Transcript Ch.4 Notes Powerpoint Version

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?
• 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.
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.
▫ Only 1 proton and 1 electron.
▫ Can gain or lose an electron.
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
melting & boiling points
IE, Atomic Radius, and E-neg
• Ionization energy: energy needed to remove an
electron (forms an ion- atom with a charge).
• Atomic Radius (size): 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.
IONIZATION ENERGY
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 an e- when bonded
with another atom.
• 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 Effect: inner electrons shield/block the
valence electrons from the positive nucleus. This
results in less attraction.
▫ Increases going down a group. Why?
▫ Stays the same going across a period. Why?
• Nuclear Charge: positive charge of the nucleus.
Increases as the number of protons increases.
▫ Increases down a group. Why?
▫ Increases across a period. Why?
• Combine these two and you get effective
nuclear charge.
The Basics Cont.
• Effective Nuclear Charge: how well valence
electrons can feel the attraction to the nucleus’
positive charge.
▫ Takes shielding (inner e-) and nuclear charge
(#protons) into account .
 More shielding = less effective nuclear charge.
▫ Decreases going down. Why?
 Shielding increases.
▫ Increases going across. Why?
 Shielding stays the same (because the number of inner
electrons stays the same).
 Nuclear charge increases because protons increase.
IE, Atomic Radius, e-neg
Across a Period
• Effective nuclear charge increases.
SO…
• Ionization energy INCREASES.
• Atomic radius DECREASES.
• Electronegativity INCREASES.
IE, Atomic Radius, e-neg
Down a Group
• Effective nuclear charge decreases.
SO…
• Ionization energy DECREASES.
• Atomic radius INCREASES.
• Electronegativity DECREASES.
Atomic Radius Cont.
http://intro.chem.okstate.edu/1314f00/lecture/chapter7/ATRADIID.DIR_PICT0003.gif
Explaining Reactivity &
Demonstration
• 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.
▫ Electron affinity is very similar to
electronegativity- it indicates 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.