The Periodic TableChapter 4

What information can be determined from the periodic table?

How have elements been organized into the periodic table used today?

Section 1: How Are Elements Organized?

• 1865- when placed in order of atomic mass, scientists noticed that properties repeated every 8 elements.

• Dmitri Mendeleev used this discovery and placed all 63 known elements into the first periodic table.▫When chemical properties repeated he started

another row. Each column had elements with similar properties.

▫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.• Periodic: happening or recurring at certain

intervals.▫Ex: waves are periodic:

• Periodic law: properties of elements are repeated at certain intervals when ordered by atomic number.

More on the Periodic Law• Properties of elements change as electron

configurations change.▫Since the # e- = atomic #, you can also think of the

periodic table being ordered by electron configurations. ▫Can easily determine the # of valence electrons , which

are the outermost e- (comes from the column #).▫Changing e- changes the elements’ properties/behavior!

• Group = column (also known as families).▫Recall that these elements all have similar properties!▫This is because they have the same # of valence

electrons!• Period = row (energy level).

▫Recall that each row begins when properties begin repeating again.

How are elements grouped on the periodic table?

Green = metals

Blue = metalloids

Yellow = nonmetals

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 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. Some used to produce color on TV

screens.▫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.

• Main group elements – s & p block elements▫Groups 1,2 and 3-8

• Group 1 = Alkali Metals▫H is NOT included

• Group 2 = Alkaline Earth Metals• Group 7 = Halogens• Group 8 = Noble Gases

• Remember: the group/column number tells you how many valence electrons those elements have!

Groups

Alkali Metals• VERY REACTIVE !

▫React with water to make alkaline 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.

• Found as compounds, rather than pure substances.

• Harder and higher melting points than group 1.

Halogens• Most reactive nonmetals.• 7 valence electrons.

▫Only need to gain one more electron to 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

Exception: He, which has s2. But 1st energy level does not have 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.

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

The Basics.•Shielding Effect: inner electrons shield/block

the valence electrons from the positive nucleus.▫Results in less attraction.▫Increases going down, stays the same going across.

•(Effective) Nuclear Charge: how well valence electrons can feel the nucleus’ positive charge.▫More shielding (inner electrons) = less nuclear

charge felt by the valence electrons.▫Decreases going down (due to shielding).▫Increases going across.

Inner electrons stay the same, so does shielding. Protons increase.

Ionization Energy• Energy needed to remove an e- (forms an ion-

atom with a charge).• Decreases down a group.

▫More energy levels between nucleus and valence e-.

▫Shielding effect increases, effective nuclear charge decreases.

• Increases across a period.▫Energy levels between the nucleus and valence e-

stay the same. Increased nuclear charge (more p+ and e-, but e-

are being added to the same energy level). Noble gases have the highest values (very

stable).

Increases

Decreases

Atomic Radius (Atomic Size)

• Half the distance between 2 bonded atoms’ nuclei.• Hard to measure exact size due to e- cloud.• Bond radius is easier to measure- then cut in half.• Increases down a group.▫ Energy levels added farther and farther from the

nucleus.▫ Shielding effect increases.• Decreases across a period.▫ Shielding remains the same, effective nuclear charge

increases.▫ Electrons feel the attraction more and are pulled closer.

Inc reases

Decreases

Atomic Radius Diagram

distance between two bonded atoms’ nuclei

2

Where should we

consider the outside of

the atom to be?

Measured in picometers (pm) or Angstroms (Å).

Atomic Radius Cont.

http://intro.chem.okstate.edu/1314f00/lecture/chapter7/ATRADIID.DIR_PICT0003.gif

Electronegativity• Ability of an atom to attract 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 highest ability to attract e- when bonded).

Noble gases don’t have a value (don’t need to form bonds- they are stable).

Electronegativity

• Decreases down a group.▫Shielding effect increases.▫Effective nuclear charge decreases.▫Thus it’s tougher to attract electrons in a

bond.• Increases across a period.

▫Shielding stays the same.▫Effective nuclear charge increases.▫Thus it’s easier to attract electrons in a bond.▫Another way to look at it: as more e- are

added, elements get closer to noble gas configuration which is very stable! The closer to stability the better.

Increases

Decreases

Other Trends- Honors• Ionic size

▫ Ions: atoms that have lost or gained e- (have a charge).

▫Same trend as atomic size. Same for both positive and negative ions.

• Electron affinity▫Ability to gain e- (when not bonded).▫Same trend as electronegativity.

• Melting & boiling points (see pg. 140)▫Very different trend from ones previously discussed.

Peaks appear in the p and d blocks across a period. Tend to increase as atoms approach half filled p

and d sublevels, and then decrease.▫More opportunities for atoms to bond together.

How have the elements been 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,O,N,P, & S.▫Compounds that are carbon-based are called

organic compounds. Found in living things.

•Big Bang Theory: these 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.

Supernovas

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.