Chapter 6.1-6.3 Periodic Table Lecture. Do members of the same family, generally behave the same?...
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Transcript of Chapter 6.1-6.3 Periodic Table Lecture. Do members of the same family, generally behave the same?...
![Page 1: Chapter 6.1-6.3 Periodic Table Lecture. Do members of the same family, generally behave the same? Yes.](https://reader036.fdocuments.in/reader036/viewer/2022081504/5697bfa01a28abf838c9512c/html5/thumbnails/1.jpg)
Chapter Chapter 6.1-6.36.1-6.3
Periodic Table LecturePeriodic Table Lecture
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Do members of the same family,
generally behave the same?Yes
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The Periodic TableThe Periodic Table
The Alkali MetalsLithium, Sodium, Potassium, Rubidium, Cesium, and francium very reactive 1 valence electron s1 sublevel is filled
Alkali Earth Metals
Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium2 valence electronss2 sublevel is filled
The Transition Metalsmetals with atomic numbers 21-112 highest s & d sublevels have electrons
MetalloidsLike metals & nonmetals
Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium
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Nonmetals• Consists of Carbon, Nitrogen, Oxygen, Phosphors, Sulfur, Selenium• poor conductors of heat and electricity compared to metals• dull and brittle
Halogens• Consists of Fluorine, Chlorine, Bromine, Iodine, Astatine• nonmetals• have 7 valence electrons• very reactive• want one more electron (octet rule)
Noble Gases• consists of Helium, neon, Argon, Krypton, Xenon, Radon• unreactive stable inert because they already have 8 valence electrons
Inner Transition metals• consists of elements with atomic numbers 58 through 71 and 90 through 103•F sublevels partially filled
• the Lanthanide Series has atomic numbers 58 -71 and the Actinide Series has atomic numbers 90-103
Other Metals
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Define the term inert gas?
noble gas –unreactive & stable
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Group 1A
Group 2A
Group 3A
Group 4A
Group 5A
Group 6A
Group 7A
Representative Elements #1 – Group IA-VIIA outer s & p orb partially filledAlkali Metals
Alkaline Earth
Nonmetals/Metalloids
Halogens
ns1
ns2
ns2 np1
ns2 np2
ns2 np3
ns2 np4
ns2 np5
Group 0
8 or 18
Noble Gases ns2 np6
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Representative Elements #1 Lewis dot structure
1s2 2s2 2p6 1s2Na
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Group B Transition Metals
Filling the “d” orbital
Group 58-71 Lanthanides Filling the “4f” orbital
Group 90-103Actinides Filling the “5f” orbital
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A. Ionic SizeA. Ionic Sizemetals (group 1A-3A)
lose electrons to become stable cation
non-metal (group 5A-7A) gain electrons to become stable anions.
1A =
2A =
3A =
5A =
6A =
7A =
Loses 1 e-
Loses 2 e-
Loses 3 e-
Gains 3 e-
Gains 2 e-
Gains 1 e-
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7PERIODS
! !v
A Family is a Group living between Columns
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Periodic Table Song by Tom Lehrer abovePeriodic Table Song by Tom Lehrer above
End of Lecture 6.1End of Lecture 6.1
Next Lecture 6.2Next Lecture 6.2
http://www.privatehand.com/flash/elements.html
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Who designed the 1st periodic table in 1869?
Dmitri Mendeleev
grouped w/ similar chemical and physical properties & ordered by atomic mass.Ex:
Co Ni
Ar K
Te I
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http://www.youtube.com/watch?v=y7dmRtlXaYQ
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http://www.youtube.com/watch?v=zUDDiWtFtEM
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Lecture 6.3Lecture 6.3Periodic TrendsPeriodic Trends
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I. Periodic Trends - Atomic Size
Atomic Radii:
NucleusDistance between
nuclei
Atomic Radius
Measured as 1/2 distance between nuclei 2 atoms
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Atomic Size generally INCREASES as you move down a group on the periodic table.
Why? down a group
increases # of energy levels
Example:Ca atom larger than a Mg atom. Why?
An energy level is added!
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Atomic Size generally DECREASES across a row on the periodic table.Why? adding more p+ pulls in extra
electrons
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Na < ionization energy than O because less protons pull.
RELATIVE ELECTRONEGATIVITY, IONIZATION RELATIVE ELECTRONEGATIVITY, IONIZATION ENERGY, RADII, SHIELDING ETC…ENERGY, RADII, SHIELDING ETC…
Hydrogen
2.1
Oxygen
3.5
Carbon
2.5
Sodium
0.9
Electro negativities:Hydrogen has the smallest atomic radius
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B. Ionization EnergyB. Ionization Energyenergy needed to pull an electron away from an atom.
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B. Ionization EnergyB. Ionization Energy
Example : Na Na+1 + e-
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Ionization energy decreases as you move down a group.increased distance from protons
reduces attractive force
Why?
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Period TrendPeriod Trend: : Ionization energy generally increases as you move across a period.
nuclear charge increases (more protons)
which increases attractive forces
Why?
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energy required to remove the energy required to remove the 1st1st outermost electron is outermost electron is 1st ionization 1st ionization energy.energy.
What is the second ionization energy?
Which is harder to remove?
Why?
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What happens to the shielding of the nucleus as you move across a period?
•ONLY adding electrons, NOT a new energy level.
Remains constantWhy?
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What happens to the shielding of the nucleus as you move down a group?
another energy level that shields those valence electrons.
IncreasesWhy?
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CaCa++ions ions –– smallersmaller than the original than the original atomatom
When electrons lost,
a whole energy level lost
decreases radius.
Why?
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Negative anions grow Negative anions grow largerlarger
there are more e- than p+
(increased electron repulsion),
Why?
Natom N-3
anion
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from group 5A to the right,from group 5A to the right,aannions ions gradually decreasegradually decrease in sizein size
groups 6A &7A only gain 1 or 2 e-
Have Same # of e-, but increased # of p+
Why?
N-3 O-2 F-1
anion anion anion.
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B. ElectronegativityB. Electronegativity
Noble gases no electronegative #
Why?inert / don’t form compounds.
Can’t force a noble gas to take an electron – they have s2 p6
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3. Period Trends3. Period Trendsleft to right electronegativity increases. Why?High ionization energy = high electronegativity
Resists electron
loss
Attracts electrons
Fluorine is the most electronegative!
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4. Group Trends4. Group TrendsElectronegativity decreases down a group.
Why?
Increased energy levels and shielding
Cs has the lowest electronegativity
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3 alkalis.MOV
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