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Ch 5
| Question | Answer |
|---|---|
| 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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