periodic table the most awesome chemistry tool ever!
TRANSCRIPT
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