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New topic The Periodic Table
The how and why
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The Modern Table A. Organization of the P.T.
Elements are still grouped by properties Similar properties are in the same column Order is in increasing atomic number Added a column of elements Mendeleev didn’t know about. The noble gases weren’t found because they didn’t react with anything.
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Horizontal rows are called periods
There are 7 periods
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Vertical columns are called groups.
Elements are placed in columns by similar properties. Also called families
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The elements in these groups are called the representative elements
18 1 2 13 14 15 16 17
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Other Systems IA IIA IIIB IVB VB VIB VIIB VIIIB IIIA IVA VA VIA VIIA
VIIIA IB IIB 1 2 13 14 15 16 17 18 3 4 5 6 7 8 9 10 11 12 1A 2A 3A 4A 5A 6A 7A 8A 3B 4B 5B 6B 7B 8B 1B 2B
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B. Metals
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Metals Luster – shiny. Ductile – drawn into wires.
Malleable – hammered into sheets. Conductors of heat and electricity.
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C. Transition metals The D and F Block elements
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D. Non-metals Dull Brittle Nonconductors- insulators
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Metalloids or Semimetals
Properties of both Semiconductors
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These are called the inner transition elements and they belong here
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Group 1 are the alkali metals
Group 2 are the alkaline earth metals
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Group 17 is called the Halogens
Group 18 are the noble gases
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Why? The part of the atom another atom sees is the electron cloud.
More importantly the outside orbitals The orbitals fill up in a regular pattern The outside orbital electron configuration repeats So.. the properties of atoms repeat.
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H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p63d104s24p65s1 1s22s22p63s23p63d104s24p64d105s2 5p66s1 1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p67s1
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Identifying the patterns
Periodic trends Identifying the patterns
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What we will investigate
Atomic size how big the atoms are Ionization energy How much energy to remove an electron Electronegativity The attraction for the electron in a compound Ionic size How big ions are
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What we will look for Periodic trends-
How those 4 things vary as you go across a period Group trends How those 4 things vary as you go down a group Why? Explain why they vary
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The why first The positive nucleus pulls on electrons
Periodic trends – as you go across a period The charge on the nucleus gets bigger The outermost electrons are in the same energy level So the outermost electrons are pulled stronger
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The why first The positive nucleus pulls on electrons Group Trends
As you go down a group You add energy levels Outermost electrons not as attracted by the nucleus
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Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus +
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Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus A second electron has the same shielding In the same energy level (period) shielding is the same +
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Shielding As the energy levels changes the shielding changes
Lower down the group More energy levels More shielding Outer electron less attracted + Three shields Two shields No shielding One shield
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Atomic Size First problem where do you start measuring
The electron cloud doesn’t have a definite edge. They get around this by measuring more than 1 atom at a time
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Atomic Size } Radius Atomic Radius = half the distance between two nuclei of molecule
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Trends in Atomic Size Influenced by two factors Energy Level
Higher energy level is further away Charge on nucleus More charge pulls electrons in closer
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Group trends H Li Na K Rb As we go down a group
Each atom has another energy level More shielding So the atoms get bigger Li Na K Rb
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Periodic Trends Na Mg Al Si P S Cl Ar
As you go across a period the radius gets smaller. Same shielding and energy level More nuclear charge Pulls outermost electrons closer Na Mg Al Si P S Cl Ar
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Rb K Overall Na Li Atomic Radius (nm) Kr Ar Ne H 10 Atomic Number
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Ionization Energy The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a +1 ion The energy required is called the first ionization energy
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Ionization Energy The second ionization energy is the energy required to remove the second electron Always greater than first IE The third IE is the energy required to remove a third electron Greater than 1st or 2nd IE
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Symbol First Second Third
HHeLiBeBCNO F Ne
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What determines IE The greater the nuclear charge the greater IE.
Increased shielding decreases IE Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE
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Group trends As you go down a group first IE decreases because of
More shielding So outer electron less attracted
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Periodic trends All the atoms in the same period Same shielding.
Increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 full orbitals
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He has a greater IE than H same shielding greater nuclear charge
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 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 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 F N First Ionization energy H C O Be B Li Na Atomic number
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Web elements First Ionization energy Atomic number
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Driving Force Full Energy Levels are very low energy
Noble Gases have full orbitals Atoms behave in ways to achieve noble gas configuration
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Ionic Size Cations are positive ions Cations form by losing electrons
Cations are smaller than the atom they come from Metals form cations Cations of representative elements have noble gas configuration.
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Ionic size Anions are negative ions Anions form by gaining electrons
Anions are bigger than the atom they come from Nonmetals form anions Anions of representative elements have noble gas configuration.
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Configuration of Ions Ions of representative elements have noble gas configuration Na is 1s22s22p63s1 Forms a 1+ ion - 1s22s22p6 Same configuration as neon Metals form ions with the configuration of the noble gas before them - they lose electrons
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Configuration of Ions Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them.
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Group trends H1+ Li1+ Na1+ K1+ Rb1+ Cs1+ Adding energy level
Ions get bigger as you go down Li1+ Na1+ K1+ Rb1+ Cs1+
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Periodic Trends N3- O2- F1- B3+ Li1+ C4+ Be2+
Across the period nuclear charge increases so they get smaller. Energy level changes between anions and cations N3- O2- F1- B3+ Li1+ C4+ Be2+
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Size of Isoelectronic ions
Iso - same Iso electronic ions have the same # of electrons Al3+ Mg2+ Na1+ Ne F1- O2- and N3- 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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Electronegativity
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Electronegativity The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How “greedy” Big electronegativity means it pulls the electron toward it.
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Group Trend The further down a group More shielding
more electrons an atom has. Less attraction for electrons Low electronegativity.
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Periodic Trend Metals - left end Low nuclear charge Low attraction
Low electronegativity Right end - nonmetals High nuclear charge Large attraction High electronegativity
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How to answer why questions
Trend Periodic Group Reason Nuclear charge Energy level and shielding Result What happens to which electron
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Nuclear Charge Energy Levels & Shielding
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Ionization energy, electronegativity
INCREASE
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Atomic size increases, Ionic size increases
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Other topics Reactivity Alkali metals most reactive
Nobel gases least reactive Noble gases found as free elements Gr 1,2,17 - never found free Metallic properties Group Period
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Properties of Groups Hydrogen By itself Can gain or lose electrons
Also shares electrons 1 and 2 Metallic characteristics Lose electrons,low electroneg and IE Never found in atomic state (compounds) Reactivity inc as you go down gr BP/MP no trend
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Properties of Groups 15 Nonmetallic to metallic change
N and P gain 3 electrons-nonmetal As and Sb metalloids Bi loses elecrons N-fixing bacteria, less reactive than P P- 4 bonds, more reactive than N
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Properties of Groups 16 O & S gain 2 electrons nonmetals
S loses 2 or 4 electrons O most important Very reactive Diatomic except photosynthesis Se & Te metalloids Po metal
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Properties of Groups 17 Halogens in free state
Gain 1 electron called halides Nonmetals, w/ metal characteristics All three states of matter at RT Occur in nature as compounds F most reactive halogen BP/MP increases from top to bottom
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Properties of Groups 18 Monatomic molecules He 2 valence electrons
The rest = 8 valence electrons Some react w/ Florine BP/MP increase from top to bottom
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