Lecture and Chapter 1 Notes
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Department
Chemistry
Course
CHM136H1
Professor
Thavarajah/ Woolley/ Nitz
Semester
Summer

Description
CHM138H1 Jasmyn Lee Chapter 1: Structure And Bonding 1.1 The Atomic Structure: The Nucleus Atomic Number and Atomic Mass  The atomic number (Z) is the number of protons in the atom’s nucleus  The mass number (A) is the number of protons plus neutrons  All the atoms of a given element have the same atomic number  Isotopes – atoms of the same element that have different numbers of neutrons and therefore different mass numbers  The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an elements naturally occurring isotopes 1.2 Atomic Structure: Orbitals  Quantum mechanics: describes electron energies and locations by a wave equation o Tries to describe the position of electrons in atoms o Shows most probable area of electrons in atoms Shapes of Atomic Orbitals for Electrons  Four different kinds of orbitals for electrons based on those derived for a hydrogen atom  Denoted s, p, d and f  s and p orbitals most important in organic and biological chemistry  s orbitals: spherical, nucleus at center  p orbitals: dumbbell-shaped, nucleus at middle  d orbitals elongated dumbbell-shaped, nucleus at center Orbitals and Shells  Orbitals are grouped in shells of increasing size and energy  Different shells contain different numbers and kinds of orbitals  Each orbital can be occupied by two electrons  First shell contains one s orbital, denoted 1s, holds only two electrons  Second shell contains one s orbital (2s) and three p orbitals (2p), eight electrons  Third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons Quantum Numbers 1. Principle (n) – main energy level 2. Angular (l) – shape and type of orbital; s=0, p=1, d=2, f=3 3. Magnetic (m) l electrons position in orbital 4. Spin (m s – spin of e- p Orbitals  In each sell there are three perpendicular p orbitaxs,yp , pzand p of equal energy  Lobes of a p orbital are separated by region of zero electron density, a node CHM138H1 Jasmyn Lee 1.3 Atomic Structure: Electron Configurations  Ground-state electron configuration (lowest energy arrangement) of an atom lists orbitals occupied by its electrons  Rules 1. Lowest-energy orbitals fill first: 1s  2s  2p  3s  3p  4s  3d (Aufbau “build up” principle) 2. Electrons act as if they were spinning around an acid  Electron spin can have only two orientation, up  and down   Only two electrons can occupy an orbital, and they must be a=of opposite spin (Pauli exclusion principle) to have unique wave equations 3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund’s rule) 1.4 Development of chemical Bonding Theory  Kekule and Couper independently observed that carbon always has four bonds  van’t Hoff and Le Bel proposed that the four bonds of carbon have specific spatial directions o Atoms surround carbon as corners of a tetrahedron  Atoms form bonds because the compounds that results is more stable than the separate atoms  Ionic bonds in salts form as a result of electron transfers  Organic compounds have covalent bonds from sharing electrons  Lewis Structures (electron dot) show valence electrons of an atom as dots o Hydrogen has one dot, representing its 1s electron o Carbon has four dots (2s 2p )  Kekule structures (line-bond structures) have a line drawn between two atoms indicating a 2 electron covalent bond  Stable molecule results at completed hell, octet (eight dots) for main-group atoms (two for hydrogen)  Atoms with one, two, or three valence electrons form one, two or three bonds  Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet  Carbon has four valence electrons (2s 2p ),rming four bonds (CH )4  Nitrogen has five valence electrons (2s 2p ) but forms only three bonds  Oxygen has six valence electrons (2s 2p ) but forms two bonds (H O) 2 Non-Bonding Electrons  Valence electrons not used in bonding are called nonbonding electrons or lone-pair electrons o Nitrogen atom in ammonia (NH ) 3  Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair CHM138H1 Jasmyn Lee 1.5 The Nature of Chemical Bonds: Valence Bond Theory  Covalent bonds forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom  Two models to describe covalent bonding Valence Bond Theory  Electrons are paired in the overlapping orbitals and are attracted to nuclei to both atoms o H-H bonds results from the overlap of two singly occupied hydrogen 1s orbitals o H-H bonds is cylindrically symmetrical, sigma (σ) bond  Valence Bond Theory o How closely two atoms should approach each other in order to form a bond Bond Energy  Reaction 2H  H 2eleased 436kj/mol  Product has 436kj/mol less energy than two atoms: H-H has bond strength of 436kj/mol o I kJ = 0.2390 kcal; 1 kcal = 4.184 kJ Bond Length  Distance between nuclei that leads to maximum stability  If too close, they repel because both are positively charged  If too far apart, bonding is weak 3 1.6 sp Orbitals and the Structure of Methane  Carbon has 4 valence electrons (2s 2p ) 
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