Electron Configuration, Periodicity, Bonding, and Molecular Geometry By XX Bonding and Molecular Geometry Table of Contents Section Topics 1. Key terms & principles Ionic versus Covalent Bonding Electronegativity Polarity 2. Octet Rule and the Exceptions to the rule Octet Rule Exception to the Octet Rule: Electron deficient molecules Exception to the Octet Rule: Expanded Octet 3. Hybrid Lewis Structures Lewis structures Formal Charge 4. Resonance structures Resonance structures 5. Molecular geometry Molecular geometry of molecules Unshared electronic pairs 6. Atomic orbitals Hybridization Sigma & Pi bonds 3 1. Key terms & principles Ionic versus covalent bonding Ionic Bond Covalent Bond Electrons transferred from one atom to another Electrons shared between atoms (single or multiple) Opposite charges hold the molecule together Sharing creates strong bond 1. Key terms & principles Electronegativity Electronegativity: tendency for atom to attract a bonding pair of electrons 1. Key terms & principles Which is the most electronegative element? ? a. Fr (Atomic number 87) b. Cs (Atomic number 55) c. Br (Atomic number 35) d. F (Atomic number 9) e. C (Atomic number 6) 1. Key terms & principles Polarity: electrons pulled towards one atom resulting in partial charge H2O molecule: Electrons are “closer” to Oxygen Opposite ends of molecules have a slight positive or negative charge because of an asymmetrical distribution of electrons 2. Octet Rule and the exceptions to the rule Octet Rule: Atoms want to fill their valence shells Octet rule: Atoms want to gain, lose or share electrons to have 8 electrons in valence shell (outer shell) where possible Electrons needed to complete valence shell 1 (no octet since can only hold 2 ) 4 3 2 2. Octet Rule and the exceptions to the rule Exceptions to the Octet Rule: Electron deficient molecules NO molecule with 11 valence electrons Some species do not follow the octet rule such as when there are odd number of valence electrons (e.g., Nitrogen has odd number of electrons in NO and NO2) Unpaired electrons are put on other atoms to get formal charge of zero Octet is NOT filled – Octet is filled – 8 7 electrons (needs 1 electrons more) BUT formal charge is 0 (remember for later) 2. Octet Rule and the exceptions to the rule Exceptions to the Octet Rule: Expanded Octet PCl5 Molecule When there are “too many” electrons, some elements can expand their octet and hold more than 8 electrons in their valence shell by utilizing the d orbitals Elements with atomic numbers greater than or equal to that of Si are able to have an expanded octets (e.g., Si, P, S, Cl) P has 5 bonds resulting in 10 electrons in the valence shell Formal charge is 0 (remember for later) 3. Hybrid Lewis Structure Hybrid Lewis Structures Diagram that show the bonding between atoms of a molecule Single lines represent single bonds, two lines represent double bonds, etc. Dots represent lone pairs and unshared electrons 3. Hybrid Lewis Structure Formal Charge – helps you decide between multiple options 𝑪𝑶𝟐 Example 3. Hybrid Lewis Structure Formal charge (FC) – an elegant technique to help you decide FC is the charge an atom would have if valence electrons distributed evenly FC assigned to each atom to help inform how to balance electrons across atoms in given molecule (try to get to “0” for each atom) Sum across all atoms to minimize formal charge across atoms (e.g., ions) Formal Charge (FC) = [# of valence electrons on atom] – [non-bonded electrons + number of bonds] 3. Hybrid Lewis Structure Steps to Follow for Drawing Lewis Dot Structure: 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) 3. Draw the skeleton structure w/ single bonds 4. Determine the remaining number of valence electrons 5. Apply octet rule 6. Formal charge - balance center atom first 7. Leftover electrons/ lone pairs go to the center atom 8. Beware of the number of electron clouds - lone pairs count as electron clouds 3. Hybrid Lewis Structure Hybrid Lewis Dot Structure: Practice Problem Draw Hybrid Lews Structure for Methane (CH4) 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons CH4 has 8 valence electrons (from periodic table) ‒ Carbon has 4 ‒ Hydrogen has 1 each (4 total) 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) The Carbon is the center atom 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) 3. Draw the skeleton structure w/ single bonds Skelton structure looks as follows: 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) 3. Draw the skeleton structure w/ single bonds 4. Determine the remaining number of valence electrons There are no remaining valence electrons, all 8 are used 2 e- 2 e- 2 e2 e- 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) 3. Draw the skeleton structure w/ single bonds 4. Determine the remaining number of valence electrons 5. Check to make sure that all atoms are abiding by the octet rule (hydrogen only needs two electrons for its valence shell): 2 e- Apply octet rule 2 e- 2 e2 e- 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) 3. Draw the skeleton structure w/ single bonds 4. Determine the remaining number of valence electrons 5. Apply octet rule 6. Formal charge - balance center atom first Check to ensure atoms are balanced via formal Charge Carbon has formal charge of 0 Hydrogen has full valence shell of 2 electrons 2 e- 2 e- 2 e2 e- 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) 3. Draw the skeleton structure w/ single bonds 4. Determine the remaining number of valence electrons 5. Apply octet rule 6. Formal charge - balance center atom first 7. There are no leftover electrons or lone pairs, hence no need to put on center atom 2 e- 2 e- 2 e2 e- Leftover electrons/ lone pairs go to the center atom 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4) 1. Determine total number of valence electrons 2. The less electronegative atoms are in the center (except hydrogen) 3. Draw the skeleton structure w/ single bonds 4. Determine the remaining number of valence electrons 5. Apply octet rule 6. Formal charge - balance center atom first There are four electron clouds/areas of density, no electron clouds from “lone” pairs 2 e- 2 e- 2 e2 e- 7. Leftover electrons/ lone pairs go to the center atom 8. Beware of the number of electron clouds lone pairs count as electron clouds 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for NOCl (Nitrosyl chloride) Use the steps to draw the lewis structure for NOCl 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for NOCl (Nitrosyl chloride) Nitrogen is the least electronegative out of the three, so you know that it will be in the middle Ensure the octet rule (which makes you realize that Nitrogen needs double bond for octet) Check that formal charges are correct 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for NO3- (Nitrate) Use the steps to draw the lewis structure for NO3- (Nitrate) 3. Hybrid Lewis Structure PRACTICE: Hybrid Lewis Dot Structure for NO3- (Nitrate) Nitrogen is the least electronegative (will be in the center) with Oxygen at the terminals All atoms satisfy the octet rule The formal charge is not zero (it is an ion) 4. Resonance Structures Drawing Resonance Structures Approach for “resonance structures” First select one correct Lewis structure for the molecule ? What happens when there are multiple Lewis Dot Structures that are possible for the electron configuration of a molecule? Make resonance structures off the selected structure - Rotate through the different positions the double, triple, etc Make sure that when you move the position of the bonds that the molecule maintains its original properties (formal charge, hybridization etc.) 4. Resonance Structures Drawing Resonance Structures 4. Resonance Structures Drawing Resonance Structures 4. Resonance Structures Drawing Resonance Structures 4. Resonance Structures PRACTICE: Create resonance structure for two molecules 1. NCO- NCO- 2. CO32- 4. Resonance Structures PRACTICE: resonance structure for NCO- NCO- 4. Resonance Structures PRACTICE: resonance structure for NCO- CO32- 5. Molecular geometry Guideline for molecular geometry of molecules Predicting the shape of molecules Once the Lewis structure has been completed, the key is to arrange the electron pairs around the central atom in way that minimizes repulsion (as far part as possible) The electronic repulsion means the electron configuration outcome may be different than the molecular geometry of the molecule The table on the left provides a guide for what the molecular geometry looks like 5. Molecular geometry Unshared pairs determine molecular geometry Even though all three molecules have 4 electron clouds (electronic configuration), the molecular geometry is different The lone pair of electrons repel bonded pairs (and hence become closer) as lone pairs are more “spread out” Bond angle (degrees) 109.5 <109.5 (slightly less) 105 6. Atomic orbitals Sigma and Pi bonds Single bonds are sigma bonds, which means their atomic orbitals overlap during the covalent bonding process (strong) The remaining bonds, are pi bonds because of unhybridized p orbitals above and below the sigma bond NOTE: one pi bond and consists of two lobes (one lobe above and one below the sigma bond) 6. Atomic orbitals Sigma and Pi bonds C3H4 (Propyne) The two triple bonded Carbons have 1 sigma bond and 2 pi bonds Total of 6 sigma bonds and 2 pi bonds 6. Atomic orbitals PRACTICE: Determine sigma and pi bonds ? How many sigma bonds and pi bonds are there? 6. Atomic orbitals PRACTICE: Determine sigma and pi bonds ? 3 sigma bonds and 1 pi bond 6. Atomic orbitals Hybridization of Oribitals for bonding Orbitals mix into new hybrid orbitals (with different energies, shapes, etc., than component orbitals), to allow electrons to pair and form bonds (sigma bonds) Example: carbon atom which forms four single bonds the valence-shell s orbital combines with three valenceshell p orbitals to form four equivalent sp3 mixtures which form tetrahedral Remaining bonds, are pi bonds because of unhybridized p orbitals above and below the sigma bond 6. Atomic orbitals Hybridization of Oribitals for bonding: oribital table and geometries Hybrid orbitals are based on number of atomic orbitals that are mixed Reminder: Hybrid orbitals are for sigma bonds and unshared pairs; extra bonds are pi bonds 6. Atomic orbitals PRACTICE: Hybridization of Oribitals Question: What is the hybridization of nitrogen in following molecules?: NO2- NH3 N2 6. Atomic orbitals PRACTICE: Hybridization of Oribitals Question: What is the hybridization of nitrogen in following molecules?: NO2- NH3 sp2 (one unshared pair, 1 single bond, 1 of the electron pairs in the double bond) N2 sp3 (one unshared pair, 3 single bonds) sp (one unshared pair, , 1 of the electron pairs in the double bond) Electron Configuration and Periodicity Table of Contents Section Topics 1. Quantum model of electrons Quantum Numbers Orbitals (s, p, d, f) Periodic element “blocks” 2. Electron configurations Aufbau Process Hunds Rule Gas Core configuration 3. Isoelectric species Isoelectric species examples 4. Periodicity Atomic size Ionization energy (first and subsequent) 46 1. Quantum model of electrons Electron behavior is more complicated than we think Bohr model (Classical Mechanics) Electrons are “particles” orbit the nucleus Electron Cloud Model (Quantum Mechanics) Electrons are “waves” that exist in certain “cloud” or “orbital” patterns 1. Quantum model of electrons Four quantum numbers describe electron behavior 1 2 4 MAIN energy level of the electrons – higher n means electrons have more energy; eletcrons are further away from the nucleus The ℓ value (Azimuthal quantum number) determines the SHAPE of the orbital (e.g., s, p, d, f) 3 “SPIN” of the electrons – clockwise or counter clockwise Each ℓ level has sub-orbitals mℓ which determines the direction or ORIENTATION in space of the electron cloud (e.g., 1, 3, 5, 7 sub orbitals) 1. Quantum model of electrons The s Orbital (ℓ quantum number of 0) ` ` ` ` The ℓ value or Azimuthal number of 0 means it is an s orbital and its SHAPE is a “sphere” ` ` ` ` As quantum number (n) increases from 1 to 3, the energy level goes up – electrons further away from the nucleus 1. Quantum model of electrons The p Orbital (ℓ quantum number of 1) Quantum number ℓ of 1 means p orbitals and its SHAPE is a “bow tie” mℓ quantum number of -1, 0, +1 means three sub-orbitals with its ORIENTATION along three axes Each sub-orbital holds 2 electrons (6 total) 1. Quantum model of electrons The d Orbital (ℓ quantum number of 2) Quantum number ℓ of 2 means d orbitals its SHAPE is a “clover” mℓ quantum number of -2, -1, 0, +1, +2 means five suborbitals with its ORIENTATION as shown here Five sub-orbitals each holding 2 electrons (10 total) 1. Quantum model of electrons f Orbital (ℓ quantum number of 3) with seven sub-orbitals Quantum number ℓ of 3 means d orbitals its SHAPE is a “tetrahedral” mℓ quantum number of -3, 2, -1, 0, +1, +2 , +3 seven sub-orbitals with its ORIENTATION as shown Seven sub-orbitals each holding 2 electrons (14 total) High energy – usually unstable 1. Quantum model of electrons Periodic table “blocks” guide which valence electron orbitals are important for reactions 2. Electron configurations Aufbau Process to determine electron configuration Aufbau Process helps determine the electron configuration of an atom keeping electrons at their lowest possible energy (close as possible to the nucleus) Algorithm for allocating electrons: ‒ Determine number of electrons present (atomic number) ‒ Fill lowest energy orbitals first and work way up as detailed in exhibit) 2. Electron configurations Hund’s rule of maximum municipality Nitrogen example – Atomic number (7) Electrons are placed into each orbital before any pairing takes place (e.g., three 2p electrons go into each sub-orbital versus pairing) Definitions for solid forms: ‒ Paramagnetic: created when unpaired electrons are present and attracted to magnetic field ‒ Diamagnetic: occurs when all electrons are paired and slightly repelled by magnetic field 2. Electron configurations PRACTICE: problems on Orbital Diagrams Hint: Make sure you know the number of electrons you have based on periodic table 2. Electron configurations PRACTICE: problems on Orbital Diagrams 2. Electron configurations Writing gas core diagrams Noble gases have full octets and therefore can be used as reference points to show orbital diagrams for subsequent elements – avoids writing long orbital notations Example: Na will has 1 more electron than Ne, hence you can write as [Ne]3s1; the one extra electron goes into next orbital 2. Electron configurations PRACTICE: Gas core diagram for Ca Draw the gas core diagram for calcium (hint: what is closest noble gas) 2. Electron configurations PRACTICE: Gas core diagram for Ca 3. Isoelectric species Isoelectronic species Examples of Isoeletronic species Isoelectronic species have same electronic number and configuration 4. Periodicity Periodicity Key Characteristic Atomic radius Periodicity refers to trends or recurring variations in element properties with their placement on the periodic table (e.g., atomic number, group, period) Ionic radius (not covered in exhibit) Ionization energy Description The size of the atom – the distance between center of atoms of an element Smaller the size, the closer and more stable are the electrons (closer to the nucleus) The size of the radius of the ions Cations have smaller radii (less electrons, so remaining are closer to nucelli) Anions have larger radii since more electrons avoiding each other (further from nucleus) Energy needed to unlock electron from gaseous atom The closer the electrons to the nuclei (lower energy states), the more energy is needed to dislodge 7/31/23 Electronegativity (covered in previous section) Measure of ability of atom to attract (pull in) electrons during covalent bodning 62 4. Periodicity Atomic size: Electrons closer to the nucleus (stronger pull) have smaller atomic size and hence are more stable The n value increases as you move down each period, and hence electrons are further away from the nucleus (higher atomic radius) 4. Periodicity Ionization energy Amount of energy needed to lose the most loosely bound e- of an atom in its gaseous state “Zig-zag” effect: as you move further right (to noble gas) higher ionization because valence shells more filled, move down less energy as electrons further out 4. Periodicity Successive ionization energies are MUCH higher Example of Aluminum Ionization Energy (Kj/mol) The energy needed to remove subsequent electrons (closer to the nucleus) jumps a LOT Note that the Y axis is logarithmic which means that to remove the 12th electron requires almost 350X the energy as the first electron ~350X
