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Transcript
```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.
• 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.
• 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.
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)
electronegativity (e- neg)
ionic size
electron affinity
melting & boiling points
• 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
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.
Across a Period
• Effective nuclear charge increases.
SO…
• Ionization energy INCREASES.
• Electronegativity INCREASES.
Down a Group
• Effective nuclear charge decreases.
SO…
• Ionization energy DECREASES.
• Electronegativity DECREASES.
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.
 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.
```
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