PERIODIC TABLEThe most awesome chemistry tool ever!

UNIT OBJECTIVES:

Understand the history of who built up the periodic table, how they did it, and what law was made

Become familiar with structure of periodic table, how the e- is related to the structure of the table, and properties of chemicals based on their location

Discover trends related to electron configuration and periodic properties; use electron configuration to predict location on periodic table.

HISTORY• Explain the roles of Mendeleev and Moseley in

the development of the periodic table.

• Describe the modern periodic table.

• Explain how the periodic law can be used to predict the physical and chemical properties of elements.

• Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number.

MENDELEEV AND PERIODICITY

• In the late 1800s, scientist Mendeleev noticed:

• when the elements were arranged in order of

increasing atomic mass, certain similarities in

their chemical properties appeared at regular

intervals and formed a pattern.

• Repeating patterns are referred to as periodic. • Something that occurs periodically occurs at

regular/fixed intervals

MENDELEEV AND PERIODICITY

• Not all elements had been discovered, so

when arranging the table by atomic mass,

• He left spaces for future elements to be

discovered and placed into the table based on

their mass

• He even predicted what their properties would

likely be

MENDELEEV AND PERIODICITY

Where Mendeleev’s work left off there were some questions:

Why could most elements be arranged in the order of increasing atomic mass and other not?

What was the reason behind chemical periodicity?

MOSELEY AND THE PERIODIC LAW

Moseley continued research on the periodic table and was able to propose an answer to the first question.

Elements follow a clearer pattern when arranged by their nuclear charge (# of protons) instead of their nuclear mass!

Why do you think this is? Hint: remember Hydrogen-3 and Helium-3 and

their similarities and differences?

MOSELEY AND THE PERIODIC LAW

A big example of what Moseley did: he realized that Tellurium (which has an atomic mass of about 127.6 amu) should be before Iodine (which has an atomic mass of 126.9 amu) on the periodic table because Tellurium has the atomic number of 52 and Iodine has the atomic number of 53.

MOSELEY AND THE PERIODIC LAW

Moseley led to the editing of Mendeleev’s principle of chemical periodicity:

Periodic law: the physical and chemical properties of the elements are periodic functions of their atomic numbers.

(When elements are arranged by increasing atomic number, similar properties occur in elements at regular intervals)

MODERN PERIODIC TABLE

After all the elements had been discovered (some synthesized as well) our modern periodic table was been established

Periodic table: an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall into the same column or “group”.

SPECIFIC GROUP: NOBLE GASES

Discovered in the late 1800s, the noble-gases were slowly discovered, often due to indirect observation Helium was discovered due to its emission

spectrum in sunlight

Argon was not-reactive and was deduced when mass of an unknown gas could not be accounted for when all other gases had been reacted out.

LANTHANIDES AND ACTINIDES HAVE THEIR OWN BLOCK

Lanthanides: The elements between Cerium (Atomic # 58)

and Lutenium (Atomic # 71) did not follow the properties of the elements around them

Whereas Lanthanum (atomic # 57) and Hafnium (atomic # 72) do fit in their groups

Actinides: The elements between atomic #s: 90 and 103 Many are synthetic and they do not follow the

properties of the metals around them in the periodic table

PERIODICITY

We will begin to explore the similar properties groups of the periodic table exhibit.

Groups to begin thinking about:Alkali Metals (s-block)Alkaline Earth Metals (s-block)Main-group elements (p-block

Halogens (p-block) Noble Gases (p-block)

Transition metals (d-block)

ELECTRON CONFIGURATION AND PERIODIC PROPERTIESObjectives:• Explain the relationship between electrons in

sublevels (shapes) and period length in the periodic table.

• Locate and name the four blocks of the periodic table. Explain the reasons for these names.

• Discuss the relationship between group configurations and group numbers.

• Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, and the noble gases.

PERIODS AND BLOCKS OF THE PERIODIC TABLE

Periods are the horizontal rows of the periodic table.

Elements are arranged vertically in the periodic table in groups that share similar chemical properties.

The four blocks: the s, p, d, and f blocks each block corresponds to the electron sublevel

being filled in that block.

PERIODS & BLOCKS OF THE PERIODIC TABLE

Period # Sublevels (shapes) in order of filling

# of elements in period

1 1s 2

2 2s 2p 8

3 3s 3p 8

4 4s 3d 4p 18

5 5s 4d 5p 18

6 6s 4f 5d 6p 32

7 7s 5f 6d 7p 32

S-BLOCK ELEMENTS• The elements of Group 1 of the periodic

table are known as the alkali metals. • lithium, sodium, potassium, rubidium, cesium, and

francium

• In their pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife.

• The elements of Group 2 of the periodic table are called the alkaline-earth metals.

• beryllium, magnesium, calcium, strontium, barium, and radium

• Group 2 metals are less reactive than the alkali metals, but are still too reactive to be found in nature in pure form.

HYDROGEN AND HELIUM Hydrogen is NOT a metal;

Its properties are very different from the metals in the alkali metal group

Helium fits in with the Noble Gases as it has a full valence shell and is not at reactive Alkaline earth metals have a full s-sublevel, but

their whole valence shell is not full (their p-sublevel is empty)

A valence shell is the highest energy level that electrons occupy. Elements are more stable/less reactive when their valence shells are full of electrons

THE d-BLOCK ELEMENTS (GROUPS 3-12)

Transition metals: the d-block elements are metals with typical metallic properties Ductility Malleability Conductivity Having luster

Groups and periods within the transition metals have varied properties and it is therefore hard to group metals by individual properties

p-BLOCK ELEMENTS (GROUP 13-18)

Main-group elements: the elements in the p-block of varying properties

All non-metals are in the p-block (aside from hydrogen)

All metalloids (boron, silicon, germanium, arsenic, antimony, and tellurium) are in the p-block

RELATIONSHIPS BETWEEN GROUP #, BLOCKS, AND ELECTRON CONFIG.

Group #

Group configuration

Block Comments

1,2 ns1, ns2 s 1 or 2 electrons in ns sublevel

3-12 (n - 1)d1-10ns0-2 dThe sum of electrons in ns and (n-1)d equals group number

13-18 ns2np1-6 p

Number of electrons in np sublevel equals group number minus 12

HALOGENS

Elements in group 17 are the halogens Fluorine, chlorine, bromine, iodine, and astatine

They are highly reactive in similar ways All form salts with group 1 and 2 metals

They are volatile ( gas forms easily)

METALLOIDS

Metalloids have properties of both metals and non-metals boron, silicon, germanium, arsenic, antimony,

and tellurium Semiconducting, brittle, harder than most

metals, have an opalescent luster

Relatively reactive (Bi only found in elemental form)

f-BLOCK ELEMENTS LANTHANIDES AND ACTINIDES

Lanthanides are shiny metals and have similar reactivity to group 2 metals

Their position reflects the fact that they involve the filling of the 4f sublevel

Actinides: Only 4 are found in nature (thorium through

neptunium)

Alkali Metals

Other Metals

Inner transition Metals

Alkaline Earth Metals

Metalloids

Other Non-Metals

Halogens

Noble-Gases

DO NOW:• List properties of the

following elemental groups:• Table 1: Halogens• Table 2: Noble Gases• Table 3: Alkali Metals• Table 4: Alkaline Earth

Metals• Table 5: Transition metals• Table 6: Actinides

• Share with your table

• Describe what you know about an element by looking at its position in the periodic table. • Table 1: Br• Table 2: Ar• Table 3: K• Table 4: Ca• Table 5: Cu• Table 6: Fe

• Identify any noticeable trends.

DO NOWIf you were absent Friday: How did Mendeleev organize

the periodic table?

Why didn’t this always work?

How did Moseley organize the periodic table?

What are the groups of the periodic table? Include at least 1 property

for each group

If you were here Friday: Please take out your

practice problem sheet and complete the problems for Atomic radius and ionization energy using your graphic organizer

Look over your SRQ Answers

Help your neighbor with their do now if they were absent Friday

ATOMIC RADII (THE DISTANCE BETWEEN THE NUCLEUS & THE END OF THE CLOUD)

One half the distance between the nuclei of identical atoms that are bonded together

Atomic radii: 200 pm

Bond length400 pm

ATOMIC RADII Increases in size from right to left Increases in size from up to down

IONIZATION ENERGY (HOW TO KNOCK AN ELECTRON OFF!) An ion is an atom that

has more or less electrons than it has protons.

is either positively (cation) or negatively (anion) charged

Ionization is anything that gives matter a charge

Ionization Energy (IE) is the amount of energy needed to remove an electron from an atom

IONIZATION ENERGY CONTINUED

Ionization often occurs because The valence electrons are more stable in the

ionic configuration To complete a valence shell

Gaining an electron can cause a neutral atom to lose or gain energy depending on the atom A metal gaining an electron would absorb energy

and become less stable A non-metal gaining an electron would release

energy and become more stable

IONIZATION ENERGY (HOW TO KNOCK AN ELECTRON OFF!) Ionization energy trends:

Increases from left to right Increases from down to up

ADDITIONAL IONIZATION ENERGIES

Removing a second electron is the 2nd Ionization Energy, and it typically greater than the first ionization energy

With each additional electron removed, the Ionization Energy will increase.

ELECTRON AFFINITY The change in energy associated with the addition

of an electron to a neutral atom

Green means no electron affinity

Yellow means high electron affinity (trends by groups)

IONIC RADII Positive ion = cation Negative ion = anion

IONIC RADII

VALENCE ELECTRONS

Valence electrons are available to be lost, gained, or shared during chemical reactions and compound formation.

Group #

Group e- Config

# of valence e-

1 ns1 1

2 ns2 2

13 ns2p1 3

14 ns2p2 4

15 ns2p3 5

16 ns2p4 6

17 ns2p5 7

18 ns2p6 8

ELECTRONEGATIVITY Electronegativity how much an atom in a

compound attracts electrons from another atom in the compound (how much the atom is greedy for electrons)

PERIODIC PROPERTIES OF THEd-BLOCK AND f-BLOCK

Atomic radii decrease from left to right

Ionization energy increases from left to right

Electronegativity slightly increase from left to right

COMMON ELEMENTAL IONS+1

+2

+3 -1-2-3

8.2

WHY?

A lot of these trends are related to electron shielding

Electron shielding is how the inner electrons effect how the outer electrons act due to charge different and repulsion