Chapter 1 Atomic Structure and Bonding PDF

Chapter 1 Atomic structure and bonding Jose Cobo, Ph.D  Principles of atomic bonding Nucleus: protons (+) and neutrons Atomic number: Number of protons in the nucleus Atomic mass: Number of protons + number of neutrons – Usually the number of protons = nu...

Chapter 1 Atomic structure and bonding Jose Cobo, Ph.D  Principles of atomic bonding Nucleus: protons (+) and neutrons Atomic number: Number of protons in the nucleus Atomic mass: Number of protons + number of neutrons – Usually the number of protons = number of neutrons. Isotope: same number of protons but different number of neutrons – Different atomic mass Electrons: found in orbitals Mass= protons + neutrons Atomic number  Electron shell and orbitals n=2 Electrons: are found in orbitals n=1 Orbitals: Energy states where an electron can be found (2 electrons per orbital) Shells: Orbitals are grouped into different shells at different distances from the nucleus. Lowest-energy shells are closest to the nucleus. Each shell is identified by a principal quantum number (n) – n = 1 lowest energy → closes to the nucleus – As (n) number increases, the shells are farther from the nucleus Higher in energy, More orbitals More electrons  Most of the common elements in organic compounds are found in the first two rows of the periodic table – Their electrons are found in the first two electron shells. The first electron shell (n = 1) – One orbital (s) → 1s n=2 – Two electrons Second electron shell (n=2) – Four orbitals: n=1 2s 2px, 2py, 2pz degenerate orbitals (same energy) – 8 electrons  Electronic Configurations Pauli exclusion principle: each orbital can only hold maximum 2 electrons Aufbau principle: Place electrons in lowest energy orbital first. Hund’s rule: Equal energy orbitals are half-filled, then filled. Electrons = 6 Electronic Configuration: n=2 1s22s22px12py1 n=1  Atomic Electronic configurations Pauli exclusion principle Aufbau principle Hund’s rule n=2 n=1  Valence electrons Valence electrons: electrons in the outermost shell n=2 Valance electrons = 4 n=1 Column number = number of valance electrons  Bond formation Octet Rule: An atom with a filled shell is more stable. – Noble gas configuration n=2 n=1 Bond Formation: atoms transfer or share electrons in such a way as to attain a filled shell of electrons.  Ionic bonding Transfer of electrons. – One atom looses an electron and one gains it, to satisfy the octet rule. – Common in inorganic compounds. – Very polar bonds  Covalent bonding Electrons are shared – Two atom share an electron to satisfy the octet rule ( fill outer shell) – Common in organic compounds. Noble gas configuration  Lewis Structures To represent covalent bonding in molecules Each valence electron is dot 4 valance e- A bonding pair is represented as a dash or pair of dots – Carbon contributes 4 valence electrons – Each hydrogen contributes 1 valance electron Carbon: gets a total of 8 electrons ( noble gas configuration) Each hydrogen: gets a total of 2 electrons ( noble gas configuration) Lewis structures examples  Lewis structures Lone Pairs Non-bonding electrons Valence electrons NOT shared Usually reactive sites Lewis structures must always show lone pairs  Multiple Bonds Single bond – One pair of electrons is shared between two atoms Double bond – Two pairs of electrons are shared between two atoms Triple bond – Three pairs of electrons are shared between two atoms  Common Bonding Patterns The number of bonds an atom usually forms to obtain noble gas configuration is its valence  Electronegativity and bond polarity Nonpolar covalent bond: – Electrons shared equally between the two atoms Two atoms with similar electronegativity Polar covalent bond: – Electrons shared unequally between the two atoms Electrons are attracted more strongly to the more electronegative atom Electronegativity is used to predicts bond polarity Dipole moment (μ): Measurement of bond polarity. – Partial charge ( ) X bond length (d)  Electronegativity and the periodic table Electronegativity: is a measure of an atom's ability to attract electrons to itself. Electronegativities increase from left to right across the periodic table. Electronegativities increase from bottom to top in the periodic table. Electronegativity Electronegativity Predict the direction of the dipole moments  Formal charge Formal charge: difference between the number of valence electrons of each atom and the number of electrons the atom is associated with Formal charge  Ionic bonds in organic compounds Organic compounds carrying a formal charge can form ionic bonds with charged ions – Methylammonium chloride (CH3NH3Cl) N has a +1 formal charge, and it can form an ionic bond to Cl ion (-1 charge)  Resonance Charged organic compounds can have electron delocalization to achieve a more stable hybrid molecule The two resonance forms do not exist individually but rather exits as a single resonance hybrid resonance hybrid  Major and Minor Resonance Contributors Not all resonance forms always contribute equally to the main structure (hybrid) – The more stable from contributes more to the structure of the hybrid The actual hybrid structure more closely resembles the more stable form nitromethane C follows octet rule C does not follow octet rule N follows octet rule N follows octet rule Number of C-N bonds formed = 2 Number of C-N bonds formed = 1  Major and Minor Resonance Contributors formaldehyde resonance hybrid The more stable form has: – More or all atoms following the octet rule – More number of bonds – Less or no charge separation  Rules to resonance structures All the resonance forms must be valid Lewis structures for the compound. – Second-row elements (B, C, N, O, F) can never have more than eight electrons in their valence shells. Only electrons (in pairs) may be shifted from one structure to another – Usually non-bonding electrons and electrons in double bonds. Nuclei (for example H) cannot be moved All bond angles must remain the same Sigma bonds (single bonds) are very stable, and they are rarely involved in resonance The major resonance contributor is the one with the lowest energy  Types of representation of organic compounds Lewis structures: – Symbolizes a bonding pair of electrons as a pair of dots or as a dash (-). – Lone pairs of electrons are shown as pairs of dots. Condensed structure: – Written without showing all the individual bonds – Each central atom is shown together with the atoms that are bonded to it. – Double or triple bonds may or may not be drawn  Types of representation of organic compounds Line–angle formula (stick figure structure): – Bonds are represented by lines – Carbon atoms are assumed to be present wherever two lines meet – Hydrogen atoms are not shown – Nitrogen, oxygen, and halogen atoms are shown  Atomic orbitals Electrons behave as waves rather than as particles Atomic orbitals: – Wave function: mathematical description of location of the electron S orbitals: Standing wave function Two phases Spherical shape P orbitals: Harmonic wave function Two phases – “Lobes” Opposite phases separated by a node – The nucleus separates the “lobes”  Molecular Orbitals (MOs) When orbitals on different atoms interact, they produce molecular orbitals (MOs) that lead to bonding Sigma Bonding: Sigma bonding of (S) orbitals: – When two electrons in the same phase of different (s) orbital approach each other, their s wave functions can add constructively Creates a sigma bond between two atoms For example Hydrogen-Hydrogen bond +  Molecular Orbitals (MOs) Sigma bonding of (p) orbitals: – When two p orbitals in the same phase overlap along the same axis, a bonding molecular orbital(MO) is formed – Creates a sigma bond between two atoms. For example, a F-F bond Sigma bonding of an (s) orbital with a (p) orbital: – When an (s) and a (p) orbital in the same phase overlap along the same axis, a bonding molecular orbital(MO) is formed – For example, a C-H bond +  Molecular Orbitals (MOs) Pi Bonding: Results from overlap between two p orbitals oriented perpendicular to the line connecting the nuclei. Usually involved in double and triple bonds – Double bonds require sharing two pairs of electrons (4 e) – First pair makes a sigma bond – Second pair makes a Pi bond – For example C=C double bond  Multiple Bonds A double bond (2 pairs of shared electrons) consists of a sigma bond and a pi bond. A triple bond (3 pairs of shared electrons) consists of a sigma bond and two pi bonds. 32  Antibonding MO When two orbitals of opposite phase overlap – Destructive addition – Higher energy (less stable) Sigma antibonding: – s – s orbitals – p – p ortitals – s – p orbitals Pi antibonding:  Hybrid atomic orbitals Overlapping of two orbitals in the same atom sp hybrid orbital: – Overlapping of the s and p orbitals in the same atom The orbitals in the same phase add constructively – 2 hybrid sp orbitals are formed – It can make bonds to two atoms Linear geometry (180 degrees) + Example: BeH2 :  sp2 Hybrid Orbitals: – Overlapping of one s orbital and 2 p orbitals in the same atom The orbitals in the same phase add constructively Creates three sp2 orbitals It can make bonds to three atoms Trigonal geometry (120 degrees) Example: Borane (BH3)  sp3 Hybrid Orbitals: – Overlapping of one s orbital and 3 p orbitals in the same atom The orbitals in the same phase add constructively Creates four sp3 orbitals It can make bonds to four atoms Tetrahedral geometry (109.5 degrees) Example: Methane (CH4) sp3 Hybrid Orbitals: Sp Hybrid Orbitals  Angles and bonding Carbons: sp3 sp2 sp Not 90° degrees because electrons repel each other!!! They try to be as far from each other as possible Lone pairs take more space than bonding angles, compressing angles Double and triple bonds are not made of hybridized orbitals Only sigma and lone pairs occupy hybridized orbitals  Lone pairs hybridization state: – A lone pair can occupy hybrid orbitals, just like sigma bonds – For example NH3. Three sigma bonds One lone pair of electrons Four orbitals – Sp3 hybridization – Tetrahedral geometry What’s the hybridization state? a b c d sp2 - sp2 Sp-sp2 a) Tetrahedral geometry (109.5) = sp2 c) Linear geometry (180) = sp b) Tetrahedral geometry (109.5) = sp2 d) Tetrahedral geometry (109.5) = sp2  3D Drawing Straight lines: indicates bonds in the plane of the page Heavy wedge-shaped lines: indicates bonds that come forward, toward the reader Dashed lines: indicate bonds that go backward, away from the reader  Bond rotation Single bond freely rotate creating different conformations. – Sigma bond is flexible, can rotate! Double bonds cannot rotate unless bond is broken – Sigma bond unaffected – Pi bond looses its overlap – Cis and trans conformations  Isomerism Isomers: molecules with the same molecular formula but differ in the arrangement of their atoms Structural isomers (Constitutional isomers) : differ in bonding sequence – DIFFERENT CONNECTIVITY C4H10 Stereoisomers: differ in atom spatial orientation – SAME CONNECTIVITY C4H8  Structural Isomers CH3 O CH3 and CH3 CH2 OH CH3 and CH3 => Chapter 2 45  Stereoisomers Cis-trans isomers are also called geometric isomers. There must be two different groups on the sp2 carbon. Br Br Br CH3 C C and C C H3C CH3 H3C Br Cis - same side Trans - across 46  Isomerism Not all double bonds show cis/ trans isomerism There must be two different groups on the sp2 carbon  Homework Problem Set Send me an email with 2 problems you would like to go through over recitation