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chem notes

by: Hannah Czajkowski

chem notes CHE 1101

Hannah Czajkowski

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chapter 3, 4, and 7
General Chemistry1
Jennifer Cecile
Class Notes
Chem genchem general chemistry
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This 13 page Class Notes was uploaded by Hannah Czajkowski on Thursday March 31, 2016. The Class Notes belongs to CHE 1101 at Appalachian State University taught by Jennifer Cecile in Spring 2016. Since its upload, it has received 13 views. For similar materials see General Chemistry1 in Chemistry at Appalachian State University.

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Date Created: 03/31/16
Chapter 3 The Mole and Stoichiometry  Stoichiometry o Greek root of two words meaning measure elements o Study of mass relations in chemistry o Remember the law of constant composition, conservation of mass, and multiple proportions  Chemical equation o An abbreviated way to describe a chemical reaction o Reactant + reactant --> product o The physical state of each reactant and product is given  Gas g  Solid s  Liquid l  Aqueous aq o The equation is balanced as the same number of atoms of a given element appear on both sides of the equation o Sometimes a temperature or the addition of other chemicals to the reaction is placed over the arrow  Coefficients give the number of complete molecules in a reaction  Subscripts are used in the chemical formula of each reactant or product molecule and give the number of atoms each element present in formula  Balancing an equation o Start balancing with the most complicated formula first o Balance atoms that appear in only two formulas (one in reactant, one in product) o Make sure that all the coefficients are in the lowest possible whole number  Reaction Types o We can classify reactions in order to help us write equations o Combination reactions- reaction in which two or more substances to form one product o Decomposition reaction- reaction in which one substance breaks down into two or more substances o Combustion reaction- rapid reaction that produces a flame, usually have hydrocarbons that react with oxygen in the air; produces carbon dioxide and water, incomplete will make carbon monoxide  The mole- a number equal to the number of atoms in exactly twelve grams of carbon atoms o Avogadro's number  Number of atoms, molecules, or particles in one mole/mol  6.022x10 23  Percent composition o The percent by mass contributed by each element in the substance o % element= (number of atoms * atomic mass) / (molar mass of the compound) *100  Empirical formulas o The lowest whole number ratio possible which gives the relative number of atoms of each element it contains o Simplest formula  Mass % elements  Grams of each element  Moles of each element  Empirical formula  Molecular formula o The subscripts in the molecular formula of a substance are always a whole number multiple of the corresponding subscripts in its empirical formula o Whole number multiple= (molecular weight)/(empirical formula weight)  Quantitative stoichiometry o Find quantities of moles from coefficients of balanced chemical reaction equations o Use moles rations from balanced chemical equations TEST THREE  Reactant quantities o Limiting reactants  Reactant that is completely used up in the reaction  Present in lower number f moles  It determines the amount of product produced o Excess reactant  Reactant that has some amount left over at end  Present in higher number of moles  Limiting reactants- finding the least abundant reactant based on the equation for a reaction that limits the maximum product formed o Find amount of product formed if first reactant is completely consumed o Find amount of product formed if second reactant is completely consumed o Choose the smaller of the two amounts of product formed, reactant that yields smallest product is limited  Reaction yields o Theoretical yield  Amount of product that must be obtained if no losses occur  Amount of product formed if all of limiting reagent is consumed o Actual yield  Amount of product that is actually isolated at end of reaction  Amount obtained experimentally  How much is obtained in mass units or in moles  Percent yield o Percent yield= (actual yield)/(theoretical yield)*100 Chapter 4 Molecular View of Reactions in Aqueous Solutions  Aqueous reactions o Aqueous solutions are solution that have water as the dissolving medium o Present has a homogeneous mixture  Two or more components mix freely  Molecules or ions completely intermingled o Types of aq reactions  Precipitation reactions  Acid base reactions  Oxidation reduction reactions  Solution- homogeneous mixture  Solvent-medium that dissolves solutes o Component present in large amount o Can be gas, liquid or solid o Aq solution--water is solvent  Solute- substance dissolved in solvent o Solution is named by solute o Can be gas liquid or solid  Concentration- the amount of solute dissolved in a given quantity of solution  The greater the amount of solute the more concentrated the solution  Dilute solution- small solute to solvent ration  Concentrated solution- large solute to solvent ration  Electrolytes o Ionic compounds conduct electricity o Molecular compounds don’t conduct electricity  Dissociation- defined as the process in which an ionic substance dissolves in water; the solvent pulls individual ions from the crystal and solvates them o An electrolyte includes substances that dissociate into ions when dissolved in water o Electrolytes conduct electricity o Strong electrolytes are strong acids and bases as well as ionic compounds  Non-electrolytes o Solute particles are surrounded by water but the molecules do not dissociate  Concentration units o Solute to solvent ration o G solute / g solvent o G solute / g solution o Percent concentration- g solute / 100g solution  Electrical conductance o Strong electrolytes  Dissociate 100% in water  Good electrical conduction  Ionic compounds (NaCl, KNO3)  Strong acids and bases (HClO4 and HCl) o Weak electrolytes  Small percentage of molecules ionize in water  Weak conductors of electricity  Weak acids and bases  Dissociation Reactions o Strong electrolytes o Ionic compounds dissolve to form hydrated ions  Hydrated- surrounded by water molecules indicated with aq  Precipitation Reactions o Solid product formed from reaction in which an insoluble product forms and separates out of solution  Writing equations o Chemical equations  Molecular equations list the reactant and products in their molecular form  Ionic equations all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions  Net ionic equation only lists substances actually involved in the reaction  Write a balanced molecular equation  Dissociate all strong electrolytes  Cross out anything that remains unchanged from the left side to the right side of the equation ( spectator ions)  Write the net ionic equation with the species that remain  Solubility o Solubility = g solute needed to make saturated solution / 100 g solvent o Temperature dependent o Saturated solution- solution in which no more solute can be dissolved at a given temperature o Unsaturated solution- solution containing less solute than maximum amount  Acid and Bases o Arrhenius defined these species  Acid- produces H+ ions in water solution  Base- produces OH- ions in water solution o Bronsted-Lowry Definition  Acid- proton donor  Base- proton acceptor  Acids o Acids are found in industrial, household environments and in our bodies o Have a sour tastes and affect color of organic dyes known as acid base indicators (turn litmus paper pink) o Substances that ionize in aq solutions to form H+ ions o Polyprotic acids  Acids can form different number of H+ ions per molecule  Mono-protic- produce 1 H+ ion per acid molecule; HCl, HNO3  Diprotic- produce 2 H+ ions per acid molecule; H2SO4  Tri-protic- produces 3 H+ ions per acid molecule; H3PO4  Sulfuric Acid o H2SO4 is a strong electrolyte but only the first ionization is complete o The aqueous solution contains H+, HSO4 and SO4 ions  Bases o Bases are found in industrial, household environments and in our bodies o Bases have a bitter taste and turn litmus paper blue  Strong acids- dissociates completely forming H+ ions and anions (memorize) o HClO 4 perchloric acid HClO 3 Chloric acid HCl hydrochloric acid HBr Hydrobromic acid HI Hydroiodic acid HNO 3 Nitric acid H 2O 4 Sulfuric acid  Strong bases- compeltely ionized to OH- ions and cations (memorize) o Group 1A metal hydroxides  LiOH, NaOH, KOH, RbOH, CsOH o Group 2A metal hydroxides  Ca(OH) 2 Sr(OH) 2 Ba(OH) 2  Weak acids- only partially dissociate to H+ ions in water, they are molecules containing an ionizable hydrogen atom. Any acid that isn't one of the seven strong acids  Weak bases- only partially dissociated to form Oh- in water  Electrolyte strength o Ionic? strong electrolyte o Molecular?  Acid?  Strong acid? Strong electrolyte  Weak acid? Weak electrolyte  Base?  Weak base? Weak electrolyte o Other? Non electrolyte  Metathesis or exchange reaction o Metathesis (exchange) comes from a Greek word that means to transpose o It appears the ion in the reactant compounds exchange, or transpose, ions o Well look at precipitation reactions and neutralization reactions  Neutralization reactions o Generally, when solutions of an acid and a base are combined, the products are a salt and water o Salt formed is an ionic compound formed from the cation of the base and the anion of the acid o Molecular, ionic, and net ionic equations o All strong electrolytes  Gas-Forming acid-base reactions o Some bases other than OH- react with acids to form gases that have low solubility in water o Carbonates and bicarbonates react with acids to form CO gas2and are used as neutralizers  Precipitation reaction- formation of an insoluble slid when two electrolyte solutions are mixed  Precipitate- solid that forms when solutions are mixed  Concentration units o Two solutions can contain the same compounds but be quite different because the proportions of those compounds are different o Molarity is one way to measure the concentration of a solution o Molarity (M) = moles of solute / volume of solution in liters  Dilution- solution made in a concentrated form called stock solutions are often diluted o Mc*Vc=Md*Vd o We can use other units of volume for this equation  Titrations o To determine the concentration of a solute in solution, chemists use titrations o Titrations involve combining a solution with a reagent of known concentration (standard) o Can do acid base, precipitation, and redox reactions this way o We use an indicator to indicate the endpoint of the reactions  Phenolphthalein T est 2: Chapter 7  Electromagnetic Radiation helps us understand the structure of atoms. It describes energy carried through space and has wave like characteristics  Wavelength (λ): distance between corresponding points on adjacent waves  Amplitude- maximum and minimum height describes wave intensity  Node- points of zero amplitude  Frequency (ν) is the number of waves passing a giben point per unit of time. Unit is hertz. Hz=cycles/sec=1/sec=s -1 o Inversely proportional to wavelength o Short wavelength = high frequency o Long wavelength = short frequency  Speed of light relationship o Electromagnetic radiation travels at the same velocity (speed m/s) the speed of light is c=3.00x10 m/s o The relationship between wavelength, frequency and the speed of light is c=λν  Electromagnetic Spectrum o Comprised of all frequencies of light o Divided into regions according to wavelengths of radiation o Radio waves- communication o Microwaves- cooking, molecular rotation o Infrared- heat sensing equipment, night scopes, burglar alarms o Visible- vision o Ultraviolet- bacterial sterilization o X rays- medical imaging o Γ rays- nuclear radiation  History of Atomic Theory o Late 1800's  Matter and energy believed to be distinct  Matter- made up particles  Energy- light waves o Beginning of 1900's  Several experiments proved this idea incorrect  Experiments showed that electrons acted like tiny charged particles in some experiments and waves in other experiments o Additional model  The wave model doesn’t explain all radiation such as  Black body radiation- how an object can glow when its temperature increases  Photoelectric effect- how electrons are emitted from matter after energy absorption  Emission spectra- the intensity of radiation of each frequency emitted when heated  Energy is quantized o Max Planck and albert Einstein (1905)  Electromagnetic radiation is stream of small packets of energy  Packets were called quanta of energy of photons  Each photon travels with velocity=c  Waves with frequency=ν o Energy of photon of electromagnetic radiation is proportional to its frequency  Energy of photon E=hν  h=Planck's constant= 6.626x10 -34Js o We can relate the wavelength of light to the energy in one photon or packet of that light  c=λν  E=hν o Einstein's photoelectric effect  Shine light on metal surface  Below certain frequency (ν)  Nothing happens even with very intense light (high amplitude)  Above certain frequency  Number of electrons ejected increases as intensity increases; kinetic energy (KE) of ejected electrons increases as frequency increases  Photosynthesis o If you irradiate plants with infrared and microwave radiation (lower energy than visible light) then no photosynthesis regardless of light intensity o If you irradiate plants with visible light photosynthesis occurs, more intense light now means more photosynthesis  Radiation Sources o Monochromatic- radiation at a single wavelength (laser) o Continuous- radiation through whole array of wavelengths (the sun) o Line spectra- line for a given gas (like hydrogen) ak a atomic spectrum of emission spectrum  Electronic Structure of the Atom o Understood by studying light o Study of light absorption  Electrons absorbs energy  Moves to higher energy, excited state o Study of light emission  Electron loses photon of light  Drops back down to lower energy, ground state  Bohr's Model of the Atom o Historical plum pudding model of the atom o Next, a new atom model with protons, electrons, and neutrons (Rutherford and Chadwick) o Finally Bohr realized electrons are moving at great speeds within the atom. Electrons are fixed in orbitals. Transitions in orbitals describes absorption and emission of light o Electrons in an atom can only occupy certain orbits (corresponding to certain energies) o Electrons in permitted orbits have specific allowed energies; these energies will not be radiated from the atom o Energy is only absoorbed or emitted in such a way as to move an electron from one allowed energy state to another; the energy is defined by E=hν o Bohr's model only worked for atoms and ions that had a single electron (hydrogen) o Energies of electron in atoms were quantized and characterized by the principal quantum number, n o Energy is involved in moving electron from one level to another  Wave Behavior of Matter o De Broglie suggested particles have wavelike properties o An electron moving about a nucleus has associated with it a particular λ which depends on mass and velocity o λ=h/mν  The uncertainty principle o Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known  (���x)(���mv)≥h/4Π o In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself o To see the electron position we must probe with radiation and change its position or velocity or both o We can only specify the highest probability of an electron in an orbital  Quantum Mechanics o Schrödinger developed a method to incorporate both the wave and particle mature of matter o Called a wave function equation it describes the energy of electrons that are quantized o The probability of finding an electron in a point in space (the orbital) is found by the square of the wave function o Each orbital can hold two electrons  Electron Cloud descriptions o Electron dot picture=snapshots  Lots of dots shown by large amplitude of wave function gives high probability of finding electrons o Electron density is the electrons charge packed into given volume  High probability occurs with high electron density or large electron density  Low probability occurs with low electrons density or small electron density  Quantum Numbers o Solving the wave equation gives a set of wave functions or orbitals and their corresponding energies. Quantum numbers are the shorthand for keeping track of electrons characteristics o Each orbital describes a spatial distribution of electron density or probability o An orbital is described by a set of four quantum numbers  n=principle quantum number  All orbitals with same n are in same shell  Allows values: positive integers from n=1 to infinity  Describes the size of orbital (distance of electron from nucleus as n increases the electron moves further from nucleus) and total energy of orbital (as n increases energy increases and becomes less stable)  Orbitals with the same n are called a shell  l=angular momentum quantum number  Divides shells into smaller groups called subshell  Allowed values: 0-infinity (n-1)  Letters s ,p, d, f correspond to numerical values of l  0 and s  1 and p  2 and d  3 and f  Describes the shape of the orbital where the electron is found  A set of orbital with the same values of n and l are called subshells  M=lagnetic quantum number  Divides subshells into individual orbitals  Allowed values from -l to +l  There are 2*l+1 different values  Describes the orientation of orbital in space  To designate a specific orbital you need three quantum numbers  M sspin quantum number  Describes the behavior of an electron in a magnetic field  From behavior of electron in magnetic field  Electron acts like a top  Spinning charge makes electrons behave like a magnet  Two possible directions of electron spin  Up and down  North and south  Orbital Designations o Based on the first two quantum numbers o Two electrons can go in each orbital  Pauli Exclusion Principle o No two electrons in the same atom can have exactly the same energy. This means no two electrons in the same atom can have identical sets of the four quantum numbers. Only two electrons can occupy a single orbital and they must have opposite spins o Two electrons in the same orbital have different spins  Spins paired- diamagnetic  Sample not attracted to magnetic field  Magnetic effects tend to cancel each other o Two electrons in different orbital with same spin  Spins unpaired- paramagnetic  Sample attracted to a magnetic field  More unpaired electrons lead to a stronger attraction to a magnet  Ground state electron arrangement notation o Orbital diagram represents electrons in orbitals o Electron configuration-distribution of all electrons in an atom  List subshells that contain electrons  Indicate their electron population with superscripts  Hund's Rule o Lowest energy is attained when the number of electrons with the same spin is maximized in each orbital of same energy o So, place one electron in each orbital before pairing up electrons in a single orbital  RULES TO FOLLOW o Aufbau principle o Pauli exclusion principle o Hund's rule  Chemical reactivity and electron configuration o Periodic table arranged by chemical reactivity  Depends on outer shell electrons (highest n)  Each row is different n o Core electrons  Inner electrons are those with n<n(max)  Buried deep in atom  Abbreviated electron configurations o Truncate core electrons by representing them with the noble gas of the row before o Only outer shell electrons are described in our electron configuration notations o This reminds us the outer shell electrons react with atoms of other elements o Valence shell- the outer shells, these electrons are involved in bonding  Exceptions o Cr, Cu, Ag, Au o Exceptions occur as exactly filled and exactly half filled subshells have extra stability o Promote one electron into ns orbital to gain this extra stability Chapter 7 Part 2 Periodic Trends  Electron cloud o Electron dot picture=snapshot  Lots of dots shown by large amplitude of wave function  High probability of finding electrons o Electron density- how much of electron charge packed into given volume  High probability- high electron density or large electron density  Low probability- low electron density or small electron density  Periodic trends o Chemical and physical properties of elements  Vary systematically periodic table location o To explain, must first consider amount of positive charge felt by outer electrons (valence electrons)  Core electrons spend most of their time closer to nucleus than valence (outer shell) electrons o Effective Nuclear Charge  Z =Z-S eff  Z is atomic number  S is the number of electrons between nucleus and electron in question  Highest is top right, lowest bottom left o Atomic Size  The closest distance separating the nuclei of two atoms that are non-bonding to the nonbonding atomic radius  The distance separating the nuclei of atoms when they are chemically bonded to the bonding atomic radius  Bottom left is the largest, top right is the smallest o Ionic Size  Ion size dependent on sign of charge and magnitude of charge  Increase down column, decreasing across row  Anions(-) large than parent atom  Same Zeff, more electron  Radius expands  Cations(+) smaller than parent atom  Same Zeff, less electons  Radius contracts  Isoelectronic series- ions possess the same number of electrons; the radius of the ion decreases with increasing nuclear charge o Ionization energy  The amount of energy required to remove an electron from the ground state of a gaseous atom or ion  First ionization energy- energy required to remove first electron  Second ionization energy- energy required to remove second electron  Greatest energy in the upper right, smallest in the lower left  As ionization energy increases, it is more difficult to remove an electron  Every successive removal pulls an electron away from an increasingly more positive ion  A large increase in ionization energy occurs when electrons are removed from the noble gas core of elements  Electrons are always removed from the shell with the largest "n" first  When electrons are added they are added to the orbital with the lowest "n" o Electron Affinity  Energy change that accompanies addition of electrons to a gaseous atom  Energy is typically released when an electron is added  The greater the attraction, the more negative the EA value  So, electron affinity measures the ease with which an atom gains an electron  Greatest in upper right, smallest in lower left  Added electrons must go in p-orbital, not the s. electron is farther from nucleus and feels repulsion from the s-electrons  Group A has a half filled p subshell, adding an electron increase the electron-electron repulsion.


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