Presentation on theme: "Ch 5.3 Electron Configuration and Periodic Properties"— Presentation transcript:
1 Ch 5.3 Electron Configuration and Periodic Properties
2 Science Content Standards for California Public Schools Atomic and Molecular Structure1. The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept:b. Students know how to use the periodic table to identify metals, semimetals, non-metals, and halogens.c. Students know how to use the periodic table to identify alkali metals, alkaline earth metals and transition metals, trends in ionization energy, electronegativity, and the relative sizes of ions and atoms.f.* Students know how to use the periodic table to identify the lanthanide, actinide, and transactinide elements and know that the transuranium elements were synthesized and identified in laboratory experiments through the use of nuclear accelerators.g.* Students know how to relate the position of an element in the periodic table to its quantum electron configuration and to its reactivity with other elements in the table.
3 Atomic Size}RadiusMeasure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.
4 #1. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller.Electrons are in the same energy level.Outermost electrons are pulled closer.NaMgAlSiPSClAr
5 #1. Atomic Size - Group trends HAs we increase the atomic number (or go down a group), each atom has another energy level.Valence electrons get further from the nucleus.So the atoms get bigger.LiNaKRb
7 IonsAn ion is an atom (or group of atoms) that has a positive or negative chargeAtoms are neutral because the number of protons equals electronsPositive and negative ions are formed when electrons are lost or gained between atoms
8 #2: Ionic Group trends Li1+ Na1+ K1+ Rb1+ Cs1+ Ions therefore get bigger as you go down, because of the additional energy level.Na1+K1+Rb1+Cs1+
9 Ionic Period TrendsAcross the period from left to right, they get smaller.N3-O2-F1-B3+Li1+Be2+C4+
10 Trends in the Periodic Table Atomic Size vs. Ion Size
11 #3. Trends in Ionization Energy Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).The energy required to remove 1 electron is called the first ionization energy.
12 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.
13 3. Trend in Ionization Potential The energy required to remove the valence electron.
14 Ionization Energy - Group trends As you go down a group, the first IE decreases because...The electron is further away from the attraction of the nucleus, andThere is more shielding.
15 Ionization Energy - Period trends All the atoms in the same period have the same energy level.So IE generally increases from left to right.
16 He has a greater IE than H. Both elements have the same shielding since electrons are only in the first levelBut He has a greater nuclear chargeHFirst Ionization energyAtomic number
17 These outweigh the greater nuclear charge Li has lower IE than Hmore shieldingfurther awayThese outweigh the greater nuclear chargeHFirst Ionization energyLiAtomic number
18 greater nuclear charge HeBe has higher IE than Lisame shieldinggreater nuclear chargeFirst Ionization energyHBeLiAtomic number
19 greater nuclear charge HeB has lower IE than Besame shieldinggreater nuclear chargeBy removing an electron we make s orbital half-filledFirst Ionization energyHBeBLiAtomic number
20 First Ionization energy HeFirst Ionization energyHCBeBLiAtomic number
21 First Ionization energy HeNFirst Ionization energyHCBeBLiAtomic number
22 HeOxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbitalNFirst Ionization energyHCOBeBLiAtomic number
23 First Ionization energy HeFNFirst Ionization energyHCOBeBLiAtomic number
24 Ne has a lower IE than He Both are full, Ne has more shielding Greater distanceFNFirst Ionization energyHCOBeBLiAtomic number
25 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
27 Electron AffinityIs the energy change that occurs when a neutral atom acquires an electron.Most atoms release energy when they gain an electron. (negative value)However, some must be forced to gain an electron which requires energy. (positive value)Halogens gain electrons most readily.
28 Metals tend to LOSE electrons, from their outer energy level IonsMetals tend to LOSE electrons, from their outer energy levelSodium loses one electron.There are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation”The charge is written as a number followed by a plus sign: Na1+Now named a “sodium ion”Lost an electron, so a decrease in size from atom to ion.
29 Nonmetals tend to GAIN one or more electrons IonsNonmetals tend to GAIN one or more electronsChlorine will gain one electronProtons (17) no longer equals the electrons (18), so a charge of -1Cl1- is re-named a “chloride ion”Negative ions are called “anions”Gained an electron so increase in size from atom to ion.
30 Valence ElectronsThe electrons available to be lost, gained, or shared in the formation of chemical compounds.Main group elements: valence electrons are in the outermost s and p orbitals.
31 Electronegativity is the tendency for an atom to attract electrons. It decreases as it goes down a group because electrons get farther away from the nucleus.
32 Electronegativity Period Trend Metals (left side)They let their electrons go easilyLow electronegativityNonmetals (right side).They want more electrons.Try to take them away from othersHigh electronegativity.
33 Summary of Trends Ionization Energy and Electronegativity Atomic and Ionic Radius
34 Additional Assessment Questions Topic5Question 1For each of the following pairs, predict which atom is larger.a. Mg, Srd. Ge, Brb. Sr, Sne. Cr, Wc. Ge, Sn
35 Answers a. Mg, Sr Sr b. Sr, Sn Sr c. Ge, Sn Sn d. Ge, Br Ge e. Cr, W W Additional Assessment QuestionsTopic5Answersa. Mg, SrSrb. Sr, SnSrc. Ge, SnSnd. Ge, BrGee. Cr, WW
36 Additional Assessment Questions Topic5Question 2For each of the following pairs, predict which atom or ion is larger.a. Mg, Mg2+d. Cl–, I–b. S, S2–e. Na+, Al3+c. Ca2+, Ba2+
37 Answers a. Mg, Mg2+ Mg b. S, S2– S2– c. Ca2+, Ba2+ Ba2+ d. Cl–, I– I– Additional Assessment QuestionsTopic5Answersa. Mg, Mg2+Mgb. S, S2–S2–c. Ca2+, Ba2+Ba2+d. Cl–, I–I–e. Na+, Al3+Na+
38 Additional Assessment Questions Topic5Question 3For each of the following pairs, predict which atom has the higher first ionization energy.a. Mg, Nad. Cl, Ib. S, Oe. Na, Alc. Ca, Baf. Se, Br
39 Answers a. Mg, Na Mg b. S, O O c. Ca, Ba Ca d. Cl, I Cl e. Na, Al Al Additional Assessment QuestionsTopic5Answersa. Mg, NaMgb. S, OOc. Ca, BaCad. Cl, ICle. Na, AlAlf. Se, BrBr