the periodic table and physical properties
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SONG. The Periodic Table and Physical Properties. Get a periodic table out. Topics 3.1 - 3.3 and 12.1.1 - 12.1.2. Dmitri Mendeleev 8 February 1834 – 2 February 1907. Russian chemist and teacher - PowerPoint PPT PresentationTRANSCRIPT
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The Periodic Table and Physical Properties
SONG
Topics 3.1 - 3.3 and 12.1.1 - 12.1.2
Get a periodic
table out.
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Dmitri Mendeleev 8 February 1834 – 2 February 1907
• Russian chemist and teacher• given the elements he knew
about, he organized a “Periodic Table” based on increasing atomic mass (it’s atomic # now)
• he even left empty spaces to be filled in later (TOK– he was a “scientist” and “risk taker”!)
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At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic
table. He predicted their discovery and estimated their properties.
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Z
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Design of the Table• Groups are the vertical columns.
– elements have similar, but not identical, properties• most important property is that
they have the same # of valence electrons
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valence electrons- electrons in the highest occupied energy level
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http://images.google.com/imgres?imgurl=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/imgper/econfig.gif&imgrefurl=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/perlewis.html&h=267&w=512&sz=22&tbnid=__EXctBwlG0J:&tbnh=66&tbnw=128&hl=en&start=1&prev=/images%3Fq%3DElectron%2BDot%2BDiagrams%26svnum%3D10%26hl%3Den%26lr%3D
2
2,3 2,5 2,8
2,8,2
Electron arrangement (SL level – 3.1.3)
2,1
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• B is 1s2 2s2 2p1;– 2 is the outermost energy level – it contains 3 valence electrons, 2 in
the s and 1 in the p• Br is [Ar] 4s2 3d10 4p5
How many valence electrons are present?
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• Periods are the horizontal rows– do NOT have similar properties– however, there is a pattern to their properties as
you move across the table that is visible when they react with other elements
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Definitions• atomic radii
– the distance from the nucleus to the outermost electron
• ionic radii– same distance, but for ions (atoms that have lost
or gained valence electrons)• first ionization energy (kJ mol-1)
– the energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gaseous phase
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• electronegativity– measures the attraction for a shared pair of
electrons• melting point• chemical properties
– how elements react with other elements
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But first, the electron shielding effect
• electrons between the nucleus and the valence electrons repel each other
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• ATOMIC RADII – McGraw Hill video – groups (alkali metals and halogens)
• increases downwards as more levels are added– periods across the periodic table (period 3)
• radii decreases – the number of protons in the nucleus increases
» increases the strength of the positive nucleus and pulls electrons closer to it
H
Li
Na
K
Rb
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• IONIC RADII – decreases across periods for same reason
as atomic radii (nucleus becomes stronger)– alkali metals
• cations are smaller that the parent atom– have lost an electron (actually, lost an entire level)– therefore have fewer electrons than protons
• radii still increases downwards as more levels are added on
Li 0.152 nm3e and 3p
Li+, 0.078 nm
2e and 3 p
+forming a
cation
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–halogens• anions are larger than parent atom
– have gained an electron to achieve noble gas configuration
• radii still increases downwards as more levels are added on
F 0.064 nm9e- and 9p+
F- 0.133 nm10 e- and 9 p+
- forming an
anion
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– IONIZATION ENERGY• decreases down a group
– outer electrons are farther from the nucleus and therefore easier to remove
– inner core electrons “shield” the valence electrons from the pull of the positive nucleus
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• increases across a period– extra electrons are just filling up the same level– the nucleus is becoming more powerful and
therefore the electrostatic force increases making it harder to remove an electron
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12.1.1
• Evidence for levels and sub-levels – First ionization energy
• electrons are harder to remove…– when there are more protons to attract them– a sub-level (s,p,d,f) is completely filled– a sub-level (s,p,d,f) are half filled
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12.1.1
• Evidence for levels and sub-levels – successive (1st, 2nd, 3rd) ionization energy
• as more electrons are removed, the electrostatic pull of the protons holds the remaining electrons closer
• therefore, more energy is required to remove them (even have to use a logarithmic scale to show this)
• large “jumps” are when the electrons are being removed from the next, lower level that are much closer to the nucleus
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4s1 removed
starting to remove the 3p sub-level
starting to remove the 2p
sub-level
starting to remove the 3s
sub-level
starting to remove the 2s sub-level
starting to remove
the 1s sub-level
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–ELECTRONEGATIVITY• as you go down a group electronegativity
decreases – the size of the atom increases
» the bonding pair of electrons (-) is increasingly distant from the attraction of the nucleus (+)
» the valence electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons
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• as you go across a period– electronegativity increases
• the atoms become smaller so the positive nucleus can hold onto the electrons better
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–MELTING POINT• group 1 (alkali metals)
– decreases as “sea of negative electrons” are farther away from the positive metal ions
• group 7 (halogens)– increases downwards as the van der Waals’ forces increase
» larger molecules have more electrons which increases the chance that one side of the molecule could be negative
Element Melting Point (K)
Li 453Na 370K 336Rb 312Cs 301Fr 295
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increases decreases
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• across the table (period 3)– from left to right
• bonding goes from strong metallic to very strong macromolecules (network covalent) to weak van der Waals’ attraction
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• CHEMICAL PROPERTIES–groups
• alkali metals– react vigorously with water and air
» 2Na (s) + H2O (l) 2Na (aq) + 2OH- (aq) + H2 (g)» (Li, Na, K… all the same equation)
» reactivity increases downwards» because the outer (valence) electron is in higher energy
levels (farther from the nucleus) and easier to remove – react with the halogens
» halogens’ reactivity increases upwards» smaller size can attract electrons
better» (see next slide)
1+ charge
1- charge
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most reactive
least reactive
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• halogens–diatomic molecules such as F2, Cl2, Br2, I2
» can react with halide ions (Cl -, Br -, and I -)» the single bond is broken and each atom can gain one
electron to form halide ions (F1-, Cl1-, Br1-, I1-)» the most reactive ends up as an ion (1- charge) and is not
visible (molecules F2, Cl2, Br2, I2 are a visible gas)
» Cl > Br > I Cl-(aq) Br-(aq) I-(aq)
Cl2 Colorless- no reaction
turns red due to formation of Br2
turns brown due to formation of I2
Br2 no reaction no reaction turns brown due to formation of I2
I2 no reaction no reaction no reaction
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–periods• from left to right in period 3
–metals…metaloids…nonmetals–when oxides react with water
»basic…amphoteric (either basic or acidic)…acidic
» Na2O(s) + H2O (l) 2 NaOH (aq) strong base» MgO (s) +H2O (l) Mg(OH)2 (aq) weaker base» P4O10 (s) + 6H2O (l) 4 H3PO4 (aq) weak/strong acid» SO3(g) + H2O (l) H2SO4 (aq) strong acid