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Chem 1127 Chapter 1-9 Notes

by: Catherine Cabano

Chem 1127 Chapter 1-9 Notes Chem 1127

Marketplace > University of Connecticut > Chemistry > Chem 1127 > Chem 1127 Chapter 1 9 Notes
Catherine Cabano

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Book chapter notes for Chem 1127 Chapters 1 - 9!
General Chemistry
Fatma Selampinar (TC), Joseph Depasquale (PI)
Class Notes
Chemistry, Book, notes, Chapter one through nine
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This 10 page Class Notes was uploaded by Catherine Cabano on Wednesday March 2, 2016. The Class Notes belongs to Chem 1127 at University of Connecticut taught by Fatma Selampinar (TC), Joseph Depasquale (PI) in Fall 2015. Since its upload, it has received 33 views. For similar materials see General Chemistry in Chemistry at University of Connecticut.

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Date Created: 03/02/16
Chapter 1 – Matter and Measurements o 3 phases o Solid – fixed shape and volume o Liquid – fixed volume but is not rigid in shape o Gas – neither a fixed volume nor a shape o Pure substances are either elements or compounds o Element – type of matter that cannot be broken down into two or more pure substances  118 known, 91 occur naturally o Compound – pure substance that contains more than one element; ex: water is a compound of hydrogen and oxygen o Mixture – contains two or more substances combines in such a way that each substance retains its chemical identity o Homogenous – uniform mixtures; composition is the same throughout; solution o Heterogeneous – non-uniform mixtures are those in which the composition varies throughout o To separate components of a mixture  filtration (used to separate a heterogeneous solid-liquid mixture); distillation (used to resolve a homogenous solid-liquid mixture) o Standard unit of length in the metric system is the meter o Mass is most commonly expressed in grams, kilograms, or milligrams o Temperature is the factor that determines the direction of heat flow o t(F) = 1.8 (t(C)) + 32 o K = 273.15 + t(C) o Intensive – must be independent of amount o Extensive – depend on amount o Chemical properties – observed when the substance takes part in a chemical reaction; a change that converts it to a new substance o Physical properties – observed without changing the chemical identity of a substance; melting point and boiling point o Density = mass/volume o Solubility – the process by which a solute dissolves in a solvent o Saturated vs. unsaturated vs supersaturated Chapter 2 – Atoms, Molecules and Ions o Atom: the smallest particle of an element that can enter into a chemical reaction o Law of conservation of mass: no detectable change in mass in an ordinary chemical reaction o Law of constant composition: a compound always contains the same elements in the same proportions by mass o Law of multiple proportions: the masses of one element that combine with a fixed mass of the second element are in a ratio of small whole numbers o Electrons: common to all atoms, carry a unit negative charge (-1), and have a very small mass, roughly 1/2000 that of the lightest atom o Discovered by J. J. Thomson o Nucleus: small, positively charged center of atom o Proton: mass nearly equal to that of an ordinary hydrogen atom; unit positive charge (+1), equal in magnitude to that of the electron (-1) o Neutron: uncharged particle with a mass slightly greater than that of a proton o Atomic number: number of protons o In a neutral atom  the number of protons in the nucleus is exactly equal to the number of electrons outside the nucleus o Mass number: number of protons + number of neutrons o Isotopes: atoms that contain the same number of protons but a different number of neutrons o Atomic masses: relative masses of atoms of different elements; indicates how heavy, on average, one atom of that element is compared with an atom of another element o Atomic mass units (amu) = 1/12 mass C atom = mass H atom o The average atomic mass shown in the periodic table is not equal to the mass number o Ex: mass of C-12 atom = 12 amu (exactly) o Isotopic abundances: atom percent in nature  can be determined by mass spectrometry o Mass number on top, atomic number on bottom, element symbol o Periodic Table: symbols of all the elements arranged in a particular way o Avagadro’s Number: 6.022 x 10 23 o It represents the number of atoms of an element in a sample whose mass in grams is numerically equal to the atomic mass of the element o Periods: horizontal rows in the table o Groups: vertical columns o Main-group elements: groups 1, 2, 13, 14, 15, 16, 17, 18 o Transition metals: ten elements in the center of each of periods 4 through 6 o The periodic table is an arrangement of elements, in order of increasing atomic number, in horizontal rows of such a length that elements with similar chemical properties fall directly beneath on another in vertical groups o Metals  high electrical conductivities; Nonmetals  right of the stairway, 18 elements Metalloids  B, Si, Ge, As, Sb, Te o Molecule: two or more atoms may combine with one another to form an uncharged molecule o Molecular formulas: the number of atoms of each element is indicated by a subscript written after the symbol of the element o Structural formulas: show the bonding pattern within the molecule o Condensed structural formula: suggests the bonding pattern in the molecule and highlights the presence of a reactive group of atoms within the molecule o Ions: when an atom loses or gains electrons; charged particles o Cations: lose electrons to form positively charged ions o Anions: formed by nonmetal atoms; negative ions o When an ion is formed, the number of protons in the nucleus is unchanged o Ionic bonds: ionic compounds are held together by strong electrical forces between oppositely charged ions o Strong electrolyte: conducts an electric current o Nonelectrolyte: does not conduct electricity; ex: sugar and water o The total positive charge of the cations in the formula must equal the total negative charge of the anions o **** See list of polyatomic ions on page 46 o **** See charts of naming compounds on page 49 Chapter 3 – Mass Relations in Chemistry; Stoichiometry  Mole: 6.022 x 10 23  The molar mass, MM, in grams per mole, is numerically equal to the sum of the masses (in amu) of the atoms in the formula  Mass = MM x n  Molarity (M) = (moles of solute)/(liters of solution)  [ ]  commonly used to represent the molarity of a species in solution  The molarity of a solution can be used to calculate: the number of moles of solute in a given volume of solution; the volume of solution containing a given number of moles of solute  When an ionic solid dissolved in water, the cations and anions separate from each other  Percent composition  citing the mass percent of the elements present  The subscripts in a formula represent not only the atom ratio in which the different elements are combined but also the mole ratio  Simplest formula: gives the simplest whole-number ratio of the atoms present  Assume a 100-g sample and calculate the mass of each elements in that sample  Any calculation involving a reaction must be based on the balanced equation for that reaction.  You cannot write an equation unless you know what happens in the reaction that it represents  The coefficients of a balanced equation represent number of moles of reactants and products.  Theoretical yield  the maximum quantity of something that can be obtained under certain conditions, assuming the reaction goes to completion and no product is lost  Limiting reactant  the amount of product formed is determines (limited) by the amount of limiting reactant o Steps:  Calculate the amount of product that would be formed if the first reactant were completely consumed  Repeat this calculation for the second reactant; that is, calculate how much product would be formed if all of that reactant were consumed  Choose the smaller of the two amounts calculated. This is the theoretical yield of product; the reactant that produces the small amount is the limiting reactant.  Take the theoretical yield of the product and determine how much of the reactant in excess is used up in the reaction. Subtract that from the starting amount to find the amount left.  % yield = (actual yield)/(theoretical yield) x 100%  Hydrates: ionic compounds often separate from water solution with molecules of water incorporated into the solid Chapter 4 – Reactions in Aqueous Solutions - Equations that exclude “spectator ions,” which take no part in the reaction = net ionic equations - Atom balance: there must be the same umber of atoms of each element on both sides - Charge balance: there must be the same total charge on both sides - ***See figure 4.6 Flowchart - An acid is a species that produces H ions in water solution - - A base is a species that produces OH ions in water solution - Strong acids ionize completely, forming H ions and anions  ex: HCl - Weak acids: molecule containing an ionizable hydrogen atom - Strong base: completely ionized to OH ions and cations - - Weak bases: produce OH ion+ in a quite different ma-ner; react with water molecules, acquiring H ions and leaving OH ions behind - ***See Table 4.1 for strong acids and bases - Neutralization occurs in strong acid-strong base reactions - ***See Figure 4.8 about strong acids and bases - Titration: measuring the volume of standard solution (a solution of known concentration) required to react with a measured amount of sample - The objective of the titration is to determine the point at which reaction is complete = equivalence point - Oxidation reduction reactions/redox reactions – transfer of electrons between two species o One species loses (donates) electrons and is said to be oxidized; the other species gains (receives) electrons is reduced o Can be split into two half reactions - Oxidation number: 1. The oxidation umber of an element in an elementary substance is 0 2. The oxidation number of an element in a monoatomic ion is equal to the charge of that ion 3. Certain elements (“leading elements”) have the same oxidation number in all their compounds a. Group 1 elements always have an oxidation number of +1 b. Group 2 elements always have an oxidation number of +2 c. Fluorine (F) always has an oxidation number of -1 4. Hydrogen in a compound has an oxidation number of +1, unless it is combined with a metal, in which case it is -1. 5. The sum of the oxidation numbers in a neutral species is 0 and in a polyatomic ion is equal to the charge of the ion. 6. Oxygen in a compound has an oxidation number of -2, unless it is combined with a Group 1 metal (always +1) or Group 2 metal (always +2). Solve algebraically for the oxidation number of oxygen. - Oxidation = an increase in oxidation number and reduction as a decrease in atomic number - The ion or molecule that accepts electrons is called the oxidizing agent; by accepting electrons it brings about the oxidation of another species - The species that donate electrons = reducing agent - Before you can balance an overall redox equation, you have to be able to balance two half-equations, one for oxidation (electron loss) and one for reduction (electron gain) - Half-equations: 1. Assign oxidation numbers to each element 2. Balance the atoms of each element being oxidized or reduced 3. Multiply the oxidation number by the number of atoms that have that oxidation number. This gives you the “total” oxidation number 4. Balance oxidation number by adding electrons + - 5. Balance charge by adding H ions in acidic solution and OH ions in basic solution 6. Balance hydrogen by adding water molecules 7. Check to make sure that oxygen is balanced - Balancing Redox equations: 1. Split the equation into two half-equations – one for reduction, the other for oxidation 2. Balance one of the half-equations with respect to both atoms and charge 3. Balance the other half-equation 4. Combine the two half-equations in such a way as to eliminate electrons Chapter 5 – Gases o I L = 10^3 m^3 = 10^-3 m^3 o Number of moles = n o N = mass/MM o Temperatures must be expressed on the Kelvin scale o T(K) = T(C) + 273.15 o Pressure = force per unit area o Measure atmospheric pressure  barometer o Gas pressure is often expressed in millimeters of mercury (mm Hg) o Standard atmosphere = atmosphere (atm) o 1.013 bar = 1 atm = 760 mm Hg = 14.7 psi = 101.3 k Pa o 1 bar = 10^5 Pa o The Ideal Gas Law o Volume is directly proportional to amount  V = k 1 k1= constant; the slope of the line; independent of individual values of V, n, and the nature of the gas o Volume is directly proportional to absolute temperature  V = k 2 o Volume is inversely proportional to pressure o V = k / P 3 o PV = nRT o STP  P = 1 atm; V = 22.4 al; n = 1.00 mol; T = 273 K o Density of a gas is dependent on pressure, temperature, molar mass o The volume ratio of any two gases in a reaction at constant temperature and pressure is the same as the reacting mole ratio. o Dalton’s law of partial pressures: The total pressure of a gas mixture is the sum of the partial pressures of the components of the mixture. o The partial pressure of water vapor, is equal to the vapor pressure of liquid water o The partial pressure of a gas in a mixture is equal to its mole fraction multiplied by the total pressure. o Kinetic theory of gasses  all gases behave similarly as far as particle motion if concerned o Gasses are mostly empty space o Gas molecules are in constant, chaotic motion o Collisions are elastic o Gas pressure is caused by collisions of molecules with the walls of the container o P = N(mass)u / 3V o Average translational kinetic energy of a gas molecules * see formula pg 141 o Effusion = the flow of gas molecules at low pressures through tiny pores or pinholes Chapter 6 – Electronic Structure and the Periodic Table  Electron configurations: show the number of electrons in each energy level  Orbital diagrams: show the arrangement of electrons within orbitals  Wavelength: the distance between two consecutive crests or troughs, most often measured in meters or nanometers (1 nm = 10^-9 m)  Frequency: the number of wave cycles (successive crests or troughs) that pass a given point in unit time; v  (wavelength in meters)(frequency in Hz) = speed of light  c = 2.998 x 10^8 m/s  Light is a stream of particles called photons  E = hv = hc/wavelength  h = 6.626 x 10^-34 J s  Energy is inversely related to wavelength  Color spectrum: between 400 and 700 nm  Quantized = limited to particular values  E n -2.180 x 10^-18 J/n^2  Bohr model o E n -R /nH 6 o E n energy of the electron o R = Rydberg constant; 2.180x10^-18 J H o n = principal quantum number  Lowest energy state = ground state  When an electron absorbs enough energy, it moves to a higher excited state  Quantum mechanics: o The kinetic energy of an electron is inversely related to the volume of the region to which it is confined. o It is impossible to specify the precise position of an electron in an atom at a given instant  n = principal energy levels o n = 1, 2, 3, 4… o As n increases the energy of the electron increases and it is found farther out from the nucleus  l = second quantum number (s, p, d, f) o l = 0, 1, 2…(n-1) o Sublevels o Ex: if n=3; l=0, 1, or 2 o In the nth principal level, there are n different sublevels o ns < np < nd < nf o s = 0; p = 1; d = 2; f = 3  m l third quantum number; orbitals o m =ll…,+1, 0, -1,…,-l o Ex: d sublevel: l=2  m = l, 1, 0, -1, -2  5 orbitals o Ex: f sublevel: l=3  m = l, 2, 1, 0, -1, -2, -3  7 orbitals  M s fourth quantum number; electron spin o M = +1/2 or -1/2 s  Pauli exclusion principal: no two electrons in an atom can have the same set of four quantum numbers  As n increases the radius of the orbital becomes larger  A p orbital consists of two lobes along an axis (x, y, or z)  Electron configuration: shows the number of electrons, indicated by a superscript, in each sublevel  The abbreviated electron configuration starts with the preceding noble gas  Actinides: 14 elements in the seventh period are filling the 5f sublevel  Orbital diagrams of atoms used with arrows and parenthesis  Hund’s rule: when several orbitals of equal energy are available, as in a given sublevel, electrons enter singly with parallel spins o In all filled orbitals, the two electrons have opposed spins o In accordance with Hund’s rule, within a given sublevel there are as many half-filled orbitals as possible  When transition metal atoms form positive ions, the outer s electrons are lost first  The chemical and physical properties of elements are a periodic function of atomic number  Atomic radius  decrease across a period from left to right; increase down a group in the periodic table  Ionic radius  increases moving down a group in the periodic table o Positive ions are smaller than the metal atoms from which they are formed o Negative ions are larger than the nonmetal atoms from which they are formed  Ionization energy: a measure of how difficult it is to remove an electron from a gaseous atom o Energy must always be absorbed to bring about ionization, so ionization energies are always positive quantities o The more difficult it is to remove ionization energies, the larger the ionization energy o Increases across the period table from left to right; decreases moving down the periodic table o Electronegativity: measures the ability of an atom to attract to itself the electron pair forming a covalent bond Chapter 7 – Covalent Bonding o Lewis suggested that nonmetal atoms, by sharing electrons to form an electron- pair bond, can acquire a stable noble-gas structure o Valence electrons are the ones involved in bonding o Shared electrons are counted for both atoms o Two kinds of electron pairs: o Shared by two atoms = covalent bond; ordinarily shown as a straight line between bonded atoms o An unshared pair of electrons, owned entirely by one atom, is shown as a pair of dots on that atom o Single, double and triple bonds o Atoms in covalently bonded species tend to have noble-gas electronic structures  octet rule o Writing Lewis Structures: 1) Draw a skeleton of the species joining atoms by single bonds 2) Count the umber of valence electrons 3) Count the number of valence electrons available for distribution AE = VE – 2(number of bond in the skeleton) 4) Count the number of electrons required to fill out an octet for each atom (except H) in the skeleton (NE) o Resonance: two structures, separated by a double-headed arrow, are written with the understanding that the true structure is intermediate between them o The double-headed arrow is used to separate resonance structures o Differ only in the distribution of electrons, not in the arrangement of atoms o Formal charge: can be applied to any atom within a Lewis structure; the difference between the number of valence electrons in the free tom and the number assigned to that atom in the Lewis structure. o The charge an atom would have if valence electrons in bonds were distributed evenly. o FC = VE – ½(bonding electrons) o Isomers have the same formula but different properties o FC = VE – unshared electrons – number of bonds o Formal charge is not an infallible guide to predicting Lewis structures o Expanded octets are possible o When the number of electrons available after the skeleton is drawn is greater than the number required to give each atom an octet  distribute the extra electrons (two or four) around the central atom as unshared pairs o Bond angles = angles between bonds o Linear = bond angle of 180 degrees o Bent = bond angle less than 180 degrees o Molecular geometry  electron pair repulsion o According to the VSEPR model, the valence electron pairs surrounding an atom repel one another. Consequently, the orbitals containing those electron pairs are oriented to be as far apart as possible. o *** See table on pg 203 o Unshared pairs reduce bond angles below ideal values o Insofar as molecular geometry is concerned, a multiple bond behaves like a single bond o Polar: as a result of an unsymmetrical distribution of electrons, the bond or molecule contains a positive and a negative pole and is therefore a dipole o Nonpolar: a symmetrical distribution f electrons leads to a bond or molecule with no positive or negative poles o Nonpolar bonds – formed whenever the two atoms joined are identical, as in H 2 o Bones in which the electron density is unsymmetrical are polar bonds o All molecules, except those of elements, have polar bonds o Polar molecules: one that contains positive and negative poles; partial positive charge (positive pole) at one point in the molecule and a partial negative charge (negative pole) at a different point o There are no positive and negative poles in a nonpolar molecule o The HF molecule is called a dipole – contains positive and negative poles o The number of orbitals is shown by the superscript o Unshared as well as shared electron pairs can be located in hybrid orbitals o *** See table of pg 215 o The extra electron pairs in a multiple bond (one pair in a double bond, two pairs in a triple bond) are not located in hybrid orbitals o Sigma bond consists of an electron pair occupying a sigma bonding orbital o Pi bonding orbitals  two lobes, one above the bond axis, the other below it. Chapter 8 – Thermochemistry - Energy = the capacity to do work - Heat = a particular form of energy that is transferred from a body at a high temperature to one at a lower temperature when they are brought into contact with each other - Thermochemistry = the study of the heat flow that accompanies chemical reactions - Calorimetry = the experimental measurement of the magnitude and direction of heat flow - System = that part of the universe on which attention is focused - Surroundings = exchange energy with the system; make up in principal the rest of the universe - State properties depend only on the state of the system, not on the way the system reached that state - “q” is positive when heat flows into the system from the surroundings - “q” is negative when heat flows out of the system into the surroundings - An endothermic process (q>0), in which heat flows from the surroundings into the reaction system  ex: melting of ice - An exothermic process (q<0), in which heat flows from the reaction system into the surroundings  ex: combustion of methane - q = C x delta T - Delta T = t(final) – t(initial) - C = heat capacity = the amount of heat required to raise the temperature of the system 1 degree Celsius and has the units J/degree C - q = mass x c x delta T - “c” = specific heat capacity = the amount of heat required to raise the temperature of one gram of a substance on degree Celsius - When the mass of that substance is equal to its molar mass, then c is called the molar heat capacity - Specific heat, like density of melting point, is an intensive property that can be used to identify a substance - q(reaction) = -q(calorimeter) - q(cal) = C(cal) x delta T - Coffee-cup calorimeter  C(cal) = Mass(water) x c(water) = mass(water) x 4.18 - Bomb calorimeter: more versatile; used for reactions involving gases - Enthalpy: a type of chemical energy, sometimes referred to as “heat content” - Thermochemical equation = a chemical equation that shows the enthalpy relation between products and reactants - Rules of thermochemistry: o The magnitude of delta H is directly proportional to the amount of reactant or product o Delta H for a reaction is equal in magnitude but opposite in sign to delta H for the reverse reaction. o The value of delta H for a reaction is the same whether it occurs in one stop or in a series of steps - The standard molar enthalpy of formation of a compound (delta Hf) is equal to the enthalpy change when one mole of the compound if formed at a constant pressure of 1atm and a fixed temperature, ordinarily 25 degree Celsius, from the elements in their stable states at that pressure and temperature. - The standard enthalpy chance, delta H, for a given thermochemical equation is equal to the sum of the standard enthalpies of formation of the product compounds minus the sum of the standard enthalpies of formation of the reactant compounds - Elements in their standard state can be omitted - The coefficients of products and reactants in the thermochemical equation must be taken into account - The bond enthalpy is defined as delta H when one mole of bonds is broken in the gaseous state - Thermodynamics = deals with all kinds of energy effects in all kinds of processes - Delta H = q(p) - Delta E = q(v) - H = E + PV Chapter 9 – Liquids and Solids - Why are liquid and solids so different from gases? o Molecules are much closer to one another in liquids and solids o Intermolecular forces, which are essentially negligible with gases, play a much more important role in liquids and solids - Once equilibrium between liquid and vapor is reached, the number of molecules per unit volume in the vapor does not change with time - The pressure exerted by the vapor over the liquid remain constant - The pressure of vapor in equilibrium with a liquid is called the vapor pressure - So long as both liquid and vapor are present, the pressure exerted by the vapor is independent of the volume of the container - A liquid boils at a temperature at which its vapor pressure is equal to the pressure above its surface - If this pressure if 1atm (760 mmHg), the temperature is referred to as the normal boiling point - Critical temperature, above which the liquid phase of a pure substance cannot exist - The pressure that must be applied to cause condensation at that temperature is called the critical pressure - Above the critical temperature and pressure, a substance is referred to as a super- critical fluid - All these equlilbria can be shown in a phase diagram - Sublimation = the process by which a solid changes directly to vapor without passing through the liquid phase - Deposition = the opposite of sublimation; the phase transition from the vapor phase to solid without passing through the liquid phase


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