CHEM 111 - Study guide #2
CHEM 111 - Study guide #2 CHEM 111
Popular in Chemistry
Popular in Chemistry
This 12 page Study Guide was uploaded by Selena Blanco on Monday February 22, 2016. The Study Guide belongs to CHEM 111 at New Mexico State University taught by Ramesh Chinnasamy in Spring 2016. Since its upload, it has received 45 views. For similar materials see Chemistry in Chemistry at New Mexico State University.
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Date Created: 02/22/16
March 22, 2016 Exam Study Guide: Ch. 3-4 •surfaces when illuminatedctrons are ejected from charged metal •Wavelength and Frequency are inversely related (reciprocal) • Energy and wavelength are inversely related(reciprocal) • Energy and frequency are directly related • Wavelength (): • Distance from crest to crest or trough to trough, unit: nm (nano meter) • Frequency (): • The number of times a wave passes a point per unit of time -1 (per second, -1), also used in Hertz • (1 Hz = 1 s ) • Amplitude: • The height of the crest or depth of the trough • Black body Radiation: Absorbs light when cold and emit when it is hot.(depends on temperature) • Radiant energy is “quantized” • Having values restricted to whole-number multiples of a specific base value • Quantum = smallest discrete quantity of energy • Photon = a quantum of electromagnetic radiation(tiny packet) • Quantized states: Discrete energy levels (e.g., steps) • Continuum states: Smooth transition between levels (e.g., ramp) • Photoelectric Effect: Albert Einstein • Phenomenon of light striking a metal surface and producing an electric current (flow of electrons) • Work function () – amount of energy needed to dislodge an electron from the surface of a metal • Bohr theory: Based on Rutherford model of an atom(Recall) and Max Planck's theory of quantized energy • Electrons are revolving around the nucleus in one of orbits • Most negative energy of electron and the addition of energy to dislodge an electron or breaking the attraction between the nucleus and an electron. Energy added should be more than the attraction between electrons and nucleus of an atom. • Electronic states • Energy Level: An allowed state that an electron can occupy in an atom • Ground State: Lowest energy level available to an electron in an atom (n = 1); • Excited State: Any energy state above the ground state, n= 2, 3, 4, 5 etc • Electron Transition: Movement of an electron between energy levels • Strengths: Accurately predicts energy needed to remove an electron from an atom (ionization) • Limitations: Does not account for spectra of multi electron atoms or limited to hydrogen atom • De Broglie (1892–1987): Another evidence for stability of electrons orbiting the nuclei • If electromagnetic radiation behaves as a particle, could a particle in motion, such as an electron, behave as a wave? • Assumed that electron could behave as waves of matter and particles of matter. • Heisenberg uncertainty principle • The principle that one can not simultaneously know the exact position and momentum of an electron • Erwin Schrödinger (1925) • Developed mathematical equations to describe behavior of electron waves; became the basis of quantum mechanics or wave mechanics(provides the probability of finding an electron at any point in the space around nucleus of atom) • Quantum numbers: energy, size, shape, orientation • of orbitals, spin orientation of electrons • Mathematical solutions to wave equations (y ) identified by unique combination of 3 integers called: Quantum Numbers: • Principle quantum number (n); indicates shell and relative size of orbital(s). n = 1, 2, 3,… Like Bohr’s n value • Orbitals with same value of n is in the same shell, energy increases with increasing value of n • Angular momentum quantum number (ℓ); defines shape of orbital and subshell. ℓ = 0 →(n – 1), Orbitals with same values of n and l are in same subshell, have same energy • l= 0(s), 1(p), 2(d), 3(f), 4(g) • Magnetic quantum number (m); defines orientation of orbital l around nucleus. m = –ℓl→ + ℓ • Shell: Defined by angular momentum quantum number (l) and magnetic quantum number(m) l • Shell consists of orbitals with same principle quantum number(n) • Atomic orbital: Provides the probability of finding an electron in the space around the nucleus of atom • An orbital with in the same subshell have same principal and angular momentum quantum numbers. • principal quantum number value be never zero • s subshell:1 orbital, p subshell: 3 orbitals, d subshell:5 orbitals • f subshell: 7 orbitals g subshell: 9 orbitals • Pauli Exclusion Principle: No two electrons in an atom may have the same set of four quantum numbers • Total number of orbitals =16 • Radial Distribution Plot: • A graphical representation of the probability of finding an electron in a thin spherical layer near the nucleus of an atom • Aufbau Principle: • Method of building electron configurations by adding one electron at a time as atomic number increases • Electron configuration: The distribution of electrons among the orbitals in atom or ion • When adding electrons to an atom: Rules • Electrons always go first to lowest energy orbitals available • Maximum of two electrons per orbital • Orbital Diagram: Depiction of arrangement of electrons in an atom or ion using boxes to represent orbita-s • Up/down arrows to indicate e with +/ – spin • Hund’s Rule: • Degenerate orbitals – orbitals of the same energy (e.g., 2p) • Degenerate Orbitals: orbitals sharing the same principal and angular momentum quantum numbers • The lowest energy configuration maximizes the number of unpaired electrons, all of which has same spin, in degenerate orbitals • Core electron: Inner shell electrons, not involved in chemical reaction • Valence Electron: Outer shell electrons, involved in a chemical reaction • Valence shell: Outermost occupied shell of atom • Effective nuclear Charge: The attraction between the electron and nucleus, the positive charge reduced by the extent to which other electrons shield electrons from the nucleus • Chromium, copper do not follow pattern: 2 2 6 2 6 5 1 • 24Cr = 1s 22 2p 2s 3p63d 2s 6 10 1 • 29Cu = 1s 2s 2p 3s 3p 3d 4s • Anomalies arise from stability associated with half-filled and completely filled d-subshells • • Formation of Ions: • Gain/loss of valence electrons to achieve stable electron configuration (filled shell) • Cations: + - • Na(g) → Na (g) + e • [He]3s → [Ne] + e - • Anions: - - • Cl(g) + e → Cl (g) • [Ne]3s 3p + e → [Ne]3s 3p = [Ar] 2 6 • Main Group Elements: • Form ions by gain/loss of e to obtain noble gas configuration: 2+ - • Mg → Mg + 2e = [Ne] • O + 2e → O - 2- = [Ne] • Isoelectronic: Describes atoms/ions having identical electron configurations • Na , Mg , O ,F , Ne = 1s 2s 2p 2 2 6 (= [Ne]) + - 2+ 2 2 6 2 6 • K , Cl , Ca , Ar = 1s 2s 2p 3s 3p (= [Ar]) • Atomic Radius (i.e., covalent radius): Half the distance between identical nuclear centers in a molecule • Metallic Radius: half the distance between nuclear centers in the crystal of a metal • Ionic Radius: Derived from the distance between nuclear centers in ionic crystals • Atomic Radii: • Increase going down a family“Shielding” by inner shell electrons decreases effective nuclear charge (z ) eff • Decrease going across a row • Increased nuclear charge (z) moving across row • Increased attraction for electrons in inner orbitals → atomic size decreases • Recall: Atomic emission spectrum, theories about structure of atom and discrete energy levels in the atom • Ionization Energy: - • Amount of energy needed to remove 1 mole of e from 1 mole of the ground-state atoms or ions in the gas phase (kJ/mole) + - • X(st → X (g) + e (g) + - • 1 Ionization Energy (IE ): Mg1g) → Mg (g) + e = 738 kJ/mol • 2 ndIonization Energy (IE ): M2 (g) → Mg (g) + e =1451 kJ/mol • Note: IE > 2E 1 2+ • Total energy required to make 1 mole of Mg = 2189 kJ/mole • Electron Affinity: • Energy change that occurs when 1 mole of electrons combine with 1 mole of atoms or ions in the gas phase • Cl(g) + e → Cl (g) - EA =1–349 kJ/mol CHAPTER 4 • Types of Chemical Bonds: • Ionic bond: Chemical bond resulting from the electrostatic attraction of a cation for an anion (metals + nonmetals, transition metals + non metals), by transfer of electrons. • Energy associated is electrostatic potential energy: E el ∞ Q x1Q 2 /d • Replacing ∞ by proportionality constant E =2.31x10 el -19J.nm Q 1Q /2 • The net charge ionic compound is zero • Covalent bond: Chemical bond that results from a sharing of outermost electrons (non-metals + non metals or metalloids), sharing of electrons involved, mutual attraction between nucleus & electron • Metallic bond: Chemical bond consisting of the nuclei of metal atoms surrounded by a “sea” of shared electrons (Metals, pool of electron • Coulomb’s law • Electrostatic potential energy: The energy of a charged particle based on its position relative to another charged particle; it is directly proportional to the product of the particles and inversely proportional to the distance between them: Coulombic attraction • what is lattice energy? • The energy released when free gas phase ions combine to form one mole of crystalline solid. • Crystal Lattice: Three dimensional arrangement of atoms, ions, or molecules in a crystalline solid • Binary ionic compounds consist of cations (usually metals) and anions (usually nonmetals), e.g., MgCl , LiCl, 2a S, Li O 2 2 • Cation named first using name of element, Mg = magnesium • Anion named by adding the -ide suffix to the name of the element, Cl = chlorine → chloride • For metals that form cations with different charges, a Roman numeral is added to indicate the charge of the cation. • Polyatomic Ions • Charged group of two or more atoms joined together by covalent bonds Ex: NH 4 cation is most common • Oxoanions • Polyatomic anions containing oxygen in combination with one or more other elements, - - 2- • Examples: acetate (C H O )2 n3tr2te (NO ), carbona3e (CO ), 3 perchlorate (ClO ), 4ulfate (SO ) and 4ulfite (SO ) (See Tabl3- 4.3) • • Binary Acids: Explain ionizable hydrogen atoms • Contain hydrogen and a monoatomic anion (e.g., Cl , S ) - 2- • Most common binary acids are halogen (e.g., HI, HCl, HBr) • Hydrogen chloride, hydrogen bromide, hydrogen iodide • If the these compounds dissolves in water, solution is called as • HI hydro iodic acid, and HCl Hydrochloric acid • Acid names: the prefix “hydro” + the halogen base name + the suffix “ic” + the word acid. • • Gilbert Lewis (1916) • Proposed that atoms form chemical bonds by sharing electrons to acquire electron configuration of a noble gas • Octet Rule : Atoms tend to lose, gain, or share electrons to obtain a set of 8 valence electrons (ns np )2 6 • Lewis Symbols (Dot Structures) • Chemical symbol for an atom surrounded by one or more dots representing valence electrons • A Nonmetal and a Metal: Ionic compounds • Metals lose valence electrons to achieve noble gas electron configuration. • Nonmetal gain electrons to achieve noble gas electron configuration. • Lewis Structures: Molecular Compounds: non-metal +non metal • Pair of electrons shared between two atoms in a covalent bond • Single bond: Two atoms sharing one pair e . Ex. H:H, or H–H (duet), stable like He - • Lone pair: Pair of e that is not shared. Or (lone pairs) • Shared electrons vs unshared pair of electrons Ex H O 2 2 • Double Bond: Two electron pairs shared by two atoms • Represented as a double line ( = ). • Example: O 2 → • Triple Bond: Three electron pairs shared by two atoms • Electronegativity (EN): ability of an atom to attract electrons in a bond to itself. • EN increases from left to right and decreases from top to bottom in the periodic table. • Polar Covalent Bond: Unequal sharing of bonding pair of electrons between atoms, Results in uneven distribution of charge • Bond Polarity: A measure of the extent to which bonding electrons are unequally shared due to differences in electronegativity of the bonded atoms • Polar Bond: Partial negative and positive charges (δ+, δ–) • Direction of polarity indicated by arrow pointing to more negative end of bond, with + sign at positive end • Electronegativity Difference is • Less than 0.4: Non-Polar • Between 0.4-2.0: Polar • Greater than 2.0: ionic • Absorption of energy by atoms: • Absorb radiation if frequency matches ∆E of electron energy levels (Chapter 7) • Absorption of energy by molecules: • Absorb infrared radiation if frequency matches vibrational modes of molecular bonds = infrared active Basis of “greenhouse effect” • Allotropes: Different molecular forms of the same element • Having different physical and chemical properties • Resonance: • When two or more equivalent Lewis structures can be drawn for one compound • Resonance Structures: • Two or more Lewis structures with the same arrangement of atoms but different arrangement of bonding pairs of electrons: • Formal Charge: - • Determined by the difference between the number of valence e - in the free atom and the sum of lone pair + ½ bonding e in a molecule • Formal charge (FC): • FC = (# valence electrons) – - - • [ (# unshared e ) + 1/2(# of e in bonding pairs)] • • Free Radical: Odd-electron molecule with an unpaired e in its - Lewis structure. Very reactive! • Two important things about Expanded Shells • 1. Atoms tend form molecules with stronger electronegative elements (F, O and Cl) • Ex: SO ,42- PO 43- (Home work, try to draw the Lewis structure) • 2. Formation of molecules resulting in smaller formal charges on the atoms: FC for sulfur = 0; FC for fluorine = 0 • Bond Length Depends On: Identity of the atoms, and # of bonds between them • Bond Order: The # of bonds between two atoms • 1 for a single bond • 2 for a double bond • 3 for a triple bond • Bond Energy or Bond strength • The amount of energy required to break the bond between the atoms • Bond energies always positive, energy required to break bond • Bond order and Bond strength or bond energy is directly related • C-C (348 kJ) C=C (614 kJ) C-C triple bond (839 kJ) • Energy needed to break 1 mole of covalent bonds in the gas phase • Breaking bonds consumes energy (+); forming bonds releases energy (-) • Using bond energies to estimate IONIC COMPOUNDS -The metal will always be the first element in the name. The non- metal will always be the last element in the name COVALENT COMPOUNDS -Will only contain non-metals Fluorine (the most electronegative element)
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