Chem 3

Tellurium electron configuration Valence electrons?

6 valence electrons

Technetium electron configuration Valence electrons

7 valence electrons

Na electron configuration. Valence and core electrons?

[Ne] 3s1 or 1s2 2s2 2p6 3s1 -1 valence electron -10 core electrons

Chromium electron configuration

1s2 2s2 2p6 3s2 3p6 4s2 3d4 --> 1s2 2s2 2p6 3s2 3p6 4s1 3d5 OR [Ar] 4s1 3d5 Ar is the nearest noble gas that comes before it

Orbital diagram

-Electrons occupy orbitals so as to minimize the energy of the atom; therefore, lower-energy orbitals fill before higher-energy orbitals. Orbitals fill in the following order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s. --Orbitals can hold no more than two electrons each. When two electrons occupy the same orbital, their spins are opposite. This is another way of expressing the Pauli exclusion principle (no two electrons in one atom can have the same four quantum numbers). --When orbitals of identical energy are available, electrons first occupy these orbitals singly with parallel spins rather than in pairs (Hund's rule). Once the orbitals of equal energy are half-full, the electrons start to pair.

Group 7A elements (ns^2np^5, n>=2)

-Halogens, -most reactive non-metals, not found uncombined in nature -Highest (most negative) electron affinities. If you give it one electron, it releases a lot of energy. -With increasing reactivity, larger negative electron affinity

Quantum numbers of helium's electrons

-Helium has two electrons --Electrons like to occupy the lowest state of energy possible -Both electrons are in the first energy level, Both electrons are in the s orbital of the first energy level -Since they are in the same orbital, they must have opposite spins -first electron: n=1, 1=0, ml=0 ms=+1/2 (1s) -second electron: n=1, 1=0, ml=0 ms=-1/2 (1s, so 1s2, 2 represents second electron) - The two electrons have three quantum numbers in common (because they are in the same orbital), but each electron has a different spin quantum number (as indicated by the opposing half-arrows in the orbital diagram).

Group 8A elements (ns^2np^6, n>=2)

-Noble gasses are the most inert -Completely filled ns and np subshells -Highest ionization energy of all elements (don't want to lose electrons), Positive EA (so they hate gaining electrons) -They don't like to be compounds -Reactivity decreases down a group

Group 1A elements (ns1, n >= 2)

-Alkali metals -Most reactive metals, not found uncombined in nature -Only one electron in their valence shell and IE decreases as reactivity increases. So it is easiest to exchange electrons for Cs

Periodic properties of the elements: Ionization energy

-As you go down any groups, the ionization energy decreases because as you go down the group it becomes easier to remove an electron from the neutral atom because that electron is lying at a higher energy level and thus further away from the nucleus -As you go across a period (left to right), the ionization energy increases because electrons in the outermost principal energy level generally experience a greater effective nuclear charge (Z eff)

Periodic properties of the elements: atomic radius

-Atomic radius increases down a group because as we move down a column in the periodic table, the principal quantum number (n) of the electrons in the outermost principal energy level increases, resulting in larger orbitals and therefore larger atomic radii. -Atomic radius decreases across period (left to right) because as we move to the right across a row in the periodic table, the effective nuclear charge (zeff) experienced by the electrons in the outermost principal energy level increases, resulting in a stronger attraction between the outermost electrons and the nucleus, and smaller atomic radii. -If you go across the row/period of transition metals, there is not a consistent trend. This is because of the effective nuclear charge, which is basically the same for all of them because they both have 1 or 2 valence electrons

Electron Configuration and Ion charge

-Atoms form ions that will result in an electron configuration that is the same as the nearest noble gas -Group 1A=+1 (lose a valence electron), Group 2A = +2 (lose two valence electrons), Group 7A = -1 (gain an electron), Group 6A= -2 (gain two electrons) etc -The tendency for many main-group elements to form ions with noble gas electron configurations does not mean that the process is in itself energetically favorable. In fact, forming cations always requires energy, and forming anions sometimes requires energy as well. However, the energy cost of forming a cation or anion with a noble gas configuration is often less than the energy payback that occurs when that cation or anion forms chemical bond

Trends in electron affinity

-Electron affinity is a measure of how easily an atom accepts an additional electron and is crucial to chemical bonding because bonding involves the transfer or sharing of electrons. -Most groups (columns) of the periodic table do not exhibit any definite trend in electron affinity. Among the group 1A metals, however, electron affinity becomes more positive as we move down the column (adding an electron becomes less exothermic). -Electron affinity generally becomes more negative (adding an electron becomes more exothermic) as we move to the right across a period (row) in the periodic table because ionization energy increases

Valence electrons

-The electrons in all the subshells with the highest principal energy shell (highest n) -For main-group elements, the valence electrons are those in the outermost principal energy level. For transition elements, we also count the outermost d electrons among the valence electrons (even though they are not in an outermost principal energy level). -Helium is an exception. Even though it lies in the column with an outer electron configuration of ns2 np6, its electron configuration is simply 1s2. ---Except for helium, the number of valence electrons for any main-group element is equal to its lettered group number. For example, we know that chlorine has seven valence electrons because it is in group number 7A. -Lastly, note that, for main-group elements, the row number in the periodic table is equal to the number (or n value) of the highest principal level. For example, because chlorine is in row 3, its highest principal level is the n=3 level.

Trends in successive ionization energies

-The jump in Na is in the 1st and 2nd electron being removed, while the jump in Mg is between the 2nd and 3rd electron being removed -This can be justified by valence electrons. Na has one valence electron, but the next electron is a core electron. Meanwhile, Mg has two valence electrons but at the third ionization energy, the third electron is hard to remove because it is the a core electron -If there is a huge jump between ionization energies, it is because you must remove a core electron, which is hard to remove

Summarizing Periodic Table Organization

-The periodic table is divisible into four blocks corresponding to the filling of the four quantum sublevels (s, p, d, and f). -The lettered group number of a main-group element is equal to the number of valence electrons for that element. -The row number of a main-group element is equal to the highest principal quantum number of that element.

The Transition and Inner Transition Elements

-The principal quantum number of the d orbitals that fill across each row in the transition series is equal to the row number minus one. In the fourth row, the 3d orbitals fill; in the fifth row, the 4d orbitals fill; and so on. This happens because, the 4s orbital is generally lower in energy than the 3d orbital (because the 4s orbital more efficiently penetrates into the region occupied by the core electrons). The result is that the 4s orbital fills before the 3d orbital, even though its principal quantum number (n=4) is higher. -As we move across the f block (the inner transition series), the f orbitals fill. For these elements, the principal quantum number of the f orbitals that fill across each row is the row number minus two. (In the sixth row, the 4f orbitals fill, and in the seventh row, the 5f orbitals fill.) In addition, within the inner transition series, the close energy spacing of the 5d and 4f orbitals sometimes causes an electron to enter a 5d orbital instead of the expected 4f orbital.

Atomic radii of transition metals roughly the same size across the d block

-the atomic radius plotted as a function of atomic number for the first 57 elements in the periodic table. Notice the periodic trend in the radii. Atomic radii peak with each alkali metal. -Defective nuclear charge is the same for transition elements -across a row of transition elements, the number of electrons in the outermost principal energy level (highest n value) is nearly constant. As another proton is added to the nucleus of each successive element, another electron is added as well, but the electron goes into an n highest - 1 orbital. The number of outermost electrons stays constant, and the electrons experience a roughly constant effective nuclear charge, keeping the radius approximately constant.

Group 2A elements (ns2, n >= 2)

Alkali earth metals Also reactive (not as much as group 1A though) -reactivity increases down group

Electron configurations for multielectron atoms

Aufbau principle: The principle that indicates the pattern of orbital filling in an atom -Explains the pattern of orbital filling -Unless otherwise specified, we use the term electron configuration to mean the ground state (or lowest energy) configuration As we move to the right across a row (which is also called a period), the orbitals fill in the correct order. With each subsequent row, the highest principal quantum number increases by one.Notice that as we move down a column, the number of electrons in the outermost principal energy level (highest n value) remains the same. -Whenever you end up with a a configuration with 4 or 9 electrons in the d subshell (one away from half full or complete full), d subshell will steal an electron from the s subshell before and make itself full (4s2 3d4 → 4s1 3d5)

Periodic properties of the elements: Ionic radius

Cations smaller than neutral atom; anions bigger than neutral atom -This is due to the imbalance in charge. Cations: Too many protons pull onto electrons, which causes them to suck in close to the nucleus, causing the ion to shrink. (For isoelectronic species, same amount of electrons) Anions: Too many electrons, so the protons don't have enough force to pull the electrons in so the anion swells. (For isoelectronic species, same amount of electrons)

Write the electron configuration and determine whether the Cu atom, Cu+, and Cu2+ ions are paramagnetic or diamagnetic

Cu=[Ar]4s^23d^9--> [Ar]4s3d^10 (paramagnetic) Cu+=[Ar]4s^03d^10 --> [Ar]3d^10 (4s is valence shell, diamagnetic) Cu2+=[Ar]4s^03d^9 (paramagnetic)

Core Electrons

Electrons in lower energy shells The core electrons are those in complete principal energy levels and those in complete d and f sublevels. For example, silicon, with the electron configuration 1s^22s^22p^63s^23p^2 has four valence electrons

Multielectron Atoms

In multielectron atoms, because of electron-electron repulsions, the sublevel energies are not degenerate In multielectron atoms (anything with more than one electron, so anything besides hydrogen), energy depends on n and l -Degenerate: Describes two or more electron orbitals with the same value of n that have the same energy ---The 3s, 3p, and 3d orbitals (which are empty for hydrogen in its lowest energy state) all have the same energy—we say they are degenerate ---The orbitals within a principal level of a multi-electron atom, in contrast, are not degenerate—their energy depends on the value of l. We say that the energies of the sublevels are split. In general, the lower the value of l within a principal level, the lower the energy (E) of the corresponding orbital. Thus, for a given value of n: E(s orbital) < E(p orbital) < E(d orbital) < E(f orbital) OR S (1=0) < p (1=1) < d (1=2) < f (1=3)

Other Irregularities in the Trend

Ionization Energy generally increases from left ot right across a period -Except from 2A to 3A, 5A to 6A -The change in going from the s block to the p block causes the 2p orbital to penetrate into the nuclear region less than the 2s orbital. Consequently, the 1s electrons shield the electron in the 2p orbital from nuclear charge more than they shield the electrons in the 2s orbital. The result is that the 2p orbitals are higher in energy, and therefore the electron is easier to remove (it has a lower first ionization energy). -Although oxygen is to the right of nitrogen in the same row, it has a lower first ionization energy. This exception is caused by the repulsion between electrons when they occupy the same orbital.

Pauli exclusion principle

No two electrons in an atom may have the same set of 4 quantum numbers. Therefore, no orbital may have more than 2 electrons, and they must have opposite spins -S sublevel has 1 orbital, therefore it can hold 2 electrons --ml=0 -P sublevel has 3 orbitals, therefore it can hold 6 electrons --ml=-1,0,1 -D sublevel has 5 orbitols, therefore it can hold 10 electrons --ml=-2,-1,0,1,2 -F sublevel has 7 orbitals, therefore it can hold 14 electrons --ml=-3 - 3

paramagnetic vs diamagnetic

Paramangnetic: net magnetic field Diamagnetic: no net magnetic field

First ionization energies

Peaks at noble gasses

Effective nuclear change

The actual nuclear charge experience by an electron, defined as the change of the nucleus plus the charge of the shielding electrons -The net positive charge that is attracting a particular electron -Z effective = Z (atomic) - S (core electrons) -When the third electron is far from the nucleus, it experiences the 3+ charge of the nucleus through the screen or shield of the 2- charge of the two 1s electrons. We can think of the third electron as experiencing an effective nuclear charge of approximately 1+(3+ from the nucleus and 2− from the electrons, for a net charge of 1+). The inner electrons in effect shield the outer electron from the full nuclear charge. -In the pic, the electron outside is experiencing a charge of 3-2 because the 2 electrons inside are shielding --Any one electron in a multi-electron atom experiences both the positive charge of the nucleus (which is attractive) and the negative charges of the other electrons (which are repulsive). Consider again the outermost electron in the lithium atom: (Li=1s^22s^1) even though the 2s orbital penetrates into the 1s orbital to some degree, the majority of the 2s orbital is outside of the 1sorbital. Therefore, the electron in the 2s orbital is partially screened or shielded from the 3+ charge of the nucleus by the 2- charge of the 1s (or core) electrons, reducing the net charge experienced by the 2s electron. -But if the electrons are all inside, so they experience a charge of 3. Because it is inside, it is feeling more of the 3+ charge because it is closer.

Shielding

The effect of an electron of repulsion by electrons in lower energy orbitals that screen it from the full effects of nuclear change 0For example, consider a lithium ion (Li+). Because the lithium ion contains two electrons, its electron configuration is identical to that of helium: Li+= 1s2 -Core electrons efficiently shield electrons in the outermost principal energy level from nuclear charge, but outermost electrons do not efficiently shield one another from nuclear charge. The effective nuclear charge experienced by an atom's outermost electrons continues to become more positive as we move to the right across the rest of the second row in the periodic table, resulting in successively smaller atomic radii. The same trend is generally observed in all main-group elements.

Ionization energy

the energy required to remove an electron from the atom or ion in the gaseous state. Ionization energy is always positive because removing an electron always takes energy. The energy required to remove the first electron is the first ionization energy (IE1). For example, we represent the first ionization of sodium with the equation: Na (g) --> Na+(g) + 1e-, IE=497 kj/mol The energy required to remove the second electron is the second ionization energy (IE2), the energy required to remove the third electron is the third ionization energy (IE3), and so on. We represent the second ionization energy of sodium as: Na (g) --> Na2+(g) + 1e-, IE=4560 kj/mol -Notice that the second ionization energy is not the energy required to remove two electrons from sodium (that quantity is the sum of IE1and IE2), but rather the energy required to remove one electron from Na+. We look at trends in IE1 and IE2 separately.

Hund's rule

when filling degenerate orbitals, electrons fill them singly first, with parallel spins. Hund's rule is a result of an atom's tendency to find the lowest energy state possible. When two electrons occupy separate orbitals of equal energy, the repulsive interaction between them is lower than when they occupy the same orbital because the electrons are spread out over a larger region of space. By convention, we denote these parallel spins with half arrows pointing up. -Electrons with parallel spins have correlated motion that minimizes their mutual repulsion.