Quantum Theory of the Atom Cheat Sheet

Chapter 7 Ð Quantum Theory of the Atom Wave Nature of Light c = speed of light = 3.00 x 108 m/s Hz) λ = wavelength (m) v = frequency (/s or c = λv Electromagnetic Spectrum Ð forms of electromagnetic radiation and their ranges of wavelength/frequency Einstein explained the photoelectric effect by postulating that light had both wave and particle properties by building on the work of Max Planck. Photoelectric effect is the ejection of electrons from the surface of a metal or from another material when light shines on it. Electrons are ejected only when the frequency of light of exceeds a certain threshold value characteristic of the particular metal. (Dependent on frequency). Light ! particle ! absorbed energy PlanckÕs Quantization of Energy E = energies of vibration in a photon v = frequency h = PlanckÕs constant 6.63 x 10-34 E = hv (this equation displays wave-particle duality) Bohr Theory of Hydrogen Atom λ in the visible spectrum of hydrogen could be reproduced by the following formula: = 1.097 x 107 /m ( - Heated solids produce continuous spectrum, heated gas produces line spectrum (all vs certain wavelengths) BohrÕs Postulates Ð account for the stability and line spectrum of the H atom. 1. An electron can have only specific energy values in an atom, called energy levels, and therefore only specific total energy values. RH = 2.179 x 10-18 J E= n = 1, 2, 3, É(principle quantum number) (allowed for quantization to be applied to electrons of atom in motion) 2. An electron in an atom can change energy only by going from one energy level to another energy level, undergoing a transition. ) Energy of emitted photon = hv = -(Ef Ð Ei) Quantum mechanics and Quantum Numbers: applies to submicroscopic particles of matter, like electrons. de Brogile Relation Ð shows that photon also has momentum (mc); only applies to electrons in free force environments, not atom (subjected to force of the nucleus) v = speed (in m/s) (often c) h = 6.63 x 10-34 Jás m = mass of a particle (in kg) λ= Uncertainty Principle of Wave Function = uncertainty in the x coordinate of the particle direction = mv (v is c, mass in kg) = uncertainty in momentum in x ѱ2: probability of finding the particle within a given region of space. - Probability is largest surrounding the nucleus (since electrons are more likely to be found surrounding this region). Quantum Numbers and Atomic Orbitals Each electron in an atom is described by 4 different quantum numbers, 3 of which specify the wave function that gives the probability of finding the electron at various points in space. Atomic orbital: a wave function for an electron in an atom (region of space with high probability of finding an electron) ms: refers to a magnetic property of electrons called spin. 1. Principle quantum number (n): the energy of an electron in an atom depends on this; it can have any positive value 1, 2, 3,É a. The smaller n, the lower the energy. b. For ions besides single-electron atomic ions, energy slight depends on the l quantum number. c. The size of an orbital also depends on n. the larger the value of n, the larger the orbital. Number of orbitals = n2. 2. Angular momentum quantum number (l): distinguishes orbitals of given n having different shapes; can have any integer value from 0 to n-1; l=subshells. a. Same n but different l = different subshells. 3. Magnetic quantum number (ml): distinguishes orbitals of given n and l Ð that is, of given energy and shape but different orientations in space; values range from -l to +l. a. Within a subshell, there are 2l + 1 orbitals of quantum number l. b. The orbitals have the same shape but different orientations in space. 4. Spin Quantum Number (ms): refers to the 2 possible orientations of the spin axis of an electron; possible values are +1/2 and -1/2. a. An electron acts as if it were spinning on its axis like earth. This spin gives rise to a circulating electric charge that would generate a magnetic field. Atomic Orbital Shapes ¥ s orbitals: spherical in shape. A 2s differs from a 1s orbital because an electron in a 2s orbital is likely to be found in 2 regions, one near the nucleus and the other in a spherical shell around the nucleus (more likely to be here) ¥ p orbitals: same basic shape of 2 lobes arranged in a straight line with the nucleus between the lobes; different in their orientations in space. ¥ d orbitals: have more complicated shapes than the first 2, also on axes. Chapter 8: Electron Configurations and Periodicity Electron Spin and the Pauli Exclusion Principle Electron spin is subject to a quantum restriction on the possible directions of the spin axis, resulting in 2 possible directions, ms = +1/2 or -1/2. Electron configuration: a particular distribution of electrons among available subshells. Orbital diagram: shows how the orbitals of a subshell are occupied by electrons. Pauli Exclusion Principle: no 2 electrons in an atom can have the same four quantum numbers. In other words, an orbital can hold at most 2 electrons, and then only if the electrons have opposite spins. Building-Up Principle and the Periodic Table An atom as an infinite number of possible electron configurations. - Configuration associated with lowest energy level = ground state. - All other configurations = excited state. - Keep in mind to count the electrons in every configuration to figure out what element it represents, then decide if itÕs excited or ground state. The chemical properties of an atom are related to the configuration of its ground state. Building-up principle (Aufbau): a scheme used to reproduce the electron configurations of the ground states of atoms by successively filling subshells with electrons in a specific order (building-up order). - Increasing order of energy in subshells (produces ground state) (from 1s to 3p) - Are determined by total energies of the atoms, not strictly the order of their energies. This total energy depends on the energies of interaction among different subshells as well. For all elements with Z = 21 or more, the energy of the 3d subshell is less than the energy of the 4s subshell. Ex. Scandium = [Ar]3d24s2 When the subshell fills to a noble gas, it is stable. Electron Configuration and the Periodic Table Noble-gas core: an inner-shell configuration corresponding to one of the noble gases plus 2 outer electrons with the ns2 configuration. - Group IIA characteristic Pseudo-noble gas core: the noble-gas core together with (n-1)d10 electrons - These electrons usually are not involved in chemical reactions. Valence electron: an electron in an atom outside the noble-gas or pseudo-noble-gas core. - Involved in chemical reactions and similarities among the valence-shell configurations account for similarities of the chemical properties among elements. Exceptions to the Building-Up Principle There are some exceptions to the building-up principle not reproducing most ground-state configurations correctly. - Chromium is actually [Ar]3d5s1. - Copper is actually [Ar]3d104s2. Orbital Diagrams of Atoms; HundÕs Rule Hund discovered an empirical rule to determine the lowest-energy arrangement of electrons in a subshell. HundÕs Rule: the lowest-energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons. Magnetic Properties of Atoms The magnetic attractions from 2 electrons opposite in spin cancel each other out. Only an atom with doubly occupied orbitals has no net spin magnetism. One with unpaired electrons does exhibit net magnetism. - Paramagnetic substance: substance that is weakly attracted by a magnetic field, and this attraction is generally the result of unpaired electrons. Ex. Na - Diamagnetic substance: substance that is not attracted by a magnetic field or is very slightly repelled by such a field. This generally means that the substance has only paired electrons. Ex. Hg Some Periodic Properties Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. - 3 physical properties: atomic radius, ionization energy, and electron affinity. Atomic Radius Ð atoms do not have a definite size because the distribution of electrons does not abruptly end, but merely decreases to super small values as they get farther from the nucleus. decreases Increasess - Why? Because first, the larger the n, the larger the size of the orbitals and thus the larger the atomic radii. Second, larger effective nuclear charge reduces the size of the orbital by pulling the electrons inward. - Effective nuclear charge: the positive charge that an electron experiences from the nucleus, equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution. Ionization Energy Ð the minimum energy needed to remove the highest-energy (outermost) electron from the neutral atom in the gaseous state. Increasess Decreasess - It is characteristic of group 1A metals to lose electrons much easier than the nonmetals. The largest ionization energies are found in the noble gas elements. This is because theyÕre so stable and unreactive. - Increasing nuclear charge divided by distance between electron and nucleus = energy needed to remove an electron. Some exceptions to the trend: - IIIA element as smaller ionization energy than the preceding IIA element. - VIA element has smaller ionization energy than preceding V element. The electrons can be moved successively: second ionization energy, third, fourth, etc. Electron Affinity Ð the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion. Basically, the amount of energy released to make a neutral atom a negative ion. - The periodic trend is a little more difficult. Electron affinities get more negative as you go to the right in a period (more energy is released and more stable negative ions are formed. o Group IIA VIIIA are positive. (absorbs energy instead of releases) o Group VA is less negative than preceding group IVA. Chapter 9: Ionic and Covalent Bonding Describing Ionic Bonds Ionic Bond: a chemical bond formed by the electrostatic attraction between positive and negative ions. - Bonds form when electrons are transferred, thus creating ions. - Representation of ionic bond electron transfer: o Na + Cl ! Na+ + Clo [Ne]3s1 + [Ne]3s23p5 ! [Ne] + [Ne]3s23p6 Lewis Electron-Dot Symbols The preceding equation can be simplified by writing Lewis electron dot symbols for the atoms and monatomic ions to represent the electron transfer. Energy involved in Ionic Bonding If atoms come together and bond, there should be a net decrease in energy because the bonded state should be more stable and thus at a lower energy level. - Think of it occurring in 2 steps: 1, an electron is transferred between 2 atoms to give ions. 2, the ions then attract one another to form an ionic bond. - Not really 2 separate steps but a practically simultaneous thing; the net quantity of energy involved is the same regardless of the order. - The energy to remove the electron from Na is ionization energy (first ionization energy) o This requires energy - The energy to add the electron to Cl is the electron affinity of Cl o This releases energy The formation of ions is not energetically favorable. It requires more energy to remove an electron than to add it. CoulombÕs Law: estimates the energy released when ions bond. E = energy k=constant -8.99x109 Jám/C2 Q1 and Q2=electric charges of ions r = distance between nuclei E= The energy released when a crystalline solid forms from ions is related to the lattice energy of the solid. Lattice energy: the change in energy that occurs when an ionic solid is separated into isolated ions in the gas phase. NaCl(s)! Na+(g) + Cl-(g) - ItÕs basically looking for enthalpy change using HessÕs law. - Use the energies of the five steps to come up with an energy for the final product: sublimation of sodium and dissociation of chlorine (each basically turns to a gas). Sodium is ionized and then chlorine is ionized. Finally, the ions are combined. Properties of Ionic Substances - High-melting solids. - Large groups of small, spherical cations and anions attract one another with strong ionic bonds forming a crystalline solid. - Liquid melt of ionic solids consist of ions so they have an electric charge. Electron Configuration of Ions Ions of the Main-Group Elements No compounds are formed with ions having charges greater than the group number because the ionization energy to remove more than 1 electron from group IA is super high, the energy needed to remove 2 electrons from group IIA is really high as well, etc... - Elements besides boron can form 3+ cations. - In period 6, elements in group IIA and VA form compounds with ions charged with group number minus 2. o Ex. Thallium = Tl+ - Tin (Sn) Lead (Pb) Thallium (Tl) and Bismuth (Bi) form cations Transition-Metal Bonds Most transition elements from cations of different charges. - Ex. Fe2+ and Fe3+ - Neither have noble-gas configurations because it would require the energetically impossible loss of 8 electrons from the neutral atom. In forming compounds, they first lose ns electrons. Then, they may lose one or more (n-1)d electrons. Ionic Radii A monatomic ion, like an atom, is a nucleus surrounded by a distribution of electrons. Ionic Radius: a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found. - Arbitrary the same way atomic radius is because an electron distribution never really ends. - Similar to atomic radii, ionic radii increase down any column because of the addition of electron shells. Isoelectronic: different species having the same number and configuration of electrons. - In any isoelectronic sequence of atomic ions, the ionic radius decreases with increasing atomic number because charge is increasing and the orbitals contract. Describing Covalent Bonds As 2 atoms approach each other, their 1s orbitals begin to overlap and the electrons then occupy the space around both atoms. Basically, the electrons can be shared by the atoms. - The distance between nuclei at the minimum energy (when repulsion of positive charges on nuclei is larger than attraction of electrons) is called bond length. - Bond dissociation energy: the energy that must be added to separate the atoms in a molecule. Lewis Formulas Lewis electron dot symbols can be used to represent the formation of a covalent bond. As the two atoms approach, unpaired electrons on each atom pair up to form a covalent bond. The pair of electrons is shared by the atoms. Bonding pair: an electron pair shared between two atoms Lone pair: an electron pair that remains on one atom and is not shared Coordinate Covalent Bonds: a bond formed when both electrons of a bond are donated by one atom. Still involves the sharing of a pair of electrons between 2 atoms; example = NH3 forms a bond with H+. Octet Rule: the tendency of atoms in molecules to have 8 electrons in their valence shells ( 2 for H). Many molecules follow this rule while some do not. Single bond: a covalent bond in which a single pair of electrons is shared by 2 atoms. Double bond: a covalent bond in which 2 pairs of electrons are shared by two atoms. Triple bond: a covalent bond in which 3 pairs of electrons are shared by 2 atoms. Polar Covalent Bonds; Electronegativity When two atoms are alike in a covalent bond, the bonding electrons are shared equally. When the two atoms are of different elements, the bonding electrons need not be shared equally. Polar covalent bond: a covalent bond in which the bonding electrons spend more time near one atom than the other. - Intermediate between nonpolar covalent bond and ionic bond. - Equally shared in nonpolar covalent, unequally shared in polar covalent, not shared in ionic. Electronegativity: a measure of the ability of an atom in a molecule to draw bonding electrons to itself. - Mulliken has suggested that the electronegativity of an atom is 1/2 (ionization energyelectron affinity) - Fluorine is most electronegative - Electronegativity increases from left to right and decreases from top to bottom. Bond polarity: the absolute value of the difference in two electronegativities. Small difference= nonpolar. Big difference = polar. Writing Lewis Electron Dot Formulas - Shows how 2 atoms are bonded, shows the positions of lone pairs of electrons. - Must know skeleton structure of the molecule first; it tells you which atoms are bonded to one another. The central atom = least electronegative atom Note that molecules with symmetrical formulas often have symmetrical structures. Ex. S2Cl2 Step 1: calculate the total number of valence electrons for molecule. Add or subtract electrons to account for polyatomic ions. Step 2: write the skeleton structure of the molecule or ion, connecting every bonded pair of atoms by a dash/pair of dots. Step 3: distribute electrons to the atoms surrounding the central atom to satisfy the octet rule. Step 4: distribute remaining electrons as pairs to the central atom after subtracting the number of electrons already distributed from total found. Sometimes, the central atom doesnÕt follow the octet rule. - Less than 8 left? ThereÕs probably a multiple bond somewhere surrounding the central atom. C, N, O, and S often form multiple bonds. Delocalized Bonding; Resonance When there is more than one possible Lewis formula for a molecule, like O3, neither formula can be correct. According to theory, one of the bonding pairs in ozone is spread over the region of all 3 atoms rather than associated with a particular O-O bond. This = delocalized bonding, a type of bonding in which a bonding pair of electrons is spread over a number of atoms rather than localized between the two. - Intermediate between single and double bond. A resonance description is often used to properly describe delocalized bonding. Resonance description: you describe the electron structure of a molecule having delocalized bonding by writing all possible electron-dot formulas. - Resonance formulas have arrows between the two formulas (double-headed) - This does not mean the molecule flips back and forth between two forms. There is only 1 ozone molecule. Exceptions to the Octet Rule Some molecules, like NO, have an odd number of electrons and cannot satisfy the octet rule. Other exceptions fall into two groups: one group of molecules with an atom having fewer than 8 valence electrons around it and another group of molecules with an atom having more than 8 valence electrons around it. Typically, anything in period 3 or greater can accommodate more than 8 electrons. - Ex. of more than 8: PF5, XeF4 - As long as the atom has available d orbitals for bonding, the atom doesnÕt have to obey the octet rule. (3rd period and below) The other group of exceptions to the octet rule consists mostly of molecules containing Group IIA or IIIA atoms. - Ex. of less than 8: BF3, BeF2 Formal Charge and Lewis Structures The formal charge of an atom in a Lewis formula is the hypothetical charge you obtain by assuming that bonding electrons are equally shared between bonded atoms and that the electrons of each lone pair belong completely to one atom. Formal charge = valence electrons on free atom Ð ½ (number of bonding electrons) Ð (number of lone-pair electrons) - The sum should equal 0 unless itÕs an ion; in that case, F.C. should equal the charge on the ion. Rule A: Whenever you can write several Lewis formulas for a molecule, choose the one having the lowest magnitudes of formal charges. Rule B: when two proposed Lewis formulas for a molecules have the same magnitudes of formal charges, choose the one with the negative formal charge on the more electronegative ion. Rule C: Choose Lewis structures that donÕt have like charges on adjacent atoms. Bond Length and Bond Order Bond distance/length: the distance between the nuclei in a bond. This sometimes provides a clue to the type of bonding present. Covalent Radii: values assigned to atoms in such a way that the sum of covalent radii of atoms A and B predicts the approximate A-B bond length. - Basically, the sum of the radii = bond length. Bond order: the number of pairs of electrons in a bond. As bond order increases, the bond strength increases and the nuclei are pulled inward, decreasing bond length. - Short bonds = strong bonds. Bond Energy The bond dissociation energy is a measure of the strength of a particular bond and is essentially the enthalpy change for a gas-phase reaction in which a bond breaks. A-B Bond Energy: the average enthalpy change for the breaking of an A-B bond in a molecule in the gas phase. - The larger the enthalpy, the stronger the bond. - Short bonds = higher bond order = stronger Molecular Geometry and Directional Bonding Molecular geometry is the general shape of a molecule as determined by the relative positions of the atomic nuclei. - Valence-shell electron-pair model usually predicts the general shape of a molecule but does not explain the chemical bonding. Valence-bond theory provides further insight into why bonds form and reveals that bonds have definite directions in space. The Valence-Shell Electron Pair Repulsion (VSEPR) Model VSEPR model predicts the shapes of molecules and ions by assuming that the valence-shell electron pairs are arranged about each atom so that electron pairs are kept as far away from one another as possible, thus minimizing electron-pair repulsions. - The direction in space of the bonding pairs gives you the molecular geometry. - In the VSEPR model, each multiple bond is treated as though it were a single electron pair. - Note the difference between the arrangement of electrons pairs and the molecular geometry, or the arrangement of nuclei. What you ÒseeÓ is nuclei arrangement (molecular geometry). What is unseen, the lone pairs, occupy the definite positions about an atom according to the VSEPR model. - VSEPR allows you to predict the approximate angles between bonds in molecules. - A lone pair tends to occupy more space than a corresponding bonding air. Why? Because the lone pair is only attracted to 1 atomic core whereas bonding is attracted to 2 and thus drawn more tightly to the nuclei. - Multiple bonds require more space than single bonds because of a greater number of electrons (smaller angles) Electron pairs (T.B.L) Arrangement of Pairs Molecular Geometry Example 2 2 0 Linear 180û Linear AX2 BeF2 3 3 3 2 0 1 Trigonal Planar 120û Trigonal planar AX3 Bent/Angular AX2 BF3 SO2 4 4 4 3 0 1 4 2 Tetrahedral 109.5û;sometimes deviate Trigonal Pyramidal AX3 CH4 NH3 2 Bent H2O Trigonal Bipyramidal AX5 PCl5 SF4 ClF3 XeF2 5 5 5 5 5 4 3 2 0 1 2 3 Trigonal Bipyramidal Not all equal angles; either 180, 90, 120 6 6 6 6 5 4 0 1 2 Octahedral Tetrahedral AX4 Seesaw (distorted) AX4 T-shaped AX3 Linear AX2 Octahedral AX6 Square pyramidal AX5 Square Planar AX4 SF6 IF5 XeF4 How to predict the geometry by the VSEPR Model: Step 1: Write the electron-dot formula from the molecular formula. Step 2: Determine the 3 of electron pairs around central atom (bonding and nonbonding) Step 3: Determine the arrangement of these electron pairs about the central atom. Step 4: Obtain the molecular geometry from the directions of the bonding pair arrangements. Dipole Moment and Molecular Geometry Dipole Moment: a quantitative measure of the degree of charge separation in a molecule. - Any molecule that has a net separation of charge has a dipole moment. - δ+ or δ- partial charge on atoms Any molecule in which the distribution of electric charge is equivalent to charges +δ and Ðδ separated by a distance d has a dipole moment = δd. Dipole moments are measured in Coulomb meters, C*m, and 1 D = 3.34 x 10-30 C*m. Capacitance: the capacity of the plates to hold a charge; affected by orientation of polar molecules in which positive end points towards positive plate and vice versa - Measure of capacitance is used to obtain dipole moments of those substances. - Bond dipoles are vector quantities, meaning they have both magnitude and direction. o Bond dipoles point from the positive to the more electronegative atom. o If dipole moment is present = polar, no dipole moment = nonpolar. ItÕs all about how symmetric the charge distribution is. Some molecular geometries imply zero (canceled out) dipole moments which allows us to narrow down the choices to which have nonzero dipole moments. Linear (AX2), Trigonal planar, tetrahedral, square planar, trigonal bipyramidal, octahedral = ZERO. This does not mean the rest are guaranteed nonzero, it just means they can be nonzero. Memorizing these is just helpful because we know these for sure canÕt be nonzero. Valence Bond Theory ^ is an approximate theory to explain the electron pair or covalent bond by quantum mechanics. Basic Theory: according to valence bond theory, a bond forms between 2 atoms when the following conditions are met: 1. An orbital on one atom comes to occupy a portion of the same region of space as an orbital on the other atom. The 2 orbitals are said to overlap. o The electrons pull the atoms together. o Strength of bonding depends on the amount of overlap; the greater the overlap the greater the bond strength. 2. The total number of electrons in both orbitals is no more than 2. o This explains why it would work with 2 H atoms but not 2 He atoms. Orbitals bond in the directions in which they protrude or point, to obtain maximum overlap. Hybrid Orbitals: orbitals used to describe the bonding that are obtained by taking combinations of atomic orbitals of isolated atoms. - In the case of CH4, a set of hybrid orbitals is constructed from one s orbital and 3 p orbitals, so they are called sp3 orbitals. - The number of hybrid orbitals formed always equals the number of atomic orbitals used. A set of hybrid orbitals always has definite directional characteristics, like angle degrees. Unhybridized orbitals are perpendicular to the plane of the hybrid orbitals and to each other. Kinds of Hybrid Orbitals Hybrid Orbitals Geometric Arrangement # of orbitals Example Degree Measures sp Linear 2 Be in BeF2 180 sp2 Trigonal planar 3 B in BF3 120 sp3 Tetrahedral 4 C in CH4 109.5 sp3d Trigonal bipyramidal 5 P in PCl5 90, 120, 180 sp3d2 Octahedral 6 S in SF6 90, 180 To obtain the bonding description about any atom in a molecule, proceed as follows: Step 1: write the Lewis electron-dot formula of the molecule. Step 2: from the Lewis formula, use VSEPR to obtain the arrangement of electron pairs about the atom. Step 3: from the geometric arrangement of the electron pairs, deduce the type of hybrid orbitals on this atom required for the bonding description. Step 4: assign valence electrons to the hybrid orbitals of this atom one at a time, pairing them only when necessary. Step 5: form bonds to this atom by overlapping singly occupied orbitals of other atoms with the singly occupied hybrid orbitals of this atom. Description of Multiple Bonding Now we will consider the possibility that more than one orbital from each bonding atom might overlap, resulting in a multiple bond. - One hybrid orbital is needed for each bond (whether a single or a multiple bond) and for each lone pairs Sigma bond: has a cylindrical shape about the bond axis. Formed either when 2 s orbitals overlap, or when an orbital with directional character, such as a p orbital or a hybrid orbital, overlaps another orbital along their axis. - Basically, the first bond. Made from hybridized orbitals. Pi bond: has an electron distribution above and below the bond axis. It is formed by the sideways overlap of 2 parallel p orbitals. - Basically, any 2nd or 3rd bond made with any other atom. Made from leftover p orbitals. Isomers: compounds of the same molecular formula but with different arrangements of the atoms. To transform isomers, bonds must be broken. Molecular Orbital Theory A theory of the electronic structure of molecules in terms of molecular orbitals, which may spread over several atoms or the entire molecule. - This theory views the electronic structure of molecules to be much like the electronic structure of atoms. Bonding and Antibonding Orbitals Molecular orbitals that are concentrated in regions between nuclei are called bonding orbitals. - Example of notation is σ1s. Subscript tells us it was obtained from 1s atomic orbitals. Molecular orbitals having zero values in the region between two nuclei and therefore concentrated in other regions are called antibonding orbitals. - Example of notation is σ*1s Bond Order Bond order refers to the number of bonds that exist between two atoms. nb = number of electrons in bonding orbitals na = number of electrons in bonding orbitals Bond order = ½ (nb Ð na) Factors that Determine Orbital Interaction The strength of the interaction between 2 atomic orbitals to form molecular orbitals is determined by 2 factors: the energy difference between the interacting orbitals, and the magnitude of their overlap. Electron Configurations of Diatomic Molecules of the Second-Period Elements Homonuclear diatomic molecules: molecules composed of 2 like nuclei. (H2, He2) Heteronuclear diatomic molecules: molecules composed of two different nuclei (CO, NO) This order of molecular orbitals reproduces the known electron configurations of homonuclear diatomic molecules composed of elements in the second row of the period table. The order of filling is: σ2s σ*2s Π2p σ2p Π*2p σ*2p You can use the order of filling to obtain the orbital diagram, magnetic character, electron configuration, and bond order of a homonuclear diatomic molecule.