Unit 2&3 Study guide
Unit 2&3 Study guide Chem 1010
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This 11 page Bundle was uploaded by Sam.Moyer on Monday November 9, 2015. The Bundle belongs to Chem 1010 at Tulane University taught by Heiko Jacobsen in Summer 2015. Since its upload, it has received 61 views. For similar materials see General Chemistry I in Chemistry at Tulane University.
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Date Created: 11/09/15
Unit 1 Study Guide (Chapter 1→4) Tuesday, September 29, 201510:17 PM Chapter 1 Classificationof Matter: (AKA: Solution) Measurement of Matter Derived Quantities Metric Prefixes Length Meter (m) Quantity Unit Formula Prefix MultiplicationFactor Scientific Notation Mass Kilogram (kg) Area Kilo k 1 000. Time Second (s) Volume — 0.0 Temp. Kelvin (K) Speed Deci d .1 Amount of Substance Mole (mol) Centi c .01 Acceleration Electric Current Ampere (A) Milli m .001 Candela (cd) Force Micro μ .000 001 Luminous Intensity Work (Energy) Pressure Power Density Temperature Straight-Line Scale Conversions Y: Desired Temperature Unit X: Reference Scale (temperature you are converting from) b: Conversion Fix Point (value of Y when X=0) m: Conversion Increment Absolute Zero: 0 K TemperatureConversions Using Straight-LineConversion -273.15 C° -459.67 F° Taking Measurements Uncertainties Systematic Error: Recurrent error due to flaws in equipment or method Random Error: Error made by an experimentor Accuracy & Precision Unit 1 Page 1 *Find Standard Deviation (σ) on calculator in stats program (sx) Scientific Notation Rules 1. All non-zeros are significant 2. Zeros between nonzero #s are significant 3. Zeros @ the end of the # to the right of the decimal are significant (indicates precision) 4. Exact #s have an unlimited # of sig figs 5. Precision can be neither gained nor lost in calculations *Even at the expense of sig fig rules‼ Accuracy: # of sig figs Precision: Position of the rightmost sig fid Chapter 2 Atoms & Atomic Theory Law of Conservation of Mass The total mass of a substance present after a chemical reaction is the same as the total mass of the substance before the reaction. Law of Constant Composition All samples of a compound have the same composition (same proportions of mass & moles) Atomic Models & Theories Dalton's Atomic Theory ○ Elements are composed of unchangeable & indivisible atoms ○ All atoms of an element are identical ○ Atoms combine in simple, whole # ratios ↓ Law of Multiple Proportions If 2 elements form more than a single compound: the mass of the 1 element, compared with the fixed mass of the 2 element are in the ratio of small whole #s. ↓ Cathode Rays The discovery of electrons ↓ Thomson's Atomic Theory ○ Atoms ar breakable ○ Atoms have structure ○ Electrons are suspended in a positively charged electric field ○ Mass of atoms is due to electrons ○ Atoms are mostly "empty space" ↓ Rutherford's Atomic Theory ○ The atom contains a tiny, dense center (nucleus), which essentially contains the entire mass of the atom ○ The nucleus is positively charged ○ The electrons are negatively charged & orbit around the nucleus Protons, Neutrons,Electrons, etc. Protons Positively charged particles in the nucleus = Electrons (if it's not an isotope or ion) Chemical Symbols: = Atomic # M# Neutrons A#E Neutral particles in the nucleus = (Atomic mass #) - (Atomic #) Atomic # = Protons Atomic Mass # (amu or u) = Protons + Neutrons Unit 1 Page 2 = (Atomic mass #) - (Atomic #) Atomic # = Protons Atomic Mass # (amu or u) = Protons + Neutrons = Electrons + Neutrons Isotopes Different forms of an element having the same # of protons, but different # of neutrons. (same atomic mass, differentmass #) Determining Protons, Neutron, & Electrons: Reference the symbol→ ○ Protons: The atomic # (same as one on the periodic table) ○ Electrons: Look @ the # of protons & the charge ○ Neutrons: (Mass # of isotope) — Protons Natural Percent Abundances The percentage of that particular isotope that occurs in nature. Intro to the Periodic Table Periods = Rows Similar Properties Groups/Families = Columns Corresponds w/ principal energy level Metals 80% of Periodic Table Nonmetals Halogens, Nonmetals, & Noble Gasses Noble Gasses Group 18 (Passive Elements) Metalloids Semimetals Moles & Avogrado's # Avogrado's # The # of representativeparticles in a mole Mole particles of a substances Molar Mass (g/mol) The mass of a mole of an element, compound, molecule, etc. Chapter 3 Chemical Compounds 2 Principle Types of Atoms 1. Metals ○ Tend to lose 1+ electrons→Cations(positive ion) ○ Electricity & heat conductors 2. Nonmetals ○ Tend to gain 1+ electrons→Anions(negative ion) *Metalloids ○ Gain or lose electrons *Noble Gases ○ Tend no neither gain nor lose electrons Formulas & Compounds Molecular Compound Compound made up of molecules. Uses covalent bonds (therefor it's a neutral group of atoms) Chemical Formula Symbolic representations that @ a minimum, indicate the elements present & relative proportions of atoms. Unit 1 Page 3 proportions of atoms. Empirical Formula The lowest whole # ratio of atoms in a compound. Macroscopic Level: Molar Microscopic Level: Atomic Molecular Formula Formula of the actual molecule of that compound. Unique to each compound Is the same or a multiple of the compound's empirical formula Structural Formula Shows the order in which atoms are bonded together & what types of bonds are present. Condensed Structural Formula A molecular formula in which connectivity is preserved When finding isomers from structural formulas, condense the formula to see which ones are the same! :) Formula Mass The mass of the formula unit (empirical formula) in atomic mass units (amu or u) s Molecular Mass Mass of molicule in atomic mass units (amu or u) Molar Mass Mass of one mole of compound. Percent Composition of Molar Mass: The contributionof each element to the molar mass (or atomic mass) Determining Empirical Formula with % Composition of Mass or Molar Mass: 3 Pinciple Types of Bonds Covalent Bond Electron sharing between 2 nonmetals Ionic Bond Electron transfer from metal→nonmetal Metallic Bond Electron sharing between metals Ionic Compounds Compounds consisting of cations & anions, joined together by electrostatic forces of attraction. • AKA: Salt • Nonmetal+Metal compounds are often ionic Diatomic Molecules H 0 N I Cl F Br 2 , 2 ,2 , 2 , 2 ,2 , 2 *Make sure you specify whether one is referring to 1 mol H or 1 mol2H Naming Compounds Organic Compound Contains either carbon or carbon & hydrogen & a small # of another element such as oxygen, sulfur, & nitrogen. Inorganic Compound Any non-organic compound. Binary Compound Composed of 2 elements. Ternary Compound Composed of 3 elements. Oxidation States Oxidation State: The # of electrons lost/gained/used by an atom Rules: 1. The total OS of all atoms is equal to the charge of the species. 2. The OS of an atom in a free element (uncombined)is 0. +1 -1 3. H —Nonmetals ; H —Metals 4. In compounds, Group 1 has an OS of +1 ; Group 2 has an OS of +2 5. In compounds, Fluorine's OS is -1 6. In compounds, Oxygen's OS is -2 7. In binary compounds with Metals or Hydrogen: Group 17's OS is -1 Group 16's OS is -2 Group 15's OS is -3 Unit 1 Page 4 7. In binary compounds with Metals or Hydrogen: Group 17's OS is -1 Group 16's OS is -2 Group 15's OS is -3 Polyatomic Ions (memorize these ones!) Ammonium NH 4+ Carbonate CO 3-2 Bicarbonate HCO 3- - Hydroxide OH - Nitrate NO 3 -3 Phosphate PO 4 Sulfate SO -2 4 Naming Compounds Hydrates: Prefix(mono, di…)hydrate @ the end of the compound name Hydrocarbons: Stem/Prefix: Meth, Eth, Prop, But, Pent,… Single bond: -ane ; Double+ Bond: -ene Alcohols: (has -OH) Stem—anol or Prefix + "alcohol" Carboxylic Acids: (has -COOH) Stem—oic or Trivial Name 3 Important Trivial Names ○ Water: H 2 ○ Ammonia: NH 3 ○ Acetic Acid: CH3COOH Isomers Molecules with the same molecular formula but different chemical structures. *WILL SEE DIFFERENCE IN CONDENSED STRUCTURAL FORMULA IF ISOMER Chapter 4 Chemical Reactions Balance the equation using stoichiometric coefficients Reflects a ratio, not absolute # of molecules→FRACTIONS are OK :) Spectator Species An ion that exists as a reactant & product. (Can be added/subtracted from equation freely) Chemical Equation Lingo ;) Reactants & Products joined by: →, =, or (reversible reaction) States of Matter Place @ end of each formula unit ○ Liquid: (l) ○ Gas: (g) ○ Solid: (s) ○ Aqueous Solution: (aq) Reaction Conditions Unit 1 Page 5 Reaction Conditions Any of these will be placed above or below the arrow. The basic condition or specific values can be used ○ Heat: ○ Pressure: ○ Catalyst (in form of compound): Stoichiometric Coefficients Placed at beginning of each formula unit. Limiting Reagent The reactant that determines the amount of product that can be formed by a reaction. (due to insufficient quantity) Reaction Yields Molarity, Concentration, & Dilution Types of Chemical Reactions Unit 1 Page 6 Reactions in Aqueous Solutions Wednesday, October 7, 2015 12:49 PM Solutions • Solution: A homogeneous mixture of solvent & solute. • Solutions exist in liquid, solid, or gas phase. • Aqueous Solution: Liquid solutions with water as solvent. Electrolytes • Electrolyte: A substance that ionizes when dissolved in water. ○ Ionization: The process in which an atom/molecule acquires ±charge by gaining/losing electrons (forming ions). ○ Strong Electrolyte: A substance that essentially completely ionizes in AQ. Almost all soluble ionic compounds Most ○ Weak Electrolyte: A substance that only partially ionizes in AQ. Molecular Most molecular compounds Compounds ○ Nonelectrolytes: A substance that doesn't ionize in AQ. Nonelectrolyte Weak Electrolyte Strong Electrolyt • An aqueous solution of an electrolyte conducts electricity. Acids, Bases, & Salts • Strong acids/bases are strong electrolytes. (Vice Versa) • Acid: Produce H +(aq)in AQ. - ○ An acid is a H ion donor n ○ Weak Acid: Has COOH (carboxyl group) • Base: Produce OH -(aq)in AQ. - ○ A base is a H ion acceptor ○ Weak Base: Produces OH -(aq)by reacting with water when it dissolves in it. + Ex) Ammonia NH (its 3eaction w/ water is reversible & incomplete) • Salt: Ionic Compound (Nonmetal+Metal) ○ Salts that are soluble are strong electrolytes. Water- A Polar Molecule • The O-H bond in water is covalent (sharing electrons), but the electrons aren't shared evenly→Polar Bond→Polar Molecule ○ Polar Bond: A bond where a pair of electrons is unequally shared between two atoms. The Hydrated Proton • NO FUCKIN CLUE MAN Ion Concentrations in Solutions • [x] = Concentration in Solution • Acidic Solution: A solution that has… + - ○ [H ] > [OH ] ○ [H ] > 1∙10 M-7 ○ [H ] > [H ]+water ○ pH > 7 • Basic Solution: A solution that has… - + ○ [OH ] > [H ] ○ [H ] < 1∙10 M-7 - - ○ [OH ] > [OH ] water ○ pH < 7 Solubility Rules for Ionic Compounds Compounds 1A Salts Nitrate Salts NO 3+ Sulfate Salts SO4-2 Chloride Salts Cl Fluoride Salts F- Ammonia NH 3 Chlorate Salts ClO 3 Bromide Salts Br - - - Perchlorate Salts ClO 4 Iodide Salts I Acetates AcO - Exceptions - Pb 2+Ag Hg 2+ Pb 2+ Ag Hg 22+ Pb 2+Ag Hg 2+ Pb2+ Ba Mg 2+ 2+ 2+ + 2+ + Sr Ca Ba Sr Ca Solubility Soluble Soluble Soluble Soluble Soluble -2 - -2 Compounds Carbonates CO 3 Hydroxides OH Chromates CrO 4 Phosphates PO 4-3 Sulfides S-2 Exceptions 1A Ions Ba2+ NH + 1A Ions NH + + 2+ 2+ 2+ 2+ NH 4 Sr Ca Mg Ca Solubility Insoluble Insoluble Insoluble • Rules in order of significance: 1. All alkali metals (Group 1) and ammonium compounds are soluble. 2. All acetate, perchlorate, chlorate, and nitrate compounds are soluble. 3. Silver, lead, and mercury(I) compounds are insoluble. 4. Chlorides, bromides, and iodides are soluble. 5. Carbonates, hydroxides, oxides, phosphates, silicates, and sulfides are insoluble. 6. Sulfates are soluble except for calcium and barium. Net Ionic Equations • In a molecular equation, it is not shown when the reactants & one of the products dissociate (separate) into cations/anions when dissolved in water. • Complete Ionic Equation: An equation that shows dissolved ionic compounds as dissociated free ions. Chapter 5 Page 7 Net Ionic Equations • In a molecular equation, it is not shown when the reactants & one of the products dissociate (separate) into cations/anions when dissolved in water. • Complete Ionic Equation: An equation that shows dissolved ionic compounds as dissociated free ions. • Spectator Ions: An ion that appears on both sides of an equation & is not directly involved in the reaction. ○ Cross these out to get the net ionic equation. • Net Ionic Equation: An equation for a reaction in an AQ that shows only the particles that are directly involved in the chemical change. Aqueous Metathesis • Metathesis Reaction: A reaction in which 2 ionic compounds displace one another in AQ. ○ AX + BY → BX + AY ○ The formation of a precipitant provides the driving force for aqueous metathesis Precipitation Reactions • Precipitate: An insoluble solid that emerges from a liquid solution. ○ Sediment: The matter that settles in the bottom of a liquid Neutralization Reactions • Neutralization Reaction: A reaction between acid & base to form water & an ionic compound (salt) • The salt will have an anion from the acid & a cation from the base. CH3COOH(aq) + OH- (aq) → CH3COO- (aq) + H2O(l) Titration • Equivalence Point: When the # of moles from the 1st reactant is 0 & any further addition of the 2nd reactant will be excess. (NO LIMITIING OR EXCESS REAGANT) ○ For acid-base solutions the # of moles of H ions is equal to the # of moles of OH ions. Moles of base(added) = moles of acid(initial) • Titration: The process of adding a known amount of solution of known concentration to determine the concentration of another solution. ○ End Point (of titration): The point at which the indicator changes color.→Indicating that neutralization has occurred (Equivalence Point) Slide 8-9 • IDK WHAT THIS IS TALKING ABOUT Redox Reactions • Redox Reaction: An oxidation-reductionreaction that is comprised of 2 parts, a reduction half & an oxidation half, that always occur together. ○ Oxidation: Loses electrons→OS increase.(oxidant) ○ Reduction: Gains electrons→OS decrease.(reductant) • Redox reactions are governed by the # of electrons transferred. Balancing Redox Reactions • Balanced By Inspection ○ Most redox reactions can be determined by balanced by inspection Balanced inspection considers 2 conservation factors: □ # of Atoms (conservation of mass) □ Total Charge (conservation of charge) ○ Process: Balance the reaction in parts in terms of atoms & charge • Half Equation Method ○ Every redox reaction is made up of two half-reactions: 1. The Oxidation Process 2. The Reduction Process ○ Process: 1. Formulate the 2 half-equations. i. Achieve mass balance for each half-equation ii. Achieve charge balance by adding electrons It is the transfer of electrons that links the 2 half-equations This produces the half-reactions (HR) 2. Adjust the coefficients in both half-reactions so the same # of electrons are in each half-reaction. 3. Add together the 2 adjusted half-reactions to obtain the balanced overall equation. • Oxidation State Methods ○ Oxidation State vs Half-Equation Method Half-Equation: Electrons added last to achieve charge balance Oxidation State: Electrons added first to account for oxidation state changes of reductant & oxidant ○ Since H & O are key players in oxidation & reduction… Half-equationsare balanced for oxygen by adding O 2- + Half-equationsare balanced for hydrogen by adding H ○ Process: 1. Assign OS's to each element/group;identify the species that are oxidized & reduced 2. Write separate, unbalanced equations for the oxidation & reduction half-reactions. □ Do not combine reduction & oxidation processes in one half-reaction! 3. For each half-reaction, add electrons to one side or the other to account for the change in OS (accounting for the # of electrons produced/consumed) 4. Balance the separate half-equationsfor the elements changing OS & update the electron flow 5. Balance O by adding O ; Balance H by adding H . + 6. Finish balancing & adjust coefficients in both half-reactions so that the same # of electrons appear in each half-reaction □ Check mass & charge balance 7. Combine the half-equations Chapter 5 Page 8 7. Combine the half-equations i. Cancel electrons & spectators ii. Combine H & O →OH ; Combine H & OH →H O + - 2 2- - iii. In the presence of water, combine additional O with H O to form 2OH2 If needed add H O to both sides! 2 8. Neutralize anything left… - + □ If acid solution: neutralize OH → add H on both sides □ If basic solution: neutralize H → add OH on both sides 9. Cancel any spectators; check mass & charge balance. Overview of Balancing Methods • Chapter 5 Page 9 Thermochemistry Monday, October 12, 20151:01 PM Energy • Energy: The capacity for doing work • 2 Basic Types: ○ Potential Energy: The work within that could come out/be done. ○ Kinetic Energy: The work within that does come out/is done. • Energy is an extensive property Name Unit Symbol Formula Momentum & Force & Energy • The 2 fundamental properties of an object in motion: Mass: m & Velocity: v Distance m - ○ Mass m (kg) Mass kg - ○ Velocity v (m/s): distance traveled per time 2 2 - Joule kg∙m /s 2 Newton kg∙m/s - 2 -2 • Momentum p (kg∙m/s) : combines mass & velocity. Gravitational Constant N∙m /kg • Vector Quantities: Possessing a direction & magnitude Coulomb's Constant N∙m ∙C 2 ○ Momentum & Velocity Velocity m/s • In a closed system, the total momentum if constant 2 • Force F (kg∙m/s ): Change in momentum per time unit Momentum kg∙m/s 2 Acceleration m/s d = derivative (could be rate of change) • If a Force is applied to a particle for a time interval (Δt)… Force N If Then • Acceleration a (m/s ): The rate of change of velocity per unit of time. Kinetic Energy J Potential Energy J • Integrating the force over the length traveled results in kinetic energy ○ Kinetic Energy KE (kg∙m /s ): The energy possessed by an object due to its motion Gravity of Earth a g Gravity Force N Fundamental Forces Gravity Force on Earth N • Gravity Force (physics) Electromagnetic Force ??? ○ Gravitational Constant: G -11 2 -2 N = Newtons G = 6.67∙10 N∙m /kg ○ Law of Universal Gravitation Force of Ionic Attraction ??? r = distance between the centers of masses Gravity Force: g Pressure Pa (Pascal) P Gravity Force on Earth:GE Work J w Voltage J/C V ○ Gravity of Earth: g g = 9.81 m/s 2 • Electromagnetic Force (chemistry) ○ Coulomb's Law of Electrostatic Force: The magnitude of the electrostatic force of interaction between two point charges is directly proportional to the scalar multiplication of the magnitudes of charges and is inversely proportional to the square of the distance between them. Scalar: having only magnitude, not direction (of a quantity) The electrostatic force ( ) uses the same unit as gravity (1/r ) but is ….??? Coulomb's Constant (k) Electromagnetic Force ( ) q & Q= electric charges r= distance between the centers of charges ○ Force of Ionic Attractio+/-F ) P p ti nal t Stronger the charg→ Stronger Force Smaller the io→ Stronger Force Conservation of Energy • Energy can nether be created nor destroyed ○ Energy can be transformed into various types, but the total energy remains constant. • Potential Energy (PE) 2 g = free-fall acceleration (gravity of earth=9.81 m/s ) h = height Work • Work w (J) : The energy required to move an object against a force. A form of energy transfer that might be expressed as a force acting through a ○ distance ○ Unit: Joules • The product of Pressure (P) & Volume (V) has the units of work ○ Pressure (P): The amount of force acting per unit of area • No change in volume = no work (P∙ΔV) • Types of Work 3 Expansion: Pa∙m -Pexexternal pressur∙) ΔV Extension: N∙m F∙Δistance Elevation: kg∙(m∙s ∙ m m∙ ∙Δeight Electrical: V∙ΔQ Voltage∙Δharge Exothermic & Endothermic Processes • System: A region containing a specific amount of matter whose behavior is being observed. • Surroundings: The portions of matter external to the system. • Boundary: A real/imaginary surface covering the volume that contains the system. Chapter 7 Page 10 • Boundary: A real/imaginary surface covering the volume that contains the system. ○ The thermometer functions on this principle, evaluating the temperature (T) of a system. Heat Law & Theory • Caloric Theory: Heat consists of a subtle fluid called caloric, which can be transferred from one body to another, but cannot be created nor destroyed ○ If the heat of a system decreases, then the heat of the surroundings must increase. • Zeroth Law of Thermodynamics: If 2 systems are in thermal equilibrium with a 3rd system, they must be in thermal equilibrium with each other. ○ Heat always flows from a warmer object to a cooler object. If the object remains in contact, heat will flow until the objects are reach a thermal equilibrium. Units of Heat Measurement • Heat q: Energy that transfers from one object to another because of a temperature difference between them. ○ There are 2 common units of heat measurement: calories & joules • Calorie cal: The quantity of heat required to raise 1 gram of water 1°C. • Joule J: The quantity of heat requires to raise 1 gram of water…????? Energy • Energy: The capacity for doing work Energy • Energy: The capacity for doing work Energy • Energy: The capacity for doing work Energy • Energy: The capacity for doing work Chapter 7 Page 11
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