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.

•  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.

Adjusting the Periodic Table

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 e- 1 IA 1A

18 VIIIA 8A

1

2 3 7 4 5 6

8

Numbers in bold are the number of valence e- for each group.

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 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.

•  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!

Groups

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

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).

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. j

Li

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.

Going 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?

The Basics • Nuclear charge: positive charge in the nucleus.

Increases as atomic number/protons increase. j

A1: Increases going down a group because atomic number increases. A2: Increases across a period because atomic number increases.

Going down a group.

+3

Going across a period.

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?

+11

+3

+19

Li

Na

K

Li

+4 +5

Be B

The Basics •  Effective nuclear charge: how well valence e- can feel

the positive nucleus based on shielding & nuclear charge.

A1: Decreases going down a group because shielding increases. A2: Increases across a period because atomic number increases & shielding stays the same.

Going down a group.

+3

Going across a period.

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?

+11

+3

+19

Li

Na

K

Li

+4 +5

Be B

Greatest effective nuclear charge.

• Effective nuclear charge decreases.

SO…

• Atomic radius INCREASES. •  Ionization energy DECREASES. • Electronegativity DECREASES.

IE, Atomic Radius, e-neg Down a Group 9p+

17p+

35p+

IE, Atomic Radius, e-neg Across a Period

• Effective nuclear charge increases. SO…

• Atomic radius DECREASES. •  Ionization energy INCREASES. • Electronegativity INCREASES.

F O

9p+ 8p+ 7p+

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 à 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.