General Chemistry Ch. 8 - Electron Configuration and Chemical Periodicity
Quiz yourself by thinking what should be in
each of the black spaces below before clicking
on it to display the answer.
Help!
|
|
||||
|---|---|---|---|---|---|
| Electron configuration | The distribution of electrons within the orbitals of its atoms
🗑
|
||||
| Recall: the three quantum numbers | n, l, and m_l
🗑
|
||||
| A fourth quantum number which is required if there is more than one electron | A property of the electron itself (as opposed to the orbital) called spin
🗑
|
||||
| What are the permitted values of spin? | m_s = +1/2 or -1/2
🗑
|
||||
| Review: the four quantum numbers and their permitted values | Principal: n (1, 2, 3, …); Angular momentum: l (from 0 to n-1); magnetic: m_l (from –l to 0 to +l); spin: m_s (+1/2 or -1/2)
🗑
|
||||
| Exclusion principle | No two electrons in the same atom can have the same four quantum numbers. E.g. two electrons might have the orbitals but have opposite spin
🗑
|
||||
| Because the spin number can have only two values, the major consequence of the exclusion principle is | An atomic orbital can hold a maximum of two electrons and they must have opposing spins
🗑
|
||||
| In the H atom, all sublevels of a given level, such as the 2s and 2p, have the same energy. Does this apply to other atoms as well? | No because the energy states of many-electron atoms arise not only from the nucleus-electron attractions, but also electron-electron repulsions.
🗑
|
||||
| The electrons of an atom in its ground state occupy the orbitals… | …of lowest energy
🗑
|
||||
| How do the 3 electrons of Li in its ground state fill its orbitals? | The first two electrons fill its 1s orbital. Then the third must go into the n = 2 level. The 2s is lower in energy than 2p so the 2s orbital is filled by the 3rd electron.
🗑
|
||||
| Why does the 2s orbital have lower energy? | Three reasons for the energy difference between 2s and 2p: (1) nuclear charge, (2) electron repulsions, and (3) orbital shape
🗑
|
||||
| The effect of nuclear charge (Z) on orbital energy | Higher nuclear charge gives the orbital energy a higher negative number (lowers it) which results in a more stable state where the electron Is harder to move.
🗑
|
||||
| Shielding | Caused by the repulsion of other electrons. Shielding reduces the full nuclear charge to an *effective nuclear charge* (Z_eff): the nuclear charge an electron actually experiences. The lower nuclear charge makes it easier to remove.
🗑
|
||||
| Example of shielding | It takes less than half as much energy to remove an electron from He than from He because the second electron in He repels the first.
🗑
|
||||
| Shielding effectiveness: inner vs. outer electrons | Outer electrons feel the effects of shielding much more than inner electrons. To rephrase: the inner electrons shield more effectively than do the outer electrons.
🗑
|
||||
| Why do electrons first fill the 2s orbital as opposed to the 2p orbital? | Even though the 2p orbital, on average, spends more time a little closer to the nucleus than does the 2s orbital, the 2s orbital *penetrates* the 1s region for a small amount of time. This penetration greatly mitigates shielding.
🗑
|
||||
| Penetration of the 2s orbital results in | A total higher Z_eff, in other words a lower orbital energy than the 2p orbital. And because electrons in the ground state fill lower energy states first, the 2s is filled first.
🗑
|
||||
| The order of l value from low to high orbital energy | s < p < d < f
🗑
|
||||
| Remember: orbital energy vs Z_eff | Lower orbital energy = higher Z_eff (effective nuclear charge felt by the electron)
🗑
|
||||
| Aufbau principle | The useful approach of determining electron configurations of elements: start at the beginning of the periodic table and add one electron per element to the lowest energy orbital available
🗑
|
||||
| Electron configuration notation | nl^#. Where n = principle energy level, l = sublevel, and # = number of electrons in the sublevel. Eg H = 1s^1, He = 1s^2. Electron spin is not indicated in the electron configuration notation but assumes you know them.
🗑
|
||||
| Orbital diagram | The boxes with arrows, review in book pg 309. The arrow in the box indicates electrons and their spins.
🗑
|
||||
| Be electron configuration | 1s 2 2s 2
🗑
|
||||
| B electron configuration | 1s 2 2s 2 2p 1
🗑
|
||||
| C electron configuration | 1s 2 2s 2 2p 2
🗑
|
||||
| N electron configuration | 1s 2 2s 2 2p 2 2p 3
🗑
|
||||
| Describe p sublevel | A p sublevel has l = 1 so the m_l can only be –l, 0, +l. the three orbitals in the 2p sublevel have EQUAL energy which means that the electron goes into any one of the 2p orbitals.
🗑
|
||||
| Hund’s rule | When orbitals of equal energy are available, the electron configuration of lowest energy has the maximum number of unpaired electrons with parallel spins.
🗑
|
||||
| Hund’s rule rephrased | The boxes in the p orbitals will fill up with single arrows first (e.g. 2p 1 thru 2p 3) then become double arrows with each additional electron.
🗑
|
||||
| O electron configuration | 1s 2 2s 2 2p 4 (note: the first box of the three 2p boxes now has two arrows)
🗑
|
||||
| Ne electron configuration | 1s 2 2s 2 2p 6
🗑
|
||||
| The electron configurations defined are also known as | Ground state electron configuration: the energy associated with the lowest energy level of the atom.
🗑
|
||||
| Ground state electron configurations of ions | They mimic the electron configurations of the atoms with the same number of electrons, e.g. Na+ has the same ground state configuration as Ne.
🗑
|
||||
| To review electron configuration pattern of periodic table… | Pages 311, 312, 313
🗑
|
||||
| In period four, what electron configuration notation is given to the inner transition metals? | [Ar]4s2 3d1 thru 3d10
🗑
|
||||
| Similar outer electron configurations correlate with… | Similar chemical behavior
🗑
|
||||
| Group 1A(1), the alkali metals, share ns1 orbital designation. How are they similar chemically? | All are highly reactive metals that form ionic compounds with nonmetals with formulas such as MCl, M2O, and M2S, and all react vigorously with water to displace H2.
🗑
|
||||
| Group7A(17), the halogens, share ns2 np5 orbital designation. How are they similar chemically? | They are reactive nonmetals that occur as diatomic molecules. All form ionic compounds with metals.
🗑
|
||||
| Group 8A(18), the noble gases, share ns2 np6 orbital designation. How are they similar chemically? | They have “filled” energy levels and are very unreactive monatomic gases.
🗑
|
||||
| What’s unique about the orbital order in period 4? | 4s is filled before 3d. This switch in order is due to shielding and penetration effects. The 4s penetrates close to the nucleus, spending more time closer to the nucleus than 3d, and thus has a lower energy level and filled first.
🗑
|
||||
| Do periods 5, 6, etc fill up in a similar pattern to period 4? | Yes; 5s fills before 4d, 6s fills before 5d, etc. In general, the ns sublevel fills before the (n-1)d sublevel. There are exceptions to this rule.
🗑
|
||||
| Electron configurations of Cr and Cu: how are they unique? | Rather than having 4s being completely filled, they’re only half filled, which adds an extra electron to the 3d orbital. Review page 313 if needed. Apparently orbitals are more stable if they’re full/half full.
🗑
|
||||
| Three categories of electrons | Inner (core) electrons: lower energy level electrons; Outer electrons: highest energy level electrons; and Valence electrons: those involved in forming compounds, i.e. among the main-group elements, they are the outer electrons
🗑
|
||||
| Valence electrons in d-block elements | d-block = transition elements (not including inner transition, those are the f-block). In the d-block all of the (n-1)d electrons are counted as valence electrons also, even though the elements Fe through Zn utilize few d electrons in bonding
🗑
|
||||
| Among the main group elements, the group number equals… | The number of outer electrons
🗑
|
||||
| The period number is the n value of… | The highest energy level. E.g. in period 2, the n=2 level has the highest energy
🗑
|
||||
| The n-value squared (n^2) gives… | The total number of orbitals in that energy level.
🗑
|
||||
| Orbitals of the inner transition elements | f-orbitals
🗑
|
||||
| Period 6 inner transition metals are called | Lanthanides (rare earths)
🗑
|
||||
| Inner transition metals period 7 are called | Actinides
🗑
|
||||
| How many different s orbitals for every shell? p orbital? d orbitals? f orbitals? | s: 1, p: 3, d: 5, f: 7. Note: every orbital can hold 2 electrons
🗑
|
||||
| If n^2 gives you the number of orbitals, how to determine the number of max electrons? | 2*n^2
🗑
|
||||
| Why is helium placed in group 8 of the periodic table rather than group 2? | Because it behaves like a noble gas, not like an alkali earth metal.
🗑
|
||||
| p block | The block on the right side of the periodic table after the transition metals: 2p, 3p, 4p, 5p, 6p
🗑
|
||||
| d block | Transition metals (except the inner transition elements): 3d, 4d, 5d
🗑
|
||||
| f block | Lanthanides and actinides: 4f, 5f
🗑
|
||||
| How does shell 3 have 9 different orbitals? (In accordance with n^2) | 3s, 3p, 3d
🗑
|
||||
| How does period 4 have 16 different orbitals? | 4s, 4d, 4p, 4f
🗑
|
||||
| Sulfide ion electron configuration | (2 extra electrons). [Ne]3s2 3p6 (rather than 3p4)
🗑
|
||||
| The rule for removing electrons when determining electron configurations for cations | Remove them from the orbitals with the highest energies first. E.g. Zn is [Ar]4s2 3d10. Zn^2+ is [Ar]3d10. Why? 4s2 has a higher energy level so those electrons were removed first.
🗑
|
||||
| What is the following neutral element: [Ar]4s1 3d10 6p1 | Zinc in excited (not ground) state. Excited state because it absorbed an electrons with higher energy so instead of having 4s2 it has 4s1 where the displaced electron is in 6p1. If you count up the electrons you can tell it’s zinc.
🗑
|
||||
| p orbitals come in groups of | 3
🗑
|
||||
| d orbitals come in groups of | 5
🗑
|
||||
| All physical and chemical behavior of the elements is based ultimately on… | The electron configurations of their atoms
🗑
|
||||
| Two common definitions of atomic size | (1) Metallic radius and (2) covalent radius
🗑
|
||||
| Atomic radius | One-half the distance between nuclei of adjacent atoms in a crystal of the element. This definition is typically used for metals.
🗑
|
||||
| Covalent radius | One-half the distance between nuclei of identical covalently bonded atoms
🗑
|
||||
| Trends among the main-group elements: changes in n | as the principal quantum number (n) increases, the probability that the outer electrons will spend more time farther from the nucleus increases as well; thus the atoms are larger.
🗑
|
||||
| Trends among the main-group elements: changes in Z_eff | As the effective nuclear charge—the positive charge “felt” by an electron—increases, outer electrons are pulled closer to the nucleus, thus, the atoms are smaller
🗑
|
||||
| The net effect of the two changes (n and Z_eff) depends on ____. What is the trend on the periodic table? | Shielding of the increasing nuclear charge by inner electrons. Down a group, n dominates (atomic radius increases). Across a period, Z_eff dominates (atomic radius decreases).
🗑
|
||||
| The trends in n and Z_eff hold true for main group elements, but not for… | The transition elements
🗑
|
||||
| Ionization energy (IE) | The energy (in kJ) required for the complete removal of 1 mole of electrons from 1 mole of gaseous atoms or ions
🗑
|
||||
| Pulling an electron away from a nucleus _____ energy. | Requires energy (to overcome the attraction)
🗑
|
||||
| First ionization energy (IE_1) and second ionization energy (IE_2) and so on | The first ionization energy (IE_1) removes an outermost electron (highest energy sublevel): Atom(g) -> ion+(g) + e-; deltaE = IE_1 > 0. The second ionization energy removes a second electron. deltaE = IE_2 > IE_1
🗑
|
||||
| Why is the second ionization energy greater than the first? | Because the second electron is pulled away from a positively charged ion which is more attracted to it.
🗑
|
||||
| Key fact about the first ionization energy and how it relates to chemical reactivity | Atoms with a low IE_1 tend to form cations during reactions, whereas those with high IE_1 (except the noble gases) often form anions.
🗑
|
||||
| Periodic pattern with regard to first ionization energy | There is roughly an inverse relationship between IE_1 and atomic size: as size decreases, it takes more energy to remove an electron.
🗑
|
||||
| Summarize the trends down a group | As we move down a main group, the orbital’s n value increases and atomic size increases. As distance from nucleus to outer electrons increases, the attraction lessens (coulomb’s law) and as a result ionization energy decreases
🗑
|
||||
| Exceptions to trends down a group | The only significant exception occur in group 3A(1), right after the transition series, and is due to the effect of the series on atomic size. The newly filled d orbitals result in a greater-than-expected Z_eff.
🗑
|
||||
| Summarize the trends across a period | As we move left to right across a period, the orbital’s n value stays the same, so Z_eff increases and atomic size decreases. The attraction between the nucleus and outer electrons increases, thus ionization energy increases.
🗑
|
||||
| As the number of times you ionize an atom increases… | It requires more and more energy. And once you take out an outside orbital and move to a core orbital, e.g. moving from p to s, the ionization energy increases by a very large amount.
🗑
|
||||
| Electron affinity (EA) | The energy change (in kJ) accompanying the addition of the 1 mol of electrons to 1 mol of gaseous atoms or ions. As with ionization energy, there is a first electron affinity, a second, and so forth.
🗑
|
||||
| The first electron affinity (EA_1) | Refers to the formation of 1 mol of monovalent (1-) gaseous anions: Atom(g) + e- -> ion-(g) deltaE = EA_1
🗑
|
||||
| In most cases, energy is ____ when the first electron is added because… | Released; because it is attracted to the atom’s nuclear charge. Thus EA_1 is usually negative.
🗑
|
||||
| During the second electron affinity (EA_2), energy is _____ because | Usually always positive because energy must be absorbed in order to overcome the electrostatic repulsions and add another electron to a negative anion.
🗑
|
||||
| Trends in EA | Not as regular as IE.
🗑
|
||||
| General trends: reactive nonmetals | The elements in groups 6A(16) and those in Group 7A(17) have high IEs and highly negative (exothermic) EAs. They lose electrons with difficulty but attract them strongly. Thus, they form negative ions in their ionic compounds.
🗑
|
||||
| General trends: reactive metals | The elements in Groups 1A(1) and 2A(2) have low IEs and slightly negative (exothermic) EAs. They lose electrons readily but attract them only weakly, if at all. Therefore, in their ionic compounds, they form positive ions.
🗑
|
||||
| General trends: Noble gases | The elements in Group 8A(18) have very high ionization energies and slightly positive (endothermic) electron affinities. Therefore, these elements tend not to lose or gain electrons. They rarely form compounds.
🗑
|
||||
| Trends and characteristics: metals | Located in the left/lower 3/4 of the table. They are shiny solids with moderate to high melting points. They’re good thermal and electrical conductors, are ductile, malleable, and tend to LOSE electrons to non-metals.
🗑
|
||||
| Trends and characteristics: nonmetals | Located in the upper right quarter of the table. They are typically not shiny, have relatively low melting points, are poor thermal and electrical conductors, are mostly crumbly solids or gases, and tend to GAIN electrons from metals.
🗑
|
||||
| Trends and characteristics: metalloids | Located in the region between the other two classes. Have properties of both.
🗑
|
||||
| Trends and characteristics: relate all three classes | Metallic behavior decreases left to right and increases top to bottom.
🗑
|
||||
| Exceptions to the trends/characteristics of the three classes (metals, metalloids, nonmetals) | Carbon in the form of graphite is a good electrical conductor. Iodine is a shiny solid. Gallium and cesium (metals) melt at temperatures below body temperature. Mercury is liquid at room temperature. Iron is brittle.
🗑
|
||||
| Why do metals lose electrons during reactions? | Because they have low IEs compared to nonmetals.
🗑
|
||||
| Oxides | A compounds containing oxygen.
🗑
|
||||
| Distinguish metals from nonmetals with regard to their behavior of their oxides in water | Most metals TRANSFER electrons to oxygen, so their oxides are IONIC. In water, these oxides ACT AS BASES. Nonmetals SHARE electrons with oxygen, so nonmetal oxides are COVALENT. In water they act as ACIDS.
🗑
|
||||
| Amphoteric oxides | Some metals and many metalloids form oxides that are amphoteric: they can act as acids or as bases in water
🗑
|
||||
| Periodic trend relating to oxide creation | As elements become more metallic down a group, their oxides become more basic. As elements become less metallic across a period, their oxides become more acidic.
🗑
|
||||
| The ions of the elements in groups 1, 2, 6 and 7 are said to be _____ with the nearest noble gas. What does that mean? | Isoelectronic. It means they have the same number and configuration of the electrons of another species (the nearest noble gas)
🗑
|
||||
| When an alkali metal atom loses its single valence electrons, it becomes… | Isoelectronic with the previous noble gas
🗑
|
||||
| So, how where does the energy come from to produce the ions that are isoelectronic with the noble gases? And why can’t you form Na^2+ in a reaction? Why only Na+? | The energy comes from the exothermic reactions with nonmetals. The energy in that reaction is enough to remove the outer electrons of the metal. To change Na to Na^2+ would require removing core electrons, which is too much energy.
🗑
|
||||
| How about metals of groups 3, 4, 5? They would have to lose a lot of electrons to attain noble gas configuration, what happens when they form bonds? | They empty enough electrons so that they attain stability of a noble gas configuration. This is known as pseudo-noble gas configuration.
🗑
|
||||
| Pseudo-noble gas configuration: tin | Tin empties its outer energy level and attains the stability of an empty 5s and 5p sublevel and a filled inner 4d sublevel: [Kr]5s2 4d10 5p2 -> [Kr]4d10 + 4e-. Alternatively it will lose two 5p electrons: [Kr]5s2 4d10 5p2 -> [Kr]5s2 4d10 + 2e-
🗑
|
||||
| What are the retained ns2 electrons called? | They are sometimes called an inert pair because they seem difficult to remove. Thallium, lead, and bismuth, commonly form ions that retain the ns2 configuration
🗑
|
||||
| Do transition metals attain noble gas configuration when bonding? | Very rarely, and for the same reason as the group 3/4/5 metals: it’s energetically impossible to lose that many electrons. The only exceptions are scandium and titanium in some compounds.
🗑
|
||||
| The typical behavior of transition element? | To form more than one cation by losing all of its ns and some of its (n-1)d electrons.
🗑
|
||||
| Why (beginning in period 4) are the ns orbitals first added but also first taken out when increasing or decreasing energy levels?? | Using period 4 as an example: when after 4s electrons have been added 3d is then added. But 3d fills the inner orbitals and thus is not shielded, so now 4s is more “vulnerable”. This is called “crossover in orbital energy”. So apply FIFO to 4s.
🗑
|
||||
| Paramagnetism | A species with unpaired electrons exhibits paramagnetism: it is attracted by an external magnetic field. A species with all electrons paired exhibits diamagnetism: it is not attracted by a magnetic field.
🗑
|
||||
| Ionic size: cations | Cations are smaller than their parent atoms. When a cation forms, electrons are removed from the outer level. The resulting decrease in electron repulsions allows the nuclear charge to pull the remaining electrons closer.
🗑
|
||||
| Ionic size: Anions | Anions are larger than their parent atoms. When an anion forms, electrons are added to the outer level. The increase in repulsions causes the electrons to occupy more space.
🗑
|
||||
| Ionic size, periodic trends | Increase down a group. For periods: decreases down a period until it hits the first anion where it increases tremendously, then starts decreasing again.
🗑
|
||||
| How is charge related to size? | In an isoelectronic series, the larger the charge, the smaller the radius: 3<2<1<1-<2-<3-
🗑
|
Review the information in the table. When you are ready to quiz yourself you can hide individual columns or the entire table. Then you can click on the empty cells to reveal the answer. Try to recall what will be displayed before clicking the empty cell.
To hide a column, click on the column name.
To hide the entire table, click on the "Hide All" button.
You may also shuffle the rows of the table by clicking on the "Shuffle" button.
Or sort by any of the columns using the down arrow next to any column heading.
If you know all the data on any row, you can temporarily remove it by tapping the trash can to the right of the row.
To hide a column, click on the column name.
To hide the entire table, click on the "Hide All" button.
You may also shuffle the rows of the table by clicking on the "Shuffle" button.
Or sort by any of the columns using the down arrow next to any column heading.
If you know all the data on any row, you can temporarily remove it by tapping the trash can to the right of the row.
Embed Code - If you would like this activity on your web page, copy the script below and paste it into your web page.
Normal Size Small Size show me how
Normal Size Small Size show me how
Created by:
Intellex_
Popular Chemistry sets