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Chemistry 1127Q Final Review

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by: Caitrín Hall

Chemistry 1127Q Final Review CHEM 1127Q 001

Marketplace > University of Connecticut > Chemistry > CHEM 1127Q 001 > Chemistry 1127Q Final Review
Caitrín Hall
GPA 3.9

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This study guide includes information from all of the chapters we've covered this semester (1-6, and 8-10). I was less selective with chapters 8, 9, and 10 since we have not been tested on them ye...
General Chemistry
Fatma Selampinar (TC), Joseph Depasquale (PI)
Study Guide
final, Chemistry, study, guide
50 ?




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"Clutch. So clutch. Thank you sooo much Caitrín!!! Thanks so much for your help! Needed it bad lol"
Ms. Audrey Lubowitz

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Popular in Chemistry

This 30 page Study Guide was uploaded by Caitrín Hall on Saturday April 30, 2016. The Study Guide belongs to CHEM 1127Q 001 at University of Connecticut taught by Fatma Selampinar (TC), Joseph Depasquale (PI) in Spring 2016. Since its upload, it has received 184 views. For similar materials see General Chemistry in Chemistry at University of Connecticut.


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Clutch. So clutch. Thank you sooo much Caitrín!!! Thanks so much for your help! Needed it bad lol

-Ms. Audrey Lubowitz


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Date Created: 04/30/16
Chapter 1 The scientific method – path of discovery that answers a question with experimental verification and modification 1. Question and observation 2. Hypothesis – tentative explanation of observations that guides experiment 3. Experimental verification 4. Theory – satisfactory, testable explanation of natural observation; can be modified! Phases of matter – occupies space and has mass  Solid – rigid object with fixed shape and volume  Liquid – variable shape that flows to fill container; fixed volume  Gas – indefinite shape and volume expand and contract to fill container  Plasma – gaseous state of matter with charged particles; high temp environment Law of conservation of matter – no matter is created nor destroyed during chemical reactions or phase changes Classifying matter:  Pure substance – constant composition; can be either compounds or elements o Elements – cannot be broken down by chemical changes; see periodic table o Compounds – can be broken down by chemical changes; consist of two or more atoms of different elements  Mixtures – composed of two or more types of matter present in varying amounts; can be separated by physical change o Heterogeneous – composition varies from point to point o Homogeneous (solution) – visibly but not chemically uniform throughout Properties:  Physical property – characteristic of matter not associated with change in chemical composition; ex) density, color, hardness, melting/boiling pts., conductivity  Chemical property – change (or inability to change) from one type of matter into another; ex) flammability, toxicity, acidity, reactivity, heat of combustion  Extensive property – depends on & is directly proportional to amount of matter; ex) mass, volume, heat  Intensive property – does not depend on amount of matter; ex) temperature Measurement  Exact numbers and defined quantities can be counted and do not change during the counting process; free from uncertainty  Significant figures – all digits in a measurement including the last uncertain digit o All nonzero numbers are significant o Captive zeros—between nonzero numbers—are significant o Leading zeros – NOT significant o Trailing zeros – NOT significant unless # is a decimal or in scientific notation  Significant figures in calculations  o Multiplication/division – result contains same # significant figures as does the factor with the least # significant figures o Addition/subtraction – result contains same # decimal places as the factor with the least # decimal places  Accuracy – results are vey close to true or accepted value  Precision – results are very similar to each other  Dimensional analysis uses conversion factors o Tips:  Begin with the initial unit  Multiply by the conversion factor  The desired unit should be the numerator of the conversion factor  The unit of the initial quantity should be the denominator  SI base units: Chapter 2 Law of definite proportion (constant composition) – all samples of a pure compound contain the same elements in the same proportion by mass Law of multiple proportions – when 2 elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers Subatomic particles  Electron – a negatively charged, subatomic particle with a mass more than one thousand-times less than th-31ass of an atom o Mass of electron = 9.107 x 10 kg  Proton – positively charged, subatomic particle in the nucleus  Neutrons – uncharged, subatomic particles with mass similar to mass of protons  Isotopes – atoms of the same element that differ in mass o Chemically identical atoms with the same # protons but different # neutrons Symbolism  Atomic number (Z) = # protons in nucleus of an atom; defining trait of an element  Mass number (A) = # protons + # neutrons in nucleus o A – Z = # neutrons  Ions – electrically charged atoms resulting from different # of protons and neutrons o Anions – atoms that gain 1 or more electrons and exhibit – charge o Cations – atoms that lose 1 or more electrons and exhibit + charge  Chemical symbol – abbreviation used to indicate an element or atom of an element o Ex) mercury’s symbol = Hg o Usually contain 1 or 2 letters (only the first letter is capitalized) Mass # Atomic #  Isotopes X o Symbol is written by placing the mass # as a superscript to the left o Ex) Mg has 3 isotopes: 24Mg; 25Mg; 2Mg  Atomic mass – almost (but not always exactly) equal to mass number because each proton and neutron contributes about 1 amu to the mass, but isotopes exist; average mass accounts for all isotopes present in a naturally occurring sample of an element Average mass = ∑ (fractional abundance × isotopic mass) o Mass spectrometers experimentally determine natural isotopic abundance Chemical formulas  Molecular formula – representation of a molecule that uses chemical symbols to indicate the types of atoms followed by subscripts to show the # of atoms of each type in the molecule  Empirical formula – indicates types of atoms present as the simplest whole-number ratio of the # atoms (or ions) in the compound o Divide the molecular formula by the greatest common factor to derive the empirical formula  Isomers – compounds with the same chemical formula but different molecular structures  majorly affects chemical properties Periodic table  The modern periodic table arranges elements in increasing order of their atomic numbers and groups atoms with similar properties in the same column  There are seven periods (horizontal rows) and 18 vertical columns called groups  Classes: o Metals – shiny, malleable, good conductors of heat and electricity o Nonmetals –dull, poor conductors of heat and electricity o Metalloids – elements that conduct heat and electricity moderately well and possess some properties of metals and others of nonmetals  Further classifications: o Main-group (representative) elements – found in columns 1, 2, and 13-18 o Transition metals are in columns 3-12 o Inner transition metals are the 2 rows at the bottom of the table  Lanthanides – top row  Actinides – bottom row Compounds  Ionic bonding involves transfer of electrons (proton and neutron #s don’t change)  Monatomic ions are formed from only 1 atom o Many main-group metals lose enough electrons to achieve the same number of electrons as the preceding noble gas o Nonmetals often gain electrons to achieve the same number of electrons as the next noble gas  Common polyatomic ions – electrically charged molecules; memorize these!!   Covalent (molecular) bonds – attractive forces between the positively charged nuclei of bonded atoms and 1 or more pairs of electrons between the atoms Chemical nomenclature Ionic Compounds  Compounds with only monatomic ions o Name of cation (metal) followed by name of anion (nonmetal) but with the ending replaced by –ide  Compounds with polyatomic ions o Named similarly but there is no need to change the ending because the suffix is in the name of the anion  Compounds containing a metal ion with variable charge o Most transition metals form 2 or more cations with different charges o The charge of the metal ion is written in parentheses Molecular (Covalent) Compounds  Compounds composed of 2 elements o The name of the more metallic element (farther left and/or bottom of periodic table) is first prefaced by a Greek prefix:  Mono- is usually excluded as a prefix for the more metallic element o Followed by the name of the more nonmetallic element (farther right and/or top of periodic table) with its ending changed to –ide  The Greek prefixes also apply to naming the less metallic element o Examples:  SO = sulfur trioxide 3  NO 2 nitrogen dioxide  N2O 5 dinitrogen pentoxide  P4O10 = tetraphosphorus decaoxide  Binary acids are composed of hydrogen plus one other nonmetallic element 1. Hydrogen is changed to hydro- 2. Add suffix –ic to the other nonmetallic element 3. The word “acid” is added as a second word o Ex) HF as a gas is “hydrogen fluoride”  as an acid it is “hydrofluoric acid”  Oxyacids – compounds that contain hydrogen, oxygen, and at least one other element, and are bonded to create acidic properties 1. Omit “hydrogen” 2. Start with root name of anion 3. Replace –ate with –ic, or –ite with –ous 4. Add “acid” Chapter 3  Formula mass – sum of average atomic masses of all atoms in the substance’s formula  Covalent formulas represent # and types of atoms in a molecule; formula mass = molecular mass  Ionic compounds contain cations and anions but do not represent the composition of a discrete molecule; formula mass is not molecular mass  Avogadro’s number (N ) = A.02 x 10 23entities composing a mole  Molar mass – mass in grams of 1 mole of that substance (g/mol) Percent Composition – percentage by mass of each element in the compound  To find % comp by each element, divide the experimentally derived mass of each element by the overall mass of the compound, then convert to a percentage  To determine % comp from formula mass, consider 1 mol of given compound and use its molar mass to calculate the percentage of each of its elements  To determine empirical formula, use the given masses to find moles of each element, divide each by the lesser number of moles, multiply ratio (if necessary) to get the smallest possible whole number subscripts Derivation of Molecular Formulas  Compare compound’s molecular or molar mass to its empirical formula mass Molar mass/empirical formula mass = n formula units/molecule  Multiply each subscript of the empirical formula by n (A B ) = A B x y n nx nx Solutions  Solvent – medium in which other components are dissolved; has a significantly greater concentration that that of other components  Aqueous solution – a solution in which water is the solvent  Solute – component of a solution present at a lower concentration than solvent  Molarity (M) – number of moles of solute in exactly 1 liter of solvent M = mol solute/L solution o Moles and volumes can be determined from molar concentrations o Molar concentrations can be determined from mass of solute o Mass of solute in given volume of solution can be determined from molarity  General dilution equation:C1V 1C V 2 2 *C = concentration & V = volume* Mass percentage – the ratio of the component’s mass to the solution’s mass (Mass of component/mass of solution) x 100%  Percent mass %mass, percent weight %weight, weight/weight percent (w/w)%  Calculation of percent by mass o Divide mass of chemical formula by mass of sample  Calculations using Mass Percentage o Use solution density given to find solution’s volume and mass, then use the  given mass percentage to calculate solute mass Chapter 4 Chemical equation – symbolic representation of a chemical reaction 1. The substances undergoing reaction are called reactants, and their formulas are placed on the left side of the equation 2. The substances generated by the reaction are called products, and their formulas are placed on the right side of the equation 3. Plus signs separate individual reactant and product formulas, and an arrow separates the reactant and product sides of the equation 4. The relative numbers of reactant and product species are represented by coefficients (numbers placed immediately to the left of each formula); it is common to use the smallest possible whole-number coefficients Equations for Ionic Reactions  Molecular equation – doesn’t explicitly represent the ionic species that are present in the solution CaCl (aq) + 2AgNO (aq) ⟶ Ca(NO ) (aq) + 2AgCl(s) 2 3 3 2  Complete ionic equation – explicitly represents all dissolved ions 2+ − + − 2+ − Ca (aq) + 2Cl (aq) + 2Ag (aq) + 2NO 3 (aq) ⟶ Ca (aq) + 2NO3 (aq) + 2AgCl(s)  Spectator ions – presence is required to maintain charge neutrality but are not chemically nor physically changed by the process  Net ionic equation – complete ionic equation MINUS spectator ions Cl (aq) + Ag (aq) ⟶ AgCl(s) Acid-Base Reactions  An acid-base reaction is one in which a hydrogen ion is transferred from one chemical species to another o Acid – any substance that dissolves in water to yield hydronium + ions, H3O o Base – a substance that will dissolve in water to yield hydroxide ions, OH-  Neutralization reaction – acid-base reaction in which the reactants are an acid and a base, while the products are often and salt and water (neither reactant is water) Oxidation-Reduction Reactions  Oxidation = loss of electrons = increase in oxidation number  Reduction = gain of electrons = decrease in oxidation number  Oxidation number – the charge an element’s atoms would possess if the compound was ionic o The oxidation # of an atom in elemental form is zero o The oxidation # of a monatomic ion = ion’s charge o Oxidation numbers for common nonmetals:  Hydrogen: +1 when combined with nonmetals, −1 when combined with metals  Oxygen: −2 in most compounds, sometimes −1, very rarely -1/2, positive values when combined with F  Halogens: -1 for F, -1 for other halogens except when combined with oxygen or other halogens (then they are varying positive #s) o The sum of oxidation numbers for all atoms in a molecule or polyatomic ion equals the charge on the molecule or ion Balancing Redox Reactions via the Half-Reaction Method 1. Write the two half­reactions 2. Balance all elements except oxygen and hydrogen 3. Balance oxygen atoms by adding H O 2olecules +  4. Balance hydrogen atoms by adding H ions 5. Balance charge by adding electrons 6. If necessary, multiply each half­reaction’s coefficients by the smallest possible  integers to yield equal numbers of electrons in each 7. Add the balanced half­reactions together and simplify by removing species that  appear on both sides of the equation 8. For reactions occurring in basic media (excess hydroxide ions), carry out these  additional steps:  −  +  o Add OH ions to both sides of the equation to equal the number of H ions o On the side of the equation containing both H and OH ions, combine  these ions to yield water o Remove any redundant water molecules    9. Check to see that both the number of atoms and the total charges are balanced Reaction Yields  Limiting reactant – the reactant present in an amount lower than required by the reaction stoichiometry, thus limiting the amount of product generated  Excess reactant – the reactant present in an amount greater than required by the reaction stoichiometry Percent yield – the extent to which a reaction’s theoretical yield is achieved  Theoretical yield – the amount of product that may be produced by a reaction under specified conditions  Actual yield – the amount of product obtained in practice Percent yield = (actual yield/theoretical yield) x 100% Chemical analysis  Titration analysis – quantitative chemical analysis method that involves measuring the volume of a reactant solution required to completely react with the analyte—substance whose concentration must be measured o Titrant – substance whose concentration is known o Equivalence point – volume of titrant solution required to react completely with the analyte; provides a stoichiometric amount of titrant for the sample’s analyte according to the titration reaction  Gravimetric analysis – quantitative analysis in which a sample is subjected to some treatment that causes a change in the physical state of the analyte that permits its separation from the other components  Combustion analysis – gravimetric method of analysis in which the weighed sample of a compound is heated to a high temperature under a stream of oxygen gas, resulting in complete combustion to yield gaseous products of known identities Chapter 5 Thermochemistry  Energy – the capacity to supply heat or do work  Work (w) – the process of causing matter to move against an opposing force  Potential energy – the energy an object has because of its relative position, composition,  or condition  Kinetic energy – the energy an object possesses because of its motion  Thermal energy – kinetic energy associated with the random motion of atoms and  molecules  Temperature – a quantitative measure of “hot” or “cold” o Increasing energy will increase temperature and the substance expands o Decreasing energy will decrease temperature and the substance contracts  Heat – the transfer of thermal energy between two bodies at different temperatures o Exothermic process – a reaction that releases heat (ex: combustion); q is ­ o Endothermic process – a reaction that absorbs heat; heat is absorbed leading to  the sensation of cold; q is +  A calorie is the amount of energy required to raise one gram of water by 1 kelvin o Depends on the atmospheric pressure and starting temperature of the water  A joule is the amount of energy used when a force of 1 newton moves an object 1 meter o SI unit of heat, work, and energy  Heat capacity (C) – the quantity of heat (q) a body of matter absorbs or releases when it  experiences a temperature change (T) of 1 kelvin C = q/T o Determined by both the type and amount of substance that absorbs/releases heat  Specific heat capacity (c) – the quantity of heat required to raise the temperature of 1  gram of a substance by 1 kelvin c = q/(mT) o Depends only on type of substance absorbing/releasing heat (intensive property) Calorimetry  Calorimetry – process of measuring the amount of heat involved in a chemical or  physical process  System – the substance or substances undergoing the chemical or physical change  Surroundings – the other components of the measurement apparatus that serve to either  provide heat to the system or absorb heat from the system  Calorimeter – a device used to measure the amount of heat involved in a chemical or  physical process o When an exothermic reaction occurs in solution in a calorimeter, heat produced  by the reaction is absorbed by the solution (solution temperature increases) o When an endothermic reaction occurs, heat required is absorbed from the thermal  energy of the solution (solution temperature decreases)  Heat produced or consumed in the system plus heat absorbed or lost by the surroundings  must add up to zero q + q = 0 reaction solution q reaction = -(q solution)  Bomb calorimeter – type of calorimeter that operates at constant volume to measure  energy produced by reactions that yield large amounts of heat and gaseous products (ex:  combustion reactions) o Require calibration to the heat capacity of the calorimeter and to ensure accuracy o Accomplished using a reaction with known q and m o Temperature change produced by the known reaction is used to determine the heat capacity of the calorimeter Enthalpy  Enthalpy (H) – the sum of a system’s internal energy (U) and the mathematical product  of its pressure (P) and volume (V) H = U + PV  Enthalpy change (ΔH) – enthalpy changes for chemical or physical processes can be  determined, but enthalpy values for specific substances cannot be measured directly ΔH = H ­ H products reactants  The following conventions apply: 1. The ΔH value indicates the amount of heat associated with the reaction involving  the number of moles of reactants and products as shown in the chemical equation 2. The enthalpy change of a reaction depends on the physical state of the reactants  and products of the reaction 3. ΔH > 0 for endothermic reactions ΔH < 0 for exothermic reactions Hess’s Law: If a process can be written as the sum of several stepwise processes, the enthalpy  change of the total process equals the sum of the enthalpy changes of the various steps 1. ΔH is directly proportional to the quantities of reactants or products 2. ΔH for a reaction in one direction is equal in magnitude and opposite in  sign to ΔH for the reaction in the reverse direction 3. Elements in their standard states have enthalpies of formation of zero ΔH° reaction = ∑ n × ΔH° (froducts) − ∑ n × ΔH° (feactants) Chapter 6 Electromagnetic radiation – light is the visible part of a vast spectrum of electromagnetic waves; kilometers (10 m) to picometers (10 -1m)  Light behaves like a wave and a particle (wave-particle duality) Waves  A wave is a periodic movement that can transport energy from one point to another  Wavelength (λ) – the distance between two consecutive peaks or troughs in a wave  Frequency (v) – wave cycles that pass a specified point in space in a specified amount of time; measured in hertz (Hz) or seconds inverse -1 (s )  Amplitude – corresponds to the magnitude of the wave’s displacement; one-half the height between the peaks and troughs; related to intensity (brightness or loudness)  Electromagnetic spectrum – the range of all types of electromagnetic radiation o The human eye sees only between 400 and 700 nm 8 −1 c = 2.998 × 10 ms = λν  Short wavelength means high frequency and energy  Long wavelength means low frequency and energy  Continuous spectrum – electromagnetic radiation given off in an unbroken series of wavelengths  Quantization – only discrete values from a more general set of continuous values of some property are observed; hence why n = 1, 2, 3…  Photons – smallest possible packet of electromagnetic radiation, a particle of light whose energy depends on its frequency Planck’s formula: E = hv OR E = (hc)/ λ Line spectra  Line spectrum – electromagnetic radiation emitted at discrete wavelengths by a specific atom (or atoms) in an excited state  Each element has a unique energy shell system The Bohr model  Models the hydrogen atom only but does not account for electron- electron interactions in atoms with more than one electron  Introduce important features of all models o An electron in its lowest energy orbit is in its ground electric state o If the atom receives energy from an outside source, an electron can move to an orbit of a higher n value; excited electronic state o When an electron moves from an excited state to a less excited state, the difference is the energy emitted as a photon o Energies of electrons are quantized, described by quantum numbers; integer #s having only specific allowed value o Electron’s energy increases with increasing distance from the nucleus o Discrete energies (lines) in the spectra result from quantized energies Quantum theory  Heisenberg uncertainty principle: It is fundamentally impossible to determine simultaneously and exactly both the momentum and the position of a particle; consequence of wave-particle duality  Atomic orbital – general region in an atom within which an electron is most probable to reside  The Pauli Exclusion Principle: no two electrons have the same set of all four quantum #s; two electrons in the same orbital must have opposite spins  Electron configuration – the arrangement of electrons in the orbitals of an atom 1. The # principal quantum shell, n 2. The letter that designates the orbital type (subshell, l) 3. A superscript number that designates # electrons in that subshell  The Aufbau Principle – each added electron (across a period) occupies the subshell of lowest energy available  Hund’s Rule – the lowest-energy configuration for an electrons within a set of degenerate (same energy) orbitals is having max # unpaired electrons Electron Configurations and the Periodic Table  These classifications determine which orbitals are counted in the valence shell 1. Main group (representative) elements – the last electron added enters and s or p orbital in the outermost shell; valence electrons are those with highest n; completely filled d orbitals count as core 2. Transition elements or transition metals – metallic elements in which the last electron added enters d orbital; valence electrons include ns & (n – 1) d electrons 3. Inner transition elements – metallic elements in which the last electron added occupies f orbital; valence shells consist of (n – 2)f, (n – 1)d, and ns subshells Variation in Covalent Radius  Covalent radius – ½ the distance between the nuclei of two identical atoms when joined by a covalent bond (possible because atoms within molecules retain identity)  Size increases from top to bottom and decreases from left to right o Effective nuclear charge, Z eff the pull exerted on a specific electron by the nucleus, taking into account electron-electron repulsions o Z increases from left to right across a period; stronger pull eff experienced by electrons on the right side of the periodic table draws them closer to the nucleus, making the radii smaller Variation in Ionic Radii  Cation is smaller than its atom because when electrons are removed from the outer valence shell, the remaining core electrons experience a greater Z eff  Down the groups of the periodic table, cations of successive elements with the same charge have larger radii  Anion is larger than its atom because addition of one or more electrons to the valence shell increases repulsion among electrons and decreases Z effper electron  Isoelectric atoms and ions have the same electron configuration Variation in Ionization Energies  The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy (IE ) 1  Endothermic process (IE values are always positive)  IE1increases with increasing Zacross a period and decreases with increasing Z down a group  Exceptions: o IE 1roup 2 > IE gro1p 3; IE group15 > IE group 61 o Because it is easier to remove electrons paired electrons from higher orbitals Variation in Electron Affinities  Electron affinity [EA] – the energy change for the process of adding an electron to a gaseous atom to for an anion  Endo/exothermic depending on element  It becomes easier to add an electron across a series of atoms as Z eff increases  From left to right across a period, EAs tend to become more negative  From top to bottom of each group, EA is less clear Chapter 8 Advanced Theories of  Chemical Bonding 8.1 Valence Bond Theory  Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms  Orbitals on two different atoms overlap when portions of two orbitals occupy the same region of space  2 conditions must be met: 1. An orbital on one atom overlaps an orbital on a second atom 2. The single electrons in each orbital combine to form an electron pair  More overlap  stronger covalent bond  Energy of system depends on how much the orbitals overlap  Number of covalent bonds depends on number unpaired electrons 8.2 Hybrid Atomic Orbitals  Hybridization – the process of combining the wave functions for atomic orbitals; mathematically accomplished by the linear combination of atomic orbitals (LCAO) o Hybrid orbitals are produced o Sigma (σ) bond – a covalent bond in which the electron density is concentrated in the region along the internuclear axis; single bonds o Pi (π) bond – type of covalent bond that results from the side- by-side overlap of two p orbitals; regions of orbital overlap lie on opposite sides of the internuclear axis 1. Hybrid orbitals do not exist in isolation 2. Hybrid orbitals have shapes and orientations different from those of the atomic orbitals in isolated atoms 3. A set of hybrid orbitals is generated by combining atomic orbitals (# hybrid orbitals per set = # atomic orbitals that were combined to produce the set 4. All orbitals in a set of hybrid orbitals are equivalent in shape and energy 5. The type of hybrid orbitals formed depends on its electron-pair geometry (VSEPR) 6. Hybrid orbitals overlap to form sigma bonds; unhybridized orbitals overlap to form pi bonds  sp hybridization – one s orbital and one p orbital hybridize to make two sp orbitals; the central atom is surrounded by two regions of valence electron density 2  sp hybridization – one s orbital and two p orbitals hybridize to make three sp orbitals; the central atom is surrounded by three regions of valence electron density 3  sp hybridization – one s orbital and three p orbitals hybridize to make four sp orbitals; the central atom is surrounded by four regions of valence electron density  sp hybridization – one s orbital and three p orbitals hybridize to make four sp orbitals; the central atom is surrounded by four regions of valence electron density  sp d hybridization – one s orbital, three p orbitals, and one d orbital 3 hybridize to make five sp d orbitals; the central atom is surrounded by five regions of valence electron density  sp d hybridization – one s orbital, three p orbitals, and two d orbitals 3 2 hybridize to make six sp d orbitals; the central atom is surrounded by six regions of valence electron density 8.3 Multiple Bonds  Hybridization involves only sigma bonds, lone pairs of electrons, and single unpaired electrons; the arrangement of pi bonds involves only unhybridized orbitals o Single bond = 1 sigma bond o Double bond = 1 sigma bond and 1 pi bond o Triple bond = = 1 sigma bond and 2 pi bonds Chapter 9 Gases 9.1 Gas Pressure   Pressure – the force exerted on a given area P = F/A o Directly proportional to force and inversely proportional to area o The SI unit is the pascal (Pa); 1 Pa = 1 N/m o 1 kPa = 1000 Pa; 1 bar = 100,000 Pa o In the U.S., pressure is measured in pounds per square inch (psi) o Atmosphere (atm) – originally represented the average sea level air pressure at the approximate latitude of Paris (45 degrees) o 101.3 atm = 101,325 Pa = 760 mm Hg = 760 torr = 14.7 psi o Barometer –the atmosphere exerts pressure on the liquid outside the test tube  the column of liquid exerts pressure inside the tube  the pressure at the liquid surface is the same inside and outside the tube; the height of the liquid in the tube is proportional to the pressure exerted by the atmosphere o Manometer – a device similar to a barometer used to measure the pressure of a gas trapped in a container 9.2 Relating Pressure, Volume, Amount, and Temperature:  The Ideal Gas Law  Pressure and Temperature: Amonton’s or Gay-Lussac’s Law  Directly proportional at constant volume P1/T1 = P2/T2 Volume and Temperature: Charles’s Law  Directly proportional at constant pressure V1/T1 = V2/T2 Volume and Pressure: Boyle’s Law  Inversely proportional at constant temperature  The graph of P vs. V is a parabola  The graph of 1/P vs. V is linear P1 x V1 = P2 x V2 Moles of a Gas and Volume: Avogadro’s Law  Directly proportional at constant pressure and temperature V1/n1 = V2/n2 The Ideal Gas Law = PV = nRT  R is the ideal gas constant = 0.08206 L atm mol K and 8.314 kPa L mol K1 -1  An idea gas is a hypothetical construct that may be used along with kinetic molecular theory to effectively explain the gas  If moles of an ideal gas are kept constant: (P1 V1)/T1 = (P2 V2)/T2 Standard Conditions of Temperature and Pressure  STP = 273.15 K and 1 atm  Standard molar volume = 22.4 L 9.3 Stoichiometry of Gaseous Substances, Mixtures, and  Reactions  Density of a Gas  Mass to volume ratio Molar Mass of a Gas  Grams per mole of a substance  Combined with the molar mass equation:  M = (mRT)/(PV) The Pressure of a Mixture of Gases: Dalton’s Law  Dalton’s law of partial pressures: The total pressure of a mixture of ideal gases is equal to the sum of the partial pressures of the component gases  The partial pressure of gas A is related to the total pressure of the gas mixture via its mole fraction (X) – a unit of concentration defined as the number of moles of a component of a solution divided by the total number of moles of all components P =X ×P where X = n n A A Total A A/ Total Collection of Gases over Water  Simple way to collect gasses that do not react with water: capture the gas in a bottle filled with water and inverted into a dish filled with water  The pressure of the gas in the bottle can be made to equal the air pressure by raising/lowering the bottle  When the water level is the same inside and outside the bottle, the pressure of the gas is equal to the atmospheric pressure  The pressure of the pure gas is equal to the total pressure minus the pressure of the water vapor; vapor pressure of water – pressure exerted by water in EQ with liquid water in a closed container; depends on temperature Chemical Stoichiometry and Gases Avogadro’s Law Revisited  Gases combine, or react, in definite and simple proportions by volume, provided all gas volumes are measured at the same temperature and pressure Chapter 10Liquids and Solids  Unlike with gases, the properties of liquids and solids depend on chemical identity 10.1 Intermolecular Forces   Intermolecular interaction refers to attractive forces between the particles of a substance, regardless of whether these particles are molecules, atoms, or ions  Intermolecular forces (IMFs) – the various forces of attraction that may exist between the atoms and molecules of a substance due to electrostatic phenomena; serve to hold particles close together  KE provides the energy required to overcome attractive forces  Phase changes occur when conditions of temperature or pressure favor the associated changes in IMF  Increased pressure brings molecules closer, while increased temp increases KE  If temp becomes sufficiently low or pressure becomes sufficiently high, the molecules don’t have enough KE to overcome the IMF  solid forms Forces between Molecules  Intermolecular forces occur between particles, while intramolecular forces are those within the molecule that keep it together (ex: bonds between atoms)  IMF determine many physical properties  IMFs between small molecules are often weak compared to the intramolecular forces that bond atoms together  Van der Waals forces – all attractive forces between neutral atoms and molecules Various types of IMFs:  Dispersion forces, dipole-dipole forces, and hydrogen bonding Dispersion Forces  The London dispersion force is present in all condensed phases  Dispersion force – attraction between two rapidly fluctuating, temporary dipoles;  significant only when particles are very close together   Because electrons are in constant motion, an atom/molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically  An induced dipole results when an instantaneous dipole distorts the electrons of a neighboring atom/molecule o Both result in weak, electrostatic dispersion forces  Polarizability – the measure of how easy or difficult it is for another electrostatic charge to distort a molecule’s charge distribution (its electron cloud) o If a charge cloud is easily distorted, it is very polarizable and will have large dispersion forces o In larger atoms, valence electrons are farther from the nuclei  less tightly held  more easily form temporary dipoles that produce the attraction  Shapes of molecules also affect magnitudes of dispersion forces o Greater surface area available for contact between molecules  stronger dispersion forces  higher boiling point Dipole-Dipole Attractions  Dipole-dipole attraction – the electrostatic force between the partially positive end of one polar molecule and the partially negative end of another  Present in polar molecules only  Stronger than dispersion forces  Two different substances with the same MM can have different boiling points if one substance is polar and the other isn’t o Presence of dipole-dipole attraction  higher boiling point Hydrogen Bonding  Strongest van der Waals force, but much weaker than covalent bonds  Hydrogen bonding – strong, type of dipole-dipole attraction that occurs when a molecule contains an H atom bonded to F, O, or N (most electronegative atoms)  Intermolecular attractive force  The large difference in electronegativity between hydrogen and F, O, or N combined with the very small size of an H atom and the relatively small sizes on F, O, or N atoms leads to highly concentrated partial charges  Hydrogen bonds are denoted by dots connecting atoms  The effect of increasingly stronger dispersion forces down a group dominates that of increasingly weaker dipole-dipole attractions  boiling points increase steadily  Hydrogen bonding molecules exhibit anomalously high boiling points  Hydrogen bonding in DNA – A and T share two H bonds while C and G share three 10.3 Phase Transitions Vaporization and Condensation  Gas  liquid = condensation  Liquid  gas = vaporization  In a closed container, gas molecules collide with the surface of the condensed phase; some collisions result in molecules re-entering the condensed phase  When rate of condensation = rate of vaporization, neither the amount of liquid nor the amount of vapor in the container changes  Dynamic equilibrium – the status of a system in which reciprocal processes occur at equal rates  Vapor pressure – the pressure exerted by the vapor in EQ with a liquid in a closed container at a given temperature o Does NOT depend on surface area of contact with container o Does depend on IMF o Strong IMF impede vaporization and favor the recapture of gas- phase molecules  low vapor pressure o Weak IMF prevent less of a barrier to vaporization and a reduced likelihood of gas recapture  high vapor pressure o As temp increases, VP of a liquid also increases due to increased average KE o Escape of more molecules per unit of time and greater average speed of molecules that escape both contribute to higher VP Boiling Points  Boiling point – the temperature at which a liquid’s EQ VP = the pressure exerted on the liquid by its gaseous surroundings  Normal boiling point – boiling point of a liquid when surrounding pressure = 1 atm  Temp remains constant throughout boiling process Enthalpy of Vaporization  Vaporization is endothermic  Energy change associated with vaporization is enthalpy of vaporization, ΔHvap Melting and Freezing  Melting – energy becomes large enough to overcome molecules in their fixed positions, and the solid transitions to liquid state o Temp of solid stops rising despite continual input of heat and remains constant until all of the solid is melted  Freezing – the reciprocal process of melting  The temp at which the solid and liquid phases of a given substance are in EQ is called the melting point of the solid or the freezing point of the liquid Sublimation and Deposition  Sublimation – solids transition directly into the gaseous state; ex: dry ice  Deposition – the reverse of sublimation; ex: formation of frost


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