Chapter 6:The Periodic Table

General Chemistry

http://www.ccsdualsnap.com/miscellan.htm

Objectives

• Periodicity of physical and chemical properties relates to atomic structure and led to the development of the periodic table.

• The periodic table displays the elements in order of increasing atomic number.

• Explain the relationship of an element’s position on the periodic table to its atomic number and mass.

Objectives

• Use the periodic table to identify metals, nonmetals, metalloids, families (groups), periods, valence electrons, and reactivity with other elements in the table.

• Relate the position of an element on the periodic table to its electron configuration.

• Identify trends on the periodic table (ionization energy, electronegativity, electron affinity, and relative size of atoms and ions).

Review/Link to Previous Learning

• In Chapter 4, we learned about electrons configurations of elements.

• Discovered there is a pattern of electron configurations on the Periodic Table.

• Are there other patterns on the Periodic Table? (yes)

• In Chapter 5 we will learn how the Periodic Table is organized.

CollectionsDo you like to play cards?Do you have a stamp, baseballcard, or comic book collection?How do you organize your collection?

Attempts at Organizing Elements

• Early scientists knew about some properties of elements.

• Is there a characteristic of elements that can organize them?

JOHAN DOBEREINER(1780-1849)

Dobereiner’s TriadsTHE LAW OF TRIADS:The atomic mass of the middle element of the triadis equal to the mean of the atomic masses of theother two elements.

EXAMPLE: Lithium Atomic Mass of 7Sodium Atomic Mass of 23Potassium Atomic Mass of 39

According to Dobereiner’s Law, the atomic mass of sodium Should equal the arithmetic mean of lithium and potassium.(7+39)/2 = 23, which is the mass of sodium.

Problems with Dobereiner’s Law of Triads.

1) All the elements known at that time could not bearranged in triads.

2) The law did not work for very low or very highmassed elements such as F, Cl, and Br.

3) As techniques improved for measuring atomicmasses accurately, the law became obsolete.

Dobereiner’s research made chemists look at groups ofelements with similar chemical and physical properties.

JOHN A.R. NEWLANDS (1837-1898)

Newland’s Law of Octaves

When placed in increasing order of their atomicmasses, every eighth element showed similar physicaland chemical properties.

Li Be B C N O FNa Mg Al Si P S ClK Ca

Problems with Newland’s Law of Octaves

1) It was not valid for elements that had atomicmasses higher than Ca.

2) When more elements were discovered(Noble gases) they could not be accommodated in histable.

However, the modern periodic table does draw from theconcept of periods of eight.

DMITRI MENDELEEV (1834-1907)

Julius Lothar Meyer (1830-1895)

Mendeleev and Meyer

• Published nearly identical schemes for classifying elements

• Arranged elements by increasing atomic mass• Mendeleev generally given more credit

– Published first– More successful at demonstrating value of table– Predicted discovery of new elements, properties

of new elements

Properties of Some Elements Predicted By Mendeleev

Mendeleev’s Table: the first periodic table of the elements.He arranged the table so that elements in the same columnhave similar properties.

Problems with Mendeleev’s Table:

1) The positions of isotopes could not beaccommodated within the table.

2) In order to make the elements fit the requirements,Mendeleev was forced to put an element of slightlyhigher atomic weight ahead of one of slightly lower atomic weight.

Henry Moseley (1887-1915)

• Developed concept of atomic number

– amount of positive charge in the nucleus

• Later determined that arranging periodic table according to increasing atomic number eliminated problems seen in Mendeleev’s table

Why is it the “periodic” table?

• Periodic Law: when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern

Study Buddy Review

• Describe the contribution each person below made to the development of the periodic table:– Johan Dobereiner– John Newland– Dmitri Mendeleev– Julius Meyer– Henry Moseley

Parts of the Periodic Table

Parts of Periodic Table

• Groups/families: vertical columns– Alkali metals: 1A– Alkali earth metals: 2A– Boron, carbon families– chalcogens (oxygen family).– pnictogens (nitrogen family)– Halogens (fluorine family): 7A– Noble gasses: 8A/0

• Horizontal rows are called periods

• There are 7 periods

1A

2A 3A 4A 5A 6A7A

8A0

• The elements in the A groups are called the representative elements

outer s or p filling

Parts of Periodic Table

• Metals: left of staircase– Luster, malleable, conduct, ductile

• Nonmetals: right of staircase– Dull in appearance, nonconductor, brittle

• Metalloids: elements adjacent to staircase (except Al, Po)– Some properties of both metals and nonmetals

The group B are called the transition metals

These are called the inner transition metals and they belong here

Study Buddy Review

• Identify the follow parts of the periodic table:– Halogens

– family

– Alkali metals

– Metals

– Inner transition metals

– Noble gases

– Metalloids

– Period

Periodic Properties of Elements

Periodic Trends

•Atomic Radius

•Ionic Radius

•Ionization Energy

•Electron Affinity

•Electronegativity

Atomic Radius

•Atomic Radius = half the distance between two nuclei of a diatomic molecule.

}Radius

Trends in Atomic Radius

• Influenced by three factors:1. Charge on nucleus– More charge pulls electrons in closer.

2. Energy Level– Higher E level is further away from nucleus

3. Shielding effect– The number of electrons between electrons and

nucleus affects the pull felt by the outer electrons

Atomic Radius Group trends

• As we go down a group...

• each atom has another energy level

• so the atoms get bigger.

HLi

Na

K

Rb

Atomic Radius Periodic Trends• As you go across a period, the radius gets

smaller.• The increasing number of protons in the

nucleus pulls the electrons in more tightly

Na Mg Al Si P S Cl Ar

Atomic RadiusAtomic Radii

0.000

0.050

0.100

0.150

0.200

0.250

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20atomic number, Z

Ato

mic

rad

ius,

nm

Atomic Radii

Ionic Size

• Ion: electrically charged atom• Cation: positively charged ion• Anion: negatively charged ion

• Ions aren't the same size as the neutral atoms they come from. – Compare the sizes of sodium and chloride ions with

the sizes of sodium and chlorine atoms.

• Positive ions are smaller than the atoms they come from.

• The sodium ion loses a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons.

• Negative ions are bigger than the atoms they come from.

• Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.

Study Buddy Review-A.R., I.R.

• Describe the pattern for atomic radius – As you move across a period– As you move down a column

• What charge does a cation have?

• What charge does an anion have?

• Which is larger than its parent atom, a cation or an anion?

First Ionization Energy

• Ionization energy is the energy required to remove the first electron from an atom of an element

• Elements want to have the e- configuration like that of a noble gas (filled)– Column 1A elements have need to LOSE one electron

to have noble gas configuration so it is EASY to remove electron

– Column 7A element need to GAIN one electron to have noble gas configuration, so it is HARD to remove electron

First Ionization Energy vs. Atomic Number

Ionization Energy

Ionization Energy• As you move down

a group ionization energy decreases…

• The electrons that are further away from the nucleus are easier to remove and thus require less energy to remove

• As you move across

a period ionization energy increases…

• Elements on left of table want to lose electrons to have full energy level (requires low energy to remove electron)

Successive Ionization Energies• more than one electron can be removed

from atoms

• Second Ionization energy: when a second electron is removed from an atom that has already lost one electron

• Third Ionization energy: when a third electron is removed from an atom that has already lost two electrons

Symbol First Second ThirdHHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11810 14840 3569 4619 4577 5301 6045 6276

Sucessive Ionization Energies

Relationship Between Common Charge and I.E.

• Consider Beryllium:

• Electron config: [He] 2s2

• Low energy to remove 1st and 2nd electrons

• MUCH higher energy to remove 3rd electron because it would be removed form a noble gas configuration

Study Buddy Review-I.E.

• What is ionization energy?

• Describe the pattern for ionization energy as you – Move down a family– Move across a row

• What does “first” ionization energy mean?

Electron AffinityElectron affinity is:

• the energy change associated with the adding an electron to a gaseous atom

• the more attraction for an electron the energy is released when the atom gains the electron – Released energy is negative (-350 kJ)

Electron Affinity

Electron AffinityGeneral Trend:

• Halogens (s2p5 configurations) are most negative electron affinities. They are most likely to want to gain electrons to obtain noble gas configuration

• As you go down a family, electron affinity is less negative (harder to gain electrons with increasing atomic size)

Study Buddy Review-E.A.

• What does it mean when an energy is negative?

• Which elements generally have a very negative electron affinity?

Electronegativity

• As you move down a group, electronegativity decreases

• As you move across a period, electronegativity increases

Electronegativity: the tendency for an atom to attract electrons to itself when it is chemically combined with another element.

Which element is the MOST electronegative?

Electronegativity

0

0.5

1

1.5

2

2.5

3

3.5

4

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

atomic number, Z

Electronegativity

Study Buddy Review-Electroneg

• Define electronegativity.

• Describe the pattern for electronegativity as you – Move down a group– Move across a period

• Which element is the most electronegative?

Resources• http://www.chemguide.co.uk/atoms/properties/

atradius.html• http://wine1.sb.fsu.edu/chm1045/notes/Periodic/

Affinity/Period05.htm• http://www.webelements.com/• http://www.public.asu.edu/~jpbirk/CHM-

113_BLB/Chpt07/sld017.htm• Jeanette Boles• Tina Lula• Dr. Stephen L. Cotton, Charles Page High School