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Gen. Chem. I chap 8
| Question | Answer | ||
|---|---|---|---|
| Degeneration | We call orbitals with the same energy degenerate | ||
| For multielectron atoms, the energies of the sublevels are split | 1) caused by charge interaction, shielding and penetration. 2) The lower the value of the l quantum number, the less energy the sublevel has: s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3) | ||
| (Chapter 7) The Principal Quantum Number (n) | determines the overall SIZE and ENERGY of an orbital. Possible values are n= 1,2,3… and so on. | ||
| (Chapter 7) The Angular Momentum Quantum Number (l) | determines the shape of the orbital. The possible values of l from 0,1,2,…, UP TO (n-1). | ||
| (Chapter 7) The magnetic Quantum Number (m1) | specifies the orientation of the orbital. The possible values of m1 are ranging from –L to +L and includes 0. For example: 1) is L=0, then m1=0. 2) L=1, then m1= -1,0, and +1. 3) L=2, then m1= -2,-1,0,+1,+2. | ||
| Coulomb’s Law | describe the attractions and repulsions between charged particles. (equation written on flashcard). 1) For the SAME charges, the potential energy (E) is POSITIVE and DECREASES as the particles get father apart (as ‘r’ increases). | Since systems go toward the lower potential energy level, the like charges repel each other (like magnets). | 2) For OPPOSITE charges, the potential energy is NEGATIVE and becomes MORE NEGATIVE as particles get closer together (as r decreases). OPPOSITE CHARGES ATTRACT! 3) The MAGNITUDE of the interaction between charged particles INCREASES as the CHARGES of the |
| Nucleus | is positively charged and is attractive. It attracts electrons to itself. | ||
| Shielding | The repulsion of one electron by another electron is SHIELDING that electron from the full effects of the nuclear charge. . | An electron far from the nucleus is partly shielded by the electrons in the 1s orbital, reducing the effective net nuclear charge it experiences | |
| Shielding continued | For MULTIELECTRIC atoms, one electron experiences both the positive charge of the nucleus (which is attractive) and the negative charges of the other elements (which are repulsive). | ||
| Effective nuclear charge | The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron | ||
| Penetration | 1) Imagine that the third electron (which was farther away), now gets closer to the nucleus. | As the electron PENETRATES the electron cloud of the 1s electrons it begins to experience the 3+ charge of the nucleus more fully | because it is less shielded by the intervening electrons. 2)*In other words: as the electron undergoes penetration, it experiences a greater nuclear charge and therefore a LOWER ENERGY. |
| Electron spatial distributions and sublevel splitting= the splitting is caused by the spatial distribution of electrons within a sublevel. | |||
| Comparing 2p and 2s orbitals= an electron in the 2p orbital is more likely to be found closer to the nucleus than an electron in a 2s orbital. BUT the 2s orbital is LOWER in ENERGY ONLY when the 1s orbital is OCCUPIED. | |||
| Electron configuration | used to mean the GROUND STATE (or lowest energy) configuration. | ||
| The number of electrons in a neutral atom is = to | to its atomic number. | ||
| Filling the Orbitals with Electrons | 1) Energy levels and sublevels fill from lowest energy to high. s -> p -> d -> f. (Aufbau Principle) | 2) Orbitals that are in the same sublevel have the same energy. No more than two electrons per orbital. (Pauli Exclusion Principle). | 3)When filling orbitals that have the same energy, place one electron in each before completing pairs. (Hund’s Rule). |
| Aufbau Principle | the pattern of orbital. For example: Lithium has 3 electrons. The electron configuration (the lowest energy) is: 1s^2 2s^1. (1s has 2 electrons with opposite spins). Example: carbon has 6 electrons: 1s^2 2s^2 2p^2 | ||
| Hund’s rule | when filling degenerate orbitals, electrons fill them SINGLY FIRST, with parallel spins (opp). | ||
| Summarizing Orbital filling | 1) lower energy orbitals fill before higher energy orbitals. 2) Orbitals fill in the following order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s. | 3) Orbitals can hold NO MORE than 2 ELECTRONS EACH with opposite spins. | 4) When ORBITALS of the SAME ENERGY are available, electrons FIRST occupy these orbitals SINGLY with parallel spins rather than in pairs. Once the orbitals of equal energy are half-full, the electrons start to pair. |
| Process on how to write an electron configuration for an element | 1) First find its ATOMIC NUMBER from the periodic table. This number= the number of ELECTRONS. 2) Then use the order of filling to distribute the electrons in the appropriate orbitals. *Remember each orbital can only hold 2 electrons maximum. | ||
| Process on how to write an electron configuration for an element Continued | 1) the ‘s’ sublevel has only one orbital and can therefore hold only 2 electrons. 2) the ‘p’ sublevel only has 3 orbitals and can only hold 6 electrons. | 3) the ‘d’ sublevel has 5 orbitals and has 10 electrons. 4) the ‘f’ sublevel has 7 orbitals and 14 electrons. | |
| valence electrons | 1) The electrons in all the sublevels with the HIGHEST PRINCIPLE energy shell. 2) Therefore, valence electrons are on the OUTERMOST PRINCIPLE ENERGY LEVEL. 3) Important in chemical bonding. | ||
| Why do the elements in a column of the periodic table have similar chemical properties | they have the SAME NUMBER of VALENCE ELECTRONS. | ||
| Core electrons | Electrons in lower energy shells are called core electrons. | ||
| the most important factors in the way an atom behaves, both chemically and physically, is the | number of valence electrons. | ||
| ***(important) Electron configuration and the periodic table | 1) **The GROUP NUMBER corresponds to the NUMBER OF VALENCE ELECTRONS. 2) The length of each “block” is the MAXIMUM NUMBER OF ELECTRONS THE SUBLEVEL CAN HOLD. 3) The Period number corresponds to the principal energy level of the valence electrons. | ||
| Because of the filling order of orbitals, the periodic table can be divided into | 2 blocks which represent the filling of particular sublevels. | ||
| The first two columns on the left side of the periodic table | are the ‘s’ block. With outer electron configuration (lowest energy level) of ns^1 (the alkali metals) and ns^2 ( the alkaline earth metals). | ||
| Inner electron configuration | Ex: take element Se. The inner electron configuration would be Ar. It is the noble gas that goes BEFORE the element. | ||
| Periodic table: n= | row number. | ||
| Outer electron configuration | 1) Ex: take the element Se. The noble gas BEFORE it is Ar. Therefore, the outer electron configuration can be found by tracing the elements between Ar and Se and assigning electrons to the appropriate orbitals. | 2) Because Se is in row 4, ADD two 4s electrons as you trace across the ‘s’ block (n= row number). | 3) Then add ten 3d electrons as you trace across the ‘d’ block (n=row number -1). 4) lastly, add four 4p electrons as you trace across the ‘p’ black to Se, which is in the fourth column of the ‘p’ block (n= row number). |
| D block | is the row minus 1. | ||
| F black | is the row number minus 2. | ||
| The chemical properties of elements are largely determined by | the number of valence electrons they contain. | ||
| Elements with electron configuration CLOSE to those of the noble gases are | most reactive because they can attain noble gas electron configuration by losing or gaining a small number of electrons. | ||
| Alkaline metals are the most reactive. And halogens are the most reactive non metals. | |||
| Irregular Electron Configurations | We know that because of sublevel splitting, the 4s sublevel is lower in energy than the 3d; and therefore the 4s fills before the 3d. . | But the difference in energy is not large. | Some of the transition metals have irregular electron configurations in which the ns only partially fills before the (n−1)d or doesn’t fill at all. Therefore, their electron configuration must be found experimentally |
| The properties of electron configuration | The properties of the elements follow a periodic pattern. Elements in the same column have similar properties. The elements in a period show a pattern that repeats. | ||
| The noble gases | have eight valence electrons. EXCEPT for He, which has only two electrons. We know the noble gases are especially non-reactive. | He and Ne are practically inert. The reason the noble gases are so non-reactive is that the electron configuration of the noble gases is especially stable. | |
| Halogens | gain one electron. They are one electron away from being stable like the noble gases. | ||
| In their reactions, the alkali metals tend to | lose one electron, resulting in the same electron configuration as a noble gas: forming a cation with a 1+ charge. | ||
| 8 valence electrons | Quantum mechanical calculations show that eight valence electrons should result in a very unreactive atom: an atom that is very stable. The noble gases have eight valence electrons and are all very stable and unreactive | ||
| Conversely, elements that have either one more or one less electron from 8 valence electrons: | should be very reactive. The halogen atoms have seven valence electrons and are the most reactive nonmetals. The alkali metals have one more electron than a noble gas atom and are the most reactive metals as a group. | ||
| Electron Configuration of Anions in Their Ground State | Anions are formed when nonmetal atoms gain enough electrons to have eight valence electrons: filling the s and p sublevels of the valence shell. The sulfur atom has six valence electrons. | ||
| Electron Configuration of Cations in Their Ground State | Cations are formed when a metal atom LOSES ALL its VALENCE ELECTRONS. Resulting in a new lower energy level valence shell. However the process is always endothermic. The magnesium atom has two valence electrons. | ||
| There are several methods for measuring the radius of an atom, and they give slightly different numbers | 1) van der Waals radius = nonbonding. Represents the radius of an atom when it is not bonded to another atom. 2) covalent radius = bonding radius. Another way to define the size of an atom. | 3) atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds. | |
| Atomic Radius Increases | down group. | ||
| Atomic Radius Decreases | across period (left to right). Adding electrons to same valence shell. | ||
| In general, atomic radii INCREASES as we move DOWN a COLUMN and DECREASES as we move to the RIGHT across a period in the periodic table. | |||
| Outer electrons are shielded from nucleus(3+) by the | core electrons. Screening or shielding effect. 1)Outer electrons do not effectively screen for each other. 2)The shielding causes the outer electrons to not experience the full strength of the nuclear charge. | ||
| Effective nuclear charge | The effective nuclear charge is net positive charge that is attracting a particular electron. | ||
| Summarizing Atomic Radii for Main-Group Elements: | 1) As we move DOWN a COLUMN in the periodic table, the PRINCIPLE QUANTUM NUMBER (n) of the electrons in the outermost principle energy level increases, | resulting in LARGER orbitals and therefore LARGER ATOMIC RADII. 2) Move to the RIGHT across a row in the periodic table, the effective nuclear charge (Z effective) experienced by the electrons in the OUTERMOST PRINCIPLE ENERGY LEVEL | INCREASES, resulting in a stronger attraction between the OUTERMOST ELECTRONS and the nucleus, and SMALLER ATOMIC RADII. |
| But it is different when moving to the right across a row in TRANSITION ELEMENTS | instead of decreasing in size the RADII OR TRANSITION ELEMENTS STAYS ROUGHLY THE SAME across each row. | ||
| Quantum-Mechanical Explanation for the Group Trend in Atomic Radius | Quantum-mechanics predicts the atoms should get larger down a column. Quantum-mechanics predicts the atoms should get smaller across a period. | 1)The size of an atom is related to the distance the valence electrons are from the nucleus. 2)The larger the orbital an electron is in, the farther its most probable distance will be from the nucleus and the less attraction it will have for the nucleus. | 3)Traversing down a group adds a principal energy level. 4)The larger the principal energy level an orbital is in, the larger its volume. 5)The stronger the attraction the valence electrons have for the nucleus, the closer their average distance will be t |
| Electron Configurations of Transition Metal Cations in Their Ground State | 1)When transition metals form cations, the first electrons removed are the valence electrons, even though other electrons were added after. 2)Electrons may also be removed from the sublevel closest to the valence shell after the valence electrons. | ||
| PARAMAGNETISM | 1)will be ATTRACTED to a magnetic field. 2)Electron configurations that result in UNPAIRED electrons mean that the atom or ion will have a NET MAGNETIC FIELD– this is called paramagnetism. | ||
| diamagnetism | 1) slightly REPELLED by a magnetic field. 2)Electron configurations that result in all PAIRED electrons mean that the atom or ion will have NO MAGNETIC FIELD. | ||
| Cations smaller than neutral atoms; anions larger than neutral atoms | Larger positive charge = smaller cation. Larger negative charge = larger anion. | ||
| Trends in Ionic Radius | 1)Ions in same group have same charge. 2)Ion size increases down the column. 3)higher valence shell, larger. | 4)Cations smaller than neutral atoms; anions larger than neutral atoms. Cations smaller than anions. Except Rb+ & Cs+ bigger or same size as F− and O2− . 4)Larger positive charge = smaller cation. 5)Larger negative charge = larger anion | |
| isoelectronic | = same electron configuration | ||
| Pauli’s exclusion principle: how many electrons can exist in an orbitals | 1) s sublevel has 1 orbital, therefore it can hold 2 electrons. 2) p sublevel has 3 orbitals, therefore it can hold 6 electrons. | 3) d sublevel has 5 orbitals, therefore it can hold 10 electrons. 4) f sublevel has 7 orbitals, therefore it can hold 14 electrons | |
| Quantum Numbers of Helium’s Electrons | 1) Helium has two electrons. 2) Both electrons are in the first energy level. 3) Both electrons are in the s orbital of the first energy level. 4) Because they are in the same orbital, they must have opposite spins. |
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