Presentation on theme: "Periodic Trends OBJECTIVES:"— Presentation transcript:
1 Periodic Trends OBJECTIVES: Interpret periodic trends in atomic radii, ionic radii, ionization energies, and electronegativities.
2 Trends in Atomic SizeFirst problem: Where do you start measuring from?The electron cloud doesn’t have a definite edge.We get around this by measuring more than 1 atom at a time.
3 Atomic Size}RadiusAtomic Radius = half the distance between two nuclei of a diatomic molecule. Diatomic means “two atoms” or a two atom molecule
4 Trends in Atomic Size Influenced by three factors: 1. Energy Level Higher energy level is further away.2. Charge on NucleusMore charge pulls electrons in closer.3. Shielding EffectElectrons between nucleus and valance electrons
5 Group Trends As we go down a group... HAs we go down a group...each atom adds another energy level,so the atoms get bigger.LiNaKRb
6 Periodic Trends As you go across a period, the radius gets smaller. Electrons are in same energy level.More nuclear charge.Outermost electrons are closer.NaMgAlSiPSClAr
11 Ion FormationAn atom can easily lose or gain electrons. The resulting ion is an atom that has an imbalance of charge or carries either a positive charge (a loss of electrons) or a negative charge (a gain of electrons).A positive ion is called a cation.A negative ion is called an anion.
12 Trends in 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 to remove the first electron is called the first ionization energy.
13 Ionization EnergyThe 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.
20 What Determines IE The greater the nuclear charge, the greater IE. Greater distance from nucleus decreases IEFilled and half-filled sublevel have lower energy, so achieving them is easier, lower IE.Shielding effect on valence electrons
21 ShieldingThe electron on the outermost energy level (valence electron) has to look through all the other inner electrons to see the nucleus.Shielding electrons hinder the nucleus form holding on to the valence electrons
22 Group Trends As you go down a group, first IE decreases because... the valence electron is further awayand there is more shielding.
23 Periodic Trends As you go across a period, All the atoms in the same period have the same highest energy level.They have similar shielding.But, increasing nuclear charge reduces atomic radius (bringing valence closer)so IE generally increases from left to right.Exceptions are at full and 1/2 full sublevels.
24 He has a greater IE than H. same shielding greater nuclear charge First Ionization energyAtomic number
25 outweighs greater nuclear charge HeLi has lower IE than Hmore shieldingfurther awayoutweighs greater nuclear chargeHFirst Ionization energyLiAtomic number
26 greater nuclear charge HeBe has higher IE than Lisame shieldinggreater nuclear chargeFirst Ionization energyHBeLiAtomic number
27 greater nuclear charge HeB has lower IE than Besame shieldinggreater nuclear chargeBy removing an electron we leave the s orbital fullFirst Ionization energyHBeBLiAtomic number
28 First Ionization energy HeFirst Ionization energyHCBeBLiAtomic number
29 First Ionization energy HeNFirst Ionization energyHCBeBLiAtomic number
30 First Ionization energy HeNBreaks the pattern, because removing an electron leaves 1/2 filled p orbitalFirst Ionization energyHCOBeBLiAtomic number
31 First Ionization energy HeFNFirst Ionization energyHCOBeBLiAtomic number
32 First Ionization energy HeNeFNNe has a lower IE than HeBoth are full,Ne has more shieldingGreater distanceFirst Ionization energyHCOBeBLiAtomic number
33 Na has a lower IE than Li Both are s1 Na has more shielding HeNeNa has a lower IE than LiBoth are s1Na has more shieldingGreater distanceFNFirst Ionization energyHCOBeBLiNaAtomic number
37 2nd Ionization EnergyFor elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected.True for s2Alkaline earth metals form 2+ ions.
38 3rd Ionization EnergyUsing the same logic s2p1 atoms have an low 3rd IE.Atoms in the aluminum family form 3+ ions.2nd IE and 3rd IE are always higher than 1st IE!!!
39 Trends in Electron Affinity The atoms ability to acquire an additional electronCl + 1e Cl1-The energy change associated with adding an electron to a gaseous atom.Easiest to add to group 7A.Gets them to a full energy level.Increase from left to right: atoms become smaller, with greater nuclear charge.Decrease as we go down a group.
40 Trends in Ionic Size 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.NaNa+1
41 Ionic Size 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.ClCl-1
42 Configuration of Ions Ions always have noble gas configuration. Na is: 1s22s22p63s1Forms a 1+ ion: 1s22s22p6Same configuration as neon.Metals form ions with the configuration of the noble gas before them - they lose electrons.
43 Configuration of IonsNon-metals form ions by gaining electrons to achieve noble gas configuration.They end up with the configuration of the noble gas after them.
48 Size of Isoelectronic Ions Iso- means the sameIso electronic ions have the same # of electronsAl3+ Mg2+ Na1+ Ne F1- O2- and N3-all have 10 electronsall have the configuration: 1s22s22p6
49 Size of Isoelectronic Ions Positive ions that have more protons would be smaller.N3-O2-F1-NeNa1+Al3+Mg2+
50 Electronegativity is a measure of the ability of an atom in a molecule to attractelectrons to itself.Concept proposed byLinus Pauling
51 ElectronegativityThe ability of an atom that is bonded to another atom or atoms to attract electrons to itself.It is related to ionization energy and electron affinity.It cannot be directly measured.The values are unitless since they are relative to each other.The values vary slightly from compound to compound but still provide useful qualitative predictions.
52 ElectronegativityThe tendency for an atom to attract electrons to itself when it is chemically combined with another element.How fair is the sharing?Large electronegativity means the atom pulls the electrons toward it.Atoms with large electron affinity have larger electronegativity.
53 Group TrendThe further down a group, the farther the valence electrons are away from the nucleus and the more electrons an atom has shielding.The larger the atom, the more willing it is to share electrons.Low electronegativity.
54 Electronegativities Electronegativity is a periodic property. Atomic number
55 Periodic Trend Metals are at the left of the table. They let their electrons go easilyLow electronegativityAt the right end are the nonmetals.They want more electrons.Try to take them away from othersHigh electronegativity.
56 ElectronegativityRelative ability of atoms to attract electrons of bond.At1.9I2.2Br2.7Cl2.8Po1.8Te2.0Se2.5S2.4Bi1.7SbAsP2.1Pb1.5SnGeSiF4.1O3.5N3.1Tl1.4Na1.0Cs0.9RbKBaMg1.2SrCaInGaAlHLiBeBC