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Chapter 6: The Periodic Table
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JW Dobereiner Early 1800s Classified elements into triads based on similar properties Ca, Sr, Ba Li, Na, K Cl, Br, I
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J.A.R. Newlands Law of Octaves – every 8 elements begin to repeat in patterns Wasn’t taken seriously because he likened chemistry to music
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Dmitri Mendeleev Father of modern periodic table
Arranged the elements according to increasing atomic mass Left blanks for elements that hadn’t been discovered but that he predicted existed Made some exceptions (Te and I) because he knew that iodine was more closely related to Br based on its properties
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Henry Moseley Assigned atomic numbers to elements (number of protons)
Arranged the elements in order of increasing atomic number No need to make exceptions like Mendeleev had to
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Periodic Law There is a periodic repetition of chemical and physical properties of elements when they are arranged by increasing atomic number.
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Groups – Vertical families
Group 1 or 1A-Alkali Metals Group 2 or 2A-Alkaline Earth Metals Group 13 or 3A-Boron Family Group 14 or 4A-Carbon Family Group 15 or 5A-Nitrogen Family Group 16 or 6A-Oxygen Family Group 17 or 7A-Halogen Family Group 18 or 8A-Noble Gas Family
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Groups Note: All of the Group A elements are called the “Representative Elements.” Group A # = # of valence electrons The other elements (Group B elements) are known as the transition and inner transition elements. The Octet Rule- atoms tend to gain, lose, or share electrons to acquire a full set of eight valence electrons.
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Some classifications Metals – grey, typically lustrous or shiny, good conductors of heat and electricity, and are generally solids at R.T Nonmetals – pink, have opposite properties. They can be solid, liquid or gas at R.T Metalloids/Semi-metals – blue, along the zigzag
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METALS Non-METALS METALLOIDS
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Periodic Trends Electron Shielding – when the presence of core electrons diminishes the effect the nuclear charge has on the valence electrons Atomic Radii – 1/2 the distance between the nuclei of two atoms of the same element when the atoms are joined
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Atomic radii: 200 pm Bond length 400 pm
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Periodic Trends Atomic Radii Across a period – tends to decrease
Down a group – atomic radii tends to increase Why? Electrons enter higher principle energy levels Across a period – tends to decrease # of protons increase, electrons enter same principal energy level, no increase in shielding but increase in nuclear charge
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Atomic Radii
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Periodic Trends Ionization energy – the energy required to remove an electron from a gaseous atom Down a group – tends to decrease Why? Electrons are farther from nucleus therefore easier to pull off Across a period – tends to increase More difficult to remove an electron as you get closer to noble gas configuration
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There are lots of electrons to remove!
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Ionization Energy
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First Ionization energy
He He has a greater IE than H. same shielding greater nuclear charge H First Ionization energy Atomic number
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outweighs greater nuclear charge First Ionization energy
He Li has lower IE than H more shielding further away outweighs greater nuclear charge H First Ionization energy Li Atomic number
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greater nuclear charge First Ionization energy
He Be has higher IE than Li same shielding greater nuclear charge First Ionization energy H Be Li Atomic number
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greater nuclear charge By removing an electron we make s orbital full
He B has lower IE than Be same shielding greater nuclear charge By removing an electron we make s orbital full First Ionization energy H Be B Li Atomic number
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First Ionization energy
He First Ionization energy H C Be B Li Atomic number
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First Ionization energy
He N First Ionization energy H C Be B Li Atomic number
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First Ionization energy
He Breaks the pattern because removing an electron gets to 1/2 filled p orbital N First Ionization energy H C O Be B Li Atomic number
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First Ionization energy
He F N First Ionization energy H C O Be B Li Atomic number
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First Ionization energy
He Ne Ne has a lower IE than He Both are full, Ne has more shielding Greater distance F N First Ionization energy H C O Be B Li Atomic number
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Na has a lower IE than Li Both are s1 Na has more shielding
He Ne Na has a lower IE than Li Both are s1 Na has more shielding Greater distance F N First Ionization energy H C O Be B Li Na Atomic number
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First Ionization energy
Atomic number
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Periodic Trends Electronegativity – the relative ability of an atom to attract an electron when the atom is in a compound Down a group – tends to decrease Why? The shielding effect is greater Across a period – tends to increase (excluding noble gases) Greater ‘desire’ to achieve noble gas configuration
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Electronegativity
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Periodic Trends Ionic Size metals - Cations
smaller than their respective atoms Why? # of electrons is less # of protons Example: Na+ is smaller than Na
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Periodic Trends Ionic Size Non-metals - Anions
larger than their respective atoms Why? # of electrons is larger than # of protons Example: Cl- is bigger than Cl Comparing ions to ions regarding size Down a group – ions tend to increase in size Across a period – tend to decrease
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Size of Isoelectronic ions
Iso - same Iso electronic ions have the same # of electrons Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3 all have 10 electrons all have the configuration 1s22s22p6
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Size of Isoelectronic ions
Positive ions have more protons so they are smaller. N-3 O-2 F-1 Ne Na+1 Al+3 Mg+2
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