Unit 3 – Periodic Table of Elements

Unit 3 Key Terms

• Energy Level – discrete regions of space around the nucleus in the electron cloud where electrons can reside

• Lewis dot structure -A model that uses electron-dot structures to show how electrons are arranged in molecules. Pairs of dots or lines represent bonding pairs

• Noble gas configuration -An electron structure of an atom or ion in which the outer electron shell contains eight electrons, corresponding to the electron configuration of a noble gas, such as neon or argon

• Orbital notation (diagram) -A way to show how many electrons are in an orbital for a given element. They can either be shown with arrows or circles

Unit 3 Key Terms (cont.)• Atomic Mass – physical property indicating the mass of an atom• Alkali Metals – extremely reactive metals found in Family 1A• Alkaline Earth Metal – very reactive metals found in Family 2A• Atomic Radius – one half the distance between two adjacent

atoms of the same element• Electronegativity – the ability of an element to attract electrons

from a neighboring atom• Ionization Energy – the energy required to remove an electron

from the electron cloud of an element• Halogens – extremely reactive non-metals found in Family 7A• Ion – charged atom resulting from the loss or gain of electrons• Ionic Radius – the resulting atomic radius found in an element

when it has either lost or gained electrons to become an ion

Unit 3 Key Terms (cont.)• Mendeleev – Russian chemist that placed elements into a

Periodic Table with periods based on atomic mass and families (groups) based on similar chemical and physical properties

• Mosley – English physicist that used X-ray spectroscopy to find the Atomic Number of elements and place elements into a Periodic Table with periods based on atomic number and families (groups) based on similar chemical and physical properties

• Noble Gases – stable, non-reactive non-metals in Family 8A• Periodic Trends – repeating patterns of chemical and physical

properties of elements within the periodic table that correspond to the Law of Periodicity

• Reactivity – chemical property describing the ability and speed with which elements react

• Transition Metals – metals in the d-block, generally with multiple valence states; most of the commonly known metals are in this category of metals

Quantum Mechanical Model of Atomic Structure• 1900: Max Planck – Develops law correlating energy to

frequency of light• 1905: Albert Einstein – Postulates dual nature of light as both

energy and particles• 1924: Louis de Broglie – Applies dual nature of light to all

matter• 1927: Werner Heisenberg – Develops Uncertainty Principle

stating that it is impossible to observe both the location and momentum of an electron simultaneously

• 1933: Erwin Schrodinger – Refines the use of the equation named after him to develop the concept of electron orbitals to replace the planetary motion of the electron

Energy Levels

• Energy levels correspond to the energy of individual electrons. Each energy level has a discrete numerical value.

• Different energy levels correspond to different numbers of electrons using the formula 2n2 where “n” is the energy level

Energy Level Number of electrons (2n2)

1 2(12) = 2

2 2(22)= 8

3 2(32)= 18

4 2(42)= 32

n 2n2

OrbitalsImpossible to determine the location of any single electron

Orbitals are the regions of space in which electrons can most probably be found

Four types of orbitalss – spherically shapedp – dumbbell shapedd – cloverleaf shapedf – shape has not been determined

Each additional energy level incorporates one additional orbital type

Each type of orbital can only hold a specific number of electrons

Orbital Types

Orbital Type

General Shape

OrbitalSublevels

# of electrons

per sublevel

Total # of electrons

per orbital type

s Spherical 1 2 2

p Dumbbell 3 2 6

d Clover leaf 5 2 10

f unknown 7 2 14

Electron Configuration

Energy Level

Orbital Type

OrbitalSublevel

# of orbitals

per energy level (n2)

# of electrons

per orbital type

# of electrons

per energy level (2n2)

1 s 1 1 2 2

2sp

13

426

8

3spd

135

926

1018

4

spdf

1357

16

26

1014

32

Electron Configuration Notation

• Find the element on the periodic table• Follow through each element block in order by stating the

energy level, the orbital type, and the number of electrons per orbital type until you arrive at the element.

1s

2s     2p

3s     3p

4s   3d 4p

5s   4d 5p

6s 4f 5d 6p

7s 5f 6d 7p

Samples of e- Configuration

• Element Electron Configuration• H 1s1

• He 1s2

• Li 1s2 2s1

• C 1s2 2s2 2p2

• K 1s2 2s2 2p6 3s2 3p6 4s1

• V 1s2 2s2 2p6 3s2 3p6 4s2 3d3

• Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap)• Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2

Noble Gas Electron Configuration Notation

Find element on the Periodic Table of Elements• Example: Pb for Lead

Move backward to the Noble Gas immediately preceding the elementExample: Xenon

Write symbol of the Nobel Gas in bracketsExample: [Xe]

Continue writing Electron Configuration Notation from the Noble GasExample: [Xe] 6s2 4f14 5d10 6p2

Valence Electrons

• The electrons in the highest (outermost) s and p orbitals of an atom

• The electrons available to be transferred or shared to create chemical bonds to form compounds

• Often found in incompletely filled energy levels

Valence ElectronsShortcut to finding valence electrons for main group elements

Family 1A (1) 1 valence electronFamily 2A (2) 2 valence electronsFamily 3A (13) 3 valence electronsFamily 4A (14) 4 valence electronsFamily 5A (15) 5 valence electronsFamily 6A (16) 6 valence electronsFamily 7A (17) 7 valence electronsFamily 8A (18) 8 valence electrons

Family 3-12 have multiple possibilities and shortcuts do not work

Electron Dot NotationElectron configuration notation using only the valence electrons of an

atom.

The valence electrons are indicated by dots placed around the element’s symbol.

Used to represent up to eight valence electrons for an atom. One dot is placed on each side before a second dot is placed on any side.

Valance Electrons: Sodium Magnesium Chlorine Neon

1 2 7 8

Electron Dot Notation: • • •• ••

Na Mg : Cl : : Ne : • • ••

Oxidation Numbers:+1 +2 -1 0

Early Development of the Periodic Table of Elements• Antoine Lavoisier (France 1778)

• Produced the first extensive list of elements showing 33 elements• Separated metals from non-metals

• John Dalton (England 1803) • Developed postulates of atomic theory with a list of elements

and symbols• Jacob Berzelius (Sweden 1828)

• Systematized letters to symbolize elements• Provided a table of atomic weights

• Johann Dobereiner (German 1828)• Discovered patterns between elements in groups of 3• Elements in triads formed similar chemical compounds• Published list of these groups called “triads”

Early Development of the Periodic Table of Elements (cont.)

• John Newlands (England 1864)• First person to devise a periodic table of elements• Expanded the concept of triads into octaves. Elements were said

to exhibit similar chemical and physical properties to the eighth element following it in the table (Law of Octaves)

• His table did not include all of the known elements

Development of Modern Periodic Table of Elements

• Dmitri Mendeleev (Russia 1869)• Produced the first Periodic Table to arrange elements in periods

(rows) and families (columns) showing all 66 known elements• Periods arranged elements in order of increasing atomic mass• Families arranged by similar chemical and physical properties• Method of arrangement left gaps for elements believed to exist and

not yet discovered.• William Ramsay (England 1894)

• Discovered a new family of gases that resisted chemical reactions• Noble gases added to the Periodic Table

• Henri Becquerel (France 1903)• With Pierre and Marie Curie credited with discovering radioactivity• Opened a new window on understanding atomic structure and

properties

Mendeleev’s Periodic Table

Development of Modern Periodic Table of Elements (cont.)• Frederick Soddy (England 1912)

• Suggested existence of isotopes• Since elements could have multiple masses they could occupy

multiple positions on the Periodic Table arranged by mass• J.J. Thomson (England 1913)

• Confirmed experimentally the existence of isotopes with Neon• Since Neon is unreactive its multiple masses could not be explained

by unidentified chemical compounds• Discovery of isotopes meant that the Mendeleev Periodic Table

could not be the final answer

Development of Modern Periodic Table of Elements (cont.)• Henry Moseley (England 1913)

• Demonstrated through x-ray spectroscopy that the characteristics of the x-rays emitted by different atoms are incremental and can be listed in numerical order (Atomic Number)

• Put forward the theory that chemical and physical properties are periodic functions of this atomic number (Law of Periodicity)

• Refined Rutherford’s theory of the atomic structure indicating a correlation between the positive charge of the nucleus and atomic number

• Developed the basic structure of the Periodic Table used today.• Rutherford (England 1917)

• Produced the first experimentally based nuclear reaction transforming nitrogen into oxygen using alpha particles

• The other product of the reaction was a Hydrogen nucleus (Proton)• Confirmed atomic number was equal to the number of protons

Moseley’s Periodic Table

Image courtesy of http://corrosion-doctors.org/Periodic/Periodic-Moseley.htm

Development of Modern Periodic Table of Elements (cont.)• Glenn Seaborg ( United States 1941-1944)

• Discovered discrepancies in Moseley’s table through the identification of new elements while conducting research as part of the Manhattan Project

• Created the lanthanide and actinide series referred to as transuranium elements

• Discoveries disclosed at the end of World War II• Continuing research

• Research laboratories use particle accelerators to identify new elements

• Recent discoveries have completed the 7th period of the table• Research is continuing to discover more new elements in an 8th

period

Seaborg’s Periodic Table

Periodic Table of Elements 2012

Image used courtesy of https://proteabio.com/resources/tools/Periodic+Table+of+Elements

Dynamic Periodic Table

• Courtesy of ptable.com

Image courtesy http://www.fanpop.com

Families of Particular ImportanceFamily 1A (1) – Alkali Metals

Soft metals and silver gray in color Extremely reactive – do not exist in elemental form in nature 1 valence electron

Family 2A (2) – Alkaline Earth Metals Soft metals and silver in color Very reactive – can exist in nature, but oxidize rapidly 2 valence electrons

Family 7A (17) – Halogens (non-metals) Very reactive Lighter halogens are gases at room temperature while heavier halogens are solids

at room temperature; Bromine is a liquid at room temperature 7 valence electrons

Family 8A (18) – Noble Gases (Non-metals) Generally non-reactive and do not form compounds Extremely Stable 8 valence electrons

Types of ElementsMetals

Good conductors of electricity and heatVast majority of the elements – Alkali metals, alkaline earth metals,

transition metals, post-transition metals, and inner transition metals

Non-MetalsPoor conductors of electricity and heatIncludes the Nobel Gases, Halogens, and only a few of the lightest

elements – hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, and selenium

MetalloidsMany are semiconductors of electricityExhibit properties of both metals and non-metalsOnly 7 elements are metalloids (some scientists include different ones

depending on perspective)

Periodic Law Created by Henry Moseley

The chemical and physical properties of the elements are periodic functions of their atomic numbers

Properties of the elements occur at repeated intervals called periods (rows on Periodic Table)

This defines the property of periodicity

Periodic TrendsAtomic Radius – half the distance between the nuclei of atoms of

the same materialDecreases generally across periods – increased positive nuclear

charge (protons) pulls electrons in tighter to the nucleusIncreases generally down families – increased number of energy

levels where electrons may reside

Electronegativity – the measure of the ability of the nucleus of an atom to attract electrons of a neighboring atomIncreases generally across periods – increased positive nuclear charge

(protons) more strongly attracts electrons from neighboring atomsDecreases generally down families – increased number of energy

levels means the nucleus is less able to overcome the distance between atom

Periodic TrendsIonization Energy – the energy required to remove an electron

from an atomIncreases generally across periods – increased positive nuclear charge

(protons) pulls electrons in tighter to the nucleus making them harder to remove

Decreases generally down families – increased number of energy levels where electrons may reside making them easier to remove

Ionic Radius – as atoms gain and lose electrons, the radius of the charged atom changesIncreases when an atom accepts electrons– the more electrons there

are the greater the overall repulsive forces between the electrons pushing them further apart

Decreases when an atom loses electrons – the fewer the electrons the greater the effectiveness of the nuclear charge (protons)

Periodic Trends