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Ch 5

Noble Gases All of them have completely filled electron shells. n2 np6 Since they have similar electronic structures, their chemical reactions are similar.
Representative Elements Are the elements in A groups on periodic chart. These elements will have their “last” electron in an outer s or p orbital. These elements have fairly regular variations in their properties.
d-Transition Elements Elements on periodic chart in B groups. Each metal has d electrons. ns (n-1)d configurations These elements make the transition from metals to nonmetals. Exhibit smaller variations from row-to-row than the representative elements
f - transition metals "inner transition metals" Electrons being added to f orbitals, two shells below the valence shell! Slight variations of properties Outermost electrons have the greatest influence on the chemical properties of elements. (n quantum #)
Outer most electrons have the greatest influence on the properties of elements. Adding an e to an s or p orbital usually causes dramatic changes in the physical and chemical properties. Adding an e to a d or f orbital typically has a much smaller effect on properties.
Outer most electrons have the highest value of the principal quantum number, n.
Atomic Radii describe the relative sizes of atoms increase within a column going from the top to the bottom of the periodic table decrease within a row going from left to right on the periodic table.
Description of atomic radii shielding/screening effect nuclear charge experienced by e- < actual nuclear charge Inner e- block nuclear charge’s effect on outer e- Moving across a period each element increases nuclear charge and e- go into same shell Outer e- feel > nuclear charg
Atomic radii increase going down a group because e are being added to shells farther from the nucleus. Atomic radii decrease from left to right within a given period owing to increasing effective nuclear charge
First ionization energy (IE1) Atom(g)+Energy=Ion+(g)+e- Measures how tightly e are bound to atoms. Low ionization energies indicate easy removal of e and hence easy positive ion formation
it is always more difficult to remove a neg charged e from a pos charged ion than from the corresponding neutral atom
Ionization Energy IE1 decrease because the outermost e are farther from the nucleus. 1st IE1 generally increase from left to rt across the periodic table because the e are held more tightly by nucleus.
Second ionization energy (IE2) The amount of energy required to remove the second electron from a gaseous 1+ ion. Symbolically: ion+ + energy --> ion2+ + e-
Periodic trends for Ionization Energy: IE2 > IE1 always takes more energy to remove 2nd e- from an ion than from neutral atom IE1 generally increases moving from Left to Right EXCEPTIONS due to filled and half-filled subshells IE1 generally decreases moving down IE1 for Li > IE1 for Na, et
Electron Affinity amount of energy absorbed when atom(g) + e- + EA --> ion-(g) Electron affinity is a measure of an atom’s ability to form negative ions. Elements with very neg. EA gain electrons easily to form negative ions (anions)
Sign conventions for electron affinity. If electron affinity > 0 energy is absorbed. If electron affinity < 0 energy is released.
General periodic trend for electron affinity values become more - from left to right the values become more negative from bottom to top on the periodic chart. EA is a precise and quantitative term, like ionization energy, but it is difficult to measure. EA of anions always +Ionic Radii
Ionic Radii Cations (positive ions) are always smaller than their respective neutral atoms. Negative ions (anions) are always larger than their respective neutral atoms. Both cation and anion sizes increase going down a group.
Within an isoelectronic series ionic radii decrease with increasing atomic number because of increasing nuclear charge (more p+ to e-).
isoelectronic species have the same number of e. Comparing such species, the higher the nuclear charge (more protons), the smaller the radius.
anion ionic radii Anion (negative ions) radii decrease from left to right across a period. Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius.
Electronegativity a measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another element
For the representative elements, electronegativities usually... increase from left to right across periods and decrease from top to bottom within groups
High EN vs low EN Elements with high EN (nonmetals) often gain E to form anions. Elements with low EN (metals) often lose E to form cations. Cs and Fr Large difference in EN --> ionic compounds
Oxidation States of an element in a simple binary (meaning 2) ionic compound is the number of e gained or lost by an atom of that element when it forms the compound. the element with the pos oxidation state is written first, except NH3, and hydrocarbons CH4 H is +1
Electronegativity and oxidation states The most EN atom in a compound is assigned a neg oxidation state, while the less EN atom lose e and have pos oxidation states.
Guidelines for assigning oxidation numbers Of any uncombined element = 0 Of an element in a simple ion = charge on the ion Any compound sum of oxidation #s of all elements = 0 In a polyatomic ion, the sum of the oxidation numbers of the constituent elements = charge on the ion
More Guidelines for assigning oxidation numbers F oxidation number = –1 in its compounds. H oxidation number = +1 unless combined with metals = -1 Oxygen usually has the oxidation number -2. Exceptions: In peroxides O has oxidation number of –1. In OF2 O has oxidation number of +2.
Periodic table helps assign oxidation numbers IA metals have oxidation numbers of +1...+2...+3...-3...-2
Created by: tilleryc



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