Final Exam Study guide for ACS Gen Chem
Final Exam Study guide for ACS Gen Chem Chem 143
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This 46 page Study Guide was uploaded by Cassidy Zirko on Thursday May 5, 2016. The Study Guide belongs to Chem 143 at University of Montana taught by Dr. Cracolice in Spring 2016. Since its upload, it has received 153 views. For similar materials see General Chemistry 2 in Chemistry at University of Montana.
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Chapter 1: How Do Scientists Use Algebra to Reason and Calculate? Algebra thought of as arithmetic in which letters are used to represent variables Variables quantities that may have different values, can appear in equations Equations represents a mathematical relationship among variables Direct Proportionality two variable quantities in which when one variable is manipulated, the other is manipulated in the same way Ex. C ∝ d, circumference is proportional to diameter The origin of a direct proportionality is always (0,0) on a graph k proportionality constant Inverse Proportionality two quantities that are related in such a way when one variable is manipulated, the other decreased by the reciprocal of that amount. Graph of inverse proportionality is not a straight line Multiple Proportionality when one quantity is proportional to each of two or more variables then it is proportional to the product of those variables. Quantity Algebra a way to use conversion factors to fine a final answer Conversion factors a direct proportionality which can be expressed in fraction form and is used to convert from one unit to second Chapter 2: How do Scientists Express Ratios and Measured Quantities? Ratio value of a partial quantity divided by the value of the whole Percentage not associated with a specific number, Percent references a specific number ppm parts per million, ppb parts per billion Important powers of 10: Kilo 10 , Centi 10 , Milli 10 , Micro 10 , Nano10 9 Weight measure of attractive force that it experiences in a gravitation field Mass measure of resistance to change in a state of motion, expressed in grams or a subset of gram, ex. kilogram Meter distance that light travels in a vacuum Measuring Volume determining, the height, width and length and then using the correct formula to find the space that an object takes up Liter, L=1 dm , 1mL=1cm 3 Temperature in C Chapter 3: What is Science? Hypothesis proposed explanation for a pattern found in nature Parts of the scientific method: 1) as a question about the natural world, 2)take measurements, qualitative or quantitative, 3) Look for a pattern in your data that expresses some sort of relationship 4) propose a hypothesis, based on an unobservable casual entity Unobservable Casual Entity an invisible something that is known to exist which is the underlying cause of something or some pattern in nature Experiment purposeful manipulation of the natural world that wants to test a specific hypothesis Pressure the strength on an area over which it is spread, P=force/area Boyle’s Law as volume decreases, pressure increases, and as volume increases, pressure decreases as long as temperature remains constant, P V =1 V1 2 2 LawStatement that summarizes a pattern found in nature 760 mm Hg=760 torr=1.01325 bar= 1 atm Force is equal to mass times acceleration, F=MA Kinetic Molecular Hypothesis matter is composed of miniscule particles that are in constant motion Charles law as the volume increases, temperature increases, and as the volume decreases, temperature decreases, as long as the pressure remains constant, T V =2 V1, a1wa2s convert temperature to kevin Theory a hypothesis that is found to explain a broad group of related events Combined Gas Law as pressure and volume increases, temperature decreases, and as pressure and volume decrease, P 1 1= P 2 2 temperature increases, T T 1 2 Chapter 4: How do Scientists Make, Report and Work with Quantitative Measurements? Precision reproducibility of results Accuracy a measured closeness to correct values Significant figure convention universally agreed upon custom about expressing the accuracy of measure quantities Uncertain Number the last digit in a number Zeros are place holders, a number with many zeros in ambiguous because you don’t know where the uncertain digit is Rounded off a way to correctly represent the uncertainty of the measurements that they are based on Adding/ Subtracting round to least amount of decimal places Multiplying/ Dividing round to lowest number of sig figs Sig figs convention applies to only measured quantities Absolute Terms the amount of one thing that weights exactly the same as the same amount another thing Density the mass per unit of volume, D=m/v 1in 2.54 cm, 1lb= 453.59237g, 1gal= 3.785411784L, 1ft=12in, 1 yd= 3 ft, 1ml=5280ft, 1lb=16oz Chapter 5: How do Scientists Classify Matter? Petrochemical substance that is made from petroleum or other closely related substances Matter describes physical materials in general Liquid Form indefinite shape, definite volume Properties something matter has which helps categorize it Solid definite shape, definite volume Gas/vapor indefinite shape, indefinite volume Solid Phase water molecules are stuck together so you get a solid shape Liquid Phase molecules would be able to move around but would still be stuck together Gas Phase water molecules would be free to move around themselves independent of another Pure Substance a substances with no dissolved substances, a kind of matter Compound a substance that can be decomposed into simpler substances Element pure substances that can’t be decomposed into another pure substance by ordinary chemical means Physical Change When a substance changes from one state to another, a change that can be observed Physical Property attributes of a substances that can be measured or observed without changing the substance Mixture combo of two or more substances that doesn’t have a fixed composition Solution homogenous mixture or heterogeneous mixture Homogenous Mixture same throughout, both in composition and stage Heterogeneous Mixture different throughout, Bot composition and stage visibly different stages Chemical Change one substance is destroyed to create a new substance Chemical Properties chemical changes possible from any substances, list of chemical changes that are possible Chapter 6: Is Matter Continuously Divisible? Law of Conservation of Mass the mass of the reactants is a chemical change is equal to the mass of the products Law of Definite Proportions the percent by mass of each of the elements that make up a compound is definite, fixed or invariant Law of Multiple Proportions when 2 substances react to form more than one product, amount of one that completely reacts with the fixed amount of the other are in a ratio of small, whole number Atomic Hypothesis matter is not infinitely divisible (the smallest particle is an atom), Atoms will always and forever exist never changing, atoms of the same element are identical, atoms of an element could combine with atoms from other elements creating compounds, atoms could combine in many ratios creating different compounds. Atomic Theory Provided support for and intertwines with the Kinetic Molecular Theory Molecule smallest piece of a substance that retains identity of that substance Relative Atomic Masses scientists use a relative scale to express the weight of atoms Atomic Mass units (u) used by the scientists who investigate the physics of chemistry Chapter 7: How are Electricity and Matter Interrelated? Law of Combining Volumes volume ratio in which gases reacted were always a ratio of small, whole numbers Avogadro’s Hypothesis law could be explained if equal volumes of gases contained an equal number of molecules, ended up being wrong Plum Pudding Model atom was made up of the large sphere of positive charge that contained many small negatively charged particles as needed to have total charge of the atom be zero Electrons negatively charged particles Oil drop Experiment pair of electrically charged plates to provide uniform electric field – oil drop fell into space between plates – rose and fell based on electrical potential Planetary Model of the Atom idea that electrons orbited the nucleus like planets Chapter 8: What is the Nucleus Made of? Rutherford discovered proton with alpha particles Proton the first subatomic particle, has positive charge Neutron a neutrally charged subatomic particle Atomic Number the Number of Protons that an element has, Z Isotope atoms of the same element that have different masses, different number of neutrons Mass Number, A total number of protons and Neutrons in the nucleus Mass number= number of protons+ number of neutrons Nuclear Symbol the way an isotope is represented 1u= 1/12 the mass of a carbon 12 atom 1u= 6.02*10 g23 Atomic Mass of an Element average mass of all atoms of that element as they occur in nature Chapter 9: How do Chemists Express Names and Formulas of Substances? Part 1 Chemical Formulas Represent the composition of elements and compounds Chemical Nomenclature a system by which substances are named Molecule the number of atoms of the element in the smallest particles of the substances Diatomic elements: Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine Polyatomic Molecules Sulfur S , 8nd phosphorous. P 4 Compounds Made of atoms or ions of different elements Metalloids Semiconductors that border the stair step line on a periodic table Binary Molecular Compounds compounds formed by two molecules or a metalloid a nonmetal Ions when an atom gains or loses one or more electrons Cations atoms loses an electron and remains with a positive charge Anion when there is a gain of an electron and the element has a negative charge Ionic Compound attraction of oppositely charged particles that cause a compound to stick together Monatomic anion named with element modified with –ide and followed by ion Transition Metal Ions have a charge written in roman numerals following the name and then following with ion Acid molecular compound that reacts with water to yield a hydrogen ion and an ion Oxoacid acid that contained a hydrogen, a nonmetal and oxygen that has all of the hydrogen ions removed 5 Important acids: Carbonic Acid H CO 2 Ni3ric Acid HNO , Phos3horic Acid H PO , Sul3uri4 Acid H SO , Chlor2c 4 Acid HClO 3 Chapter 10: How do Chemists Express Names and Formulas of Substances? Part 2 Per, ic has one more oxygen than the –ic acid ic ending changes to –ous when the acid has one less oxygen then the –ic acid Hypo, ous, 2 fewer oxygen’s than ic ending Hydro, ic, no oxygens Acids ending in –ous create anions ending in –ite Acids ending in –ic create anions ending in –ate Polyproptic Acids acids with more than 1 hydrogen Total Dissociation acid had all of the hydrogens removed + Important: Ammonium ion NH , hydr4xide ion OH Ionic Compound compound made up of ions Hydrate ionic compound that contains water molecules Anhydrous Compound without water Chapter 11: How do Chemists Weight and Count Particles? Formula Mass based on the chemical formula of the compound Molecular Mass average mass of molecules (or formula units) comparted with the mass of a carbon12 atom which is exactly 12 atomic mass units Mole the amount of any substance that contains the same number of units in the number of atoms in exactly 12 grams of a carbon 12 atom 23 Avogadro’s Number, N a 6.02∗10 atoms, molecules or formula units Molar Mass mass in grams of one mole of a substance Percent Composition the percentage by mass of each element in the compound Chapter 12: How are Chemical Formulas Determined? Combustion Analysis a compound burned for the purpose of analyzing its composition Empirical Formula simplest formula of a compound Procedure for finding empirical formula 1. Find masses of different elements in a sample of a compound 2. Convert the masses into moles 3.Express the moles of atoms as the lowest possible ratio of whole number integers 4. Write the empirical formula, using the number of each atom in the integer ratios as the subscript Procedure for finding molecular formula 1. Determine the empirical formula 2. Calculate the molar mas of the empirical formula 3. Determine the molar mass of the compound (usually given) 4. Divide the molar mas of the compound by the molar mass of the empirical formula to get the number of empirical formula units per the molecular given 5. Multiply the empirical formula by the factor from step four and write the molecular formula Chapter 13: How is a Chemical Change Expressed Symbolically? Ways to recognize a chemical change: color change, formation of solid, formation of a gas, expulsion/absorption of heat or production of light State symbols indicate the state of each compound and is sometimes omitted Aqueous solution a substance that is dissolved into water Qualitative Equation a chemical equation that is not balanced Balanced placing a coefficient in front of one or more of the formulas indicating it is used more than once Procedure for writing an equation 1. Write a qualitative description of the reaction. In this step write the formulas based on the reactants and products 2. Quantify the description by balancing the equation. Done by adding coefficients Combination reaction, synthesis reaction where two or more substances combine to form single product, A+X AX Decomposition reaction a compound that breaks down into simpler substances, AX A+X Single Replacement Reaction where one element will replace a second within a compound, A+BX AX+B Oxidization means loss of electrons during electron transfer, products are always C2 and water Precipitate when solutions of two compounds are mixed, positive ion form one compound and the negative ion from the second compound will combine to form a solid compound that will settle at the bottom of the container. Double Replacement Reactions ions of two reactants will appear to change partners in the products, AX+BY AY +BX Base substance that contains a hydroxide ion Neutralization Reaction when an acid is added to an equal amount of base, each hydrogen ion reacts with each hydroxide ion to form water HX +(aq) H(aq) MX2 (aq) Acid + Base Water + Salt Chapter 14: How are Masses Related in a Chemical Change? Part1 Stoichiometry given a macroscopic quantity of one species before and after a chemical change, you can find the quantity of a second species Mass Molar mass in grams mole Chapter 15: How are Masses Related in Chemical Change? Part 2 Limiting Reactant the reactant that is completely used up by the reaction Excess Reaction the amount of the reaction which is still left over after the reaction is complete Procedure for Limiting Reactant 1. A) the reactant that yields the smaller amount of product is the limiting reactant B) the smaller amount of product is the amount that will be formed when all of the limiting reactant is used up 2. Calculate the amount of excess reactant that is used by using that limiting reactant and converting it to the excess reactant 3. Subtract the amount of excess reactant that was given in the problem initially form the amount calculated in step two. This difference will give you the amount of the excess reactant that is unreacted. Chapter 16: How are Masses Related in a Chemical Change? Part 3 Ideal yield amount of product formed from the complete calculation Actual Yield measured quantity determined by experiment Percentage yield actual yield expressed as a percentage of the ideal yield actual product yield yield= ∗100 ideal product yield Chapter 17: What Models Represent the Gaseous State? Law of Combining Volumes reacting volumes are always in the ratio of small whole number if the volumes are measured at the same temperature and pressure Avogadro’s Law equal volumes of all gases at the same temperature and pressure contain the same number of molecules, V ∝ n, n= number of moles Ideal Gas Equation PV=nRT Ideal Gas Law volume is directly proportional to the number of moles and the temperature while being indirectly proportional to pressure Equation could also be PV=(m/MM)RT Gases can be compresses, they can expand, gases have low densities due to the large amount of space between each of the individual molecules, gases may be fixed in volume, Gases exert constant pressure on the walls of its container uniformly in all directions Ideal Gas model allows use to be able to visualize the nature of gas by comparing it with a physical system 1. Gases consist of particulates moving at any given instant in a straight line 2. Molecules collide with each other and with the container walls without a loss of total kinetic energy 3. Gas molecules are very widely spread out 4. Actual volume of molecules is negligible to the space they occupy 5. Gas molecules are independent; attractive forces between them are negligible Ideal Gas a model of a gas constructed of identical particles that occupy no sufficient volume and exert no forces on one another Dalton’s Law of Partial Pressure total pressure exerted by a mixture of gases is the sum of the partial pressures of gases in the mixture Real gases are different than ideal gases because the actual volume of molecules are negligible and that gas molecules behave as independent particles are only approximate Real gases approach the characteristics of a real gas at low pressure and high temperatures 2 P+a n (V−nb =nRT van de Waals Equation [ (V) ] a= adjusts for the attractive forces among that particles, b corrects for the volume of molecules closer a and b are to zero, the more the gas will behave like an ideal gas Chapter 18: How are Measurable Quantities Related in a Chemical Change When One or More of the Species is a Gas? Density of a mass over volume Molar Volume (MV) the volume of a gas occupied by one mole of a gas molecule s MV ≡ V = RT n P Molar volume depends on temperature and pressure and is not constant Solved in two based steps, convert to same temperature and pressure and then solve stoichiometry or vice versa Molar volumes will always cancel if given and wanted are at the same temperature and pressure Chapter 20: How is Energy Related Systems and Chemical Change? Energy the ability to do work or transfer heat Work (w) work done to move an object is the product of force and the distances that the object moves Heat a form in which energy is transferred between substances with different temperature s Temperature a measure of average energy of particles in a substance Exothermic reactionchemical change that releases energy to its surroundings Endothermic Change chemical change that absorbs energy from its surroundings Potential Energy depends on its positon in a field where forces of attraction and/or repulsion are present There are always electrostatic forces between charged particles The minimization of energy is one driving forces that causes chemical reactions to occur Kinetic Energyany moving object, mechanical energy Calorie (Cal) used by nutritionists to measure food energy Thermodynamic system the portion of the universe under consideration Open system has imaginary boundary, matter can move in and out of the system Close System has physical barriers that keep quantity of matter in the system State function Property whose volumes are determined only by the state of the system at a given moment Internal energy energy of a system that results from sources other than the influence of external forces Law of Conservation of Energy energy can neither be created nor destroyed E, energy of a system state function only depends on the initial energy and the final energy states of the system Second Law of Thermodynamics heat flows into or out of a system, internal energy of the system changes by a quantity equal to the quantity of heat transferred First Law of Thermodynamics internal energy of a system changes by an amount equal to the sum of the heat that flows into or out of the system and the work done by the system w ΔE Δ Change of internal energy of the system , q heat flow, w wo signifies final minus initial Heat flow, q, is not a state function Enthalpy, H H≡∆E+P∆V Pressure volume of work work associated with chemical change in a liquid and gas phase, Enthalpy is a state function ∆ E=q−P∆V Heat of Reaction, Enthalpy of Reaction, H when a system gives off heat to surroundings Thermochemical Equation a chemical equation that includes change in energy by having an enthalpy value Make sure to use state symbols because enthalpy will change based on state of reactants and products Endothermic reaction, enthalpy is positive and a reactant in the equation Exothermic reaction, enthalpy is negative and a product in the equation Chapter 21: What are the Relationships Among Heat, Energy, Temperature and Mass for a Pure Substance ? Change form liquid to solid heat removed Liquid to gas heat is added ∆ H vap = enthalpy of vaporization, q ∝ m q= H vap Units for heat of vaporization, kJ/g Enthalpy or Heat of Condensation the energy required to condense a liquid into a gas at its boiling point Same values as heat of vap but negative Enthalpy of Fusion energy required to melt one gram of that substance Fusion changing a solid to a liquid by heat, melting q=m∗∆ H Heat of Solidification liquid to a solid, numerical equal to heat of fusion but negative, fus Specific Heat, c heat flow required to change the temperature of one gram of a substance by one degree centigrade, q c= , m ∆T Procedure: Calculating total Heat flow for a Change in Temperature and State 1) sketch a graph, marking starting and ending points of each stage as well as the starting and ending points for each problem 2) calculate heat flow, q for each sloped and horizontal portion of the graph between the starting and ending points 3) Add the heat flows calculation in step 2. Make sure that all of the units are either kJ or J for all the numbers being added. Chapter 22: How are Heat Energy Change Determined? Calorimetry where heat transfer is used to determine specific heat or the heat of reaction Calorimeter Constant, Kcal heat flow per unit of change in temperature Enthalpy of formation change in enthalpy for the reaction in which pure, stable elements react at a pressure of 1 bar if they are a gas to form one mole of product also at 1 bar if it is a gas Liquids at STP bromine and mercury Gases at STP hydrogen, nitrogen, oxygen fluorine, chlorine, and all 8A/18 elements Diatomic elements hydrogen, nitrogen, oxygen, fluorine, bromine, chlorine, iodine Carbon standard form graphite or diamond Hess’s Law if a reaction is carried out in a series of steps the enthalpy for the total reaction will be equal to the sum of the enthalpy for each step Chapter 23: How are Electrons Distributed within an Atom? The Bohr Model Electromagnetic Radiation form of energy that consists of both electric and magnetic fields Electromagnetic Spectrum gamma rays, x0rays, ultraviolet light, infrared light, microwaves, radio waves, visible light Speed of Light 3.00*10 meters/s Wave equation c=λν , the speed of light is equal to the wavelength multiplied by the frequency Continuous spectrum band of colors that results from the electromagnetic radiation emission over a range of wavelengths Line Spectrum spectral lines that appear when light emitted from a sample is analyzed in a spectroscope Discrete Lines indicates that elements are individually distinct Photon particle of light Wave particle Duality all matter and energy possess both wavelike and particle like properties, photons can behave like waves and particles −34 Plank’s Constant 6.626∗10 J ∙s speicifies that light is not a continuous flow of radiation but a series of individual photons E=hv, energy is equal to plank’s constant time frequency Photoelectric effect light can cause ejection of electrons from some metal surfaces Quantized limited to specific values, it may never be between 2 values Continuous can have any values between any two values, there is an infinite number of acceptable values Quantized Energy Levels at any instant, electrons may have one of several possible energies, but at no time can they have energy between them Quantum Jump/leap when an electrons moves between orbits Electron can orbit at a certain specified distances but not found in between those distances Ground state condition when all electrons in an atom occupy the lowest possible energy level Excited State condition at which one or more electrons in an atom has an energy level above the ground state, the electron is unstable −R ∗H E= 2 n R H Rydberg constant for hydrogen= 2.1798722∗10 −18 J Chapter 24: How are Electrons Distributed within an Atom? The Quantum Mechanical Model h λ= whereλ=¿ de Broglie wavelength, hPlanks constant, mmass of particles, vvelocity, describes the mν wavelength of any particle Heisenberg proved that it is impossible to simultaneously know both the position and velocity of an electron (Δx (mΔυ)≥ h Heisenberg Uncertainty Principle 4π Δ x−uncertianty∈position,Δν−uncertiantyof velocity ,h−planksconstant The relationship in the Heisenberg uncertainty principle tells us that the product of uncertainty in position multiplied by the uncertain in velocity is greater than a specific constant Quantum Mechanical Model of the atom an atomic concept that recognizes four quantum numbers by which electron energy levels may be described Probability Density probability that a particle is found in a specific three dimensional region in space Seventh level is the highest occupied by the ground state electron Sublevelthe levels into which principle energy levels are divided according to the quantum mechanical model of the atom usually specified as the letters s, p, d and f l=0,1,2,(n−1) Sublevel specified by , which lines up with a specific orbital (s, p, d, f) N= principle quantum number, the period in which the element is in. n=1, 2, 3… Electron orbital quantum numbers m l , m =ll…0…+l l=0 m =0 l , 1 orbital for every s sublevel l=1 m l−1,0,1 3 orbitals for every p sublevel l=2 m l−2,−1,0,1,2 5 orbitals for every d sublevel l=3 m =l3 , 2 1, 0, 1, 2, 3 7 orbitals for every f sublevel Pauli Exclusion Principle at any time an orbital my be unoccupied, occupied by one electron or occupied by 2 electrons M spin is either + ½ or 1/2 because electrons have to be spinning the opposite direction s Chapter 25: How are Electrons Distributed within an Atom? Electron Configuration ground state distribution of electrons among the orbitals of gaseous atoms At ground state, electrons fill lowest energy orbitals available first, no orbital can have more than 2 electrons Group 1A/1 and 2A/2 (s sublevel) are the highest occupied sublevels P orbitals fill across 3A/138A/18 D orbitals fill across groups 3B 2B F orbitals fill across lanthanide and actinide series Nobel Gas Core short hand notation used in electron configuration when the chemical symbol of a Noble gas is in square brackets which substitutes for the electron configuration of the element List sublevels in order of increasing energy, establish the number of electrons in the highest occupied energy sublevel of the atom Procedure: Writing Electron Configurations 1. Locate element in the periodic table. From position, identify and write the electron configuration of the highest occupied energy sublevel 2. To the left of the highest occupied energy sublevel, list all lower energy sublevels in order of increasing energy 3. For each filled lower energy sublevel write the superscript of the number of electrons that fill that particular sublevel 4. Confirm the total number of electrons is the same as the atomic number of the element. (adding together all the superscripts) Orbital Diagram diagram that shows how many electrons are in each orbital Hund’s rule the most stable arrangement of electrons is the one that has the maximum number of unpaired electrons Procedure: How to Write an Orbital Diagram 1. Write electron configuration using the noble gas core 2. Draw boxes to represent orbitals, 1s orbital (1 box), 3 p orbitals (3 boxes), 5 d orbitals (5 boxes), and 7 f orbitals (7 boxes) 3. Fill each with electrons. Fill boxes in each sublevel one at a time before doubling up as necessary until all electrons are accounted for Valance electron the highest energy s and p electrons in an atom, which determine the bonding characteristics of an element Lewis Dot Symbols, Electron Dot Symbols shows the number of valance electrons for an element Only the noble gases have 8 total electrons Chapter 26: What Causes Trends in Properties of Elements? Franklard discovered the patterns among the periodic table Valence combing power of an element Effective Nuclear charge the actual nuclear charge subtracted from the shielding constant, Z ieffeases across the period table, remains constant going down the periodic table Atomic Radius average distance between the nucleus of an atom of the element and the outer limits of electron cloud Atomic Orbital probability of finding an electron, not definite boundary Atomic size increases from right to left across a period and increases from top to bottom, smaller atoms are towards the upper right corner and the larger are toward the bottom left corner Isoelectric Series group of atoms and/or ions that have the same number of electrons Ionic Size size of ions increases down a group in in the periodic table, size of ion in an isoelectric series decreases with increasing atomic number. Monoatomic cations are smaller than the original atom. Monoatomic anions are large than the original atom Ionization energy energy required to removed one electron from a neutral gaseous atom of an element Increasing effective nuclear charge negatively charged valance electrons are attracted more strongly to the nucleus meaning that the atom has more energy to remove electrons First Ionization energy increases from left to right across a period and from the bottom to the top. The highest first ionization energy is in the upper right hand corner Electron Affinity change in energy that occurs when an electron is added to a gaseous atom or ion Trend of electron affinity generally increase from right to left on any row, atoms with the most electron affinities are on the right side Metal can lose one or more electrons and become positively charged Nonmetal lack the metal qualities Metalloids/semimetals properties of both metals and nonmetals Summary: Metallic Character increases from right to left across the periodic table Chemical families groups with properties in common Reactivity tendency to react with other elements to form compounds Alkali metals group 1A/1, easily valance electrons, density increases going town the table Alkaline Earth metals group 2A/2, reactivity increases going down the table Halogens group 7A/17, very easily gains electrons, reactivity decrease down the table Noble gases group 8A/18, very unreactive, high ionization energies, full outer orbital Chapter 27: How do Atoms Combine to Form Ionic Compounds and Molecules? Chemical bond forces that hold atoms together to form molecules or polyatomic ions, holds atoms together in metals, or that hold oppositely charge ions together to form ionic compounds The outer most electrons are the valance electrons Cation atom loses one or more electrons, creating a positive charge Anion atom loses one or more electrons, creating a negative charge Ionic Compounds compounds made up of ions, or solutions of ionic compounds Crystal ions arranged so potential energy resulting from the attractions and repulsions between atoms are at a minimum Ionic Bonds strong electrostatic forces that hold the ions in a fixed position in the crystal When substances melt or dissolve, the crystal structure is destroyed Molecular compounds ultimate structural unit is an individual particle Covalent bond shares one or more pairs of electrons between 2 or more nonmetals Electron cloud/Charge density two electrons that are concentrated in the region between 2 nuclei Overlap atomic orbital of separated atoms Electrons link nuclei together they are the glue that bonds atoms to each other Lone pairs unshared electron pairs Octet rule outer most orbital that has to have 8 total elections, the stability of the noble gas formation minimizes the energy associated with the configuration Types of bonds single, double, triple, multiple Bond Strength energy required to break a bond Nonpolar bond a bond in which bonding electrons are shared equally Polar Bond a bond in which bonding electrons are shared unequally Electronegativity ability of an atom of that element in a molecule to attract bonding electrons pairs to itself High electronegativity’s elements have strong attraction for bonding electrons The greater the difference in electronegativity’s the more polar the bond is Electronegativity increases from left to right, fluorine is the most electronegative element Chapter 28: How can the Two Dimensional Arrangement of Atoms in Molecules be predicted? 1) Determine the number of valance electrons 2) Determine the central atom 3) Draw the first draft of the Lewis Diagram 4) Adjust as necessary too many electrons, add more bonds to the central atom, two few electrons, add more lone pairs to the central atom 5) Analyze formal charge Constitutional Isomer made up of the same atoms connected differently, yields different properties When drawing a Lewis diagram of an ion, surround the drawing with brackets and the charge in the upper right hand corner Better to have a charge spread around the terminal atoms Hydrogens will always attach to oxygens in acids Formal charge charge that an atom would experience if bonding electrons are shared equally Formal charge= electrons on a separate atom electrons on the atom in the molecule, (count one electron for every bonded pair and two for every lone pair Best Lewis diagram will have the least number of formal charges Negative formal charges should be on the most electronegative atom Opposite charges occur on neighboring atoms Chapter 29: How are Covalent Bonds in Molecules Optimally Arranged and What are their Lengths and Strengths? Resonance structures series of Lewis diagrams that illustrate the different possible distributions of valance electrons A molecules structure doesn’t change, only the bonds Odd electron molecules, molecules with atoms that have more or less than the octet are the only atoms that will violate the octet rule Odd electrons will go on the least electronegative atom Free radicals species with odd numbers of electrons 3 period and higher can have up to 9 bonds because of the d orbitals Atomic size increases going down a group causing larger atoms to accommodate greater number of adjoining bonds Make sure that Carbon, nitrogen, oxygen and fluorine always fulfil the octet rule After first Lewis diagram is drawn, then proceed to draw any possible resonance structures Bond order number of electrons pairs shared in a covalent bond between two atoms Bond Length average distance between nuclei of two atoms in a molecule Bond length varies with atomic size, large atom, greater length of bond from it constant second atom Bonding Enthalpy amount of energy required to break a specified bond in one mole of a compound in the gas phase Enthalpy is equal to the bond energy of the reactants subtracted from the bond energy of the products General reaction, aA+bB dD+eE, ∆ H=[a ΣBEof A )+b (ΣBE of B )]−[d (ΣBEof D )+e ΣBEof E ] Chapter 30: How can Three Dimensional Arrangement of Atoms in molecules be predicted? Molecular Geometry shape of molecules Valance shell electron pair repulsion theory (VSEPR) theory use to describe molecular geometry, applies primarily to carbon, nitrogen and oxygen, electron pairs we draw repel each other in real molecules Electron Pair geometry arrangement of electron pairs Description Lone pairs Bonded Geometrical Name Bonding Angle pairs 2 electron pairs, 2 bonded atoms None 2 Linear 180 3 electron pairs, 3 bonded atoms None 3 Trigonal Planar 120 3 electron pairs, 2 bonded atoms 1 2 Angular 120 4 electron pairs, 4 bonded atoms None 4 Tetrahedral 109.5 4 electron pairs, 3 bonded atoms 1 3 Trigonal pyramidal 107.5 4 electron pairs, 2 bonded atoms 2 2 Bent 104.5 5 electron pairs, 5 bonded atoms None 5 Trigonal bipyramidal 120 and 90 5 electron pairs, 4 bonded atoms 1 4 See saw 120 and 90 5 electron pairs, 3 bonded 2 3 T shape 120 and 90 5 electron pairs, 2 bonded 3 2 Linear 180 6 electron pairs, 6 bonded atoms None 6 Octahedral 90 6 electron pairs, 5 bonded atoms 1 5 Square pyramidal 90 6 electron pairs, 4 bonded atoms 2 4 Square planar 90 Drawing Wedge and Dash diagrams 1) 2 atoms on the same plane as the page connect with a solid uniform line 2) Atoms that are behind the plane of the page, connect to central atom by a dashed line, width increases moving away from central atom 3) Atoms that are in front of the plane of page, connect to central atom by a wedge shaped solid line, increases in width away from central atom Predicting Molecular Geometries 1) Draw Lewis diagrams 2) Count total number of electron pairs around the central atom, bonding and lone pairs, the count only bonding paris 3) Determine electron pair geometry 4) Determine molecular geometry 5) Draw wedge and dash diagram of molecule Polar molecule one which there is an asymmetrical distribution of charge Nonpolar molecule cancellation of polar bonds, no overall net regions of positive or negative charge Chapter 31: How is Bonding Explained when Atomic Orbital Overlap Model Fails? Sigma ( σ ) bond orbitals overlap end to end, so only one of the lobes of the p orbital is involved in the covalent bond Hybrid orbitals orbitals of separate atoms change into hybrid orbitals when bonding to 2 or more atoms Lowest energy orbitals mix to form hybrid orbitals with energy in between energies of atomic orbitals Sp Hybrid orbital hybrid orbitals from mixing 1 s and 1 p orbitals, looks like a leaf of a fourleaf clover, line up with a liner electron pair geometry Sp 1 s orbitals and 2 p orbitals will mix and hybridize, lines up with trigonal planar electron pair geometry 3 Sp hybrid four half filled orbitals, forming 4 single bonds, 1 s orbital and 3 p orbitals, lines up with tetrahedral electron pair geometry Sp d hybrid one s orbital, 3 p orbitals, and 1 d orbitals are hybridized, lines up with trigonal biprymidal electron pair geometry Sp d hybrid orbitals when 1 s orbital, 3 p orbitals, and 2 d orbitals overlap to form 6 total bonds, lines up with octahedral electron pair geometry Cistrans isomer molecules that have 2 isomers that result from a particular atom being on either the same side (cis) or the opposite side (trans) of the double bond Sigma bonds allow for rotation around the bond. Double bonds can rotate in any way because of pi bond present P orbitals maximize the distance from the hybrid orbital Delocalized pi electrons are not restricted to a narrowly defined location, they belong to the entire molecule Chapter 33: What Particulate Level Interactions Govern the Liquid State? liquids can be compressed because there is no space between liquid particles Strong attractions hold liquid particles together Liquids have high densities, mass per unit of volume Liquids cant be mixed in a fixed volume Vapor pressure partial pressure exerted by gaseous particles Equilibrium Vapor Pressure gas space above the liquid is closed, vapor increases to a definite volume High vapor pressure, low intermolecular forces Heat of Vaporization energy required to change 1 mole of a liquid to a gas, while at constant temperature and pressure Heat of condensation energy released in opposite processes, vapor condenses to liquid phase Boiling point average kinetic energy of liquid particles is high enough to overcome the forces of attraction that holds the particles in the liquid form Viscosity ability of a liquid to flow, internal resistance to flow, partially based on intermolecular attractions Surface tension internal energy toward minimum surface, at surface all attractive forces are downward Dipole forces the forces between polar molecules Polar molecules have higher boiling point than nonpolar substances, have stronger intermolecular attractions Induced dipole forces attractions between substances with nonpolar molecular, caused by shifting electron cloud within molecules Hydrogen bond extra strong attractions between molecules that occur because of hydrogen bonds, only occur between nitrogen, oxygen and fluorine Solubility determined by strength of intermolecular forces within the solute, solvent and between solute and solvent, partial pressure of solute and temperature Chapter 34: What are the Characteristics of Phase Equilibrium? Kinetic Energy Distribution Curve range of kinetic energies at a certain temperature Equilibrium rates of change in opposite directions are equal Dynamic Equilibrium molecules constantly switching between liquid and vapor phases Equilibrium Vapor Pressure the partial pressure exerted by a vapor in equilibrium with its liquid phase at a given temperature Reversible changes changes that can occur in either direction symbolized with double arrows Boiling vapor bubbles from throughout the liquid and rise to the surface and burst Boiling point temperature at which vapor pressure of a liquid is equal to the external pressure above the surface Solid water is less dense than liquid water Phase diagram plotting pressure and temperatures at which a substance undergoes a phase change Sublimation describes transition between solid state and vapor state without a liquid in between Deposition transition between vapor and solid phase with no liquid phase in between Triple point all three phases exist at a specific temperature and pressure Supercritical fluid at a specific temperature and pressure phase are no longer distinctly seen Critical point pressure and temperature where liquid vapor equilibrium doesn’t matter, essentially a liquid and a gas at the same time Melting points increase with increasing pressure Chapter 36: What is a Solution and How is it Concentration Expressed? Solution homogeneous mixture, uniform distribution throughout the solution components Solute substances in relatively small amount Solvent medium which solute is being dissolvent into Concentrated relatively large quantity of solute per unit of solution Dilute relatively small quantity of solute per unit of solution Solubility measure of how much solute will dissolve in a given amount of solvent at a given temperature Saturated solution whose concentration is at the solubility limit Unsaturated solution whose concentration is below the solubility limit Supersaturated solution whose concentration of solute is greater than the solubility limit Misible soluble, but only applied to liquids Immisible insoluble, only applied to liquids Hydrated ions surrounded by polar water molecules Dissolving is a reversible process Rate of crystallization depends on concentration of solute crystal at surface Rate of dissolving and crystallization are the same at equilibrium Dissolving time depends on surface area, crystallization rate and temperature Molarity, M moles of solute per liter of solution Molarity can be used as a conversion factor when going from liters to moles or moles to liters M 1 =1 V ,2M2 molarity, V= volume, used in dilutions Chapter 37: How are Chemical Changes in Solution Described? Part 1 Electrodes electrical cord attached to a battery and a light bulb A solution must contain ions so that electricity can flow through the solution Electrolysis indication of ions that make up an electrical current in a solution Strong electrolyte good conductor Weak electrolyte poor conductor Nonelectrolyte doesn’t conduct electricity Acid hydrogen bearing molecule or ion that releases a hydrogen ion in a water solution Strong acid almost all of the molecules of the original compound dissociate into the solution, major species is ions Weak acid only slightly ionized into solution, major species is molecules 7 strong acids HCl, HCl3 , HClO4, HBr, HI, HNO3, and 2 SO4 Conventional equation standard chemical equation that we have learned to write, used for stoichiometry Total Ionic Equation replaces formulas of dissolved substances with major species in solution Spectator ions ions present at the scenes of the reaction but experiences no chemical change, appears on both sides of the equation Net ionic equation equations where all spectators are removed. Chapter 38: How are Chemical Changes in Solution Described? Part 2 Activity Series able to tell which molecules are more likely to give up electrons or if they want to stay a solid, used in predicting products in a single replacement reaction Look at the order in which an element appears in the activity series, if element A has a higher positon than element B, the A will replace B in the products to form AX Look at the order in which an element appears in the activity series, if element A has a higher positon than element B, the A will replace B in the products to form AX Ion Combing Reaction cation from one reactant combines with anion from another to form a particular kind of product compound Precipitate a solid formed in a double replacement reaction, precipitation reaction Molecular Product reaction of an acid which leads to a combination of ions instead of precipitate Water is not ionized in liquid state Carbonic acid breaks apart into water and carbon dioxide gas Sulfurous acid product splits into water and sulfur dioxide gas Ammonium hydroxide (doest exist) forms water and ammonia Chapter 39: How are Solution Quantities Related in Chemical Change? Titration careful addition of one solution to another by means of a device can measure delivered volume precisely Indicator changes color to show complete neutralization of an acid or base Standardize finding concentration for use in titrations When dealing with a hydrate, water doesn’t need to be in the equation but is accounted for in weight Ideal Solutions no interaction among dissolved solute particles Quantity of dissolved solids in solution increases small decrease of vapor pressure Change in vapor pressure (VP) occurs when solute is added to solvent ∆ P change∈pressure ) VP of pure solvent VP of solution= ∆ P ∝ X solution fraction of solution made up, mole fraction Mole Fraction number of moles of a component of a mixture divided by the total number of moles of all components molescomponent1 X= =mole fractionof component1 Σmolesof all componentsaddedtogether Sum of mole fraction of all components of any mixture must be equal to 1 In dilute solutions change in some properties is ∝ to concentration of solute particles Proportionality constant is the vapor pressure of pure substance, P° ∆ P=P° X Solution, mole fraction Raoult’s Law vapor pressure of a solution is proportional to the mole fraction of the solvent, which is the solution component responsible for vapor pressure Entropy of solution is greater than the entropy of the pure solvent Vapor pressure of a solution is lower than the vapor pressure of corresponding pure solvent Boiling Point increase with increasing concentration ∆ T Boiling Point Elevation b change in boiling point, difference between a solution and the corresponding temperature for boiling the pure solvent is proportional to the molal concentration of the solute (moles of solute per kilogram of solvent), ∆ T b =K b ℃ K bmolal boiling point elevation proportionality constant in /kg, T boiling point elevation colligative property b Molality, m number of moles of a solute dissolved in 1kg of solvent, m= molesolute kgsolvent Freezing point of solution decreases with increasing concentration When solution freezes, solid forms pure solvent Solid solvent + heat liquid solvent Heat is removed from the system during the freezing process ∆ Tfchange in freezing temperature, freezing point depressing ∆ T f∝ m ∆ Tf= Kmf K – folal freezing point depression constant ∆T f (¿b) Can find molality from m= which then is expressed as mol solute/ kg solvent K f (¿b) Using the defining equation for molar mass, MM=g/mol, calculate molar mass of solute Vapor pressure of a solution is the sum of the partial vapor pressure of the solute plus that of the solvent, soln P1 1 ° X 2 2 ° … Ion Multiplier, i equal to the number of moles of ions resulting from the dissociation of one mole of solute ∆ ∆ T f iKmf(freezing) or Tb = iKbm (Boiling), i number of total compounds in a molecular formula ex. NaCl i=2 Semipermeable membrane having very ting submicroscopic spores Molecules continue to move through the membrane in both directions Dynamic Equilibriumopposite changes occur at the same rate Osmosis flow of any solvent through a semipermeable membrane from a dilute solution or pure solvent to a more concentrated solution Osmotic pressure π pressure required to restore equilibrium in the rates of flow between the two sides of a semipermeable membrane π=RTiM , mass and molar mass can be substituted in for M Equation can be sued to estimate moles or molar mass of a large complex unknown solvent molecules Same units as ideal gas equation Isotonicsolution with same concentration Hypertonic solution which has a concentration greater than the internal of the cell Hypotonic solute concentration that is less than the cell Colloid glue, substances that are partially transparent and has visible particles in it Tyball Effect sunlight scattered by colloidal particles in air or water Hydrophilic dispersed phase molecules are attracted to water Hydrophobic colloidal particles separate from water Electrical double layer surrounded colloid particles caused by large colloidal molecules having many electrons Absorb ions of the colloidal particles that adhere to the surface of the water molecules because of ions having induced dipole attraction Large outer shell surrounding the colloidal particles repelled by the outer shell of other colloidal particles Surfactants, surface active agents large nonpolar section and a polar head and at one end of the molecules Dialysis using semipermeable membrane to separate colloidal particles from solution Agglomeration, coagulation or flocculation causes dispersed particles to come into contact so that they are attracted to each other through intermolecular forces with no ionion repulsion hydrophobic colloidal particles can agglomerate and form particles that are large enough to filter out Kinetics investigates speed, or rate, of chemical reactions Reaction Mechanism closely related knowledge of events during a chemical reaction some reactions are very slow, some are fast States of reactants: depends on frequency of molecular collisions more particles, more collisions fast reaction Solids have less surface area less space for collisions so slower reaction Homogenous phase states of reactants are all the same Heterogeneous phase reactants are in different phases Concentration of reactants: higher concentration means more particles in a given space and allows for more frequent collisions Reaction temperature: chemical reactions occur faster at higher temperatures (pressure cookers) Addition of a catalyst: catalyst substance that increases speed of reaction because it lowers the energy needed for the reaction initiation Catalyst is either nonparticipant or regenerated Enzyme molecule that acts as a biological catalyst Speed quantity that states how fast something is moving Rate quantity of one thing considered in relation to another measure quantity u Reaction rates change in concentration of a species per unit of time Reaction rates expressed in terms of reactant for product Simple reaction: decomposition of reactants formation of products, RP −∆ [R ] ∆ [P] Rate= ∆T = ∆T , square brackets indicate concentration Rates expressed as positive quantities but final concentrations are always less than initial concentrations hence the use of the negative sign on the reactants Spectrophotometer measures color intensity electronically Average rate defined as the average rate of change in concentration of a species over a specified time interval ≡− ∆ R ] | [R]| average rate ∆T or ∆T Instantaneous Reaction rate at a particular instant in time, found using calculus Rate of chemical reaction is based on which reactant or product is measured |∆[A |=|∆ [B|= ∆ [ ]= ∆ [ ] aA+bB cC+dD (general equation) Rate = a∆T b∆T c∆T d ∆T Rate of decomposition of reactants is always negative without absolute value signs Use quantity algebra, don’t need the negatives or absolute value sign on reactants to go from
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