Periodic Trends. Atomic Size The electron cloud doesn’t have a definite edge. Scientists get...
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![Page 1: Periodic Trends. Atomic Size The electron cloud doesn’t have a definite edge. Scientists get around this by measuring more than 1 atom at a time. Summary:](https://reader035.fdocuments.in/reader035/viewer/2022070414/5697c01f1a28abf838cd1896/html5/thumbnails/1.jpg)
Periodic Trends
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Atomic Size
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Atomic Size• The electron cloud doesn’t have a definite
edge.
• Scientists get around this by measuring more than 1 atom at a time.
• Summary: it is the volume that an atom takes up
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Atomic Size
•Atomic Radius = half the distance between two nuclei of a diatomic molecule.
}Radius
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Group trends – Atomic Radius
• As we go down a group the atoms have more e-, therefore more energy levels and the atoms get bigger
HLi
Na
K
Rb
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Periodic Trends – Atomic Radius The atomic radius decreases as you go
from left to right across a period.
Na Mg Al Si P S Cl Ar
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Explaining this trend
When moving across a period additional p+ are in the nucleus and more e- are in the same energy level.
• The opposite charges in the nucleus and the e- cloud cause the atom to be 'sucked' together a little tighter.
Therefore the radius decreases.
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Ionization Energy
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Ionization Energy (IE)• The amount of energy required to completely
remove an e- from a gaseous atom.
• Recall: removing one e- makes a +1 ion.
• The energy required is called the first ionization energy.
X(g) + energy → X+ + e-
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• Second and third ionization energies represents losing a 2nd and then a 3rd e- from the same atom.
• It can be shown as:
X+ (g) + energy X2+ (g) + e-
X2+ (g) + energy X3+ (g) + e-
• More energy required to remove 2nd e-, and still more energy required to remove 3rd e-.
• The closer the e- is to the nucleus, the more
difficult it will be to remove.
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Group Trends (I.E.)• Ionization energy decreases down the group.
• Ex. Going from Be to Mg, IE decreases because:– Mg outer e- is in the 3s sub-shell rather than the 2s.
– This is higher in energy and further from the nucleus.
– So the 3s e- is more easily removed, requiring less energy.
• A similar decrease occurs in every group in the periodic table.
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Period Trends (IE)IE generally increases from left to right.
Why? The e- are attracted more strongly to the nucleus
(smaller radius).
It takes more energy to remove one e- from the atom with stronger attraction, therefore IE increases.
Ex. From Na to Ar (11 p+ to 18 p+), the attraction of the protons to e- within the same energy level increases.
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Why is there a decrease in IE from Mg to Al?
• Al is 1s2 2s2 2p6 3s2 3p1
It has one e- that is in a p sublevel.
• Mg is 1s2 2s2 2p6 3s2.
Mg - the ‘s’ sublevel is full – this gives it a slight stability advantage and will require more energy to let go of its e-.
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Why is there a fall in IE from phosphorus to sulfur?
• This can be explained in terms of e- pairing.• Phosphorus - 1s2 2s2 2p6 3s2 3p3
• Sulfur - 1s2 2s2 2p6 3s2 3p4
• As the p sublevel fills up, e- fill up the vacant sub levels and are unpaired.
• Phosphorus’ configuration is more energetically stable than sulfur’s because there are e- that are unpaired.
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• When e- are paired, there is some repulsion which lessens their attraction to the nucleus.
• It becomes easier to remove!
• Having a half filled sublevel is more stable than a partially filled sublevel.
• So… sulfur will break the expected trend and want to lose an e- requiring less IE.
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Why an exchange in e- ?• Noble Gases have full energy levels.
• Atoms behave in ways to achieve noble gas configuration.
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2nd Ionization Energy• The amount of energy required to remove
the 2nd e- from a gaseous atom.
• For elements that reach a filled or half filled sublevel by removing 2 e- the 2nd IE is lower than expected.
• Makes it easier to achieve a full outer shell
• True for s2 , the alkaline earth metals which form +2 ions.
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3rd IE• Using the same logic s2p1
atoms have a low 3rd IE.
• Atoms in the aluminum family form +3 ions.
• 2nd IE and 3rd IE are always higher than 1st IE!!!
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Reactivity
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Reactivity• Reactivity refers to how likely or vigorously
an atom is to react with other substances.
• This is usually determined by how easily e- can be removed and how strongly atoms want to take other atom's e- .
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Reactivity - for Metals:Period - reactivity decreases from left to
right
Group - reactivity increases going down a group
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Why?
- Elements located toward the left of the periodic table (alkali metals) and near the bottom easily lose their e-, resulting in higher reactivity.
- Within the same group, the more e- an atom has, the easier it will give it up.
(Ex. Li (3) and Fr (87))
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Reactivity -for Non-Metals
• Period - reactivity increases from left to right (not including the noble gases)
• Group - reactivity decreases going down the group. (not including the
noble gases)
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• Why?
– Atoms are most stable when they have noble gas electron configuration.
– Groups closest to the noble gases want to gain an e- to become stable therefore they have a more vigorous exchange of e-.
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– Elements within the same group vary significantly in number of e- but contain the same number of valence e-.
– The lower energy levels are found closer to the nucleus, having a stronger desire to complete their energy level and will react more violently.
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Shielding
• Electrons on the outside energy level (valence e-) have the inner energy levels blocking the positive force field (nucleus).
• These inside energy level e- shield (block) the nuclear (pos) force field from the valence e-
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• As you go across the row the nuclear charge (positive charge) gets larger because protons are being added to the nucleus.
• As you go across the row valence e- are added to the valence shell but the valence e- have the same shielding.
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• The blocking strength (shielding effect) of these inner e- is the same across the period.
• Further right in a period the valence e- will have a greater attraction to the nucleus because of the greater positive charge.
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• Shielding becomes less effective across the row; 2e- can shield +3 better
than 2e- can shield +10.
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As you move down a group the valence e- are being added to a new energy level further from the nucleus.
These new valence e have additional levels of inner shielding e- and are more effectively shielded from the positive charge.
Ex. Campfire
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Electronegativity
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Electronegativity• The tendency for an atom to attract e- to
itself when it is chemically combined with another element.
• How fair it shares.
• Large electronegativity means it has a strong pull on an e- toward itself.
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Group Trend
• The further down a group the farther the e- is away from the nucleus and the more e- an atom has.
• Going to the bottom of a group, the e-are further away from the nucleus.
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• This means they are better shielded from the nuclear (+) charge and thus not as attracted to the nucleus.
• For that reason the electronegativity decreases as you go down the periodic table.
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Period Trend• Electronegativity increases from left to
right across a period
• When the nuclear charge increases, so will the attraction that the atom has for e- in its outermost energy level.
This means the electronegativity will increase
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Electron Affinity• The energy change associated with
adding an e- to a gaseous atom.
• Easiest to add to group 17 or 7A.
• Gets them to full energy level.
• Energy is often required (+) when adding an e- to metals. Energy is given off (-) when adding an e- to non-metals.
• EA decreases as we go down a group.
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Electron affinity decreases as we go down a group because the atoms are getting bigger and the valence electrons are not attracted as strongly to the nucleus.
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Ionic Size
• Cations form by losing e- (have a positive charge).
• Cations are smaller than the atom they come from.
• Metals form cations.
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Ionic size• Anions form by gaining electrons.
• Anions are bigger than the atom they come from.
• Nonmetals form anions.
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Configuration of Ions• Ions always have noble gas configuration.
• 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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Periodic Trends• Across the period nuclear charge increases so
the attractive force gets stronger and the atoms get smaller when filling to the same energy level.
• Energy level changes between anions and cations.
Li+1
Be+2
B+3
C+4
N-3O-2 F-1
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• http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/pem3s3_2.swf
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Size of Isoelectronic ions• Iso - same
• Isoelectronic 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 have a
smaller atomic radius. • The greater the # of protons the stronger the
attraction to the same # of electrons.• This will cause the atomic radius to be smaller.
Al+3
Mg+2
Na+ Ne F- O-2 N-3
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• Organize the isoelectronic ions/atoms in order from smallest to largest
• P-3, Ar, Cl-, K+, Ca2+, S2-, Sc3+,
• http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/pem3s3_4.swf
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Atomic size decreases
Ionic size decreases
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Ionization energy, electronegativity, electron affinity INCREASE
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