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CM-UY 1004 General Chemistry for Engineers Study Guide

by: Sanbir Rahman

CM-UY 1004 General Chemistry for Engineers Study Guide CM-UY 1004

Marketplace > New York University > Chemistry > CM-UY 1004 > CM UY 1004 General Chemistry for Engineers Study Guide
Sanbir Rahman
GPA 4.0

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About this Document

A review of all of the units covered with some examples included.
General Chemistry for Engineers
Janice E. Aber
Study Guide
percent yield, MassPercentCompostion, atomic, number, acids, bases, Binary, compounds
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This 5 page Study Guide was uploaded by Sanbir Rahman on Friday September 30, 2016. The Study Guide belongs to CM-UY 1004 at New York University taught by Janice E. Aber in Fall 2016. Since its upload, it has received 10 views. For similar materials see General Chemistry for Engineers in Chemistry at New York University.


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Date Created: 09/30/16
CM-UY 1004 General Chemistry for Engineers Chemistry Unit: Matter and Measurements Phase Technical Physical Properties Chemical Properties Change Name  color  reactivity with Solid --> Liquid Melting  mass acid Liquid --> Solid Freezing  density  flammability Liquid --> Gas Evaporation  boiling point  oxidizability Gas --> Liquid Condensation  solubility Solid --> Gas Sublimation Gas --> Liquid Deposition Extensive Property: depends on the amount of substance (ex. mass, volume, weight) Intensive Property: not dependent on the amount of substance (ex. temp, density, color, pressure) Matter: anything that has mass and takes up space Atoms: the basic building blocks of all matter Classification of Matter Element: made up of only one type of atom Compound: made up of different types of atoms Mixture: separable based on physical properties Precision: measurements that are well-reproduced OR made to many decimal points; improve by repeating trials Accuracy: measurements that are true and improved by comparing with a standard Conversions: needed= given x conversion (s) Temperature: average kinetic energy kilo 1000  ̊C = (̊ - 32) (5/9)  ̊K = ̊ +273 (k) deci 0.1 addition or subtraction: think about number of (d) decimal places centi 0.01 multiplication or division: think about number of (c) sig figs Percent Error= 100 (measured - accepted)/ milli 0.001 accepted (m) micro 10-6 (µ) nano 10-9 (n) CM-UY 1004 General Chemistry for Engineers Chemistry Unit: Formulas and Reactions Elements that have single letter abbreviations: Hydrogen (H), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Phosphorus (P), Sulfur (S), Potassium (K), Iodine (I), Tungsten (W), Fluorine (F), Yttrium (Y), Vanadium (V), Uranium (U) If parentheses are in a formula, the subscript to the right tells how many sets of atoms exist in the formula. The total number of atoms is the number within the parenthesis multiplied by the subscript outside the parenthesis. Ex. Pb (NO ) 3 2 contains 1 Pb, 2 Ns, and 6 Os Hydrate: a compound that has one or more water molecules bonded to it; the coefficient in front of the 2 O shows the number of water molecules Diatomic Elements: Hydrogen (H ), 2itrogen (N ), 2xygen (O ), F2uorine (F ), 2hlorine (Cl2), Bromine (Br2), and Iodine (2 ) |Elemental Phosphorus is P a4d Elemental Sulfur is S | 8 Atomic Number: number of protons constant for an element Mass Number: number of protons + the number of neutrons Charge: number of protons – the number of electrons Atoms can gain or lose electrons to become ions. When they lose an election, they become a cation (+) and when they gain an electron, they become an anion (-) Isotopes: species that have the same number of protons but differ in their amount of neutrons In a chemical reaction, energy is absorbed when bonds break and energy is released when atoms rearrange and new bonds form. States: (s) solid; (g) gas; (l) liquid; (aq) aqueous Reactions Explanation Equation Example Synthesis elements react to form a compound A+B AB Fe+O 2Fe O 2 3 Decompositio compound  elements (or, smaller AB  A+B KCIO 3 KCIO +K4I n compounds) Combustion hydrocarbon + oxygen  CO +H O2 2 C a b  C 7 6 2O 2 CO 2 H O2 CO 2H O2 Double compound + compound  different AB+XY  ZnCl 2Na C2  3 Replacement compound + compound AX+BY NaCl+ZnCO 3 Single element + compound  different AB+X  Na+H O  2 Replacement element + compound AX+ B NaOH+H 2 Rules and Tips for Balancing Reactions: - Subscripts don’t change but coefficients do - Always reduce coefficients to the lowest whole number - Use fractions to “force balance” a reaction, then multiply - CHECK! Examples: CM-UY 1004 General Chemistry for Engineers Double Replacement: Al(NO ) + H SO3 3Al (SO2) +4HNO 2 4 3 3  Then, BALANCE: 2 Al(NO ) + 3 H S3 3 Al (SO )2+ 6 4NO 2 4 3 3 Combustion: C H S+2 6O + H2O + SO 2 2 2 BALANCED: 2 C H S+ 9 O  4 CO + 6 H O + 2 SO  2 6 2 2 2 2 Single Replacement: Cu(NO ) + Zn 3 2 (NO ) + Cu 3 2 Chemistry Unit: Stoichiometry Avogadro’s Number (mol or N ): when aAsample of an element weights its atomic mass in grams, there are 6.022x10 23 atoms present Molar Mass of C H O = 12C+22H+11O= 342 g/mol 12 22 11 1 mol of a gas= 22.4 L at STP (standard temperature and pressure: 0 ̊C, 1 atm) Mass of 1 mol of Cu (CO 3 (OH)3 2… (632546 x 3) + (12.011 x 2) + (15.9994 x 8) + (1.0079 x 2) = 344.67 g/mol Molecular Formula: shows the actual number of each element in a formula Empirical Formula: shows the lowest whole number ratio of elements in a compound When asked to calculate the empirical formula, assume that 100 g of that substance exists! Limiting Reactant: Mole/ Coefficient ratio Ex. in C 2 +4O  2CO2+ 2H O Gi2en 56.02g C H and 128 g O 2 4 2 mol O /2coefficient O = (228/32) / 3 = 4/3 mol C H2/ 4oefficient C H = 256428) / 1 = 2 O 2s smaller so it’s the limiting reactant Formula Mass is in amu Percent Composition of X= (mass of X) / (mass of molecule) x 100 Ex. % H in H2O = (2 x 1.10079)/ (18.015) x 100 = 11% Percent Yield: actual yield/ theoretical yield x 100 Chemistry Unit: Acids and Bases Strong Acids - HCl - HBr - HI H SO - 2 4 - HNO 3 - HClO 4 CM-UY 1004 General Chemistry for Engineers Strong Bases - metal (Group I or II cations) hydroxide: NaOH, LiOH, KOH - metal oxide: Na O, K 2, MgO2 Properties of Acids - dissolves in H O 2 - usually found in foods - electrolytes - tastes sour - most common weak acid: acetic acid loses a proton  (conjugate) base - - K eqvalue is called the acidity constant (K ) and onlyadepends on temperature and solvent - Monoprotic: one acidic proton - Diprotic: two acidic proton Polyprotic: more than two acidic protons - Properties of Bases - Dissolves in H O 2 - Usually found in CLEANERS - Tastes bitter - Most common weak base: NH 3 Gains a proton  (conjugate) acid - Definitions - Arrhenius o + Acid: release hydronium (H O ) ions 3 o Base: release hydroxide (OH ) ions - - Bronsted-Lowry o + Gives H o Takes H + - Lewis o Electron-poor, accepts e - o - Electron-rich, donates e Naming Acids - Anion Ending (remove ending) o –ide  hyrdo_______ic acid o –ate  _______ic acid o –ite  ________ous acid Water - amphoteric: can act as acid and/or base Equations - pH + pOH = 14 M V H = M V OH - a a b b Neutralization - Acid + Base  Salt (aq) + H O 2 - + - Add OH to cation and H to anion to neutralize Strong Acid + Weak Base  Acidic Salt - - Weak Acid + Strong Base  Basic Salt - Strong Acid + Strong Base  Neutral Weak Acid + Weak Acid  DEPENDS! - CM-UY 1004 General Chemistry for Engineers Buffers - a solution of a weak acid and its conjugate base OR a weak base and its conjugate acid maintain a constant pH - - Ex. CH C3 H/C2 CO and3NH /2H 3 4+


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