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# CH131-Chemistry for the Engineering Sciences

Marketplace > Boston University > CH131 Chemistry for the Engineering Sciences
Omaima
BU
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-Chapters 1-4 in full detail are included with visual aids and detailed explanations -Very neat and organised
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This 9 page Study Guide was uploaded by Omaima on Thursday January 7, 2016. The Study Guide belongs to a course at Boston University taught by a professor in Fall. Since its upload, it has received 67 views.

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Date Created: 01/07/16
Chapter 1 Law of Multiple Proportions:When two (2) elements form a series of compounds, the masses of one element that combine with a fixed mass of the other element are in the ratio of small integers to each other —> Typically, will divide by the smallest mass to satisfLaw of Multiple Proportions Chemical Formula: The composition of a compound is shown by this formula—it is a way of expressing information about the proportions of atoms that constitute a particular chemical formula Law of Combining Volumes:​ The ratio of the volumes of any pair of gases in a gas phase chemical reaction (at the same temperature and pressure) is the ratio of simple integers Abundance:​ Measure of the occurrence of an element in a given environment Equation 1:​Relative Mass= Sum of ((abundance)/100)​ isotope mass) —>Relative Mass does not have units Isotope:​Elements with the same atomic number, but different masses —> Remember, changing an element’s proton number means changing the whole element Equation 2:​Mass Number= Sum of Protons and Neutrons for the Isotope —> n=neutrons —> p=protons Chapter 2 Avogadro’s Number: ​6.0221420x10^23 Equation 1:Density=Mass/Volume —> p​=density (cm^3) —> m=​mass (g) —> v​= volume (mL, L) —>​ Density is not fixed; depends on the pressure and temperature during the time of measurement Equation 2:Molar Volume= Vm= Molar Mass/Density Tips: Do not forget to read and understand what you are asked to find. Molecular Formula:​The number of atoms of each element in one molecule of that substance —> Preferred because molecular formula provides more detailed information, however, it is not always easy to derive this formula Empirical Formula:This formula gives the correct relative of numbers —> Simplest formula Equation 3:Percentage Composition= (Element’s mass/Total Mass) x100 Equation 4:Mass Fraction (to determine if there any other elements in sample)Mass of the element present in one (1) mole of the compound/Mass of one (1) mole of compound Mass Relationships in Chemical Reactions —> Molecular formula is the same whole number multiple of the Empirical Formula —> To determine the molecular formula, you must know the approximate molar mass of the compound under study —> Ratio of molar masses of two (2) gaseous compounds is the same as the ratio of their densities Limiting Reactant Limiting Reactant:​eactant that is used up first in reaction In Excess: All the other reactants present after the limiting reactant has been used up —> Whichever reactant gives the smaller mass of product is the limiting reactant —> To determine how much is in excess take initial mole amount minus final mole using Stoichiometry (the amount reached) —> Total mass of reactant in the end should be equal to the original mass present **This satisfies the Law of Conservation of Mass Percentage Yield Theoretical Yield:Amounts of products calculated so far (if reaction goes smoothly) Actual Yield:The amount present after separating it from other products and reactants and purifying it **This amount is always less than theoretical yield Percentage Yield:​Ratio of actual yield to the theoretical yield multiplied by 100% Equation 3:(Actual Yield/Theoretical Yield)x100 Summary ❖ 1 Mole= 6.02x10^23 atoms or molecules ❖ The Limiting Reactant determines maximum theoretical yield ❖ Percent yield is somewhat less than theoretical yield Chapter 3-Nomenclature Electrostatics: The forces between and energies of systems of stationary charged particles Ionic Bond:​One or more electrons iscompletely​transferred from one atom to the other, and the dominant contribution to the strength of the bond is the electrostatic attraction between the resulting positive and negative ions Covalent Bond:​Electrons are shared more or less equally between the two (2) atoms comprising the bond —>Polar Covalent:​A partial transfer of charge from one atom to the other; bonds that posses a mixture of ionic and covalent character (there is a partial transfer of charge) Polar Covalent Bonding—Electronegativity and Dipole Moments: ​ The absolute value of the difference in electronegativity of Z bonded atoms tells the degree of polarity in their boond Electronegativity:The tendency of an atom in a molecule toattract electronsfrom other atoms —> Explains whether a given pair of atoms forms an ionic, covalent, or polar covalent bond Greater than 2.0 Ionic; Electrons have been transferred Smaller than 0.4 Largely covalent; Electrons in bond are evenly shared Intermediate Values (0.41-1.9) Polar Covalent Lewis Dot Diagram:​ Shows the number of valence electrons associated with each atom in a molecule and indicates whether they are bonding (shared) or non-binding —> These diagrams are helpful in predicting the structural formula—which atoms are bonded to each other in polyatomic molecules-but they do not describe the three-dimensional shapes of molecules Valence Electrons:​Outermost electrons Valence Shell:​Outermost, partially filled shell; contain electrons involved in chemical bonding —>Lewis Model​ represents the valence electrons as dots arranged around the the chemical symbol for an atom; the core electrons are not shown Core Electrons:​Electrons in the inner shell VSEPR Theory (Valence Shell Electron Repulsion Theory): ​ Predicts molecular shapes based on the electrostatic argument that electron pairs in a molecule will arrange themselves to be as far apart as possible Molecule:​A collection of atoms bonded together Condensed Structural Formula:​ Specifies which atoms are bonded to each other and by what types of bonds —> Ex: C2H60 ———> CH3CH2OH —> Easier to understand which atoms are bonded to a specific atom Electrostatic Potential Energy DiagramDisplays the electrostatic potential energy that a small positive “test charge” would experience at every position on the electron density surface that defines the space filling the model Isomers:​Different compounds found in nature or in the laboratory that have the same molecular formula but different molecular structures, therefore, different properties Z:​Atomic Number Forces and Potential Energy in Atoms Coulomb's Law​(Coulomb's​Inverse-Square ​Law): Law​of physics describing the electrostatic interaction between electrically charged particles. Equation 1:F(r)=((q1q2)/(4piE0r^2)) —> Describes the electrical force between two (2) charges q1 and q2 separated by a distance r —> E0= 8.854x10^-12 (Permittivity of the Vacuum) —> C (Coulombs)= Charge —> +q (proton) —> -q (electron) Coulomb’s Force Law:​The potential energy stored in the compressed spring measures how much force the spring can exert, and in which direction it is released. Equation 2:V(r)=(q1q2)/(4piE0R) **V(r)---->0 as r----->infinity —> Same charges: V(r)= +; V(r) dec. as r inc. —> Different charges V(r)=-; V(r) inc as r dec —> Charge of Proton: 1.602x10^-19J —> Charge of Electron:-1.602x10^-19J —> 1ev (electron vol= 1.6021764x10^-19J Ionization Energies, Shell Model of the Atom, and Shielding ➢ The shielding effect​ describes the attraction between an electron and the nucleus in any atom with more than one electron shell. Shielding Effect can be defined as a reduction in the effective nuclear charge on the electron cloud, due to a difference in the attraction forces of the electrons on the nucleus. First Ionization EnergyMinimum energy necessary to remove an electrofrom a neutral atom in the gas phase (first) X(g)-->X + e- —> Delta E for Ionization Reaction is always positive —>​Trends: Ionization Energy increases across periodic table Ionization Energy increases up a group **Noble gases require a higher energy to remove an e- Second Ionization Energy: Minimum energy required to remove a second electron ElectronA​ffinityThe ease with which an atomaccepts an extra electron to form an anion X(g) + e——> X-(g) ElectronegativityThe tendency of atoms to attract electrons in molecules —>​Trends:​Electronegativitincreases across the periodic t; elements on left are near noble gases and do not really want to lose their electrons while elements on the right are super close in attaining a noble gas electron configuration—> they are more like to attract electrons Electronegativitncreases up a group —>​Electrnegative:Electron acceptors (electrons are negative) —>​Electrpositiv: Electrondonors (when you give up electrons, you become +) Ionic Bonding:Ionic bonding form between atoms with large differences in electronegativity —> Stability occurs when an atom either loses or gains electrons —> Hydrogen/Helium are satisfied with two (2) valence electrons Covalent and Polar BondingElements of identical or comparable electronegativity Bond Length: Distance between the nuclei of two (2) atoms —> Bond Length and ​ond Energy​increases across a period —> Bond Energy: Energy required to break one mole of a particular bond (stability is determined by bond energy) —> Bonds generally grow weaker with increasing atomic number —> Increase force constant correlates with higher bond energy and smaller bond length (stiffer bonds or stronger and shorter) Lewis Dot Diagrams Octet:​Whenever possible, the electrons in a covalent compound are distributed in such a way that each main group element is surrounded by eight electrons (Hydrogen/Helium are satisfied with two (2). Tip:Draw out single atom Lewis Dot Diagrams for each element and observe the number of unmatched electrons—these are the electrons that actually do the bonding Steps: 1. Count total number of valence electrons present —> If species is a negative add​additional electrons to achieve the total charge —> if species is a positivesubtract​enough electrons to result in the total charge 2. Calculate total amount of electrons needed​atisfyeach atom 3. Subtract the number two (2) from number one (1): 2-1 4. Assemble the structure—double bonds​form only between atoms of C, N, O and S. Triple bondsare usually restricted to C, N, or O 5. Determine Formal Charge​; formal charges must add up to the correct charge on the molecule or polyatomic ion 6. If more than one diagram is possible, choose the one with the smallest magnitudes of formal charges (0, +1, -1) and with any negative formal charges placed on the most electronegative atoms —> Formal Charge: ​# of Valence electrons - lone pair electron (each dot counts as one) -(.5)(number of electrons bonding) Resonance: ​Two or more equivalent Lewis diagrams can be written denoted with “<-->” arrows in brackets ( [ ] ) Molecular Geometry Steric Number:​ Determines which geometry applies —>Steric Number: ​(number of atoms​bonded to central atom)+(number of lone pairs​on central atom) —> Electrons position themselves to minimize electron-pair repulsion —> When lone pairs are present, there can be three different types of repulsions (lone pairs tend to occupy more space than bonding pairs) ● bonding pair against bonding pair ● bonding pair against lone pair ● lone pair against lone pair Less Volatile:Have lower vapor pressure Oxidation Numbers: 1. Oxidation of neutral molecule must add up to zero and those in an ion must add up to the charge on the ion (Charge conservation) 2. Alkali-Metal atoms (group 1) are assigned the oxidation number +1 3. Alkaline Earth Metals (group 2) are assigned the oxidation number +2 4. Fluorine (F2) ialways​assigned oxidation number -1 in icompounds 5. Halogens are generally also assigned the oxidation numberexcept​ those containing oxygen and other halogens in which the halogen can have a positive number 6. Hydrogen is assigned the oxidation +1 in compounds​, except when convention two (2) takes precedence—> becomes -1 7. Oxygen is assigned oxidation number -2 in nearall compounds​except when it is in a compound with fluorine (take convention 3) —> +2 and when it is in O-O bonds (convention 2 and 4) —> -1 8. Neutral elements’ oxidation numbers are zero Chapter 4 Quantum Mechanics Theory:​The central idea behind quantum theory is that energy, like matter, is not continuous but it exists only in discrete packets Visible Spectrum:400-700nm

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