The Periodic Table Chapter 6 Notes. History of the PT Dobereiner –German Chemist –Proposed...

The Periodic Table

Chapter 6 Notes

History of the PT

• Dobereiner– German Chemist– Proposed “triads” in

1829: grouping of 3 elements with similar properties

– Observed that melting point decreased and density increased as the atomic mass of an element increased

• Mendeleev– Russian Chemist– Published a table of

elements in 1869 in which the elements were organized by increasing atomic mass

– First periodic table to use rows and columns– Left “missing spots” for elements not yet

discovered– Tellurium (mass = 128) and Iodine (mass =

127) were in the wrong order – thought the masses were wrong

Mendeleev’s Periodic Table

The Modern PT

• Very similar to the periodic table developed by Mendeleev

• Organized by horizontal rows (called periods) and vertical columns (called groups or families)

• Shows a pattern of physical and chemical properties

• Uses atomic number to order elements, NOT atomic mass

The Modern Periodic Table

Within a Period (Row)

• Add one proton and one electron each time you move one element across a row

• 1A – s1

• 2A – s2

• 3A – s2p1

• 4A – s2p2

• 5A – s2p3

• 6A – s2p4

• 7A – s2p5

• 8A – s2p6

Within a Group (Column):1

2 3 4 5 6 7

8

The number of valence electrons (outermost electrons) =

the group number

• Because the number of valence electrons within a group are the same, elements within a group have similar physical and chemical properties

• For example, what happens when an alkali metal comes into contact with water?

Alkali metals react violently with water to form

hydrogen gas and release a lot of energy!

Classification of Elements

• Metals– Have luster (shine)– Conduct heat and electricity– Malleable (can bend without breaking)– Solids (except for mercury)– High melting points– Groups 1A, 2A and all of Group B

• Group B – Transition Metals• Bottom 2 Rows – Inner Transition Metals

– Lanthanides – top row– Actinides – bottom row (also called the Rare Earth

Metals)

Classification of Elements

• Nonmetals– Do not conduct heat or electricity– Are brittle– Many are gases at room temperature– Do not have luster– Lower melting points– Located in Groups 3A – 8A

Classification of Elements

• Metalloids– Have some properties of metals and other

properties of nonmetals depending on the conditions the element is under (temperature, pressure, etc.)

– Border the metals and nonmetals on the periodic table

– Semiconductors: conduct electricity better than a nonmetal but not as well as a metal

Group Characteristics• Representative

Elements– All Group A elements– Represent the entire

spectrum of element characteristics

• Metals• Nonmetals• Metalloids• Solids• Liquids• Gases

Transition Elements

– All Group B elements

– Show a gradual transition in properties from metals to nonmetals

– Used on alloys (metal mixtures)

– Have high melting and boiling points

Inner Transition Metals

– “f block”– Lanthanides – top row

• Also called the “Rare Earth Metals”

• Found naturally in the Earth’s crust

• used in the glass and TV picture tube industry

– Actinides – bottom row• All are radioactive• U and Pu are used as

nuclear fuels• Some are used for

cancer therapy

Alkali Metals

• Group 1A• All have one valence

electron• All end in the electron

configuration s1

• Soft, silvery white metals• Good conductors of heat

and electricity• Very reactive

Alkaline Earth Metals

• Group 2A• All have 2 valence

electrons• All end in the

electron configuration s2

• Less reactive than the alkali metals

• More dense, harder, and have a higher melting point

Boron Group

• Group 3A

• All have 3 valence electrons

• All end in the electron configuration s2p1

• Not very reactive

• Silvery, fairly soft, and good conductors of heat and electricity

Carbon Group

• Group 4A• All have 4 valence

electrons• All end in the electron

configuration s2p2

• Form bonds by sharing electrons

• Mixture of metals, nonmetals, and metalloids

Nitrogen Group

• Group 5A

• All have 5 valence electrons

• All end in the electron configuration s2p3

• Nonmetals and metalloids

Chalcogens• Group 6A

• All have 6 valence electrons

• All end in the electron configuration

s2p4

• Nonmetals and metalloids

• Very reactive

Halogens

• Group 7A• All have 7 valence electrons• All end in the electron

configuration s2p5

• All are nonmetals• Exist as diatomic molecules – two

atoms of the same element bonded together (Cl2, F2)

• Very reactive

Noble Gases• Group 8A• All have 8 valence

electrons (except for Helium which has 2 valence electrons

• All end in the electron configuration s2p6

• All are gases• All have a full outer shell

which causes them to be unreactive (previously called inert)

Periodic Trends• Atomic Radius

– The radius of an atom– As you move across a period, the atomic

radius decreases because the greater number of protons in the nucleus increases the magnetic pull on the electrons making the atom smaller

3p+

4n0

3e-

4p+5n0

4e-Lithium atom

Beryllium atom

– As you move down a group, the atomic radius increases because there are more energy levels so electrons are farther away from the nucleus

Lithium atom Sodium atom

2 energy levels

3 energy levels

• Ionic Radius– The radius of an ion (an ion is an atom that

has gained or lost electrons causing it to have an overall positive or negative charge)

– As you move across a period, the ionic radius decreases because the greater number of protons in the nucleus increases the magnetic pull on the electrons making the atom smaller

– The trend is NOT smooth across a period because metals lose electrons and nonmetals gain electrons – there is a spike around the metalloids on the periodic table

– As you move down a group, the ionic radius increases because there are more energy levels so electrons are farther away from the nucleus

• Ionization Energy– The energy required to REMOVE an electron

from an atom– As you move across a period, ionization energy

increases because metals easily lose electrons and nonmetals strongly hold onto their electrons, so it would require less energy to remove an electron from a metal and more energy to remove an electron from a nonmetal

– As you move down a group, ionization energy decreases because the further the electron is from the nucleus, the less attraction there is between the protons in the nucleus and the electrons, so it is easier to remove the electron from the shell

• Electronegativity– How well an atom attracts an electron for

bonding– As you move across a period,

electronegativity increases because metals then to lose electrons and nonmetals tend to gain electrons

– As you move down a group, electronegativity decreases because the energy levels are farther away from the nucleus and the magnetic pull between the positive nucleus and the electrons decreases