REVIEW CHAPTERS 5,6, 7,8-(Only Lewis structure section 8.1-8.3)

1. REVIEW CHAPTERS 5,6, 7,8-(Only Lewis structure section 8.1-8.3) Chapter 5 Thermochemistry Energy Thermochemistry: Some Basic Terms - Open system Energy and mass exchange - Closed system Only energy wxcahge - Isolated No exchange In an exothermic process the system gives off heat to the surroundings Q (system) = -ve In an endothermic process the system absorbs heat from the surroundings (heat enters the system Q (system) = +ve W = -PDV Work done on the system W = +ve, volume decreases Work done by the system Expansion W = -ve, volume increases Internal Energy (U), State Functions, and the First Law of Thermodynamics Keep track of sign of q and w Increase in internal energy DU +ve Decrease in internal energy DU -ve Heats of Reaction and Enthalpy Change,DH –State Function and Extensive Property (depends on mass, moles etc) -Heat exchange in chemical reaction under constant T and P Exothermic DH = -ve ; Endothermic = DH =+ve

2. Calorimetry: Measuring Quantities of Heat -IMPORTANT Read definitions of heat capacity (cal/ºC) and Specific heat (cal/g ºC) Specific of heat of water = 1cal/g ºC Heat absorbed = q = Calorimetry problems: Heat lost by hot metal or reaction = Heat gained by Calorimeter or water DT = -ve !!! DT = +ve !!! Heat absorbed by calorimeter = qcalor = +ve = heat capacity X DT • Heat changes in change of state • Melting, Vaporization, sublimation, condensation, deposition • Heat exchange in chemical reactions • Hess’s Law problems • Standard enthalpy and DH for reaction = • sum of DH for products - sum of DH for reactants • Remember DH formation at standard state for elements in natural form =0 • And for compounds we must form 1 mole of compounds using elements in natural state Formation of CaCO3 (s) is Ca(s) + C (gr,S) + 3/2 O2 CaCO3 (s)

3. 6.1-6.2 Photons: Energy by the Quantum Direct conclusion from photoelectric effect – Light is discrete packet of energy called photons QUANTIZED! Ephoton = h * n One should be able to compute the energy of various wavelengths of light, by using and . Constant Chapter 6 Atomic Structure • 6.1 The Wave Nature of Light –electromagnetic radiation • -characterized by Wavelength l (length),frequency n (hertz, or /s) • -Speed C = 3X108 m/s or C = l * n • or energy order (low to high) Radio, microwave, Infrared, Visible, UV, x-ray, gamma l order (low to high, is reverse) Gamma, x-ray, UV, Visible, infrared, Microowaves, radio

4. Chapter6 Atomic Structure 6.3-6.4 Bohr’s Hydrogen Atom: A Planetary Model • The energy of each stable orbit is given by En = –B/n2 • n is an integer, 1, 2, 3, 4, 5, … It is the quantum number for the atom. • The negative sign indicates that lower (more negative) energies are more stable, which is the usual convention. In the case of a nucleus and electron, the (arbitrary) zero of energy corresponds to an electron and a nucleus separated far from each other. • 2. Light is emitted or absorbed only when an electron moves between energy levels. • a. When the electron drops in energy, from a higher to a lower level, light is given off. The atom becomes more stable. The higher level has a larger quantum number; the lower energy level has a smaller quantum number. • b. When light is absorbed by the atom, the electron moves from a lower energy level to a higher one, from a level with a small quantum number to a level with a larger quantum number. The atom becomes less stable. • c. In addition, recall that light of high energy has a short wavelength. Light of a small energy has a long wavelength. • The electron in the lowest possible energy level is in the ground state. All other possibilities are excited states. • E = Efinal -Einitial

5. 6.5 Wave Mechanics: Matter as Waves A. In 1923 Louis de Broglie proposed that particles of matter have wavelengths determined by their mass and velocity. 1. The de Broglie expression is m X v = Momentum Kinetic Energy = • Heisenberg’s uncertainty principle • The uncertainty principle states that we cannot determine both position and momentum with perfect precision • In the Bohr view, the electron follows an exact pathway, known as an orbit. But in the Schrödinger view, the electron has a high probability of occupying a specific region of space, called an orbital.

6. 6.6-6.8 Quantum Numbers and Atomic Orbitals • - An atomic orbital defines a region in space in which there is a high probability of finding an electron. Atomic orbitals are identified by three quantum numbers. These quantum numbers have limitations on their values, the most general being that they are all integers. They also have definite interpretations. • - Quantum Numbers- a) Principal Quantum number “n” • The principal quantum number is designated by the symbol n (takes values 1, 2, 3, 4, 5….) • The value of n indicates the size of the orbital, or how far the electron in it is away from the nucleus. A large value of n designates a large orbital.n also indicates energy of the electron in the orbital • b).The orbital quantum number or angular momentum quantum number is designated by the symbol l. • The values of l are restricted to non-negative integers smaller than the value of n. • Thus takes values o to n-1 and l=0 s subshell, l =1 p subshell, l =2 d subshell and l=3 f subshell • l gives shapes of orbitals /subshells • c) The magnetic (or orientation) quantum number is designated by the symbol (or ml). • The values of are positive and negative integers, ranging from through 0 to 2. -l to +l • ml determines how the energy of the orbital will change when the atom is placed in a magnetic field. • Gives Orientation of these orbitals in space (X, y, Z axes etc.) • d) Spin Quantum number s • This quantum number indicates that the electron is spinning, either spin up with (and counterclockwise spin) or spin down with with clockwise spin. S = +1/2 or s= -1/2 • The allowed values for quantum numbers dictate the relationships among shells, subshells, and orbitals. [Read Examples] are f orbitals.

7. Chapter 7 • Polyelectronic Atoms • An Introduction to Electron Configurations • There are two notations in common use for describing electron configurations, spdf notation and orbital diagrams • 7.1-7.2 The Rules for Electron Configurations • Aufbau Principle –Order of filling • 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p • Follow maximum occupancy for s, p, d and f susbshells! • No 2 electrons can have all 4 quantum numbers same (Pauli’s exclusion principle), so electrons occupy orbitals spin up and spin down • Each orbital has 2 electron maximum • - Electrons in the same subshell occupy the orbitals singly, with spins parallel, before they are paired. (Hund’s rule) Paramagnetic Valence e- = 1 Valence config. 3s1

8. the electron configurations of Cr and Cu group have anamolies • A way to remember these anomalies is to assume that half-filled or filled subshells are more stable than other partly-filled subshells. • Cr config : Cu (Z=29) • Valence e- = 1, V config = 4s1 Valence e- = 1, V config = 4s1 • 7.3 Electron Configurations of IONS: Periodic Relationships • - Remove valance electrons first to make ions • Groups and Periods, Main group elements (group A) and others (group B) • The period number and the principal quantum number of the outermost shell are the same. • All elements of same group have same valence config. • eg. Group IA has ns1,group 7A halogens have ns2np5, 8A: ns2np6 • In main group group number = number of valance electrons • Core electrons are those residing in inner shells. • Magnetic Properties: Paired and Unpaired Electrons (seen from box diagram) • Any unpaired electron Atom is Paramagnetic • All paired electrons Atom is Diamagnetic

9. 7.4 Periodic Atomic Properties of the Elements A. Atomic size 1. Decreases as you move to the right along a row (PERIOD) 2. Increases as you go down a column *GROUP) B. Ionization energy (IE) 1. Increases as you move left to right along a Period 2. Decreases as you go down a group 3. The second ionization energy is greater than the first, the third greater than the second, etc. The electron is being removed from an increasingly positive species 4. In successive IE when you start removing core electrons, extreme high IE seen C. Metallic character (group 1A, 2A highly metallic) 1. Decreases as you move left to right along a Period 2. Increases as you go down the group D. Electron Affinity (energy released when electron is added) 1. Increases as you move left to right along a Period 2. Decreases as you go down a group E. Electronegativity (tendency to pull shared electron towards it) 1. Increases as you move left to right along a Period 2. Decreases as you go down a group

10. Chapter 8 Strategies for Writing Lewis Structures • A. The first problem to be solved when writing Lewis structures is where the atoms are. The location of the atoms is known as the skeletal structure. There are some guidelines that help one choose a plausible skeletal structure. • 1.Hydrogen atoms are rarely central atoms; they are almost always terminal atoms. • 2.The central atom usually is the atom with the lowest Electronegativity (except for H that often has a relatively low Electronegativity). • 3. In oxoacids, H usually is bonded to an oxygen atom.(HNO3, H2SO4 etc) • 4.Molecules usually are clusters of atoms, rather than long chains. • 5. As is evident from the wording of these guidelines, they are not absolute, but they can be helpful. • B. Lewis structures are easier to draw if one follows a systematic plan in constructing them. One such plan is as follows. • 1.Sum the valence electrons in the structure. • a.For each atom, the number of valence electrons equals that atom’s group number. • b. Add electrons if the entire species is an anion, subtract them if the species is a cation. • c. Take off electrons ( 2 electrons per bond) to form single bonds, • 2. Arrange the atoms in the preferred skeletal structure and place a pair of electrons between each pair of bonded atoms. • 3. Place remaining electron pairs around terminal atoms so that each has an octet. • 4. Place LEFT OVER electron pairs on the central atom. • Create multiple bonds as necessary to make sure each atom follows the octet rule. • 8.3Molecules that Don’t Follow the Octet Rule (SF6, IF5, XeF4 etc.) • Formal charge calculation, resonance

11. 0 EXAM 3- 100 POINTS 10 points Bonus question!! Part 1 Multiple Choice –Show calculations for partial/full credit 15 questions 4 POINTS EACH Part 2(40 points) 1) Hess Law and Standard enthalpy 2)Electron configuration, paramagnetism, valance e-, 3) Lewis structure and Count central atom electron groups Part 2 Bonus question: 10 points –HARD! Partial credit!