Download presentation
Presentation is loading. Please wait.
1
Trends & the Periodic Table
2
Trends more than 20 properties change in predictable way based location of elements on PT some properties: - anyone know where we can find these numbers?! Atomic radius Ionization energy Electronegativity Reactivity
3
Going down column 1: 7 Fr 6 Cs 5 Rb 4 K 3 Na 2 Li 1 H # of energy levels Element Period increasing # energy levels as go down
4
Increasing number of energy levels
5
Atomic Radius Atomic radius: defined as ½ distance between neighboring nuclei in molecule or crystal Affected by 1. # of energy levels 2. Proton Pulling Power
7
Increasing Atomic Radius
Increasing number of energy levels Increasing Atomic Radius Increasing Atomic Radius
8
Cs has more energy levels, so it’s bigger
previous | index | next Li: Group 1 Period Cs: Group 1 Period 6
9
As we go from LEFT to RIGHT across a row, elements gain electrons, but they are getting smaller!
WHY?
10
Cesium is bigger because it has more energy levels than Lithium.
11
Decreasing Atomic Radius
Increasing number of energy levels Increasing Atomic Radius
12
Why does this happen.. As you go from left to right, you again more protons (the atomic number increases) You have greater “proton pulling power” Remember the nucleus is + and the electrons are - so they get pulled towards the nucleus The more protons your have, the more Proton Pulling Power
13
We can “measure” the Proton Pulling Power by determining the NUMBER of Protons
Number of protons = ATOMIC NUMBER So in any row, atoms have the same number of energy levels, atoms that have less protons will be bigger, they can not pull the e;ectrons to the nucleus with as much force.
14
Going from L to R size tends to decrease a bit because of greater PPP “proton pulling power”
15
Think of protons as MAGNETS.
More magnets more greater “proton pulling power” (PPP).
16
Ionization Energy Ionization Energy = amount energy required to remove a valence electron from an atom in gas phase Basically how easily an atom will LOSE an electron. High ionization energy means it takes a LOT of energy to remove an electron from that atom.
17
Cs valence electron lot farther away from nucleus than Li
previous | index | next Cs valence electron lot farther away from nucleus than Li electrostatic attraction much weaker so easier to steal electron away from Cs THEREFORE, Li has a higher Ionization energy then Cs
18
Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius Decreased Ionization Energy (easier to remove an electron) Increased Electron Shielding Increasing number of energy levels Increasing Atomic Radius
19
Electronegativity Ability of atom to attract electrons in bond
noble gases tend not to form bonds, so don’t have electronegativity values Atoms that have high electronegativity have a HIGH electron affinity This means they LIKE or electrons Which means they easily become NEGATIVE ions.
20
Electronegativity Electronegativity- tendency of an atom to attract an electron.
21
Electronegativity Electronegativity- tendency of an atom to attract an electron. Group Trend – As you go down a column, electron affinity decreases. As you go down, atomic size is increasing, so less attraction of electrons to the nucleus. Periodic Trend – As you go across a period (L to R), electron affinity increases. As you go L to R, atomic size is decreasing, so the electrons are more attracted to the nucleus.
22
Atoms that have high electronegativity have a HIGH electron affinity
23
Increased Electronegativity
Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Increasing number of energy levels Decreased Electronegativity Increasing Atomic Radius
24
Reactivity of Metals judge reactivity of metals by how easily give up electrons (they’re losers)
25
Most reactive metal = Fr (the most metallic)
Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius More metallic Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Increasing number of energy levels Decreased Electronegativity Increasing Atomic Radius Most reactive metal = Fr (the most metallic)
26
Reactivity of Non-metals
judge reactivity of non-metals by how easily gain electrons (they are winners)
27
Reactivity In General Reactivity – tendency of an atom to react.
Metals – lose e- when they react, so metals’ reactivity is based on lowest Ionization Energy (bottom/left corner) Low I.E = High Reactivity. Nonmetals – gain e- when they react, so nonmetals’ reactivity is based on high electronegativity (upper/right corner). High electronegativity = High reactivity
28
Most Reactive Nonmetal = F
Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Most Reactive Nonmetal = F More metallic Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Increasing number of energy levels Decreased Electronegativity Increasing Atomic Radius Most reactive metal = Fr (the most metallic) Nonreactive BACK
29
How do you know if an atom gains or loses electrons?
Think back to the Lewis structures of ions Atoms form ions to get a valence of 8 (or 2 for H) Metals tend to have 1, 2, or 3 valence electrons It’s easier to lose them Nonmetals tend to have 5, 6, or 7 valence electrons It’s easier to add some Noble gases already have 8 so they don’t form ions very easily
30
Positive ions (cations)
Formed by loss of electrons Cations always smaller than parent atom 2e 8e 8e 8e 8e 2e 2e Ca Ca Ca+2
31
Negative ions or (anions)
Formed by gain of electrons Anions always larger than parent atom
Similar presentations
