
A1: Decreases going down a group because
shielding increases.
A2: Increases across a period because atomic
number increases & shielding stays the same.
IE, Atomic Radius, e-neg
Down a Group
• Effective nuclear charge decreases.
SO…
17p+
• Atomic radius INCREASES.
• Ionization energy DECREASES.
• Electronegativity DECREASES.
35p+
9p+
IE, Atomic Radius, e-neg
Across a Period
7p+
8p+
9p+
O
F
• Effective nuclear charge increases.
SO…
• Atomic radius DECREASES.
• Ionization energy INCREASES.
• Electronegativity INCREASES.
Atomic Radius Cont.
http://intro.chem.okstate.edu/1314f00/lecture/chapter7/ATRADIID.DIR_PICT0003.gif
Explaining Reactivity
• Recall that groups 1 and 7 are the most reactive
metals and nonmetals.
• As we move down group 1, the alkali metals become
more reactive- this is because of the trend seen in
ionization energy!
• As we move down group 7, the halogens become less
reactive- this is because of the trend seen in electron
affinity (very similar to electronegativity).
▫ Electron affinity is how well an atom can gain an
electron.
How are elements created?
Section 4: Where Did the Elements Come From?
• Only 93 of the elements are found in nature.
▫ 3 of these are not found on Earth.
 Technetium, Promethium, Neptunium
 Found in stars.
• Most living things contain C,H,N,O,P, & S.
▫ Compounds that contain carbon are called organic
compounds.
 Found in living things.
• Big Bang Theory: elements were created when
universe was formed in a violent explosion.
Big Bang Theory Cont.
• VERY high temperatures existed after the big bang. This
form of energy cooled and formed matter (e-, p+, n).
 Further cooling allowed subatomic particles to join together
to form H.
 Gravity pulled H clouds together and formed stars.
 Stars worked as nuclear reactors to form He (under high
temperature and pressure).
 4 H  1 He + energy (gamma radiation)
• Other elements were formed as He and H combined
(fusion) to form even heavier elements.
• Supernovas formed all elements heavier than iron.
▫ Star collapses and blows up, releasing heavier elements into
space.
▫ This can emit more energy than the sun does in its life span!
Supernovas
http://en.wikipedia.org/wiki/Supernova_remnant
Synthetic & Superheavy Elements
• Transmutations: type of nuclear reactions
that change one element into another element
• All elements greater than number 93 (except 61)
are not naturally occurring– synthetic
elements.
▫ Particle accelerators can be used to create
them. Different types exist.
 Nuclei collide and fuse together.
• Superheavy elements are those that have an
atomic number greater than 100.
▫ Only exist for fractions of a second.