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Chem 101 Exam 1 Notes & Study Guide

by: Lauren Notetaker

Chem 101 Exam 1 Notes & Study Guide CHEM

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Lauren Notetaker
Texas A&M

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These notes cover material on the first exam (review chapter, chapter 1, chapter 2, and chapter 3).
General Chemistry (101)
Dr. Collins
Class Notes
Chem:, Stochiometry, Bonding, lewis, structures, nomenclature
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Date Created: 10/15/16
Chemistry Review1: Measurementsand Calculationsin Chemistry R-1Units of Measurement  Measurement is a quantitative observation constructed intwo parts: a number and a unit  Measurements are usedfor mass,length, time, temp., electric current, amount, etc.. o All measurements are standardized  SI System is basedon the metric system o Volume is not measured in the SI System which is why we used units becauseit is derived from length  Mass vs.Weight o Terms are similarbut NOTinterchangeable o Mass is a measure of resistanceofan object to a changein its stateof motion o Weight is a measure of the force of gravity on anobject –varies with the strength ofa gravitational field R-2Uncertainty inMeasurement  When recording measurements, numbers are often estimated after a certain point o Ex.While reading a buret, you measure 20.15 the, 5 is estimatedand can vary depending on who takes the measurement o Uncertain digits areestimated and therefore vary, usuallyall certain digits are recorded plus an additional uncertain digitfor further accuracy  A measurement always has a degree ofuncertainty – depends on the precision of the device  SignificantFigures of a measurement are all certain number and an additional uncertain number -communicates something about the uncertainty in a measurement o Amount of significantfigures needed to be recorded vary depending on the decimal the measuring devicereaches o Ex.A buret reads to the 0.01mL, the measurement must be written as 25.00mL or 25.70mL to account for the accuracy ofthe measurement  Precision and Accuracy o Accuracy is the agreement of a particular valuewith its true value o Precision is the degree of agreement among several measurements of the same quantity  Reflects reproducibility of a measurement  Indication of accuracy only ifsystematic errors are present (seefigure& explanation R-4 on page7) o Random error / indeterminate error means a measurement has an equal probability of being highor low  Occurs in estimating the lastdigitof a value o Systematic error / determinate error occurs in the samedirection each time; always highor always low R-3SignificantFigures and Calculations  Final results ofcalculations involve some type of calculation,uncertainty is important to the final result Rules for SignificantFigures  1. Non-zero integers always count as significantfigures 2. Zeros: a. Leading zeros occur before aninteger (0.003) do not count as significant figures b. Captivezeros occur between digits (90008) and ALWAYS count as significant figures c. Trailing zeros follow an integer number (900) and ONLYCOUNT IF THEREIS A DECIMAL (900.0has 4 significantfigures,900only has 1) 3. Exactnumbers (4/3pir^2) do not influence the number of significantfigures when doing operations Calculations &SignificantFigures  1. MULTIPLICATION AND DIVISION –number ofsignificantfigures inthe end result is the number inthe “leastaccuratemeasurement”, or number with the leastsig figs. a. 4.56 x 1.4 =6.4 2. ADDITIONAND SUBTRACTION– number of significantfigures inthe end is determined by the “leastaccuratemeasurement” or the number with the leastsig figs AFTER the decimal point. a. 12.11 + 18.0 +1.013 = 31.1 3. ROUNDING a. Do not round until the end ofan equation *the final result is the only number that should be rounded* b. Round final digitas normal (<5 stays thesame, >5round up) *useonly the digitto the right of the lastsig fig* R-4Learning to Solve Problems Systematically  To solveproblems know: o Where you’re going o What you already know o What you need to do to getthere R-5Dimensional Analysis  Unit Factor Method/Dimensional Analysis converts a givenresult from one system of units to another  Dimensional Analysis does not changethe measurements value,only the units  See page12-15 for example problems R-6Temperature  Temp. is measured in Celsius,KelvinandFahrenheit  Temperature Conversions: o Temp. K =Temp. C + 273.15 o Temp. F =1.8(Temp. C)+32 *Degreesymbol is not used when measuring K* o See page17 and 18 for conversion examples R-7Density  Densityis the mass of a substanceper unit volume, unique for every element o D=M/V  Densityis used to identify substances and as measurement R-8ClassificationofMatter  Matter is anything occupying spaceand having mass o Complex, many levels oforganization  States of Matter: Solid, Liquid, and Gas o Solid – rigid,fixed volume, and shape o Liquid – definitevolume, no shape o Gas –no fixedvolume or shape,takes shape and volume of its container  Mixture – variablecomposition o Homogeneous / Solution – having visiblyindistinguishable parts o Heterogeneous –having visiblydistinguishable parts o Canbe separated into pure substances byphysical methods  Ex.Seawater is NaCl +H20 ->seperatable through boiling/freezing o Physical change– changein form of substance,not the chemical composition  Ex.Boiling / freezing water o Distillation–a liquid mixture is heated in a device, the most volatilecomponent vaporizes at the lowest temperature. Thetemperature passes through a cooled tube (condenser) where it condenses into its liquid state  depends on differences in the volatility of the components o Filtration –a mixture consisting ofa solidand a liquid, is poured onto mesh which passes theliquid but leaves the solid o Chromatography –general name applied to a series ofmethods that usea system with two phases (or states)ofmatter o Stationary phase –a solid o Mobile phase –a liquid or gas o Separation occurs because components with high afinity for the mobile phasemove quickly, whereas the component with an affinityfor the solid moves more slowly  Ex.Paper chromatography  Stationary phase: strip of porous paper /filter paper  Mobile phase: liquid  A drop of mixture is placedon the paper. It is then dipped into the liquid that travels up the paper likea wick, carrying the mobile afinitecomponent up the paper with it o Absolute purity of separated components is ideal  Compound – substancewith constant composition that canbe broken down into elements by a chemical process o Ex.Electrolysis –electric current is passedthrough water to break H2O molecules into Hydrogen and Oxygen  Chemical change: givensubstancebecomes a new substancew/ different properties and composition  Element –substances that cannot be decomposed into simpler substances bychemical or physical means R-9Energy  Energy – abilityto do work or produce heat o Work – force acting over a distance o Heat –energy that flows from one object to another due to a temperature difference between the objects  Lawof Conservation of Energy – energy canbe converted from one form to another, but never created or destroyed  Kinetic Energy – energy of motion  Potential Energy –stored energy due to position o Attraction ofopposite charges is animportant PE o Energy can switchfrom kinetic to potential and vice versa  Joule – unit of energy, kilogrammeter squared per second squared R-10The Mole  6.022 x 10 ^23 molecules =1 Mole  a unit for counting atoms and molecules Chapter 1: ChemicalFoundations 1-1 Chemistry: An Atoms-First Approach  Study atoms becauseitgives us a better understanding ofthe world o Atoms are the foundation of chemistry  STM / Scanning Tunnel Microscope allows us to seeatoms – confirms matter is made of atoms  Atoms are complex o Different atoms from the microscopic world (world of atoms and molecules) come together to form different substances inthe macroscopic world (our world) o Think of atoms as the letters of the alphabet. Letters construct words as atoms construct substances o We cantake them apart, and put them together  Matter is composed of various types of atoms  One substancechanges to another by re-organizing the way atoms are attached to eachother 1-2 The Scientific Method  Scientific Method: 1. Make Observations (collectdata) 2. Suggestan explanation (hypothesis) 3. Do experiments to test hypothesis o Scientists approach to solving problems o Investigator repeats above steps until there is enough knowledge to explain the phenomenon  Scientific Models o Theory / model –a setof tested hypotheses that gives anoverall explanation of some natural phenomenon o Observation vs.theories  Observation are witnessedand recorded  Theory is aninterpretation of why nature behaves a certain way – can changeas more information becomes available o Theories are human invention – explains observed natural behavior interms of human experience o Natural law –formulated statement of generallyobserved behavior o Theory vs.natural law  Law– summarizes WHAT happens  Theory – summarizes WHY ithappens  Theories focus hypotheses by giving expectations for outcome o Human politics effectscience 1-3 The EarlyHistory of Chemistry  1000 BC – ancientpeoples processed natural ores to produce metals for ornaments and weapons, alsoused chem. for embalming fluids  400 BC –Greeks proposed all matter was composed of four fundamental substances:Fire, Earth, Water and Air. Considered ifmatter is continuous and thus infinitely divisible  460 – 370 BC –Demokritos and Leucippos – termed atomos (later atoms) to describe ultimate particles  370 BC –1400 AD – Alchemy, discovered several elements and learned to prepare mineral acids  1500 AD –Georg Bauer developed systematic metallurgy (extraction of metals from ore). Paracelsus developed medical application of minerals  1627-1691 – Robert Boyle, first“chemist” o measured relationship between pressure and volume of air o published The Skeptical Chymist – 1661 o discovered quantitative behavior of gases o contributed ideas about chemical elements – samedefinition of element we usetoday  1600-1700 – Georg Stahl suggestedcontained fire stopped burning becausethe air became contaminated with “phlogiston”  1700-1800 – Joseph Priestley discovered Oxygengas 1-4 Fundamental Chemical Laws  Lavoisierdiscovered theLawofConservation ofMass –matter cannot be created or destroyed o alsodiscovered that combustion is fueledby oxygen, and oxygen is involved in a process that supports life  Joseph Proust followed Lavoisier’s work, and discovered LawofDefiniteProportion:a givencompound always contains exactly the sameproportion of elements by mass  Dalton was inspired by Proust, and reasoned ifelements were composed oftiny individual particles,a given compound should contain the samecombination of atoms o Why relativemasses ofelements are found in a compound  Alsodiscovered two given compounds canhave the same“ingredients” yet be different compounds o CO and CO2both have Carbon and Oxygenbut they are not the same compound  Lawof Definite Proportions: when two elements form a series ofcompounds, the ratio of the second element which combines with the first atom of the firstelement can be reduced to whole numbers o EX. Carbon Dioxide: For every ONEatom of carbon, TWO atoms of Oxygen bond to form one compound of Carbon Dioxide, therefore we write the ratio as C O2 1-5 Dalton’s Atomic Theory DALTON’S ATOMIC THEORY 1. Eachelement is made of tiny particles,atoms 2. Atoms of a givenelement are identical,atoms of different elements are fundamentally different 3. Chemical compounds form when atoms of different elements combine, a givencompound always has the samenumber and types of atoms 4. Chemical reactions involve the re-organizing of atoms –changes how they are bonded but the atoms themselves are not changed  Dalton prepared firsttable of Atomic Masses –atomic weight, his assumptions were wrong but his idea started the periodic table  Gay-Lussac –experimented with volumes of gases whichreacted with eachother  Avagadro’s Hypothesis – “atthe sametemp. and pressure, equal volumes of different gases containthe samenumber ofparticles o distancebetween particles is huge compared to the sizeof particles,thus volume would be determined by number of molecules present, not size 1-6 Early Experiments to Characterize the Atom  The Electron o J.J. Thomas –studied electrical discharges inCathode-ray tubes (partially evacuated tubes)  When highvoltage is appliedto the tube, a cathode ray was produced  Raywas a streamof negativelycharged particles (electrons) that were being repelled by a negative pole of the electrical field  Determined charge-to-mass ratio e/m = -1.76 x 10^8 C/g  e represents charge on the electrons in Coulumbs per gram mass o Concluded:  Since all electrons could be produced from electrodes of various metals – all atoms must contain electrons  Since atoms have to be electricallyneutral, they must alsohavea positive charge o Atomic structure must be a diffusecloud of positive charge with negative electrons randomly embedded –plum pudding model o Robert Milikan –studied magnitude of electron charge  Calculatedmass ofelectron to be9.11 x 10^-31 kg  Radio Activity o Some elements produce highenergy radiation o Henri Becquerel – a mineral containing uranium could produce an imageon a photographic platein absenceof light o Radioactivity– spontaneuous emissionof radiation by uranium  Types of radioactivity: 1. Gamma rays (y) –high energy light 2. Beta particles (B) – high speedelectron 3. Alpha particles (a) – 2+charge, 7300 biggerby mass ofan electron  The Nuclear Atom o Ernest Rutherford – tested Thomson’s plum pudding model  Directed a particles on a thin sheetof metal foil, ifthe model was correct than massiveparticles should crashthrough the foil  Results werenot expected, most a particles passedthrough, many were deflected at largeangles,somereflected – plum pudding model could not be correct o largedeflection had to be causedby a center of positive chargethat contains most of the mass,most particles passedthrough the foil becausethe atom is mostly open space o nuclear atom – atom with a dense center ofpositive charge with electrons moving around the nucleus at a distancelarger than nuclear radius 1-7 The Modern View of Atomic Structure: An introduction  Atom: tiny nucleus with electron that move about the nucleus  Nucleus contains protons and neturons o Protons: positivecharge equal to electon’s negative charge o Neutron: same mass as a proton but no charge  The arrangement and number of electrons determines anatoms chemical properties o Electrons constitute most of the atomic volume and thus “intermingle” when molecules form  Isotopes –atoms with the same number of protons but different numbers of neutrons  Atomic number: number of protons  Mass number: total mass of protons and neutrons Chapter 2:Atomic Structureand Periodicity 2.1 Electromagnetic Radiation  Electromagnetic radiation –how energy travels through space,havewavelike behavior & travel atthe speedof light o Ex.light from sun, microwave, X-rays  Characteristics: o Wavelength – (ƛ lambda) distancebetween two consecutive peaks or troughs in a wave o Frequency –(???? nu) number ofwavesper secondthatpassa givenpoint in space  All magnetic radiation travelat the speedoflight, therefore o Shortwavelength must travel at highfrequency o Longwavesmust have low frequency o THUS: ????ƛ=speedoflight (c) =2.9979x 10^8 2.2 The Nature of Matter  1900 – Max Planckproved energy could only be gainedor lostin whole number multiples of hv o h = Planck’sconstant =6.626x 10^-34J *h is the frequency of electromagnetic radiation absorbed/emitted o charge inevery systemΔE=nhv *n is aninteger  energy is quantized and can occur only indiscrete units of sizehv calledquantums  Albert Einstein– electromagnetic radiation is itselfquantized, canbe viewed as a stream of photons E= hv = (hc)/ƛ h=Plancks constant v=frequency ƛ=wavelength  The Photoelectric Effect o Photoelectric effect –phenomenon inwhich electrons are emitted from the surfaceof a metal when light strikes it 1. When frequency of lightis varied, shows no electrons are emitted by a given metal below threshold frequency 2. When frequency of lightis lower than threshold frequency, no electrons are emitted regardless of light intensity 3. When frequency of lightis greater than threshold frequency, number of electrons increases with intensity of light 4. When frequency of lightis greater than threshold frequency, kinetic energy of emitted electrons increaselinearly o Threshold frequency is the minimum energy required to remove an electron from a metal’s surface E0 =hv- o Lightwith a frequency less than threshold frequency produces no electrons becausea photon with energy less than Ecannot remove an electron o Lightwhen v>v0 , the energy inexcess is givento the electron as kinetic energy KEelectron = 1/2mv^2 =hv –hv0 (seepage58)  Intensity of light is a measure of the number of photons present in part of the beam, greater intensity = more photons are availableto release electrons E= mc^2  Special Theory of Relativity– energy has mass  So, mass of a photon length = M = (E/c^2)= ((hc/ƛ)/c^2) =(h/ƛc) seepage58 o 1922 – Arthur Compton through x-ray experiments showed that photons do exhibit mass but not in the classical sense.Photon mass is relative,ithas no rest mass Important takeaways: -Energy is quantized, canoccur in units calledquanta -Electromagnetic radiation shows characteristics of particulate matter as well as wave properties “dual nature of light” o Louis Broglie – questioned “iflighthas particulate qualities,do particles have wavelike qualities” Particle velocity =m =h/ƛv thus, ƛ=h/mv o Diffraction –when light is scatteredfrom a regular array of points or lines  Colors resultbecausevarious wavelengths of visiblelightarenot scattered in the sameway o Diffraction pattern – pattern produced by scattered radiation, includes areas of dark and light o Diffraction occurs most efficientlywhen spacing between scattering points is the sameas the wavelength of the wave being diffracted  Electromagnetic radiation was found to possess particulateproperties  Electrons which were thought to be particles were found to have a wavelength  Energy is a form of matter, all matter shows the same types ofproperties (both eave and particulate) 2.3 The Atomic Spectrum of Hydrogen  When H2molecules absorb energy &some of the H-Hbonds arebroken, the resulting H atoms are excited o Excited =contain excess energy which they releasebyemitting lightof various wavelengths to produce the emissionspectrum  Continuous spectrum – results when white lightis passedthrough a prism, a rainbow, contains all wavelengths  Linespectrum –results when the hydrogen emissionspectrum passes through a prism, only a few lines which correspond with a different wavelength o Significance:only certain energies are allowed for the electron in the hydrogen atom, energy is quantized o Changes inenergy between discreteenergy levels in hydrogen will produce certain wavelengths of emitted light o Calculatedwith Planck’s equation *only certain energies are possible,levels arequantized. If any energy level were allowed, spectrum would be continuous 2.4 The Bohr Model  1913 – Neils Bohr developed Quantum model – electron inhydrogen atoms moves around the nucleus in certain allowed circular orbits o particle inmotion tends to stayina straightline, unless there is a force acting upon it o charged particle under acceleration should radiate energy, sincean electron is constantly changing direction (think circle),itis constantly accelerating o thus the electron should emit light& loseenergy and be drawn into the nucleus o BUT the electron doesn’t do that soBohr concluded the electron could occur only in increments o Energies availabletothe electron inthe Hatom E = -2.178 x 10^-18 (z^2/n^2)  N = aninteger  Z = nuclear charge  - signmeans energy of electron bound to nucleus is lower than ifthe electron were at infinite distance,which would be 0 o ground state –lowest possiblestateof a hydrogen atom o changes instate arecalculatedby ΔE= energy of final state – energy of initial state * -indicates the atom has lostenergy and is now more stable o wavelength can be calculatedby ΔE=h(c/ƛ) or ƛ=hc/ΔE * cannot have negative wavelength, sothe negative signis ignored  Bohr’s model is fundamentally incorrect becauseit cannot be applied to other atoms 2.5 The Quantum Mechanical Model of the Atom  Louis de Boglie –electron while being a particle, still shows waveproperties  Schrödinger –emphasized the wave properties of an electron o Paralleledan electron bound to a nucleus is likea standing wave, thus began to research o Standing wave – wave does not travel, the end points are fixed while other points vibrate atmaximum amplitude – must be a whole number of half wavelengths inany allowed motion of the string  Hypothesized the electron follows a “standing wave pattern” throughout the circumference of the atom  Schrödinger had to test whether or not the model would correctly fit the experimental data on hydrogen and other atoms, used a complicated math equation o ĤΨ=EΨ  Ψ =wave function – function of (x,y, and z), electrons 3Dposition in space  Ĥ =operator – contains mathematical terms that produce the total energy of the atom when applied to the wave function  E= total energy of the atom –sumof potential energy due to attraction of proton and electron, and kinetic energy of moving electron  Produces multiple solutions, eachsolution represents an orbital or specific wavefunction  Quantum mechanical model of the atom o Lowest energy of H atom –Ψ (wave function) = 1s orbital o We do not know the pathway ofan electron  Heisenberg Uncertainty Principle –there is a fundamental limitation to how preciselywe canknow both the position and momentum of a particle ata giventime ℎ Δ???? ⋅ Δ(????????) ≥ 4????  Δ???? =uncertainty of a particle position  Δ(????????) =uncertainty ina particles momentum  h = planck’s constant  the more accuratelywe know a particles position, the less accurately we know its momentum & viceversa  therefore, we cannot assumethe nucleus moves ina well defined orbit as inthe Bohr model  The Physical Meaning of a WaveFunction o The square of the function indicates the probability of finding an electron near a particular point in space  Ex.two points (X1, Y1, Z1)and (X2, Y2,Z2), the relative probability of finding the electron inposition 1and 2is givenby substituting into the wave function ????(????1,????1,????1) ????1 ????(????2,????2,????2)= ????2  Quotient is ratio of probabilities of finding the electron at possibility1 and 2  No indication ofwhen it will be there or how it moves o Probability distribution – square of the wave function a. intensity of color is used to indicate the probability value near a given point in space –darkness of a point indicates the probability of finding anelectron at that position can be depicted on a graph b. seepage71 o Radial probability Distribution – total probability is plotted vs.distancefrom nucleus a. calculatethe probability at various points on a line drawn outward in any direction from the nucleus and depict it in a graph b. seepage71 o Maximum of the curve occurs becausethe total probability increases to a certain radius then decreases as the electron probability ateach position becomes smaller o The sizeof anorbital is the radius of the sphere that encloses 90% of the total electron probability 2.6 Quantum Numbers  Quantum numbers are a series ofnumber used to characterize an orbital o Principal Quantum Numbers: (n),integral values:1, 2,3…  Relatedto sizeand energy of the orbital  As “n” increases,theorbital becomes larger,electron spends more time farther from the nucleus  Increasein “n” alsomeans higher energy, electron is less bound to nucleus and energy is less negative o Angular momentum quantum number: (????), integral valuesfrom0to n-1for eachvalueofn  Related to shapeofatomic orbitals  Assigneda letter  ????=0 iss, ????=1 is p,????=2 is d, ????=3is f, ????=4is g o Magnetic Quantum number: (???? ),integral values between ???? and−???? ????  Relatedto the orientation of the orbital in spacerelative to the other orbitals inthe atom 2.7 Orbital Shapes and Energies  Nodal surfaces /nodes – orbitals which contain areas of high probability separated by areas of zero probability o Increases as “n” increases o For s orbitals, the number of nodes is givenby n-1  Atomic orbitals havesigns o Function for s orbital is positive everywhere in three-dimensional space o P orbital functions havedifferent signs indifferent regions of space o Signs are simplymathmetical signs forlocation purposes, not as charges o Just as a sinealternates between positive and negative,p orbitals also alternate in different phases  3p orbitals have more complex probability distribution, represented by the sameshapes & boundaries as 2p orbitals but surfaces increaseas nincreases  d orbitals first occur in level n = 3 o two different fundamental shapes  four orbitals (???????????????? ???????? ,???????????? ???? ???? −???? 2)have four lobes centered in the plane indicated inthe orbital label  ???? ????2 has a special shapewith two lobes along x-axis anda belt centered on xy plane  f orbitals first occur in level n=4  energy of an orbital is determined by n, all orbitals with the samevalueof n are degenerate because they have the same energy  Hydrogen’s singleelectron can occupy any of the atomic orbitals, but inground stateit resides in1s orbital –can be transferred to a higher orbital with a higher energy when excited 2.8 Electron Spin and the Pauli Principle  Goudsmit and Uhlenbeck developed concept of electron spin  A fourth quantum number was needed to account for the details of the emission spectra of atoms o Indicative the electron has a magnetic moment with two possibleorientations when the atom is placedin a magnetic field  Electron spinquantum number – (???? ),+1/2 or -1/2 ???? o An electron has two spinstates which produce two oppositely directed magnetic moments  Pauli exclusionprinciple o In a given atom, no two electrons can havethe samesetof four quantum numbers: n, ????, ????????,and ???? ????  Electronsin the sameorbital havethesame n,????, ???? so the????eforeany orbitalcan onlyholdtwo electronsandthey must haveoppositespins 2.9 Polyelectronic Atoms  Polyelectronic atoms – atoms with more than one electron  How does quantum mechanical model apply to polyelectronic atoms o First consider  Kinetic energy as electrons move around nucleus  Potential energy of attraction between nucleus and electrons  Potential energy of repulsion between two electrons o Electron correlation problem: Difficultyin the repulsions between electrons, sincewe cannot know exactlyhow a singleelectron behaves,we cannot preciselycalculatetheir repulsions  Must treat eachelectron as ifitwere moving ina fieldof chargethat is the net result of the nuclear attraction and averagerepulsions of all the other electrons  Ex.Na has 11 electrons. The outermost electron is attracted to the nucleus, but the repulsions from the other 10 electron prevent it from being more tightly bound to the nucleus  Electron is screenedor shieldedfrom the nuclear charge o Orbitals of polyelectronic atoms are hydrogenlike, they havesimilarshapes but sizes andenergies are different  In the hydrogen atom. All orbitals in a givenprincipal quantum level have the same energy  In polyelectronic atoms, orbitals vary in energy ???? ???????? < ???? ???????? < ???? ???????? < ???? ????????  See page78 2.10 The History of the Periodic Table  Unsuccessful attempts o Originallyconstructed according to patterns in chemical properties  became evident that this would not work due to the amount and diversity of elements o Johann Dobereiner –1790-1849 –found groups of three elements that have similarproperties – triads o John Newlands –1864 – suggestedelements could be arranged in octaves, becausecertain properties repeated for every eighth element  Successful attempt o Julius Meyer and Dmitiri Mendeleev – formed the present periodic table o Mendeleev emphasized the table’s abilityto predict elements properties o Now universallyadopted 2.11 The Aufbau Principle and the Periodic Table  Assumption: all atoms have the same type of orbitals as the hydrogen atom  Aufbau principle –as protons areadded one by one to the nucleus to build up the elements, electrons are similarlyadded to hydrogenlike orbitals o Ex.Hydrogen has one electron inthe ground state 1  Configuration: 1???? ↿ o Helium: has two electrons in the 1 s orbital with opposite spins  1???? ↿⇂ o Lithium: th2ee1electons,2 in 1s orital, 1in 2s orbital  1???? 2???? ↿⇂ ↿  Hund’s rule – the lowest energy configuration for anatom is the one having the maximum number of unpaired electrons allowed by the pauli principle in a particular set of degenerate orbitals o There are three 2p orbitals with the sameenergy. Mutually repulsive electons will occupt separateorbitals before pairing together 2 2 2  Ex.Carbon: 1???? 2???? 2???? ↿⇂ ↿⇂ ↿ ↿ o Abbreviations canbe used soextensive electron configurations are not written out 1  Ex: Na is (Ne)3????  ValenceElectrons – the electrons in the outermost principal quantum level of a main group atom, important becausethey are involved in bonding o Nitrogen the valence electrons are 2s and 2p electrons o Sodium the valenceelectrons areis the 3s  Core electorns – inner electrons  Elements in the samegroup (vertical column) havethe samevalenceelectron configuration o Elements with the samevalenceelectron configuration displaythe same behavior  Only transition metals have valence electron configurations including “d”,the rest fall into the pattern of 2s,2p, 3s,3p, 4s,4p, ….  Rules of configuration 1. The (n+1)s orbitals always fill before the nd orbitals 2. After Lanthanum, the lanthanide series occurs – corresponds to the filling of 4f orbitals 3. After actinium, the actinide series occurs – fills 5forbitals 4. Group labels (representative elements) for 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A indicate the total number of valenceelectrons for the atoms inthis group 2.12 Periodic Trends in Atomic Properties  Atomic Radius o Size of an atom cannot be exactly specified o Atomic radii values areobtained by measuring the distancebetween atoms in chemical compounds  Radii –covalent atomic radii becauseof how they are determined o Nonmetallic atoms that do not form diatomic molecules, the atomic radii are estimated from covalent compounds o Metallic radii are obtained from halfthe distance between metal atoms in solid metal crystals o Representative element radii:  Radii decreasein going leftto right across a period  Valenceelectrons are drawn closerto the nucleus, decreasing the sizeof the atom  Ionization Energy o Ionization energy is the energy required to remove an electron from a gaseous atom or ion + − ???? ???? → ???? (???? + ????  Where the atom is assumedto be in ground state o First ionization energy: The highest-energy electron (leastbounded to nucleus) is removed first o Second ionization energy: second highest-energy electron is removed after the highestis removed, the value is significantlylargercompared to the first ionization energy  Makes sensebecause:  Charge; negativelycharged electron is removed from a neutral atom soit becomes (1+ion) o Thus the increaseinpositive energy (due to absenceof negative)binds the electrons more firmly  Electron configuration; the first electron is removed from an orbital that is higher inenergy than the second electron *strongest when removing core electrons o Generally: as you travel from leftto right, the first ionization energy increases o Electrons added in the same principal quantum level to not completely shieldthe increasing positive chargefrom the increaseinprotons o Thus,electrons inthe samePQL are more strongly bound as you move to the right of the periodic table o Generally: as you travel down a group, the ionization energy decreases o becauseelectrons being removed are farther from the nucleus o As n increases,sizeoforbital increases,electrons are easierto remove  Electron Affinity o Electron affinityis the energy change associatedwiththe addition of an electron to a gaseous atom ???? ???? + ???? → ???? − (????)  If the addition of the electron results ina lower energy, then the corresponding valuefor electron affinity will carry a negativesign o The more negative the energy, the greater quantity of energy released o Generally: become more negativefrom left to right across a period  Exceptions: o Generally: become more positivedown a group 2.13 The Properties of a Group: The Alkali Metals  Information contained in the Periodic Table 1. Groups of representative elements exhibit similarchemical properties that changeregularly. Quantum mechanical model allows us to understand why they aresimilar– number and type of valenceelectrons that determine an atoms chemistry 2. Representative elements caneasilybeelectronically configurated 3. certain groups are reffered to in names 4. most basic divisionon tableis metals to nonmetals. Metals: Most important chemical property of a metal atom is the tendency to giveup electrons to form a positive ion; metals tend to have low ionization energies.Most chemically reactivemetals are inthe lower- left with the smallestionization energies Nonmetals: abilityto gainone or more electrons to from ananion when reacting with a metal. Elements with largeionization energies and most negativeelectron affinities.Locatedin right sideof the table, most reactive in upper-right corner except for noble gas elements Metalloids/semimetals:exhibit properties of both. Si,Ge, As,Te, Po, and At.  The Alkali Metals –metals of group 1A o Most chemically reactive o Hydrogen is not a metal o Elements agreewith general trends previously discussed o Increasein density as you go down the group  Becauseatomic mass increases more than atomic size o Smooth decreasing in melting point is unique to this group o Most characteristic property: ability to losevalenceelectrons – very reactive o Low ionization energies Chapter 3Bonding:GeneralConcepts 3.1 Types of Chemical Bonds  Dalton – Firstrecognized that chemical compounds are collections of atoms, but didn’t understand the structure of the atom or the need to bond  Chemical bonds –forces that hold groups of atoms together and make them function as a unit o Covalent bonds- chemical bonds formed by sharing electrons, results ina molecule (electrons are shared by nuclei) o Ionic bonds – chemical bonds formed when anatom that loses electrons relatively easilyreacts withan atom that has a highaffinity for electrons  Ionic compounds forms when a metal reacts with a nonmetal o Chemical formula – symbols indicate the types of atoms present and subscripts indicate the relative number of atoms  Structural Formula – individual bonds are shown (indicated by lines),may or may not represent actual shape of the molecule represented  Space-filling model – shows the relative sizes ofthe atoms as well as their relative orientation inthe molecule  Ball andstickmodel – used to represent molecules  In a compound, individual molecules move around as independent units  We canuse experiments to testthe properties of materials o Study physical properties such as melting point, hardness, and electrical and thermal conductivity o Solubility and properties of solutions o Electric charge by placing itin a magnetic field o Strength of bonding by measuring bond energy – energy required to break a bond  Coulumb’s lawcalculates theenergy of interaction between a pair of ions −19 ???? 1 2 ???? = (2.31 × 10 ????∙ ????????)( ) ????  Eis in Joules o If Eis negative,that is an attractive force, the ion pair has a lower energy than when separated o If Eis positive,the force is repulsive  r = distancebetween ion centers in nm  ????1&???? 2re numerical ion charges  Bonding between two identical atoms o When two Hatoms are brought together, there are two unfavorable potential energies  Proton – proton repulsion  Electron –electron repulsion o And one favorable energy  Proton – electron attraction o Favorable conditions for bond formation?  Any systemin nature has a strong tendency to achievethe lowest possibleenergy  A bond will form ifthe system canlower its total energy in the process  The systemwill act to minimize the sum of the positive (repulsive)energy and maximize the negative(attractive) energy  Bond length –the distancewhere the energy is minimal between the hydrogen nuclei *net potential energy results from attraction and repulsions among the charged particles *kinetic energy results from the motions of the electrons *zero point of energy is defined with the atoms at infiniteseparation *at short distance,the energy rises steeplybecausethe importance of repulsive forces when the atoms are closetogether *bond length is the distanceat which the systemas minimum energy o In ???? the electrons reside inthe spacebetween the two nuclei,where they are 2 simultaneously attracted by both protons –leads to optimal stabilityof the atom o ???? i2 more stablethan two separated Hatoms o in terms of forces, simultaneous attraction of eachelectron by the protons generates a force that pulls the protons towards eachother and balances the proton – proton and electron – electron repulsions at bond length  Ionic bonds – electrons are transferred to form oppositely charged ions which then attract each other  Covalent bonds – two identical atoms share electrons equally,results from mutual attraction of two nuclei for shared electrons  Polar covalent bonds – intermediate bond of ionic and covalent, atoms end up at opposite poles in a molecule due to the unequal sharing of incompletely transferred atoms o Results infractional charges on different atoms o One atom has a stronger attraction for shared electrons than the other atom o Ex.HF 3.2 Electronegativity  Electronegativity is the ability of anatom in a molecule to attract shared electrons to itself  Linus Pauling (1901-1995) – won nobel prize in chemistry and peace, o Consider hypothetical molecule HX o Relativeelectronegativity is determined by comparing measure H-X bond energy with “expected” H-X bond energy o Resultis the averageof H-H and X-X bond energies ???? − ???? ???????????????? ????????????????????????+ ???? − ???? ???????????????? ???????????????????????? ???????????????????????????????? ???? − ???? ???????????????? ???????????????????????? = 2 o The difference (Δ)between actual (measured) and expected bond energies is Δ = (???? − ????) ????????????− (???? − ????) ???????????? o IF H and X have identical electronegativities are the sameand ∆ = 0 o IF X has a greater electronegativity then, shared electrons will be closerto the X atom and the molecule will be polar  Bond has ionic and covalent components – attraction between partially and oppositely charged Hand X atoms will lead to a greater bond strength  The actual bond energy difference will be greater than the expected bond energy difference  The greater distanceof bond energies,the greater the ionic component and the greater ∆ value  Electronegativity generally increases fromleft to right across a period and decrease going down a group  For identical atoms, (electronegativity of 0),the electrons inthe bond are shared equally, no polarity  For two atoms with very different electronegativities interact, electron transfer can from ions  Intermediate cases givepolarcovalent bonds with unequal electron sharing 3.3 Ions: Electron Configuration and Sizes  Electron Configuration of Compound o Two nonmetals - form covalent bonds, they share electrons sothat the valence shell is complete for both atoms – both atoms have nobel gas configuration o Nonmetal and representative group metal – form binary ionic compound, ions form so the valenceelectron configuration of the nonmetal is filled,andthe metal is emptied –both atoms have noble gas electron configuration  Predicting formulas of ionic compounds o Ionic compound –solidstate of the compound, ions are closeand contain a large collection of positiveand negative ions closetogether that minimize repulsions and maximizes attractions  In a gaseous state,ions are far apart, a pair of ions may getcloseto interact but largecollections do not exist  Assume when you hear “ionic compound” it is inthe solidform  Ex.electron configuration of stable,solidionic compounds of calciumand oxygen ???????? ???????? 4???? 2 2 4 ???? ???????? 2???? 2???? o Electronegativity of oxygen (3.5) is largerthan Calcium(1.0)*known from table o 2 Electrons will be transferred from Calciumto Oxygen ???????? +???? → ???????? 2+ + ???? 2− o chemical compounds are always electricallyneutral, samequantities of positive and negative charges o will from CaO  Ex.Aluminum and Oxygen 2 1 ???????? [????????]3???? 3???? o Aluminum loses three electrons to form ???????? 3+becauseas a metal, itmust giveall its electrons to the nonmetal (O) o Aluminum achieves Neon configuration o Results inthe ions ???????? 3+???????????? ???? 2−,sincecompounds must be neutral there must be three Oxygenions for every two Aluminum ions o Achieve the empirical formula ???????? 2 3  Exceptions: Tinforms ???????? 2+ ???????????? ???????? , Leadforms ???????? 2+ ???????????? ???????? ,+ 3+ 5+ + 3+ Bismuth forms ???????? ???????????? ???????? ,and Thalliumforms ???????? ???????????? ????????  Sizes of Ions o Various factors influence size:  Relativesizeof the ion and its parent atom  Positivecharge removes an electron – the cation will be smaller than parent atom  Negativecharge adds an electron –anion will be largerthan parent atom  Position of the parent elements in the periodic table  Ratioof protons and number of electrons in the ion  ions can have the samenumber of electrons but if they have a greater number of protons the force on electrons will be greater which makes the sizeof the ion smaller 3.4 Energy Effects inBinary Ionic Compounds  LatticeEnergy – the change inenergy that takes placewhen separated gaseous ions are packed together to form an ionic solid o Energy releasedwhen an ionic solidforms from its ions, signis determined by the change inenergy that occurs  Negativeif process results in a lower energy  Positiveif process results in a higher energy o See page112 for example /demonstration  LatticeEnergy Calculations Latticeenergy = ????( ????1 2) ????  K is proportionality constant that depends on the structure of the solid and electron configuration of the ions  Q1 and Q2are the charges of the ions  r is the shortest distance between the centers of the cations and anions o Latticeenergy is negativewhen Q1 and Q2have opposite sings o Releases moreenergy as the ionic charges increaseand as the distances decrease o See page114 for example /demonstration o 3.5 Partial Ionic Character ofCovalent bonds  Probably no totally ionic bonds between discrete pairs of atoms o Canbe calculatedbasedon comparisons of measured dipole moments for molecules of the type X-Y  Percent ionic character of a bond = (???????????????????????????? ???????????????????????? ???????????????????????? ????????+????−????)x 100% ???????????????????????????????????????? ???????????????????????? ???????????????????????? ???????? ???? ????  Ionic character increases with electronegativity difference, bonds never reach 100% ionic character, no bonds canbe completely ionic  All compounds with more than 50% ionic character are considered to be ionic solids  Another complication arises incompounds with polyatomic ions which are bonded with covalent bonds  Any compound that conductsan electric currentwhen melted willbeclassified as ionic 3.6 The Covalent Chemical Bond: Amodel  Bonds occur when collections of atoms are more stable(inlower energy) than separate atoms (in higher energy)  Ex.Methane consistof four hydrogen atom arranged at the corners of a tetrahedron around a carbon atom o Four individual reactions of hydrogen and carbon calledbonds o Energy stabilizationis divided equally among the four parts 1642 ????????/???????????? = 413 ????????/???????????? 4 o Methyl Chloride o 1578 kJ of energy is required to bread down one mole of gaseous methyl chloride o it is lower in energy by 1578 kJ than separate gaseous atoms,thus one molecule is held together by 1578 kJ of energy o seepage117-118 o bonds area human invention, they provide a method for dividing energy  Models: an Overview o Models – attempts to explainhow nature operates on the microscopic level, basedon experiences in the macroscopic world o Originatefrom out observations of the properties of nature o Bonds form due to a tendency toward lower energy – collections of atoms aggregatebecausetogether they have a lower energy o It is correct to assumethat bonds hold atoms together in every type of structure, including polyatomic ions o Bonding model provides framework to systematize chemical behavior o Bonding molecules physicallysensible,makes sensethatatoms can form stable groups by sharing electrons becausesharing gives a lower stateof energy becausethey are simultaneouslyattracted to two nuclei 3.7 Covalentbond energies and chemical reactions  bonds aresensitive to their molecular environments  energy required to break bonds in elements is not systematic  useaverageof individual bond dissociationenergies even though itonly approximates the bond energies inparticular molecule o singlebond – bonds inwhich one electron pair is shared o double bond – bonds in which two pairs of electrons are shared  bond energy is not twice of a singlebond o triple bond –bonds inwhich three pairs of electrons are shared  bond energy is not three times the singlebond  relationship exists between number ofbonds and bond length  Bond Energies o Bond energy values canbe used to calculateapproximate energies for reactions o Energy must be added to a bond for itto break (positive ΔE)  ΔE= sum of the energies required to break old bonds (positivesigns)+ sum of energies releasedin the formation of new bonds (negativesigns ΔE= Σn x D(bonds broken) – Σn x D(bonds formed) Energy required energy released  Σ = sum of terms  D= bond dissociationenergy per mole of bonds  N represents moles of a bond type 3.8 The LocalizedElectron Bonding Model  Bond strength and polarity canbe assignedto individual bonds  LocalizedElectron (LE)model –assumes a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms (covalent bonds) o Electron pairs are assumedto be localized on a particular atom or in the space between the two atoms  Lone pairs – pairs localizedon an atom  Bonding pairs – pairs found in the spacebetween atoms LEmodel has 3 parts: 1. Description of valenceelectron arrangement inthe molecule using Lewis structures 2. Prediction of geometry of the molecule using valenceshell electron-pair 3. Description of the type of atomic orbitals used by atoms to share electrons or hold lone pairs 3.9  Lewis Structure – shows how valenceelectrons are arranged among the atoms inthe molecule  Most important guidelines for the formation of a stablecompound is that the atoms achievenoble gas electron configurations  Rule: only valenceelectrons are included  Principles of achieving noble gas configuration o DUET RULE: stablemolecules are formed when itshares two electrons  Ex.H2 o Helium does not form bonds because its valenceorbital is already filledand therefore itis a noble gas o Second-row nonmetals (C through F) form stablemolecules when they are surrounded by enough electrons to fill the valenceorbitals –Octet rule: surrounded by 8 electrons o Neon does not form bonds becauseit already has anoctet of valenceelectrons  Problem-Solving strategies o Sum valenceelectrons from all atoms. Don’t worry about keeping track of which electrons come from which atoms o Usea pair of electrons to form a bond between eachpair of bound atoms o Arrange remaining electrons to satisfyduet rule for hydrogen and octet for second-row elements 3.10 Exceptions to the Octet Rule  Second row elements C,N, O and F always obey the octet rule  Second row element B and Be often have fewer than 8 electrons –make them very reactive  Second row elements never exceedthe octet rule, valenceorbitals only accommodate 8 electrons  Third row and heavier elements satisfyoctetrule but can exceed  Satisfyoctet rule atoms first,if electrons remain then placethem on any element period 3 and beyond 3.11 Resonance  More than one validLewis structure is possiblefor a givenmolecule o There is no reason to choose a particular atom for the double bond, thus all situations must be satisfied o The correct description of NO3 is not given by any one ofthe three Lewis structure but by the superposition of all three  Resonance– condition in which more than one valid Lewis structure canbe written for a particular molecule, resulting electron structure is givenby the averageof resonance structures o Necessarybecauselocalizedelectron models postulate that electrons are localizedbetween a givenpair of atoms, but nature does not operate this way – resonance is necessarytocompensate for defectiveassumption of the localizedelectron model  Odd-Electron molecules o Rare o A more sophisticated model is needed  Formal Charge o Molecule or polyatomic ions that canexceed the octet rule have many nonequivalent Lewis structures o Formal charge –the difference between the number of valence electrons on the free atom and the number of valenceelectrons assignedtothe atom in the molecule o Need two things:  Number of valenceelectrons on the free neutral atom  Number of valenceelectrons “belonging” to the atom in a molecule o Formal Charge= (number of valenceelectrons on free atom) –(number of valenceelectrons assignedto the atom in the molecule)  To compute we assignvalenceelectrons inthe molecule to various atoms:  Lone pair electrons belong entirely to the atom in question  Shared electrons are divided equally between the two sharing atoms o Atoms inmolecules try to achieve formal charges as closeto zero as possible o Any negativeformal charges areexpected to reside on the most electronegative atoms 3.12 Naming Simple Compounds  Common names (i.e.sugar) were used before the system was invented  Binary compounds –compounds composed oftwo elements  Binary compounds (Type I) o Binary ionic compounds contain a positive ion (cation) -always written first- and a negative ion (anion) o Rules:  Cation is named first  Monatomic cation takes its name from the name of the element  Monatomic anion is named by taking the root of the element and adding


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