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ISU / Chemistry / CHEM 177 / How many significant figures are in the number 4001?

How many significant figures are in the number 4001?

How many significant figures are in the number 4001?

Description

School: Iowa State University
Department: Chemistry
Course: General Chemistry I
Professor: Anderson
Term: Fall 2016
Tags: General Chemistry, final exam, Chem177, and Iowa State
Cost: 50
Name: Final Exam Chem 177_Cumulative
Description: It’s that time of year… FINALS This study guide was created and designed to assist you somewhat with the online review on Bb (Blackboard). I suggest you utilize both at the same time. It was not created to give you the answers and exact exam questions. It was designed to help you study effectively. Thank you & Good luck Happy Studying
Uploaded: 12/11/2016
37 Pages 278 Views 5 Unlocks
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It’s that time of year… FINALS


How many significant figures are in the number 4001?



Before we begin, let me start by saying that this study guide was created and designed to assist you somewhat with the online review on Bb. I suggest you utilize both at the same time. Thank you & Good luck

Let’s start off with some helpful hints and tips for studying:

✵ Even though finals are stressful it is

important to remember to sleep and eat

your brain food. ​ Sleep deprivation is


The symbol for the element sodium is what?



We also discuss several other topics like Transmission genetics means what?

detrimental to your brain's ability to

formulate and generate memories

✵ Practice old exams! Practice makes

perfect

✵ Dr. Anderson explained that even

though the final is all multiple choice

there is still a need for calculations to

be done, so do not skip out

✵ ​ There. Is. No. Curve. BUT recitation

points will be used to bump people up.

✵ If you wish to review the exams and

you do not have your own, Dr.

Anderson still has all of the exam keys taped up outside of his office in Howe Hall.

NOW, LET’S BEGIN 


1 kilogram is how many milligrams?



Recitation Quiz One Review

Don't forget about the age old question of What are the levels of organization in the body?

Recitation Quiz Two Review

Recitation Quiz Three ReviewDon't forget about the age old question of Why does methylation turn dna into heterochromatin?

Recitation Quiz Four Review

Recitation Quiz Six Review

Recitation Quiz Seven ReviewIf you want to learn more check out What is the diffusion of urbanization?
Don't forget about the age old question of What is fascism?

Recitation Quiz Eight Review

Chapter Reviews

Chapter One 

Chemistry is the study of matter and the changes that matter undergoes. The first chapter is an overview of what chemistry is about and chemists do. If you want to learn more check out What is true about habituation?

1.1 The study of chemistry

1.2 classifications of matter: fundamental ways to classify matter and distinguish the differences between pure substances and mixtures and between elements and compounds ● Homogenous vs. Heterogeneous mixtures

● Compounds are composed of two or more elements

1.3 properties of matter: different characteristics, or properties used to describe solids, liquids and gases. Distinguishing between chemical and physical properties

Physical Properties: weight, dimensions, area, mass, density, surface, etc. Chemical Properties: reactivity, electronegativity, etc.

Liquids = viscous, turbid, etc.

Gases = atoms spread farther apart, vapor, etc.

Solids = how hard, brittle, texture

1.4 units of measurement: many properties rely on quantitative measurements involving numbers and units; metric system

1.5 significant figures: Rules to SigFigs

1. Non-zero digits are always significant.

2. Any zeros between two significant digits are significant.

3. A final zero or trailing zeros in the decimal portion ONLY are significant. 1.6 dimensional Analysis

Chapter Two 

Introduction of the basic structure of atoms and the formation of molecules and ions. 21. The atomic theory of matter: matter is composed of infinitely small particles called atoms. 2.2 The discovery of the atomic structure: remember plum pudding?? - J.J. Thompson

2.3 The modern view of the atomic structure: the ideas of atomic numbers, mass numbers and isotopes

● What is an isotope again? - each of two or more forms of the same element that contain equal numbers of protons but dif erent numbers of neutrons in their nuclei. 

○ Uranium ring a bell? 

2.5 The periodic table:

2.6 Molecules and molecular compounds: the assemblies of atoms into arrangements called molecules and how their compositions can be represented by empirical and molecular formulas. ●

2.7 Ions and ionic compounds: atoms gain and lose electrons to form ions. Periodic table helps us predict the charges of ions and the empirical formulas of ionic compounds 

2.8 Naming inorganic substances: Nomenclature- rules for naming a substance

Chapter Three 

Introducing into stoichiometry

3.1 chemical equations: how to use chemical formulas to write chemical equations & 3.2 simple patterns of chemical reactivity: examining combustion, combination and decomposition reactions

● Combustion- an exothermic reaction in which something reacts with oxygen ● Decomposition- in which a single compound breaks down into its simplest form ● Combination- two simple compounds or elements form to create an entirely new substance

3.3 formula weights: adding up the atomic weights of each element in a compound

3.4 Avogadro's number and the mole:

● What is Avogadro's Number?- 6.02 • 1023 = 1 mole

● What is Molar Mass?- a physical property defined as the mass of a given substance divided by the amount of total substance

3.5 Empirical formulas from analysis: apply the mole concept to determine the empirical formula of a compound

● How to find an empirical formula

3.6 Quantitative information from balanced equations: Using balanced equations to predict the amount substances consumed and produced in a reaction

3.7 Limiting reactants:

Chapter Four 

4.1 general properties of aqueous solutions: examine whether substances dissolve or do not dissolve in water and whether they exist as ions, molecules or a mixture of the two

● Solvent: the thing doing the dissolving

● Solute: the dissolved

● Electrolytes and nonelectrolytes: a substance such as NaCl whose aqueous solutions contain ions is called an electrolyte. A substance such as C12H22O10 that does not form ions in a solution is called a nonelectrolyte. The dif erent classifications of the 2 arise largely because NaCl is an ionic compound, whereas C12H22O10 is a molecular compound

4.2 Precipitation reactions: when subtle reactants yield an insoluble product ● A precipitate is an insoluble solid formed by a reaction in a solution

● Solubility Rules and Guidelines

4.3 Acids, bases and neutralization reactions

● Think titrations, when you drip an acid into a base and wait for a subtle color change ● Acids form Hydrogen H+

● Bases form Hydroxide OH

Table of Strong Acids

Table of Strong Bases

4.4 Oxidation- reduction reactions: examine the reactions in which electrons transferred from one reactant to another

Practice:

4.5 Concentrations of solutions:

● Molarity (M): expresses the concentrations of a solution as a number of moles of solute in a liter of a solution

M = moles of solute / volume of solution in liters

Example:

“Suppose we had 1.00 mole of sucrose (it's about 342.3 grams) and proceeded to mix it into some water. It would dissolve and make sugar water. We keep adding water, dissolving and

stirring until all the solid was gone. We then made sure that when everything was well-mixed, there was exactly 1.00 liter of solution.”

What would be the molarity of this solution?

1. The answer is 1.00 mol/L. Notice that both the units of mol and L remain. Neither cancels.

a. A symbol for mol/L is often used. It is a capital M. So, writing 1.00 M for the answer is the correct way to do it.

2. And never forget this: replace the M with mol/L when you do calculations. The M is the symbol for molarity, the mol/L is the unit used in calculations.

Example:

Calculate the molarity of 25.0 grams of KBr dissolved in 750.0 mL.

4.6 Solution stoichiometry and chemical analysis: titration

Example of Acid-Base Titration: Calculating Molarity from Titration Data

“Titration reveals that 11.6ml of 3.0 M sulfuric acid are required to neutralize the sodium hydroxide in 25.00ml of NaOH solution. What is the molarity of the NaOH solution? “

H2SO4(aq) + 2NaOH(aq) → 2H2O(l) + Na2SO4(aq)

? mol NaOH = 11.6ml H2SO4 soln (103 ml) ( 3.0 mol H2SO4 ) (2 mol NaOH ) 1L NaOH soln 25 ml NaOH soln ( 1L ) ( 103 ml H2SO4 soln ) ( 1 mol H2SO4 ) ANS = 2.8 M NaOH

Chapter Five 

5.1 The nature of energy: the form energy takes, notably kinetic energy and potential energy and the fact that energy can be used to do work.

● Work on the system

● Exothermic vs. Endothermic Reactions

5.2 The first law of thermodynamics: energy cannot be created or destroyed but can be transformed from one form to another

● The energy possessed by a system is called internal energy

5.3 Enthalpy: state function known as enthalpy that measures the quantity of heat gained or lost by a system in a process occurring under constant pressure & 5.4 Enthalpies of a reaction: enthalpies of the products minus enthalpies of the reactants

● Some reactions give of so much energy (primarily as heat) that they are explosive, other reactions give off only a little bit of heat, still other reactions don’t take place unless we add heat from the surroundings. It’s very useful information to know how much heat a reaction will give off/absorb (if nothing else so you won’t blow yourself up in the laboratory). This quantity is called reaction enthalpy. 

Example: Consider the combination reaction between two moles of hydrogen and one mole of oxygen to create two moles of water: 

2H2(g) + O2(g) → 2H2O(g) ........................ DH = -483.6 kJ 

You can see that DH < 0 so this is an exothermic reaction (the reaction gives off heat). Of course if you were in class on the day that we ignited balloons filled with H2(g)/O2(g) mixtures you know that this reaction releases heat to the surroundings.

● Born-Haber Model 

5.5 Calorimetry: an experimental technique used to measure heat changes in chemical processes 

Example: 

“The temperature of a calorimeter increases 0.10 K when 7.52 J of electric energy is used to heat it. What is the heat capacity of the calorimeter?” 

1. Solution: 

a. Dividing the amount of energy by the temperature increase yields the heat capacity, C, 

b. C = 7.52 / 0.10 = 75.2 J/K. 

2. Discussion: We often compare the heat capacity of a calorimeter to that of a definite amount of water. The heat capacity of 75.2 J/K for the calorimeter is equivalent to the heat capacity of 1 mole (18 g) of water (18 g mol-1*1 cal (g K)-1*4.184 J cal-1 = 75.3 J (K mol)-1. 

3. Do you know that the electric energy = q * dV = i * dt * dV, where q is charge; dV, voltage; i, current; and t, time.

5.6 Hess’s Law: states that regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes. This law is a manifestation that enthalpy is a state function 

Example: 

“The heats of combustion of C, H2 and CH4 at 298 K and 1 atm are respectively -393 kJ/mol, -286 kJ/mol and -892 kJ/mol. What is the enthalpy of formation for CH4?” Solution: 

1. The three combustion reactions are: 

C + O2 ---> CO2 ΔH = -393 

kJ 

H2 + (1/2)O2 ---> H2O ΔH = -286 

kJ 

CH4 + 2O2 ---> CO2 + 2H2O 

ΔH = -892 kJ 

2. The reaction we're looking for is: 

C + 2H2 ---> CH4 

3. Discussion: 

● This is a Hess's Law problem. If you multiply the first reaction by 1, the second by 2, and the third by negative 1 (write it backwards) they add together to give the reaction you're looking for. So, the enthalpy of the reaction you're solving for is equal to 1(-393) + 2(-286) + (-1)(-892). I'll let you finish it, the critical thing is understanding where the 1, 2, and -1 came from. 

5.7 Enthalpies of formation: how to establish standard values for enthalpy change in chemical reactions and how to use them to calculate enthalpy changes for reactions ● Using the table in the back of the chemistry book, sound familiar?? 

Chapter Six 

6.1 The wave nature of light: (radiant energy, or electromagnetic radiation) wave like properties and is characterized by wavelength, frequency and speed.

● Frequency = f

● C = velocity of light

● = wavelength

6.2 Quantized energy and photons: interaction of light with metal surfaces, we recognize that electromagnetic radiation also has particle-like properties and can be described as photons, particles of light

● Planck’s Constant(h) = 6.62607004 × 10-34 m2 kg / s 

● E = hv ; energy of photon 

6.3 Line spectra and the Bohr Model: 

● Rydberg equation: 

● RH being a constant (1.096776 x 107 m -1) 

● N1 and n2 are positive integers (n2 being the larger of the two) 

The Bohr’s Model 

6.4 The wave behavior of matter 

● = h/ mv 

○ h = a constant 

○ m = mass 

○ v = velocity 

○ = wavelength 

6.5 Quantum mechanics and atomic orbitals:

6.6 Representations of orbitals 

6.9 Electron configurations and the periodic table

Chapter Seven 

Mostly centered around periodic tables and their many trends

● Zeff = Z - S

Chapter Eight 

Basic concepts of chemical bonding

8.1 Lewis structure: description of the three main types of chemical bonds 1. Ionic

2. Metallic

3. Covalent

Lewis is a shorthand way to keep track of the valence electrons

8.2 Ionic bonding: atoms are held together by electrostatic attractions between the ions of opposite charges

8.3 covalent bonding: bonding between molecules that share one or more electrons 8.4 Bond polarity and electronegativity: electronegativity is the compounds ability to attract electrons to itself

● Polar covalent:  8.5 Drawing Lewis Dot Structures

8.6 Resonance structures:

Key equations:

● Eel = kQ1Q2 / d ← the potential energy of 2 interacting charges

● µ = Qr ← the dipole moment of 2 charges of equal magnitude but opposite sign, separated by the distance (r)

● Formal charge = valence electrons - ½ (bonding electrons) - nonbonding electrons ● (delta)Hrxn = (sigma)(bond enthalpies of bonds broken) - (sigma)(bond enthalpies of bonds formed) ← the enthalpy change as a function of bond enthalpies for reactions involving gas-phase molecules

Chapter Nine 

9.1 Molecular shapes

● Linear

● Trigonal planar

● Tetrahedral

● Trigonal bipyramidal

● Octahedral

9.2 VSEPR model

9.3 Molecular shape and molecular polarity

Example:

“Decide whether the molecules represented by the following formulas are polar or nonpolar. (You may need to draw Lewis structures and geometric sketches to do so.)” 

a. CO2 

Solution: 

a. The Lewis structure for CO2 is 

The electronegativities of carbon and oxygen are 2.55 and 3.44. The 0.89 dif erence in electronegativity indicates that the C-O bonds are polar, but the symmetrical arrangement of these bonds makes the molecule nonpolar. 

If we put arrows into the geometric sketch for CO2, we see that they exactly balance each other, in both direction and magnitude. This shows the symmetry of the bonds. 

9.4 Covalent bonding and orbital overlap: 

Steps to drawing an orbital overlap: 

1. Draw the atoms in the same positions as the Lewis Dot Structure

2. Draw the s orbitals for H and the hybrid orbitals for all the other atoms 9.6 Multiple Bonds 

● Sigma bond vs. pi bond 

● Pi bonds are usually weaker than sigma bonds 

● Single bond = one sigma bond 

● Double bond = one sigma bond and one pi bond 

● Triple bond = one sigma bond and two pi bonds 

Chapter Ten

Chapter Eleven 

11.1 Molecular comparison of gasses, liquids and solids

11.2 Intermolecular forces: four intermolecular forces

1. London dispersion forces: attractive forces that arise as a result of temporary dipoles induced in atoms or molecules

2. Dipole-dipole: attractive forces between polar molecules

3. Ion-dipole: attractive forces between an ion and a polar molecule

4. Hydrogen bonds: hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a dif erent molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine, which has a partial negative charge.

11.3 Select properties of liquids : properties of liquids that are related to the nature of the atoms of a liquid. i.e. viscosity or surface tension

11.4 Phase changes: energy associated with changing phases. Gas to solid to liquid to gas etc. ● Melting is sometimes called heat of fusion or enthalpy of fusion

11.5 Vapor Pressure: dynamic equilibrium that exists between a liquid and its gaseous state and an introduction of vapor pressure

11.6 Phase diagrams: learn how to read phase diagrams

11.7 Liquid crystals:

● There are many classes and subclasses of liquid crystals

● In a nematic phase (the term means "thread-like") the molecules are aligned in the same direction but are free to drift around randomly, very much as in an ordinary liquid. ● In smectic ("soap-like") phases the molecules are arranged in layers, with the long molecular axes approximately perpendicular to the laminar planes. The only long-range

order extends along this axis, with the result that individual layers can slip over each other (hence the "soap-like" nature) in a manner similar to that observed in graphite.

Chapter Twenty One 

Nuclear Chemistry

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