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Numbers and Scientific Notation

• Created by Mitchell Foltz , last modified by Leta R Moser on May 13, 2021

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Some questions you may encounter in Quest will use numeric free response answers that you type in. The following example is a numeric type question on a learning module slide, but these questions can appear on all types of Quest assignments.

• If your instructor has allowed retries, you will be given 7 attempts
• Most solution answers are at least six digits (unless significant figures are relevant to the question/otherwise denoted) 
• For credit your answer must be within 1% of the correct answer, unless tolerance is otherwise denoted (so entering in four digits to the right of the decimal is usually sufficient)
• Start with at least four significant digits for numeric entry; your response must be within 1% of the correct answer unless otherwise designated
• Scientific notation may use the format of "e" or "x10^"
• Comma use is fine.
• Do not use symbols in solutions (ie do not use $ in monetary solutions, but you can write 'dollars'-or the specified 'answer in units of ---- '  stated in the question, after your numeric answer).
• If offered, use the function pallet. If you don't see the pallet, plan on entering an actual number calculated out or simple expression (ex: 3x-5).

For applicable practice, your professor may opt to include a short 3 part question ( sig fig practice , #222082 ) for you to get use to what is and is not acceptable:

Note that if you get a 'that response has already been entered' message, try to use another way to say the same thing (ie, if used 10^, try e). If you continue to get the 'that response has already been entered' accept that it is not correct and try again. This is a safeguard in place so you don't spend all your tries insisting on an answer that is incorrect.

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Hours Tuesday – Wednesday 10AM-4PM

• Introduction
• 1.1 Chemistry in Context
• 1.2 Phases and Classification of Matter
• 1.3 Physical and Chemical Properties
• 1.4 Measurements
• 1.5 Measurement Uncertainty, Accuracy, and Precision
• 1.6 Mathematical Treatment of Measurement Results
• Key Equations
• 2.1 Early Ideas in Atomic Theory
• 2.2 Evolution of Atomic Theory
• 2.3 Atomic Structure and Symbolism
• 2.4 Chemical Formulas
• 2.5 The Periodic Table
• 2.6 Molecular and Ionic Compounds
• 2.7 Chemical Nomenclature
• 3.1 Formula Mass and the Mole Concept
• 3.2 Determining Empirical and Molecular Formulas
• 3.3 Molarity
• 3.4 Other Units for Solution Concentrations
• 4.1 Writing and Balancing Chemical Equations
• 4.2 Classifying Chemical Reactions
• 4.3 Reaction Stoichiometry
• 4.4 Reaction Yields
• 4.5 Quantitative Chemical Analysis
• 5.1 Energy Basics
• 5.2 Calorimetry
• 5.3 Enthalpy
• 6.1 Electromagnetic Energy
• 6.2 The Bohr Model
• 6.3 Development of Quantum Theory
• 6.4 Electronic Structure of Atoms (Electron Configurations)
• 6.5 Periodic Variations in Element Properties
• 7.1 Ionic Bonding
• 7.2 Covalent Bonding
• 7.3 Lewis Symbols and Structures
• 7.4 Formal Charges and Resonance
• 7.5 Strengths of Ionic and Covalent Bonds
• 7.6 Molecular Structure and Polarity
• 8.1 Valence Bond Theory
• 8.2 Hybrid Atomic Orbitals
• 8.3 Multiple Bonds
• 8.4 Molecular Orbital Theory
• 9.1 Gas Pressure
• 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
• 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
• 9.4 Effusion and Diffusion of Gases
• 9.5 The Kinetic-Molecular Theory
• 9.6 Non-Ideal Gas Behavior
• 10.1 Intermolecular Forces
• 10.2 Properties of Liquids
• 10.3 Phase Transitions
• 10.4 Phase Diagrams
• 10.5 The Solid State of Matter
• 10.6 Lattice Structures in Crystalline Solids
• 11.1 The Dissolution Process
• 11.2 Electrolytes
• 11.3 Solubility
• 11.4 Colligative Properties
• 11.5 Colloids
• 12.1 Chemical Reaction Rates
• 12.2 Factors Affecting Reaction Rates
• 12.3 Rate Laws
• 12.4 Integrated Rate Laws
• 12.5 Collision Theory
• 12.6 Reaction Mechanisms
• 12.7 Catalysis
• 13.1 Chemical Equilibria
• 13.2 Equilibrium Constants
• 13.3 Shifting Equilibria: Le Châtelier’s Principle
• 13.4 Equilibrium Calculations
• 14.1 Brønsted-Lowry Acids and Bases
• 14.2 pH and pOH
• 14.3 Relative Strengths of Acids and Bases
• 14.4 Hydrolysis of Salt Solutions
• 14.5 Polyprotic Acids
• 14.6 Buffers
• 14.7 Acid-Base Titrations
• 15.1 Precipitation and Dissolution
• 15.2 Lewis Acids and Bases
• 15.3 Multiple Equilibria
• 16.1 Spontaneity
• 16.2 Entropy
• 16.3 The Second and Third Laws of Thermodynamics
• 16.4 Free Energy
• 17.1 Balancing Oxidation-Reduction Reactions
• 17.2 Galvanic Cells
• 17.3 Standard Reduction Potentials
• 17.4 The Nernst Equation
• 17.5 Batteries and Fuel Cells
• 17.6 Corrosion
• 17.7 Electrolysis
• 18.1 Periodicity
• 18.2 Occurrence and Preparation of the Representative Metals
• 18.3 Structure and General Properties of the Metalloids
• 18.4 Structure and General Properties of the Nonmetals
• 18.5 Occurrence, Preparation, and Compounds of Hydrogen
• 18.6 Occurrence, Preparation, and Properties of Carbonates
• 18.7 Occurrence, Preparation, and Properties of Nitrogen
• 18.8 Occurrence, Preparation, and Properties of Phosphorus
• 18.9 Occurrence, Preparation, and Compounds of Oxygen
• 18.10 Occurrence, Preparation, and Properties of Sulfur
• 18.11 Occurrence, Preparation, and Properties of Halogens
• 18.12 Occurrence, Preparation, and Properties of the Noble Gases
• 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
• 19.2 Coordination Chemistry of Transition Metals
• 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
• 20.1 Hydrocarbons
• 20.2 Alcohols and Ethers
• 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
• 20.4 Amines and Amides
• 21.1 Nuclear Structure and Stability
• 21.2 Nuclear Equations
• 21.3 Radioactive Decay
• 21.4 Transmutation and Nuclear Energy
• 21.5 Uses of Radioisotopes
• 21.6 Biological Effects of Radiation
• A | The Periodic Table
• B | Essential Mathematics
• C | Units and Conversion Factors
• D | Fundamental Physical Constants
• E | Water Properties
• F | Composition of Commercial Acids and Bases
• G | Standard Thermodynamic Properties for Selected Substances
• H | Ionization Constants of Weak Acids
• I | Ionization Constants of Weak Bases
• J | Solubility Products
• K | Formation Constants for Complex Ions
• L | Standard Electrode (Half-Cell) Potentials
• M | Half-Lives for Several Radioactive Isotopes

The temperature of 1 gram of burning wood is approximately the same for both a match and a bonfire. This is an intensive property and depends on the material (wood). However, the overall amount of produced heat depends on the amount of material; this is an extensive property. The amount of wood in a bonfire is much greater than that in a match; the total amount of produced heat is also much greater, which is why we can sit around a bonfire to stay warm, but a match would not provide enough heat to keep us from getting cold.

Heat capacity refers to the heat required to raise the temperature of the mass of the substance 1 degree; specific heat refers to the heat required to raise the temperature of 1 gram of the substance 1 degree. Thus, heat capacity is an extensive property, and specific heat is an intensive one.

(a) 47.6 J/°C; 11.38 cal °C −1 ; (b) 407 J/°C; 97.3 cal °C −1

1310 J; 313 cal

(a) 0.390 J/g °C; (b) Copper is a likely candidate.

We assume that the density of water is 1.0 g/cm 3 (1 g/mL) and that it takes as much energy to keep the water at 85 °F as to heat it from 72 °F to 85 °F. We also assume that only the water is going to be heated. Energy required = 7.47 kWh

lesser; more heat would be lost to the coffee cup and the environment and so Δ T for the water would be lesser and the calculated q would be lesser

greater, since taking the calorimeter’s heat capacity into account will compensate for the thermal energy transferred to the solution from the calorimeter; this approach includes the calorimeter itself, along with the solution, as “surroundings”: q rxn = −( q solution + q calorimeter ); since both q solution and q calorimeter are negative, including the latter term ( q rxn ) will yield a greater value for the heat of the dissolution

The temperature of the coffee will drop 1 degree.

5.7 × × 10 2 kJ

-2.2 kJ; The heat produced shows that the reaction is exothermic.

22.6. Since the mass and the heat capacity of the solution is approximately equal to that of the water, the two-fold increase in the amount of water leads to a two-fold decrease of the temperature change.

1.4 × × 10 2 Calories

The enthalpy change of the indicated reaction is for exactly 1 mol HCL and 1 mol NaOH; the heat in the example is produced by 0.0500 mol HCl and 0.0500 mol NaOH.

25 kJ mol −1

81 kJ mol −1

1.83 × × 10 −2 mol

–802 kJ mol −1

−495 kJ/mol

44.01 kJ/mol

90.3 kJ/mol

(a) −1615.0 kJ mol −1 ; (b) −484.3 kJ mol −1 ; (c) 164.2 kJ; (d) −232.1 kJ

−54.04 kJ mol −1

−2660 kJ mol −1

On the assumption that the best rocket fuel is the one that gives off the most heat, B 2 H 6 is the prime candidate.

(a) C 3 H 8 ( g ) + 5 O 2 ( g ) ⟶ 3 CO 2 ( g ) + 4 H 2 O ( l ) ; C 3 H 8 ( g ) + 5 O 2 ( g ) ⟶ 3 CO 2 ( g ) + 4 H 2 O ( l ) ; (b) 330 L; (c) −104.5 kJ mol −1 ; (d) 75.4 °C

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