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Chapter 1 Notes (Chemical Foundations)
1.1→
a.
b.
2. 1.2→
a.
1.
b.
3. 1.3→
a.
b.
Chemistry: An Overview
Matter is composed of various types of atoms
One substance changes to another by reorganizing the way the atoms are attached to each other
The Scientific Method
Steps
i.
Scientific Models
i.
A model is:
1. A set of hypotheses that agrees with various observations which explains some part of
some natural phenomenon
2. An interpretation of observations; reasoning for a natural occurrence
ii.
Models change over time even if the natural phenomenon does not, as scientists are constantly
finding new information
iii.
A natural law and a theory/model are NOT the same
1. Natural law summarizes what happens; a theory/model attempts to explain why
something happens
Units of Measurement
A measurement has two parts: a number and a unit
SI system
i.
Systeme International
ii.
System based on the metric system
iii.
Prefixes are used to change the size of a measurement
prefix
symbol
meaning
mega
M
106
kilo
k
103
hecto
h
102
deka
da
101
deci
d
10-1
centi
c
10-2
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milli
m
10-3
micro
μ
10-6
nano
N
10-9
iv.
Mass v Weight
1. Mass is a measure of the resistance of an object to a change in its state of motion;
weight is the force that gravity exerts on an object
2. Not interchangeable
4. 1.4→ Uncertainty in Measurement
a. Certain v. Uncertain digits
i.
Certain: same regardless of who reads measurement; uncertain: estimated value which varies
based on who reads it
ii.
A measurement always has some degree of uncertainty
b. Significant Figures in Measurement
i.
All certain values AND...
ii.
Increment/10→ round to this number UNLESS it has a
decimal, then do increment/5
c. Precision and Accuracy
i.
Precision and accuracy are not the same!
1. Precision: how repeatable is the experiment
2. Accuracy: is there a systematic error
ii.
We often average many measurements of the same value which should provide a more accurate
result IF there is only random error present (does not work if there is a systematic error)
5. 1.5→ Significant Figures and Calculations
a. Rules for Counting Significant Figures
i.
Nonzero integers: significant figures
ii.
Zeros
1. Leading: not significant figures
2. Captive: significant figures
3. Trailing
a. With decimal point: significant figures
b. Without decimal point: not significant figures
iii.
Exact Numbers
1. Exact numbers DO NOT count in calculations of significant figures (they have infinite
significant figures)
2. Counted numbers, exact conversions
a. Note: conversions include Fahrenheit→
Celsius, cm→ in
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iv.
6. 1.6→
a.
7. 1.7→
a.
8. 1.8→
a.
b.
Rules for significant figures in mathematical operations
1. Multiplication/ Division: number of sig figs in result is the same as the number in the
least precise measurement used in the calculation
2. Addition/Subtraction: result has the same number of decimal places as the least
precise measurement used in the calculation
v. Rules for rounding:
1. Round at the end!
2. Round up at 5, down at 4
a. Only round using the first number to the right of the last significant figure (eg.
3.48 to one significant figure is 3 not 4)
Learning to Solve Problems Systematically
Look at goal, starting point, and method to finding solution when trying to solve a problem
Dimensional Analysis
Multiply by different forms of one to get into the units you want to be in
Temperature
3 types used are Fahrenheit, Celsius, and Kelvin
Conversions:
i.
℃→ K
1.
ii.
℃+273.15=K
℉→ ℃
1.
℃ (9/5)+32=℉
9. 1.9→ Density
a. Formula
i.
Mass/Volume
b. Often used to identify a substance
10. 1.10→ Classification of Matter
a. Matter is anything with mass. It is very complex
b. Matter has 3 states:
i.
Solid: Fixed volume and shape
ii.
Liquid: Fixed volume, not fixed shape
iii.
Gas: Not fixed volume or shape
c. Mixtures
i.
A mixture has variable composition
ii.
Types of mixtures:
1. Homogenous (aka solution)→ wine, brass, air
2. Heterogenous→ sand in water, iced tea
iii.
Mixtures can be separated into pure substances by physical methods
d. Pure Substances
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i.
Constant composition
ii.
Compounds or elements
e. Methods of separating mixtures
i.
Distillation:
1. most volatile component is vaporized, goes through tube, re condenses
2. Good for simple mixtures
ii.
Filtration:
1. Pour mixture into mesh
2. Separates solid from liquid
iii.
Chromatography
1. Uses two phases of matter (mobile-- liquid/gas; stationary-- solid)
2. Separation happens because the components have different affinities for the two
phases so they move through the system at different rates
3. Paper Chromatography
a. Stationary Phase: Paper, Mobile Phase: Solvent (usually water)
b. Put drop of mixture on paper
c. Put solvent in beaker, but not enough to cover drop on paper
d. Put paper in, don’t let dot touch liquid
e. Liquid travels up the paper through capillary action, takes some dyes w it
f. Different spots that you end up with signify separate components
i.
Move differently because of attraction to solvent or paper
11. TERMS
a. Scientific Method: series of steps followed by scientist to answer questions about scientific nature
b. Measurement: quantitative observation
c. Qualitative Observation: Observation not pertaining to amounts or numbers
d. Quantitative Observation: an observation involving numbers
e. Hypothesis: educated prediction of the result to an experiment
f. Theory: a set of assumptions put forth to explain some aspect of the observed behavior of matter.
Chemistry theories usually involve assumptions about the behavior of individual atoms or molecules
g. Model: a theory
h. Natural Law: a statement that expresses generally observed behavior
i. Law of Conservation of Mass: mass is neither created nor destroyed
j. Mass: the quantity of matter in an object
k. Weight: the force exerted on an object by gravity
l. Significant Figures: the digits of a measurement which are not placeholders
m. Precision: the degree of agreement among several measurements of the same quantity, the reproducibility
of a measurement
n. Accuracy: the agreement of a particular value with the true value
o. Random error: measurement has equal probability of being high or low
p. Systematic error: error that occurs in the same direction each time
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q. Unit Factor Method: an equivalence statement between units used for converting from one unit to another;
dimensional analysis
r. Dimensional Analysis: Using ratios of units to convert from one unit to another
s. Matter: the material of the universe
t. States: the three different forms in which matter can exist (solid, liquid, and gas)
u. Mixture: 2 or more substances mixed together in a way that allows them to be separated without breaking
chemical bonds or a chemical reaction
v. Solution: a homogenous mixture
w. Homogenous Mixture: a mixture that has consistent properties throughout; evenly dispersed
x. Heterogenous Mixture: a mixture that does not have consistent properties throughout; not evenly
dispersed
y. Pure Substance: a substance with constant composition; a substance which cannot be broken apart by
physical means
z. Physical Change: a change in the form of a substance, but not in its chemical composition; chemical bonds
are not broken in a physical change
aa. Distillation: a method for separating the components of a liquid mixture that depends on differences in the
ease of vaporization of the compounds
bb. Filtration: a method for separating the components of a mixture containing a solid and a liquid
cc. Chromatography: the general name for a series of methods for separating mixtures by using a system with
a mobile phase (liquid or gas) and a stationary phase (solid)
dd. Volatile: the ease with which a substance can be changed to its vapor
ee. Compound: a substance with constant composition that can be broken down into elements by chemical
processes
ff. Chemical Change: the change of substances into other substances through a reorganization of the atoms; a
chemical reaction
gg. Element: a substance which cannot be decomposed into a simpler substance by chemical or physical
means
Chapter 2 Notes (Atoms, Molecules, and Ions)
2.1→ The Early History of Chemistry
a. Boyle
i.
First real chemist
ii.
1600s
iii.
Relationship between pressure and volume
2. 2.2→ Fundamental Chemical Laws
a. Lavoisier
i.
verified the Law of Conservation of Mass
ii.
Showed that combustion involved oxygen
iii.
Named Oxygen
iv.
Guillotined because of French Revolution
b. Proust
1.
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i.
Law of Definite Proportion (Proust’s law): A given compound always contains exactly the same
proportion of elements by mass
c. Dalton
i.
Thought that atoms were the particles that composed elements
ii.
Law of Multiple Proportions
3. 2.3→ Dalton’s Atomic Theory
a. Dalton’s Atomic Theory
i.
Each element is made up of tiny particles called atoms
ii.
The atoms of a given element are always identical; the atoms of different elements are different
in some fundamental way or ways
iii.
Chemical compounds are formed when atoms of different elements combine with each other. A
given compound always has the same relative numbers and types of atoms
iv.
Chemical reactions involve reorganization of atoms--changes in the way they are bound together.
The atoms themselves are not changed in a chemical reaction
b. Dalton made assumptions about the relative masses of different elements based on how they formed
compounds. His assumptions were not correct, but his idea of putting the elements into a table was a good
one.
c. Joseph Gay Lussac and Amedeo Avogadro
i.
Gay Lussac→ volumes of gases under same pressure,
found the ratios in which they reacted
ii.
Avogadro→ at the same temperature and pressure, equal
volumes of different gases contain the same number of
particles (aka Avogadro’s hypothesis)
1. Determines volume using amount of molecules present, not the size of the molecules
themselves
2. Not accepted by most chemists
4. 2.4→ Early Experiments to Characterize the Atom
a. The Electron
i.
JJ Thomson
1. Cathode Ray Experiment (around 1900)
a. Used cathode ray tubes to study electrical discharges
b. Ray forms when the tube gets energy, ray is made of negative particles called
electrons
2. Atoms must have some sort of positive particle inside of them to neutralize the
negativity of the electrons
3. Plum Pudding Model
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ii.
5. 2.5→
a.
b.
c.
Radioactivity
1. Certain elements produce high energy radiation
2. Types of radioactive emission: gamma rays, beta particles, and alpha particles
iii.
The Nuclear Atom
1. Rutherford
a. Directed particles at a thin sheet of metal foil
i.
Trying to prove plum pudding model
ii.
Atoms were deflected at large angles--knew that previous model
didn’t work (had to have positive center and negative outer part)
The Modern View of Atomic Structure: An Introduction
Atom consists of very small nucleus and electrons that move about the nucleus
i.
Chemistry of an atom mostly relies on electrons
ii.
Nucleus is very dense compared to the rest of the atom
Isotopes have almost identical chemical properties
Symbols
6. 2.6→ Molecules and Ions
a. The forces that hold atoms in compounds together are chemical bonds, the result is a molecule
i.
Atoms can share electrons
b. Ways to represent molecules
i.
most simply represented by using a chemical formula
ii.
structural formula
iii.
Ball and stick model
iv.
Space filling model
c. In a compound made of molecules, individual molecules move in independent ways
d. Types of Bonds
i.
Covalent: electrons are shared
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7.
2.7→
a.
b.
c.
d.
e.
f.
ii.
Ionic: attraction between oppositely charged ions
An introduction to the Periodic Table
Periodic Table is a chart that shows all of the known elements
Metals
i.
Efficiently conduct heat and electricity, can be hammered into thin sheets, ductility (can be put
into wires)
ii.
Usually lose electrons form positive ions
Nonmetals
i.
Usually gain electrons to form negative ions
Diatomic Elements
i.
H, N, O, F, Cl, Br, I
ii.
Bond with themselves to form “diatomic” gases
Groups (Columns)→ same number of valence electrons
i.
Alkali Metals: Group 1, very reactive
ii.
Alkaline Earth Metals: Group 2
iii.
Halogens: Group 17
iv.
Noble Gases: Group 18, do not react since they are stable
Periods (Rows)→ same number of electron “shells”
8. 2.8→ Naming Simple Compounds
a. Ionically Bonded Substances
i.
Cation is first, anion is second
ii.
Cation
1. Groups 1, 2, 13, Ag+, Cd2+, Zn2+ → elemental name
iii.
2. All other metals → elemental name (oxidation state in Roman
Numerals)
a. One with lower oxidation state is given the root ous, the one with the higher
charge is give the root ic
3. Polyatomic Ions→ given name
Anion
1. elemental→ root+ide
2. polyatomic→ given name
b. Ionic Compounds with Polyatomic Ions
i.
Oxyanions: anions of a given element and different numbers of oxygen atoms
1. One with less oxygen ends in ite, one with more oxygen ends in ate. If there are more
than 2, hypo goes to the one with the fewest oxygen atoms, per goes to the one with the
most
c. Covalently Bonded Substances
i.
(prefix--not mono) elemental name+ prefix+ ide
ii.
Prefixes
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1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
d. Acid
i.
ii.
1: mono
2: di
3: tri
4: tetra
5: penta
6: hexa
7: hepta
8: octa
9: nona
10: deca
A molecule that produces a solution containing free H+ ions; a molecule in which one or more H+
ions are attached to the anion
Naming
1. Anion doesn’t contain oxygen
a. Hydro+ elemental name+ ic acid
2. Anion contains oxygen
a. Ends in ate→ polyatomic ion+ic acid
b. Ends in ite→ polyatomic ion+ous acid
3. Notes:
a. Sulfur is root for sulfur
b. Phospor is root for phosphorus
9. Terms
a. Law of Definite Proportion: a given compound always contains exactly the same proportion of elements by
mass
b. Law of Multiple Proportions: when two elements form a series of compounds, the ratios of the masses of
the second element to combine with 1g of the first element can always be reduced to small whole numbers
c. Atomic Mass: the mass of an atom of a chemical element expressed in atomic mass units
d. Atomic Weight: the weighted average mass of the atoms in a naturally occurring element
e. Avogadro’s Law: equal volumes of gases at the same temperature and pressure contain the same number
of particles
f. Cathode Rays: the “rays” emanating from the negative electrode in a partially evacuated tube; a stream of
electrons
g. Electron: a negatively charged particle that moves around the nucleus of an atom
h. Radioactivity: the spontaneous decomposition of a nucleus to form a different nucleus
i. Nuclear Atom: an atom having a dense center of positive charge with electrons moving around the outside
j. Nucleus: the small, dense center of positive charge in an atom
k. Protons: a positively charged particle in an atomic nucleus
l. Neutrons: a particle in the atomic nucleus with mass virtually equal to the proton’s but with no charge
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m. Isotopes: atoms of the same element with different numbers of neutrons. They have identical atomic
numbers but different mass numbers
n. Atomic Number: the number of protons in the nucleus of an atom
o. Mass Number: the total number of protons and neutrons in the atomic nucleus of an atom
p. Chemical bonds: the force, or more accurately, the energy that holds two atoms together in a compound
q. Covalent Bonds: a type of bonding in which electrons are shared by atoms
r. Molecule: a bonded collection of two or more atoms of the same or different elements
s. Chemical Formula: the representation of a molecule in which the symbols for the elements are used to
indicate the types of atoms present and subscripts are used to show the relative numbers of atoms
t. Structural Formula: the representation of a molecule in which the relative positions of the atoms are shown
and the bonds are indicated by lines
u. Space filling model: a model of a molecule showing the relative sizes of the atoms and their relative
orientations
v. Ball and Stick Models: a molecular model that distorts the sizes of atoms but shows bond relationships
clearly
w. Ion: an atom or group of atoms that has a net positive or net negative charge
x. Cation: a positive ion
y. Anion: a negative ion
z. Ionic bonding: the electrostatic attraction between oppositely charged ions
aa. Ionic Solid: a solid containing cations and anions that dissolves in water to give a solution containing the
separated ions, which are mobile and thus free to conduct an electrical current
bb. Polyatomic ion: an ion containing a number of atoms
cc. Periodic Table: a chart showing all the elements arranged in columns with similar chemical properties
dd. Metals: an element that gives up electrons
ee. Nonmetals: an element not exhibiting metallic characteristics; an element which accepts electrons from a
metal
ff. Binary Compounds: a two element compound
gg. Oxyanions: an anion containing one or more oxygen atoms bonded to another element
hh. Acids: a substance that produces hydrogen ions in solution; a proton donor
Chapter 3 Notes (Stoichiometry)
3.1→ Counting by Weighing
a. Average mass
i.
Total mass of objects/number of objects
b. Counting by Weighing
i.
If you know how much an object weighs, you can measure out a certain number of that object
using its weight
2. 3.2→ Atomic Masses
a. Modern system of atomic masses is based on Carbon 12, which is exactly 12 amu
b. Mass spectrometer
i.
Most accurate method currently available for comparing the masses of elements
1.
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ii.
c.
3. 3.3→
a.
b.
c.
4. 3.4→
a.
b.
5. 3.5→
How it works:
1. Ionization: atoms/molecules are passed into a beam of high speed electrons, which
knock electrons off the atoms or molecules being analyzed and change them into
positive ions
2. Acceleration: Electric field accelerates these ions
3. Deflection: An accelerating ion produces its own magnetic field, so an interaction with
the applied magnetic field occurs which changes the path of the ion
a. The most massive ions are deflected the least
4. Detection: A metal plate detects the masses of the atoms/compounds, thus allowing
you to figure out what they are
iii.
Peaks indicate isotopes and their relative abundances
Weighted Averages
i.
Atomic masses are weighted averages of all of the isotopes of that element's (average atomic
mass)
ii.
Isotope A (%A)+ Isotope B (%B)+...=average atomic mass
The Mole
Avogadro’s Number
i.
6.022 x 1023
Not all moles have the same mass, but they are all composed of the same number of whatever the unit may
be
A mole is defined such that a sample of a natural element with a mass equal to the element’s atomic mass
expressed contains 1 mole of atoms
Molar Mass
The molar mass of a substance is the mass in grams of one mole of the compound
Ionic compounds are NOT molecules, instead they are known as formula units
Learning to Solve Problems
6. 3.6→ Percent Composition of Compounds
a. Describing composition of compound
i.
Numbers of constituent atoms
ii.
Percentages by mass
b. Mass Percent
i.
Aka weight percent
ii.
Mass of element in 1 mole of compound/molar mass of compound (100%)
7. 3.7→ Determining the Formula of a Compound
a. Empirical Formula Determination
i.
Since mass percentage gives number of grams of a particular element per 100 g of compound,
base the calculation on 100 g of compound. Each percent will then represent the mass in grams of
that element
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ii.
b.
c.
8. 3.8→
a.
b.
c.
9. 3.9→
a.
b.
c.
Determine the number of moles of each element present in 100 g of compound using the atomic
masses of the elements present
iii.
Divide each value of the number of moles by the smallest of the values. If each resulting number
is a whole number, these numbers represent the subscripts of the elements in the empirical
formula
iv.
If the numbers obtained in the previous step are not whole numbers, multiply each integer so
that the results are all whole numbers
Determining Molecular Formula from Empirical Formula
i.
Obtain the empirical formula
ii.
Compute the mass corresponding to the empirical formula
iii.
Molar mass/ empirical formula mass
iv.
The integer from the previous step represents the number of empirical formula units in one
molecule. When the empirical formula subscripts are multiplied by this integer, the molecular
formula results
Determining Molecular Formula from Mass Percent and Molar Mass
i.
Using the mass percentages and the molar mass, determine the mass of each element present in
one mole of compound
ii.
Determine the number of moles of each element present in 1 mole of compound
iii.
The integers from the previous step represent the subscripts in the molecular compound
Chemical Equations
Chemical Reactions
i.
The process in which old bonds are broken and new ones are formed, but no atoms are created or
destroyed
1. The number of one type of atoms must match on both sides of the equation
The Meaning of a Chemical Equation
i.
Put physical states next to each reactant
State
Subscript
Gas
(g)
Liquid
(l)
Solid
(s)
Aqueous
(aq)
After balancing an equation, the integers before each compound are known as coefficients
Balancing Chemical Equations
Unbalanced equations are not very helpful
The formulas of the compounds must never be changed in balancing a chemical equation
Writing and Balancing the Equation for a Chemical Reaction
i.
Determine what reaction is occurring
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ii.
Write the unbalanced chemical equation
iii.
Balance the equation by adding coefficients to the different compounds and seeing if they work
10. 3.10→ Stoichiometric Calculations: Amounts of Reactants and
Products
a. Calculating Masses of Reactants and Products in Chemical Reactions
i.
Balance the equation
ii.
Convert the known mass of the reactant or product to moles
iii.
Do a mole to mole ratio using the chemical reactant
iv.
Convert back from moles to grams if required
11. 3.11→ The Concept of Limiting Reactant
a. To determine how much product can be formed from a given mixture of reactants, we have to look for the
reactant that is limiting--the one that runs out first and thus limits the amount of product that can form
i.
Don’t assume a reaction is stoichiometric
ii.
Reactant that runs out first is the limiting reactant
b. Determination of Limiting Reactant Using Reactant Quantities
i.
Mol sub A required/mole sub B required
1. If actual mol A/mol B is more than this, A is limiting reactant
2. If actual mol A/mol B is less than this, B is limiting reactant
c. Determination of Limiting Reactant Using Quantities of Product Formed
i.
Determine which reactant would make less product→
that is the limiting reactant
d. Percent Yield
i.
Actual Yield/Theoretical Yield (100%)=percent yield
e. Solving a Stoichiometric Problem Involving Masses of Reactants and Products
i.
Write and balance equation
ii.
Convert the known masses to moles
iii.
Determine limiting reactant
iv.
Compute number of moles of desired product
v. Convert from moles to grams using molar mass
12. Terms
a. Chemical Stoichiometry: the calculation of the quantities of material consumed and produced in chemical
reactions
b. Mass Spectrometer: an instrument used to determine the relative masses of atoms by the deflection of
their ions on a magnetic field
c. Average Atomic Mass: the sum of the masses of its isotopes, each multiplied by its natural abundance
d. Mole: the number equal to the number of carbon atoms in exactly 12 g of Carbon 12
e. Avogadro’s Number: 6.022 x 1023
f. Molar Mass: the mass in grams of one mole of molecules or formula units of a substance
g. Mass Percent: the percent by mass of a component of a mixture or of a given element in a compound
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h. Empirical Formula: the simplest whole number ratio of atoms in a compound
i. Molecular Formula: the exact formula of a molecule, giving the types of atoms and the number of each type
j. Chemical Equation: a representation of a chemical reaction showing the relative numbers of reactant and
product molecules
k. Reactants: a starting substance in a chemical reaction
l. Products: a substance resulting from a chemical reaction
m. Stoichiometric Quantities: quantities of reactants mixed in exactly the correct amounts so that all are used
up at the same time, leaving no limiting or excess reactant
n. Limiting Reactant: the reactant that is completely consumed when a reaction is run to completion
o. Theoretical Yield: the maximum amount of a given product that can be formed when the limiting reactant is
completely consumed
p. Percent Yield: the actual yield of a product as a percentage of the theoretical yield
Chapter 4 Notes (Types of Chemical Reactions and Solution Stoichiometry)
1.
4.1→ Water, the Common Solvent
a. Water (H2O)
i.
Bent shape
ii.
Polar covalent bonds
1. Oxygen has partial negative charge, Hydrogens have partial positive charges
2. Water’s polarity is the reason that it is such a strong solvent
b. Polar molecules dissolve in water because of hydration
c. Not all ionic substances have the same solubility in water
i.
Depending of relative attractions of ions for each other and relative attractions of the ions for
water molecules
d. Water can dissolve many non ionic substances→ substances
which are polar can also be dissolved
2. 4.2→ The Nature of Aqueous Solutions: Strong and Weak
Electrolytes
a. Conductivity Properties
i.
High conductivity→ solution contains strong
electrolyte
ii.
Low conductivity→ solution contains weak electrolyte
iii.
No conductivity→ solution contains nonelectrolyte
b. Strong Electrolytes
i.
Substances that are completely ionized when they are dissolved in water
ii.
Soluble Salts
1. Become hydrated when salt dissolves
iii.
Strong Acids
1. An acid is substance that produces H+ ions (protons) when it is dissolved in water
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c.
d.
3. 4.3→
a.
b.
2. A strong acid is an acid for which virtually every molecule ionizes
a. Completely dissociates into its ions
iv.
Strong Bases
1. Soluble ionic compounds containing the hydroxide ion (OH-)
2. When dissolved in water, cation and OH- ion separate completely
Weak Electrolytes
i.
Small degree of ionization in water
1. Produce relatively few ions when dissolved in water
ii.
Weak acids
1. Only ionizes to a slight extent in aqueous solution
iii.
Weak base
1. Some OH- ions produced but not many; most molecules stay in tact
Nonelectrolytes
i.
Substances that dissolve in water but don’t produce any ions
ii.
Molecules are spread out within water, but they stay in tact
iii.
Eg. ethanol, sugar
The Composition of Solutions
Chemical reactions often take place when two solutions are mixed. Need to know:
i.
Nature of reaction
ii.
Amount of chemicals
Molarity
i.
Explains concentration
ii.
iii.
Can use to find concentration of ions in solution
c. How to make a solution
i.
From a solid
1. Put weighboat on balance, zero balance
2. Add solid to weigh boat, record mass
3. Put solid in x mL volumetric flask
4. Fill up volumetric flask almost to line
5. Cover volumetric flask with parafilm, swirl and invert flask to dissolve solid
6. Fill up volumetric flask to line
7. Repeat step 5
ii.
From a solution
1. Measure y mL of stock solution in a graduated cylinder, record volume
2. Put measured stock solution into x mL volumetric flask
3. Fill up volumetric flask almost to line
4. Cover volumetric flask with parafilm, swirl and invert flask to dissolve solid
5. Fill up volumetric flask to line
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6. Repeat step 4
4.
5.
6.
7.
8.
d. Dilution
i.
Start with stock solution (solution with higher concentration than you want), add water to dilute
the solution and get the M you want
ii.
M1V1=M2V2
1. M1V1= moles solute before dilution
2. M2V2=moles solute after solution
4.4→ Types of Chemical Reactions
a. Precipitation Reactions
b. Acid Base Reactions
c. Redox Reactions
4.5→ Precipitation Reactions
a. A reaction between to solutions which produces an insoluble solid
b. AB+CD→ AD+CB
c. Solubility Rules
i.
NO3- is soluble
ii.
Group 1 cations and NH4+ are soluble
iii.
Halides are soluble EXCEPT WITH Ag+, Pb2+, Hg22+
iv.
SO42- is soluble EXCEPT with Ba2+, Pb2+, Hg22+, and Ca2+
v. OH- is insoluble
vi.
S2- is insoluble
vii.
CO32- is insoluble
viii.
PO43- is insoluble
ix. CrO42- is insoluble
4.6→ Describing Reactions in Solution
a. Formula equation gives overall reaction but not necessarily the actual forms of the reactants and products
in solution
b. Complete ionic equation represents as ions all reactants and products that are strong electrolytes
c. Net ionic equation includes only solution components undergoing a change
i.
Leaves out spectator ions
4.7→ Stoichiometry of Precipitation Reactions
a. ID the species present in the combined solution, and determine what reaction occurs
b. Write the balanced net ionic equation for the reaction
c. Calculate the moles of reactants
d. Determine which reactant is limiting
e. Calculate the moles of product (s)
f. Convert back to grams if required
4.8→ Acid Base Reactions
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a. In acid base reactions, one reactant includes an acid, the other a base. These components come together
during the reaction to form (usually) H2O
b. Performing Calculations for Acid Base Reactions
i.
List the species present in the combined solution before any reaction occurs, and decide what
reaction will occur
ii.
Write balanced net ionic equation
iii.
Calculate moles of reactants
iv.
Determine limiting reactant
v. Calculate moles of required reactant or product
vi.
Convert to units needed
c. Acid Base Titrations
i.
Add titrant (measured volume of solution of known concentration) in buret to analyte (substance
being analyzed)
ii.
Reach a point (equivalence point) when exactly enough
titrant has been added to react with analyte→ marked
by change in color of indicator
9. 4.9→ Redox Reactions
a. A reaction that involves the transfer of electrons
b. Oxidation States
i.
Way to keep track of electrons in redox reactions
ii.
In covalent bonds, arbitrarily assigned
iii.
Sum of charges of individual atoms in a compound must equal the total charge of that compound
iv.
Rules:
1.
v.
Oxidation state of:
Charge:
Example(s)
Atom in element
0
Na(s), O2 (g)
Monatomic ion
Charge of ion
Na+
Fluorine
1-
HF
Oxygen (EXCEPT WHEN
PEROXIDE: O22-)
2-
H2O
Hydrogen with a nonmetal 1+
HCl
Group 1 cations
NaCl
1+
2. Can have non integer oxidation states because of how oxidation states are arbitrarily
assigned
a. Average of all atoms of that element within the compound
Characteristics of Redox Reactions
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1. Oxidation: increase in oxidation state
2. Reduction: decrease in oxidation state
3. Oxidizing Agent: reactant that includes element that gets reduced
4. Reducing Agent: reactant that includes element that gets oxidized
10. 4.10→ Balancing Redox Reactions
a. Balancing Redox Reactions by Oxidation States
i.
Write unbalanced equation
ii.
Determine oxidation states of all products and reactants
iii.
Separate into half reactions
iv.
Balance all elements other than H and O
v. Balance H and O using H+ ions and H2O
vi.
Balance electrons within half reactions
vii.
Add half reactions together
viii.
Cancel similarities
ix. IF IN BASIC SOLUTION: get rid of H+ ions using OH11. Terms
a. Polar Molecule: a molecule that has a permanent dipole moment
b. Hydration: the interaction between solute particles and water molecules
c. Solubility: the amount of a substance that dissolves in a given volume of solute at a given temperature
d. Solute: a substance dissolved in a liquid to form a solution
e. Solvent: the dissolving medium in a solution
f. Electrical Conductivity: the ability to conduct an electric current
g. Strong Electrolyte: a material that, when dissolved in water, gives a solution that conducts an electric
current very efficiently (dissociates completely)
h. Weak Electrolyte: a material that, when dissolved in water, gives a solution that conducts only a small
electric current (only dissociates a little)
i. Nonelectrolyte: a substance that, when dissolved in water, given a non conducting solution (does not
dissociate at all)
j. Strong Acid: an acid that completely dissociates to produce a H+ ion and the conjugate base
k. Strong Base: a metal hydroxide salt that completely dissociates into its ions in water
l. Weak Acid:an acid that dissociates only slightly in aqueous solution
m. Weak Base: a base that reacts with water to produce hydroxide ions to only a slight extent in aqueous
solution
n. Molarity: moles of solute per liter of solution; a measure of concentration
o. Dilution: the process of adding solvent to lower the concentration of a solution
p. Precipitation reaction: a reaction in which an insoluble substance forms and separates from the solution
q. Precipitate: insoluble solid formed in a precipitation reaction
r. Formula Equation: an equation representing a reaction in solution showing the reactants and products in
undissociated form, whether they are strong or weak electrolytes
s. Complete Ionic Equation: the equation that shows all substances that are strong electrolytes as ions
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t.
u.
v.
w.
x.
y.
z.
aa.
bb.
cc.
dd.
ee.
ff.
gg.
hh.
ii.
Net Ionic Equation: an equation for a reaction in solution, where strong electrolytes are written as ions,
showing only those components that are directly involved in the chemical change
Spectator Ion: ions present in solution that do not participate directly in a reaction
Acid: a substance that produces hydrogen ions in solution; a proton donor
Base: a substance that produces hydroxide ions in aqueous solution; a proton acceptor
Neutralization reaction: an acid base reaction
Volumetric Analysis: a process involving titration of one solution with another
Titration: a technique in which one solution is used to analyze another
Equivalence Point/Stoichiometric Point: the point in a titration when enough titrant has been added to
react exactly with the substance in solution being analyzed
Indicator: a chemical that changes color and is used to mark the end point of a titration
Endpoint: the point in a titration at which the indicator changes color
Redox reaction: a reaction in which one or more electrons are transferred
Oxidation states: a concept that provides a way to keep track of electrons in redox reactions according to
certain rules
Oxidation: an increase in oxidation state (loss of electrons)
Reduction: a decrease in oxidation state (gain of electrons)
Oxidizing Agent: the reactant which contains the substance that gets reduced
Reducing Agent: the reactant which contains the substance that gets oxidized
Chapter 5 Notes (Gases)
1.
5.1→ Pressure
a. Barometer→ a device that measures atmospheric (barometric)
pressure in mmHg
i.
made by filling tube all the way with Hg, then flipping it over, thus creating a vacuum at the top of
the tube
b. Units of Pressure
i.
1 atm=760 mmHg=760 torr=101,325 Pa=14.7 psi
ii.
Pressure= force/area
2. 5.2→ The Gas Laws of Boyle, Charles, and Avogadro
a. Boyle’s law
i.
PV=k
1. Hyperbolic
ii.
b. Charles’ Law
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i.
V=bT
ii.
1. Linear
0 K is absolute zero, the place at which there is 0 volume
i.
b. Avogadro’s Law
i.
V=an
ii.
c. Amonton’s Law
i.
P=Tk
ii.
3. 5.3→ The Ideal Gas Law
a.
i.
4. 5.4→ Gas Stoichiometry
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a. Molar volume
i.
22.4 L/mol at STP
b. Molar mass of a gas
i.
5. 5.5→ Dalton’s Law of Partial Pressures
a. Dalton’s Law of Partial Pressure
i.
if there is a mixture of gasses, the total pressure exerted is the sum pressure each gas would
exert if it was alone
ii.
1.
iii.
gas collection over water
1. use Dalton’s law to account for the vapor pressure, which is directly correlated with T
iv.
Since there are no IMF, the type of atom doesn’t matter
b. Mole Fraction: ratio of the number of moles of a given component in a mixture to the total number of moles
in the mixture
i.
1. X is called a chi
6. 5.6→ The Kinetic Molecular Theory of Gases
a. Postulates of KMT
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i.
b.
c.
d.
e.
f.
The particles are so small compared with the distances between them that the volume of the
individual particles can be assumed to be negligible (0)
ii.
The particles are in constant motion. The collisions of the particles with the walls of the container
are the cause of the pressure exerted by the gas.
iii.
The particles are assumed to exert no forces on each other (no intermolecular forces)
iv.
The average kinetic energy of the gas is directly proportional to the temperature of the gas (in K)
v. All collisions are perfectly elastic (no energy is lost)
Ideal Gas Law
i.
PV=nrT
ii.
r is the ideal gas constant
1. .0821 (L)(atm)/(mol)(K)
Relationships between parameters
i.
P and V (Boyle’s Law)
1. Inverse relationship
2. gas particles hit the wall of a smaller container more often, causing an increase in
pressure
ii.
P and T
1. Direct relationship
2. When temperature increases, the speeds of the particles of the gas increase, so the
particles hit the walls of the container with greater force and frequency, causing an
increase in pressure
iii.
V and T (Charles’s Law)
1. Direct relationship
2. As temperature increases, the speeds of the particles of the gas increase, so in order to
keep the pressure constant the volume must increase
iv.
V and n (Avogadro’s Law)
1. Direct
2. An increase in moles means that particles collide with the walls of the container more
frequently (because there are more particles), unless the volumes also increases to
keep the pressure constant
Mixture of Gases (Dalton’s Law)
i.
Dalton’s law makes sense because KMT assumes all gas particles are independent, and it is
unimportant which gas the particles are
The Meaning of Temperature
i.
KE=3/2 RT
ii.
Index of the random motions of the particles of a gas
Root Mean Square Velocity
i.
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1.
ii.
Square root of the average velocity squared
1.
M is molar mass in kJ/mol
a. As molar mass increases (all else constant), velocity decreases
2. R is 8.314 J/mol (K)
3. T is temperature in Kelvin
4. U rms is given in m/s
iii.
Mean free path
1. The path a particle travels between collisions
a. Different particles have different energies based on collisions
7. 5.7→ Effusion and Diffusion
a. Diffusion: mixing of gases
i.
More complicated than diffusion because gas particles hit one another, causing them to move
differently than they would in a vacuum
b. Effusion: passage of a gas through a tiny orifice into an evacuated chamber
i.
Graham’s law of Effusion
1.
a. Predicted by kinetic molecular theory
8. 5.8→ Real Gases
a. A real gas exhibits most ideal behavior at low pressures and high temperatures
b. Van der waals equation
i.
ii.
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9. 5.9→ Characteristics of Several Real Gases
a. Lower a means lower IMF (since a corrects for IMF)
10. 5.10→ Chemistry in the Atmosphere
a. Atmosphere is mostly N2 and O2
b. Upper atmosphere depends a lot on radiation from other things in space
i.
Protects us from this radiation→ breaking down
because of pollution (mostly CO2)
c. Ozone created, process called photochemical smog
d. Burning coal produces acid rain
11. Terms
a. barometer: a device that measures atmospheric (barometric) pressure in mmHg
b. manometer: a device that measures the pressure of a specific gas in mmHg
c. pressure: the force exerted by gaseous particles on the walls of the container
d. Absolute zero: 0 K
e. temperature: a measurement of the average kinetic energy of molecules in a system
f. volume: the amount of space a substance occupies
g. pressure: the force exerted by gaseous particles on the walls of the container
h. inverse relationship: a relationship between two numbers in which an increase in the value of one number
results in a decrease in the value of the other number.
i. direct relationship: a relationship between two numbers where an increase or decrease in one variable
causes the same change to occur in the second variable
j. elastic collisions: encounters between two bodies in which the total kinetic energy of the two bodies after
the encounter is equal to their total kinetic energy before the encounter (no energy is lost in the collision
as heat, noise, etc)
k. Mole Fraction: ratio of the number of moles of a given component in a mixture to the total number of moles
in the mixture
l. Effusion:the movement of a gas through a pore or capillary into a vacuum
m. Diffusion: the mixing of gases
Chapter 6 (Thermochemistry)
1.
6.1→ The Nature of Energy
a. Energy is the capacity to do work or to produce heat
b. Law of conservation of energy
i.
Energy cannot be created or destroyed, only transformed
c. Types of Energy
i.
Potential: energy due to position or composition
ii.
Kinetic: energy due to the motion of the object
d. Heat and work
1.
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i.
Heat: the transfer of energy between two objects due to a temperature difference
1. q=mCΔT, where C is the specific heat capacity of the substance
ii.
Work: force acting over a distance
iii.
Heat and work are path functions, but energy is a state function
e. Chemical energy
i.
System v surroundings
1. System: part of the universe on which we wish to focus attention
2. Surroundings: everything else in the universe
ii.
Endothermic v exothermic
1. Exothermic
a. some of the potential energy stored in the chemical bonds is being converted
to thermal energy via heat
f. First law of thermodynamics
i.
The energy of the universe is constant
ii.
Energy can neither be created nor destroyed, only transferred from one object to another (law of
conservation of energy)
g. Types of systems
i.
Open: can exchange matter and energy with surroundings
ii.
Closed: can only exchange energy with surroundings (no transfer of matter)
iii.
Isolated: cannot exchange matter or energy with surroundings
h. Internal energy
i.
the sum of the kinetic and potential energies of each particle in the system
ii.
State function
i.
iii.
Work (with a piston)
i.
1.
Because work= force (distance), pressure=force/area
a. work=pressure (area) (Δh)
b. w=-pΔv
2. 6.2→ Enthalpy and Calorimetry
a. Enthalpy
i.
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ii.
iii.
State function
Under constant pressure
iv.
1.
Change in enthalpy
1.
v.
a. ΔH is positive→ exothermic
i.
Heat on product side
b. ΔH is negative→ endothermic
i.
Head on reactant side
Enthalpy and Energy
1. ΔH is approximately ΔE when a small amount of work is done
(usually: when the moles of gas stay roughly constant)
b. Calorimetry
i.
The science of measuring heat
ii.
Heat capacity
1.
2. Specific v molar
a. specific→ given per gram→ J/(℃ x g)
iii.
iv.
v.
b. molar→ given per mole→ J/(℃ x mol)
Calorimeter should be adiabatic
1. Eg. coffee cup
2. 0=qrxn+qcalorimeter
a. Remember that q=(mass)(C)(ΔT)
Constant pressure calorimetry (coffee cup calorimetry)
1. Pressure doesn’t change
2. Can measure temperature change in order to measure change in enthalpy
3. qp=ΔH=ΔE
Constant volume calorimetry (bomb calorimetry)
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1.
2. Inside steel, usually used for combustion reactions
3. Start w large amount of O2 gas
4. Make sure to always set up the same way, because Ccal is a combination of mass and C
a. qv=CcalΔT
vi.
Extensive property
3. 6.3→ Hess’s Law
a. Since enthalpy is a state function, the change of enthalpy for a reaction is not dependent upon the order in
which the steps of the reaction occur
i.
b. Characteristics of Enthalpy Changes
i.
If a reaction is reversed, the sign of ΔH is reversed
ii.
If a reaction is multiplied by a factor, ΔH is also multiplied by that
factor
iii.
If two reactions are added, add ΔHs
4. 6.4→ Standard Enthalpies of Formation
a. ΔH°f means enthalpy of formation under standard conditions
i.
the enthalpy change that accompanies the formation of one mole of product at 1 atm and 25°C
from its elements in their standard states
b. ° symbolizes standard conditions
i.
Gaseous state: 1 atm
ii.
Solution: 1.00 M
iii.
For a condensed state, the standard state is pure liquid/solid
iv.
USUALLY 25°C (can assume unless told otherwise)
v. AN ELEMENT’S STANDARD STATE: the element’s state at 1 atm and 25°C
1. We define an element’s standard state as having a ΔH of 0. All
other ΔH are based on this standard.
c. Enthalpy Calculations
i.
Steps
1. Break reactants into elemental form (standard state)
a. Σ -n reactant ΔH°f, reactant
2. Form products using standard states of elemental form
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a. Σ n product ΔH°f, product
ii.
ΔHrxn=Σ n product ΔH°f, product-Σ n reactant ΔH°f, reactant
1.
5. 6.5→ Present Sources of Energy
6. 6.6→
7. Terms
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
m.
n.
o.
p.
q.
New Energy Sources
System: area of interest
Surroundings: everything not in system
Energy: the capacity to do work or to produce heat
Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed
Potential energy: energy due to position or composition
Kinetic energy: energy due to the motion of the object
Heat: the transfer of energy between two objects due to a temperature difference
Work: force acting over a distance
State function: a property of a system that depends only on the system’s present state
Path function: a property of a system that depends on the path taken to get to the current state
System: part of the universe on which we wish to focus attention
Surroundings: everything else in the universe
First law of thermodynamics: The energy of the universe is constant
Internal energy: the sum of the kinetic and potential energies of each particle in the system
Enthalpy: the sum of the internal energy added to the product of the pressure and volume of the system
Calorimetry: the science of measuring heat flow
Calorimeter: an apparatus for measuring the amount of heat involved in a chemical reaction or other
process
r. Heat Capacity: the amount of heat (usually expressed in calories, kilocalories, or joules) needed to raise
the system's temperature by one degree
s. Extensive Property: dependent on amount
t. Intensive Property: not dependent on amount
u. Adiabatic: no heat transfer w surroundings
v. Hess’s Law: in going from a particular set of reactants to a particular set of products, the enthalpy change is
the same whether the reaction takes place in one step or in a series of steps
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w. Standard Enthalpy of Formation: the enthalpy change that accompanies the formation of one mole of
product at 1 atm and 25°C from its elements in their standard states
Chapter 7 Notes (Atomic Structure and Periodicity)
1.
7.1→ Electromagnetic Radiation
a. parameters
parameter
symbol
unit
wavelength
λ
m
frequency
ν
1/s, Hz
energy
E
J;//
velocity
c
m/s
b. equations
i.
λν=c
1. inverse relationship
ii.
E/ν=h
1. direct relationship
iii.
Eλ=hc
1. inverse relationship
2. 7.2→ The Nature of Matter
a. Equations
i.
ΔE=nhν
1. n is an integer
2. h is Planck’s constant
3. ν is frequency
ii.
Ephoton=hν
iii.
Ephoton=-2.178 x 10-18 J (1/nf2- 1/ni2)
1. Bohr equation
2. ONLY FOR HYDROGEN
b. Energy can only be gained or lost in “packets” of discrete size known as quantums
c. Einstein suggested that electromagnetic radiation can be viewed as a stream of “particles” called photons
d. The Photoelectric Effect
i.
Phenomenon in which electrons are emitted from the surface of a metal when light strikes it
ii.
observations:
1. no electrons emitted by a given metal below a specific threshold frequency, v0
2. For light with frequency lower than the threshold, no electrons are emitted regardless
of the intensity of the light
3. For light with frequency greater than the threshold frequency, the number of electrons
emitted increases with the frequency of the light
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iii.
iv.
4. For light with frequency greater than the threshold frequency, the KE of the emitted
electrons increases linearly with the frequency of the light
Photons behave as if they have mass under certain circumstances, but it has no rest mass
1. EM radiation can act either as a wave or a particle; it is therefore said to have wave
particle duality
2. Since light can act as a particle, people started to ask whether particles could act as
waves. De Broglie was one of these people. Scientists now believe all matter exhibits
both particulate and wave properties. His equation is below.
a. The larger the piece of matter, the more it acts like a particle
Equations
1. Minimum energy to remove an electron
a. E0=hv0
2. KEelectronwhen v>v0
a. 𝐾𝐸(𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛) = 1/2𝑚𝑣 2 = ℎ𝑣 − 𝑣 0
3. Theory of relativity
a. E=mc2
4. De Broglie’s equation
a. λ=h/mν
e. Diffraction
i.
Scattered radiation in metals/crystals can cause diffraction
1. Interfere constructively→ makes light look
brighter
2. Interfere destructively→ makes like look less
bright/like its not there at all
3. 7.3→ The Atomic Spectrum of Hydrogen
a. Line spectrum v. Continuous Spectrum
i.
hot solids, liquids, and very dense gases emit continuous spectra (white light)
ii.
energized atoms emit non-continuous spectra (discrete lines with specific wavelengths)
b. Hydrogen emits a line spectrum, which matches up with what Planck’s ideas said it should emit
4. 7.4→ The Bohr Model
a. Niels Bohr
b. electrons reside in “shells” based on how much energy they have--can only lose and gain specific amounts
of energy
c. uses quantum
d. experimentation showed certain chemicals only emit certain wavelengths of light, thus proving that they
can only emit certain amounts of energy in a specific amount of time
e. works well with Hydrogen, encounters issues with larger atoms
f. Came up with an equation which define which energy “levels” were available to an electron in the
Hydrogen atom
i.
E=-2.178 x 10-18 J (z2/n2)
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ii.
1. z=nuclear charge, n is an integer
2. An electron with more negative energy is more tightly bound to the nucleus
To find the energy consumed or emitted when an electron jumps from one energy level to
another, subtract the final energy from the original one
1. To find the wavelength of the emitted photon
a. ΔE=hc/λ
5. 7.5→
a.
b.
c.
d.
The Quantum Mechanical Model
Schrodinger, Heisenberg, and de Broglie
Thought about electrons like standing waves
Heisenberg uncertainty principle
The Physical Meaning of a Wave Function
i.
The square of the function indicates the probability of finding an electron near a particular point
in space
ii.
It is more probable to find an electron near the nucleus, but the orbitals increase in size as they
get further from the nucleus. This causes one maximum that is close to the nucleus but not the
closest to it
6. 7.6→ Quantum numbers
a. principal quantum number
i.
n
ii.
size of orbital/proximity of nucleus (energy level)
1. As n increases, the electrons energy increases
2. As n increases, the electron becomes less negative (less tightly bound to the nucleus)
iii.
n=1, 2, 3, 4, …
b. angular quantum number/subshell
i.
l
ii.
type of orbital
iii.
l=n-1, …, 0
iv.
c. magnetic quantum number
i.
ml
ii.
Orientation in space
iii.
ml=-l, ..., 0, ..., l
iv.
ml possibilities for s, p, and d orbitals
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v.
d.
7. 7.7→
a.
b.
c.
8. 7.8→
a.
9. 7.9→
a.
b.
1.
note that any orbital can be named with any ml value that makes sense with its angular
quantum number, as long as no other orbital has already been named with that value
(there are no assigned ml values for specific positions)
electron spin quantum number
i.
ms
ii.
describes one of the two possible electrons in an orbital
1. They must be spinning in opposite directions
iii.
ms=± ½
Orbital Shapes and Energies
Orbitals have specific areas of high probability, and areas of 0 probability (nodes)
As n increases, the number of nodes increases
i.
Number of nodes for an orbital is n-1
All orbitals with the same value of n have the same energy (ONLY IN HYDROGEN)
Electron Spin and Pauli Exclusion Principle
An orbital can hold only two electrons, and they must have opposite spins
Polyelectronic Atoms
When considering polyelectronic atoms instead of Hydrogen, must also consider:
i.
Kinetic energy of electrons as they move around the nucleus
ii.
Potential energy of attraction between nucleus and electrons
iii.
Potential energy of repulsion between electrons
Electron correlation problem
i.
Don't know exactly where electrons are→ cannot
calculate repulsion
c. Outer electrons are “shielded” from positive nuclear charge by other electrons
d. Ens<Enp<End<Enf
i.
S electrons have one peak closer to nucleus than p, higher peak further
1. S electrons penetrates to the nucleus more than one in a 2p orbital (penetration effect)
10. 7.10→ The History of the Periodic Table
a. Originally constructed to represent the patterns observed in the chemical properties of elements
b. Dobereiner, Newlands, Meyer, Mendeleev
i.
Mendeleev used his table to correctly predict qualities about elements which were undiscovered
at the time
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11. 7.11→ The Aufbau Principle and the Periodic Table
a. As protons are added one by one to the nucleus to build up elements, electrons are similarly added to these
Hydrogen like orbitals (Aufbau principle)
b. The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons
allowed by the Pauli exclusion principle (Hund’s rule)
c. electron configuration can be abbreviated. Instead of writing this out for all shells, we can put the noble
gas closest (but below) to the atom, then write out the electron configuration for the remaining electrons
d. The elements in the same group have the same valence electron configuration
e. Many transition metals have electron configurations that are different than their expected electron
configuration, because it is sometimes more energy efficient to have an unpaired electron in a higher
energy level than a paired one in a lower energy level
i.
Cr: [Ar] 3d54s1
ii.
Cu:[Ar] 3d104s1
f. Orbitals and the Periodic Table
1. 4 blocks (s, p, d, f)
2. think of each elemental block as one more electron
a. explains the order in which electrons “fill up” orbitals in an atom (filling sequence)
3. as electrons are further from the nucleus, they have more energy
a. as n increases, so does energy
b. s<p<d<f
12. 7.12→ Periodic Trends in Atomic Properties
a. Ionization Energy
i.
The energy required to remove an electron from a gaseous atom or ion
ii.
1.
Ionization energy increases (first<second, etc)
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1.
Positive nuclear charge remains the same,
negative charge lowers→ nucleus holds on to
iii.
iv.
remaining electrons more tightly
2. Changes based on whether the electron is core or valence (more energy to remove
core)
First ionization energy increases as you go right on the periodic table, and goes down as you go
down on the periodic table
1. Discrepancies caused by energy lost in pairing electrons
2.
Electron Affinity
1. The energy change associated with the addition of an electron to a gaseous atom
a.
2. The more negative the energy, the greater the quantity of energy released
3. Many elements do not form stable isolated 1- ions
4. Only a first electron affinity (because that electron causes the atom to be negative, and
thus to repel other electrons)
5.
v.
a. Periodic trends
i.
As you go right, electron affinities increase
ii.
As you go down, electron affinities decrease
Atomic Radius
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1. Cannot be specified exactly
2. Obtained by measuring distance between atoms in chemical compounds
3.
a. Goes down as you go right
b. Goes up as you go down
13. 7.13→ The Properties of a Group: The Alkali Metals
a. Information contained in the periodic table
i.
Number and type of valence electrons in a group are the same
ii.
The most chemically reactive metals are the ones in the lower left hand corner of the table, where
ionization energies are smallest
iii.
The most chemically reactive nonmetals are the ones with the largest electron affinities
b. The Alkali Metals
i.
Experience many of the same qualities for the reasons listed above
ii.
Most form 1+ ions
14. Terms
a. electromagnetic radiation: light (can think of it as wave and as particle)
b. wavelength: distance between two consecutive crests or troughs
c. frequency: the number of wavelengths that pass a specific point in one second
d. energy of light: energy emitted per photon
e. velocity: distance travelled over unit time
f. photon: discrete “packet” of light; the manner in which we look at light as particles
g. Hertz: the unit equivalent to 1/s
h. c: speed of light (2.9979 x 108 m/s)
i. h: planck’s constant (6.626 x 10-34 Js)
j. spectroscope: a tool which separates light into its component wavelengths
k. excited state: a state of a physical system (eg. electron) that is higher in energy than the ground state
l. ground state: the lowest energy state of an atom or other particle, the starting point of an electron
m. endothermic: absorbing energy
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n.
o.
p.
q.
r.
s.
exothermic: releasing energy in the form of heat/light
Quantization: the concept that energy can occur only in discrete units called quanta
Photon: a quantum of electromagnetic radiation
Photoelectric effect: the production of electrons or other free carriers when light is shone onto a material
Dual nature of light: the statement that light exhibits both wave and particulate properties
Diffraction: the scattering of light from a regular array of points or lines, producing constructive and
destructive interference
t. Diffraction Pattern: the distinctive pattern of light and dark fringes, rings, etc, formed by diffraction
u. Continuous Spectrum: a spectrum that exhibits all wavelengths of visible light
v. Line Spectrum: a spectrum showing only certain discrete wavelengths
w. Standing Wave: a stationary wave as on a string of a musical instrument; in the wave mechanical model,
the electron in the hydrogen atom is considered to be a standing wave
x. Wave Function: a function of the coordinates of an electron’s position in 3D space that describes the
properties of the electron
y. Wave mechanical model: a model for the hydrogen atom in which the electron is assumed to behave as a
standing wave
z. orbital: probability distribution describing a volume of space around the nucleus where up to two electrons
are likely to be found (90% of the time)
aa. Heisenberg’s uncertainty principle: we cannot know momentum (mass x velocity) and position of an
electron at the same time
bb. quantum numbers: used to describe the most probable position of an electron within an atom
cc. Paulib Exclusion Principle: no two electrons within an atom can be identified by the same set of 4 quantum
numbers
dd. Hund’s rule: the lowest energy configuration for an atom is the one having the maximum number of
unpaired electrons allowed by the Pauli exclusion principle
ee. Aufbau Principle: as protons are added to the nucleus of an atom, electrons fill orbitals from the lowest to
the highest energy
ff. Probability Distribution: the square of the wave function indicating the probability of finding an electron at
a particular point in space
gg. Radial Probability Distribution: probability that an electron in the orbital with quantum numbers n and l will
be found at a distance r from the nucleus
hh. Nodal surfaces/nodes: an area of an orbital having 0 electron probability
ii. Degenerate: a group of orbitals with the same energy
jj. Electron Spin: a quantum number representing one of the two possible values for the electron spin
kk. Polyelectronic Atoms: an atom with more than one electron
ll. valence electrons: electrons in the outermost principal quantum level of an atom
mm. Core electrons: an inner electron in an atom; one not of the outermost principal quantum number level
nn. Lanthanides: a group of 14 elements following lanthanum in the periodic table, in which the 4f orbitals are
being filled
oo. Actinides: a group of 14 elements following actinium in the periodic table, in which the 5f orbitals are being
filled
pp. Representative elements: those elements within the A groups on the periodic table
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qq. First Ionization Energy: energy required to remove the most loosely held electron from one mole of
gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.
rr. Second Ionization Energy: energy needed to remove a second electron from each ion in 1 mole of gaseous 1+
ions to give gaseous 2+ ions
ss. Atomic Radii: half the distance between the nuclei in a molecule consisting of identical atoms
tt. Metalloids: elements along the division line of metals and nonmetals (on periodic table). They exhibit
metallic and nonmetallic properties
Chapter 8 Notes (Bonding: General Concepts)
1.
8.1→
a.
b.
c.
Types of Chemical Bonds
Can study compounds in many different ways
Systems act in the way in which they achieve the lowest possible energy
Ionic Bonding
i.
Energy of interaction between ions is calculated using Coulomb’s Law
ii.
If you get a negative number, that means that the ion pair has lower energy than the separated
ions
d. Bonding between identical atoms
i.
Opposite forces
1. Unfavorable forces : electron-electron repulsion, proton-proton repulsion
2. Favorable force: electron-proton attraction
ii.
Get the right distance apart (lowest energy) → bond
length
1.
Too far→ not enough attraction, electronelectron repulsion stronger than attraction
2. Too close→ proton-proton repulsion stronger than
attraction
e. Intermediate bonding (polar bonds)
i.
Electrons not shared equally, atoms have fractional charges
2. 8.2→ Electronegativity
a. Expected bond energy
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i.
ii.
b.
i.
Expected bond energy of a bond is the average of the diatomic bonds of each element involved in
the bond
The greater the difference between the actual bond energy and the expected bond energy, the
more polar the bond is
F has the greatest electronegativity
c.
3. 8.3→ Bond Polarity and Dipole Moments
a. Dipole
i.
A molecule with sigma positive and sigma negative parts (molecule with partially charged parts)
1. Only if polar bonds are not equal and opposite
ii.
Can be shown with colors→ red is electron rich, blue
is electron poor
4. 8.4→ Ions: Electron Configurations and Sizes
a. Electron Configuration of Compounds
i.
Two nonmetals→ covalent bond
ii.
Non metal and metal→ ionic bond
b. Predicting Formulas of Ionic Compounds
i.
Ionic solids are set up in a way that minimizes repulsions (+ + and - -) and maximizes attractions
(+ -)
1. In gases, the ions are much further apart and normally don’t interact as much
ii.
Electrons are transferred in a manner that gives all atoms involved an octet (or duet for H or He)
c. Size of Ions
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i.
Hard to exactly define ionic radii (similar to atomic
radii) → too hard to measure
ii.
Cations are smaller than neutral atoms (of same element), anions are bigger than neutral atoms
and cations (of the same element)
1. The fewer electrons an atom has, the smaller it is
iii.
The isoelectronic ion with the MOST protons is the smallest, the one with the LEAST protons is the
largest
5. 8.5→ Energy Effects in Binary Ionic Compounds
a. Lattice energy defines how strongly the ions attract each other. Lattice energy is the change in energy that
takes place when separated gaseous ions are packed together to form an ionic solid.
i.
Shown with negative sign because energy is released
b. Lattice energy calculations
6. 8.6→
a.
7.
8.7→
a.
b.
i.
Partial Ionic Character of Covalent Bonds
Hard to tell difference between ionic and polar covalent bonds
i.
No totally ionic bonds between discrete pairs of atoms
ii.
Percent ionic character
1. (Measured dipole moment of X-Y/calculated dipole moment of X+Y-) (100%)
iii.
Any compound that conducts an electric current when melted will be classified as ionic
The Covalent Chemical Bond: A Model
Chemical bonds are forces that cause a group of atoms to behave as a unit
i.
Bonds result from the tendency of a system to seek its lowest possible energy
Models: An Overview
i.
Models are attempts to explain how nature operates on the microscopic level based on
experiences in the macroscopic world
ii.
Fundamental Properties of Models
1. Models are human inventions, and therefore are not always completely accurate
a. Can actually be entirely incorrect
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2. Models become more complicated as they age
3. Have to understand how complicated/accurate a model is so that your results match
with the precision of the model
8. 8.8→ Covalent Bond Energies and Chemical Reactions
a. Bonds don’t always have the same energy (even between the
same elements) → bond energies we use are averages
b. Different energies for multiple bonds
c. Bond Energy and Enthalpy
i.
ΔH= sum of energies required to break old bonds (+) plus the sum of
the energies released in the formation of new bonds (-)
9. 8.9→ The Localized Electron Bonding Model
a. Localized electron model
i.
Description of the valence electron arrangement in the molecule using Lewis structures
ii.
Prediction of the geometry of the model using the valence shell electron pair repulsion model
iii.
Description of the type of atomic orbitals used by the atoms to share electrons or lone pairs
10. 8.10→ Lewis Structures
a. Lewis structure show how valence electrons are arranged among the atoms in the molecule
b. The most important requirement for the formation of a stable compound is that the atoms achieve noble
gas electron configurations
i.
To obey octet rule
c. Ionic
i.
Just draw ions with their valence electrons and charge
d. Covalent
i.
Sum the total number of valence electrons (add another for each negative charge on the
molecule)
ii.
Use a pair of electrons to form a bond between each pair of bound atoms
iii.
Arrange the remaining electrons to satisfy the duet rule (H or He) or the octet rule
11. 8.11→ Exceptions to the Octet Rule
a. Boron usually has 6 valence electrons instead of 8
b. S and Se often have more than 8 electrons (either 10 or 12) when in compounds such as SF6
c. Comments about Octet Rule:
i.
C, N, O, F always obey octet rule
ii.
B and Be often have fewer than 8 electrons (electron
deficient) → molecules they are in are very reactive
iii.
Second row elements can never exceed octet rule (because they don’t have a d orbital)
iv.
Third row and heavier elements usually obey the octet rule but can exceed it
12. 8.12→ Resonance
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a. Resonance is invoked when more than 1 valid Lewis structure can be written for a particular molecule. The
resulting electron structure of the molecule is given by the average of the resonance structures
b. Occurs because model shows electrons as localized when they don’t necessarily have to be
c. Odd Electron Molecules
i.
Cannot be dealt with using localized electron model, because this model thinks of electrons in
pairs
d. Formal Charge
i.
The difference between the number of valence electrons on the free atom and the number of
valence electrons assigned to the atom in the molecule
ii.
Lone pair electrons belong entirely to the atom in question
iii.
Shared electrons are divided equally between the two sharing atoms (counts for one electron for
each atom)
iv.
Still estimates→ do not count as entirely accurate
e. Rules Governing Formal Charge
i.
To calculate the formal charge on an atom:
1. Take the sum of the lone pair electrons and one half the shared electrons. This is the
number of valence electrons assigned to the atom in the molecule
2. Subtract the number of assigned electrons from the number of valence electrons on the
free, neutral atom to obtain the formal charge
ii.
The sum of the formal charges of all atoms in a given molecule or ion must equal the overall
charge on that species
iii.
If non equivalent Lewis structures exist for a species, those with formal charges closest to zero
and with any negative formal charges on the most electronegative atoms are considered to best
describe the bonding in the molecule or ion.
13. 8.13→ Molecular Structure: VSEPR
a. Valence Shell Electron Pair Repulsion (VSEPR)’s main postulate is that the structure around a given atom is
determined principally by minimizing electron pair repulsions
i.
That electrons pairs will be as far apart as possible
b. Steps to apply VSEPR
i.
Draw Lewis structure
ii.
Count electron pairs
iii.
Determine position of atoms from the way the electron pairs are shared
iv.
Determine name of molecular structure from position of atoms
c. Lone pairs require more room than bonding pairs
i.
This makes sense because they move around more→ not
centered between two nuclei
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a.
d. The VSEPR Model and Multiple Bonds
i.
Multiple bonds are counted as one effective pair
ii.
When a molecule exhibits resonance, any one of the resonance structures can be used to predict
the molecular structure using VSEPR
e. Molecules Containing No Central Atom
i.
Can use VSEPR with multiple central atoms
f. Summary of VSEPR
i.
Determine Lewis structure
ii.
Sum electron pairs around central atom
iii.
Count pairs (each multiple bond counts as one pair)
iv.
Figure out molecular structure
g. How well does VSEPR work?
i.
Pretty well most of the time→ some exceptions to
every model
14. Terms
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a. Bond Energy: energy required to break a specific bond
b. Ionic Bonding: the electrostatic attraction between oppositely charged ions
c. Ionic Compound: a compound that results when a metal reacts with a nonmetal to form a cation and an
anion
d. Coulomb’s Law: the equation that defines the strength of an ionic bond
e. Bond Length: the distance between the nuclei of two atoms connected by a bond, the distance where the
total energy of the bond is minimal
f. Covalent Bonding: bonding in which electrons are shared by nuclei
g. Polar Covalent Bond: a covalent bond in which the electrons are not shared equally
h. Electronegativity: the ability of an atom in a molecule to attract shared electrons to itself
i. Dipole Moment: a property of a molecule whose charge distribution can be represented by a center of
positive charge and a center of negative charge
j. Isoelectronic Ions: ions containing the same number of electrons
k. Lattice energy: the change in energy that takes place when separated gaseous ions are packed together to
form an ionic solid
l. Single Bond: bond in which one pair of electrons is shared
m. Double Bond: bond in which two pairs of electrons are shared
n. Triple Bond: bond in which three pairs of electrons are shared
o. Localized Electron Model: assumes that a molecule is composed of atoms that are bound together by
sharing pairs of electrons using the atomic orbitals of the bound atoms
p. Lone Pairs: electrons localized on an atom
q. Bonding Pairs: electrons found in the space between two atoms
r. Lewis Structure: a diagram of a molecule showing how the valence electrons are arranged among the
atoms in the molecule
s. Octet Rule: the observation that atoms of nonmetals tend to form the most stable molecules when they are
surrounded by eight valence electrons (to fill their valence orbitals)
t. Resonance: a condition occurring when more than one valid Lewis structure can be written for a particular
molecule. The actual electronic structure is not represented by any one of the Lewis structures but by the
average of all of them
u. Formal Charge: the difference between the number of valence electrons on the free atom and the number of
valence electrons assigned to the atom in the molecule
v. Molecular Structure: the 3D arrangement of atoms in a molecule
w. VSEPR: a model whose main postulate is that the structure around a given atom in a molecule is
determined principally by minimizing electron pair repulsions
Chapter 9 Notes (Covalent Bonding: Orbitals)
1.
9.1→ Hybridization and the Localized Electron Model
a. sp3 hybridization
i.
Assume that atoms have 4 equal orbitals→ put s and p
orbitals together to combine into sp3 orbitals
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ii.
b.
c.
d.
e.
f.
Whenever a set of equivalent tetrahedral atomic orbitals is required by an atom, this model
assumes that the atom adopts a set of sp3 orbitals; the atom becomes sp3 hybridized
sp2 hybridization
i.
Whenever an atom is surrounded by three effective pairs, a set of sp2 orbitals is required
ii.
Sigma (𝜎) and Pi(𝝅) bonds
1. Sigma bond
a. Formed using sp2 orbitals
b. In plane
2. Pi bonds
a. Formed using the unused p orbitals
b. Above and below plane
3. Double bonds are one sigma and one pi bond
sp hybridization
i.
Two effective pairs around an atom will always require sp hybridization of that atom
3
dsp hybridization
i.
A set of five effective pairs around a given atom always requires a trigonal bipyramidal
arrangement, which in turn requires dsp3 hybridization of that atom
d2sp3 hybridization
i.
Six electron pairs around an atom are always arranged octahedrally and require d 2sp3
hybridization of that atom
Summary of using localized electron model
i.
Lewis
ii.
VSEPR→ molecular structure
iii.
Hybrid orbitals
1.
2. 9.2→ The Molecular Orbital Model
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a.
b. Molecular Orbital Model
i.
The electron probability of both molecular orbitals
is centered along the line passing through the two
nuclei→ called sigma orbitals
1.
MO1→ between nuclei (bonding molecular orbital)
2. MO2→ either side of nuclei (antibonding
ii.
molecular orbital)
The shape of the orbital depends on whether the interference is constructive or destructive
(whether the wave functions of the orbitals are in phase or out of phase)
1.
iii.
iv.
In phase→ constructive→ bonding orbital
2. Out of phase→ destructive→ antibonding orbital
# of molecular orbitals= # of atomic orbitals
1. Conservation of energy
2. atomics orbitals no longer exist if they are used to form molecular orbitals
MO1 is lower in energy than the original orbitals, MO2 is higher in energy than the original orbitals
1. If the energy of the electrons in the molecular orbitals is lower than their energy in their
separated state, the situation is favorable to bonding (pro bonding)
2. If the energy of the electrons in the molecular
orbitals is higher than their energy in their
separated state, the situation is antibonding→
the atoms will not bond
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v.
A bonding orbital is one that is lower in energy than the atomic orbitals it is formed from, an
antibonding orbital is higher in energy than the atomic orbitals it is formed from
1. Bonding orbitals place electrons between the
nuclei, which prevents as much proton proton
repulsion from occurring→ lower energy→ more
vi.
stable
Naming molecular orbitals
1. Use * to denote anti bonding
2. Use sigma/pi and then the orbitals they came from
a.
vii.
Molecular electron configurations can be written in the same way as atomic electron
configurations
viii.
Each molecular orbital can hold 2 electrons
c. Bond Order
i.
The difference between the number of bonding electrons and the number of antibonding
electrons divided by 2
ii.
Larger bond order means greater bond strength
3. 9.3→ Bonding in Homonuclear Diatomic Molecules
a. P orbitals can also form MOs
i.
4 form Pi MOs (2 p orbitals not along internuclear axis)
ii.
2 form Sigma MOs(1 p orbital along internuclear axis)
iii.
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iv.
1. For F2
2. Sigma below pi because sigma bonds are stronger than pi bonds
b. Paramagnetism
i.
Paramagnetism: substance is attracted into the inducing magnetic field
ii.
Diamagnetism: substance is repelled from the inducing magnetic field
iii.
Paramagnetism is associated with unpaired electrons, diamagnetism is associated with paired
electrons
c. SP mixing
i.
Occurs when the 2s and 2p have relatively similar energies
ii.
All orbitals have contributions from 2s and from 2p orbitals, the amount from each determines
the energy of the orbitals
iii.
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iv.
v. Pi bonds do not change energy
d. Orbital Diagrams, Bond Energies, Bond Lengths
i.
As bond order increases, bond energy increases, bond length decreases
4. 9.5→ Combining the Localized Electron and Molecular Orbital
Models
a. Sigma bonds can be treated as localized, pi bonds must be treated as delocalized
i.
b. Interaction of p orbitals
5. 9.6→ Photoelectron Spectroscopy
a. Can be used to determine what energies electrons have (ionization energies) and the percent abundance
of each one
i.
Percent abundance expresses how many electrons exist at a certain energy, the energy level
expresses the orbital
b. Energy decreases from left to right
i.
Core electrons are the closest to the axis
6. Terms
a. Hybridization: the mixing of the native orbitals on a given atom to form special atomic orbitals for bonding
b. Hybrid Orbitals: a set of atomic orbitals adopted by an atom in a molecule different from those of the atom
in a free state
c. sp3 hybridization: hybridization of 3 p orbitals and 1 s orbital
d. sp2 hybridization: hybridization of 2 p orbitals and 1 s orbital
e. dsp3 hybridization: hybridization of 1 d orbital, 3 p orbitals, and 1 s orbital
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f.
g.
h.
i.
j.
k.
l.
m.
n.
o.
Molecular Orbital Model: a model that regards a molecule as a collection of nuclei and electrons, where the
electrons are assumed to occupy orbitals as much as they do in atoms, but having the orbitals extend over
the entire molecule. In this model the electrons are assumed to be delocalized rather than always located
between a given pair of atoms.
Sigma Bond: a covalent bond in which the electron pair is shared in an area centered on a line running
between the atoms
Pi Bond: a covalent bond in which parallel p orbitals share an electron pair occupying the space above and
below the line joining the atoms
Molecular Orbitals: mathematical function describing the wave-like behavior of an electron in a molecule
Sigma Molecular Orbital: molecular orbital where shared electron density is directly between the bonding
atoms, along the bonding axis
Bonding Molecular Orbital: molecular orbital with lower energy than the orbitals it was formed from
Antibonding Molecular Orbital: molecular orbital with higher energy than the orbitals it was formed from
Bond Order: The difference between the number of bonding electrons and the number of antibonding
electrons divided by 2
Paramagnetism: substance is attracted into the inducing magnetic field
Diamagnetism: substance is repelled from the inducing magnetic field
Chapter 10 Notes (Liquids and Solids)
1.
10.1→ Intermolecular Forces
a. Intramolecular bonding is within molecules, intermolecular forces are between molecules
b. Phase changes are physical not chemical
c. Dipole-Dipole Forces
i.
Partially charged parts of polar molecules are attracted to each other
ii.
Solid has more ordered organization of molecules than liquid
iii.
Not dependent on size, dependent on charge
iv.
When there are dipole dipole forces, there are also London Dispersion Forces
d. Hydrogen Bonding
i.
Special dipole-dipole
1. Not actually a bond
2. H-F, N-H, O-H because these are all small and have a very high effective nuclear charge
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a.
3. Stronger than normal Dipole-Dipole
ii.
When there is hydrogen bonding, there are also Dipole-Dipole Forces and London Dispersion
Forces
e. London Dispersion Forces
i.
In every kind of molecule
ii.
Caused by a temporary dipole moment (induced dipole)
1. Electrons are momentarily on one side of an atom (because of probability distribution),
causing the electrons in an atom near that side to be repelled
2.
iii.
The strength of the bond depends of the shape of the molecules and the number of electrons
1. The larger the molecule, the stronger the force because there is a larger chance of a
temporary dipole
2. 10.2→ The Liquid State
a. Liquid exists usually in droplets
i.
Molecules in interior pulled by more molecules than those on surface, spherical shape minimizes
surface area
b. Surface tension
i.
The resistance of a liquid to increase its surface area
ii.
Liquids with higher intermolecular forces generally have higher surface tensions
c. Capillary action
i.
Spontaneous rising of liquid in a narrow tube
ii.
Cohesive forces v adhesive forces
1. cohesive→ between molecules of liquid
2. adhesive→ between liquid molecules and container
a. Occur when the container a liquid is in has polar bonds (eg. glass)
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iii.
polar→ concave meniscus
a. adhesive>cohesive
2. nonpolar→ convex meniscus
a. cohesive>adhesive
1.
d. viscosity
i.
The measure of a liquid’s resistance to flow
ii.
As IMF increase, viscosity increases
e. Structural model for liquids
i.
Liquids have non negligible intermolecular forces and molecular movement, so they are much
harder to model than gases or solids
ii.
Can use spectroscopy to model them better
3. 10.3→ An Introduction to Structures and Types of Solids
a. Crystalline v Amorphous solids
i.
Amorphous: solid with considerable disorder in its structure
ii.
Crystalline: solid with a highly regular arrangement of its components
1. Shown in lattice, made up of unit cells
b. X Ray Analysis of Solids
i.
Structures of solids determined by X Ray Diffraction
1. Diffraction occurs when beams of light are scattered from a regular array of points in
which the spacing between the components are comparable with the wavelength of the
light
2. Used to determine intermolecular spacings
c. Bragg Equation
i.
d. Types of Crystalline Solids
i.
Ionic v molecular v atomic solids
1. ionic→ ions at the points of the lattice that
describes the structure of the solid
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2. molecular→ discrete covalently bonded molecules
at each of its lattice points
3. atomic→ have atoms at the lattice points that
describe the structure of the solid
4.
e.
4. 10.4→ Structure and Bonding in Metals
a. Closest packing: placing the spherical atoms of a metal as close together as possible
i.
Hexagonal closest packing: body centered cubic (aba)
ii.
Cubic closest packing: face centered cubic (abca)
iii.
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iv.
1. Coordination number is the number of atoms a single atom in the structure is closest to
b. Density of a unit cell
i.
ii.
Volume of unit cell
1.
2. V=a3
iii.
Packing efficiency
1. Fraction of volume occupied (higher packing efficiency means more dense)
c. Bonding Models for Metals
i.
Electron sea model
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ii.
1. Regular array of metal cations in a sea of valence electrons
Band model
1. Electrons are assumed to travel around the metal crystal in molecular orbitals formed
from the valence atomic orbitals of the metal atoms
2. The large number of molecular orbitals become more closely spaced and form a virtual
continuum of levels, called bands
3. Metals conduct electricity and heat very well
because of the availability of highly mobile
electrons (can easily move to conduction band→
no band gap)
a.
d. Metal Alloys
i. Alloy: a substance that contains a mixture of elements and has metallic properties
1. Substitutional Alloy: a type of alloy in which host metal atoms are replaced by other
metal atoms of similar size
2. Interstitial Alloy: a type of alloy in which holes in the packed metal structure are filled
with small atoms
a. Eg. steel (Carbon in holes of iron crystal)
ii.
5. 10.5→ Carbon and Silicon: Network Atomic Solids
a. Network solid→ giant molecule
b. Diamond v graphite
i.
Graphite has weak bonding between layers, but strong bonding within layers
ii.
Conductivity
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1.
Diamond has all sigma bonds→ very large band gap
a.
2. Graphite has pi bonds→ very small band gap
a.
iii.
c. Silicon
i.
Silica (SiO2)
1. Quartz
a. Strong long range order
b. Each silicon is connected to 4 Carbons
2. Glass
a. Amorphous
ii.
3.
silicates: substances that contain silicon oxygen anions
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d. Ceramics
i.
Nonmetallic materials that are strong, brittle, and resistant to heat and attack by chemicals
ii.
Heterogenous
iii.
Made by the weathering action of water and carbon dioxide on the mineral feldspar, which is a
mixture of silicates
e. Semiconductors
i.
Between conductors and insulators
ii.
1.
Doping
1. Replacing main atom (usually Si) with an atom with a different number of valence
electrons
a. N type: replace using atom with more electrons
i.
ii.
Extra electron
iii.
Electrons closer to conduction band
b. P type: replace using atom with less electrons
i.
ii.
iii.
Hole forms
Electrons move around to fill hole in valence band, hole moves
around
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c.
2. Junctions
a. PN junction
i.
ii.
Holes move from p to n, electrons move from n to p
b. Reverse and Forward Bias
i.
Hooking PN junction up to battery
ii.
1.
In reverse, no current flows. In forward, current flows
6. 10.6→ Molecular Solids
a. Strong bonds within molecules, weak forces between molecules
i.
Stronger as the molecule gets bigger because of London Dispersion Forces
7. 10.7→ Ionic Solids
a. Stable, high melting substances held together by the strong electrostatic attraction forces that exist
between oppositely charged ions
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b. Usually packed by doing closest packing with larger ion (usually anion) and filling holes with smaller ion
(usually cation)
i.
Types of holes:
1. Trigonal
2. Tetrahedral
3. Octahedral
ii.
4.
Type of hole depends on relative radii of ions
1. Smallest hole→ trigonal→ if bigger ion is a lot
bigger than smaller ion
2. Medium hole→ tetrahedral→ if bigger ion is kind
of bigger than smaller ion
3. Biggest hole→ octahedral→ if bigger ion is
pretty much the same size as smaller ion
iii.
Not all holes are filled→ depends on ratio of ions
c. Types and Properties of Solids
i.
8. 10.8→ Vapor Pressure and Changes of State
a. Vapor Pressure
i.
Liquid
1. Evaporation occurs at a constant rate at a given temperature
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2. Eventually, enough vapor molecules are present above the liquid so that the rate of
condensation equals the rate of evaporation
a. Equilibrium (no further net change occurs in the amount of liquid or vapor
because they balance each other)
b. Pressure of vapor present at equilibrium=vapor pressure
3. Liquids with high vapor pressures are volatile
a. Evaporate quickly
4. Determinants of Vapor Pressure
a. Vapor pressure is determined by the size of the intermolecular forces in the
liquid
i.
Large intermolecular forces→ low
vapor pressure because molecules need
high energies to escape to the vapor
phase
ii.
Larger molar mass usually means lower vapor pressure because it
also means larger dispersion forces
b. Vapor pressure increases with temperature
i.
As a liquid gets hotter, it has more energy and thus more molecules
have enough energy to jump to the vapor phase
ii.
Can be plotted on a line
c.
5. Boiling point
a. When vapor pressure is equal to atmospheric pressure
b. Normal boiling point→ vapor pressure is 1
atm
6. Equations
a.
i.
Enthalpy of vaporization: energy it takes to convert a substance
from liquid to vapor, usually in kJ/mol
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b.
i.
If you know VP at one temperature and you want to know it at
another
If you know VP and want to know enthalpy of vaporization
ii.
Solid
1. Sublimation is when a substance changes from a solid to a gas
b. Changes of State
i.
Physical changes
ii.
ii.
iii.
Melting and boiling points are determined by the vapor pressures of the solid and liquid states
1. Melting point: Solid and liquid phase have the same vapor pressures
2. Case 1: Solid vapor pressure>liquid vapor pressure
a. Conversion from solid to liquid through vapor phase
b. Only the liquid state can exist in this situation at equilibrium
3. Case 2: Solid vapor pressure<liquid vapor pressure
a. Liquid requires greater pressure than solid to be at equilibrium, liquid turned
into solid
4. Case 3: Solid vapor pressure=liquid vapor pressure
a. Liquid and solid can coexist
Phase changes can occur at different temperature if a liquid is supercooled or superheated
1. supercooled→ not enough organization to form
solid
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2. superheated→ bubble forms, vapor
pressure>atmospheric pressure, liquid is
splashed
9. 10.9→ Phase Diagrams
a. A way to represent the phases of a substance as a function of temperature and pressure
i.
In a closed system
ii.
1.
Points
a. Triple point: the point at which all three phases of water are present
b. Critical pressure: the pressure required to produce liquefaction at the critical
temperature
c. Critical temperature: the temperature above which vapor cannot be liquefied
no matter what pressure is applied
d. Critical point: (critical temperature, critical pressure)
e. Normal melting point: melting point at 1 atm
f. Normal boiling point: temperature at which vapor pressure of the liquid is 1
atm
1.
Fusion curve: solid→liquid, liquid→ solid
iii.
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2. Sublimation curve: solid→ gas, gas→ solid
3. Vaporization curve: liquid→ gas, gas→ liquid
iv.
Make sure to note whether pressure is to scale or not
b. Phase Diagram of Water
i.
ii.
Fusion curve has negative slope
1. Because solid is more dense than liquid
c. Phase Diagram of Carbon Dioxide
i.
ii.
10. Terms
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
m.
At 1 atm, Carbon Dioxide sublimes (dry ice)
Intramolecular force: a force within a molecule
Intermolecular force: a force between molecules
Condensed States: liquid or solid
Surface tension: The resistance of a liquid to increase its surface area
Capillary action: Spontaneous rising of liquid in a narrow tube
Viscosity: The measure of a liquid’s resistance to flow
Crystalline Solids: solid with a highly regular arrangement of its components
Amorphous solids: solid with considerable disorder in its structure
Lattice: a 3D system of points designating the positions of the components that make up the substance
Unit Cell: smallest repeating unit of the lattice
Ionic solid: ions at the points of the lattice that describes the structure of the solid
Molecular solid: discrete covalently bonded molecules at each of its lattice points
Atomic solid: have atoms at the lattice points that describe the structure of the solid
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n. Band Model: Electrons are assumed to travel around the metal crystal in molecular orbitals formed from the
valence atomic orbitals of the metal atoms
o. Conduction Band: a delocalized band of energy partly filled with electrons in a crystalline solid. These
electrons have great mobility and are responsible for electrical conductivity.
p. Valence Band: The outermost electron shell of atoms in an insulator or semiconductor in which the
electrons are too tightly bound to the atom to carry electric current
q. Alloy: a substance that contains a mixture of elements and has metallic properties
r. Substitutional Alloy: a type of alloy in which host metal atoms are replaced by other metal atoms of similar
size
s. Interstitial Alloy: a type of alloy in which holes in the packed metal structure are filled with small atoms
t. Network solid: giant molecule
u. silicates: substances that contain silicon oxygen anions
v. Ceramics: Nonmetallic materials that are strong, brittle, and resistant to heat and attack by chemicals
w. Semiconductor: a solid substance that has a conductivity between that of an insulator and that of most
metals, either due to the addition of an impurity or because of temperature effects
x. Insulator: a substance that doesn’t conduct
y. Doping: the intentional introduction of impurities into a lattice in order to make it a better semiconductor
z. N type Semiconductor: the addition of an atom into a lattice with more valence electrons than the “central
atom”
aa. P type Semiconductor:the addition of an atom into a lattice with less valence electrons than the “central
atom”
bb. PN Junction: a boundary or interface between two types of semiconductor material, p-type and n-type,
inside a single crystal of semiconductor
cc. Reverse Bias Diode: The application of a reverse voltage to the p-n junction which causes a transient
current to flow as both electrons and holes are pulled away from the junction.
dd. Forward Bias Diode: a voltage applied to a circuit or device, usually a semiconductor device, in the direction
that produces the larger current
ee. Vaporization: conversion of a liquid into a gas
ff. Evaporation: the process of a substance in a liquid state changing to a gaseous state due to an increase in
temperature and/or pressure
gg. Heat of Vaporization: the energy required to vaporize one mole of a liquid at a pressure of 1 atmosphere
hh. Condensation: Process by which a gas changes to a liquid
ii. Equilibrium: a dynamic condition in which two opposing changes occur at equal rates in a closed system
jj. Vapor Pressure: pressure exerted by a vapor in equilibrium with its corresponding liquid at a given
temperature
kk. Sublimation: change of state from a solid directly to a gas
ll. Normal Melting Point: the temperature at which the solid and liquid states have the same vapor pressure
under conditions where the total pressure is 1 atm
mm. Normal Boiling Point: the temperature at which the vapor pressure of the liquid is exactly 1 atm (and the
atmospheric pressure is 1 atm)
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nn.
oo.
pp.
qq.
rr.
ss.
tt.
uu.
vv.
Phase Diagram: A way to represent the phases of a substance as a function of temperature and pressure
Triple point: the point at which all three phases of a substance are present
Critical pressure: the pressure required to produce liquefaction at the critical temperature
Critical temperature: the temperature above which vapor cannot be liquefied no matter what pressure is
applied (above which there is supercritical fluid)
Critical point: (critical temperature, critical pressure)
Fusion curve: liquid/solid barrier, also known as freezing/melting
Sublimation curve: solid/vapor barrier
Vaporization curve: liquid/vapor barrier, also known as evaporation/condensation
Polarizability: the ability to form instantaneous dipoles (directly related to LDF)
Chapter 11 Notes (Properties of Solutions)
1.
11.1→ Solution Composition
a. Ways to define concentration
i.
Molarity (M)
1. Moles solute/L soln
ii.
Mass percent
1. (Mass solute/mass solution)(100%)
iii.
Mole Fraction (𝟀)
1.
a. if A and B are the only components of the solution
iv.
Molality (m)
1. Moles solute/kg solvent
a. Does not change with temperature
v. Normality (N)
1. # equivalents/L soln where the definition of an equivalent depends on the reaction
taking place in the solution
a. ACID BASE RXN: The mass of acid or base that can furnish or accept exactly 1
mole of protons
b. REDOX RXN: the quantity of oxidizing or reducing agent that can accept or
furnish exactly one mole of electrons
2. 11.2→ The Energies of Solution Formation
a. Steps of forming solution
i.
Separate solute
Takes energy
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Make room for solute in solvent
Takes energy
Form interactions between solute and solvent Releases energy
1.
2. Change in energy associated with the formation of the solution
is the enthalpy of solution (ΔHsoln)
3.
a. Forms if ΔHsoln is close to or less than 0. Even if it is slightly above the
solution will form because of entropy
ii.
3. 11.3→ Factors Affecting Solubility
a. Structure effects
i.
Hydrophobic→ water fearing (mostly nonpolar)
1. CH bonds
ii.
Hydrophilic→ water loving (polar or charged)
1. OH, NH, C=O bonds
b. Pressure Effects
i.
Pressure increases the solubility of a gas
ii.
Henry’s Law
1.
2. Best for low concentration, weak IMF between solute and solvent, and gases that do not
dissociate
c. Temperature Effects
i.
For solid: as temperature increases, solubility USUALLY increases
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ii.
1.
For gas: as temperature increases, solubility decreases
1.
a. Gas has more kinetic energy, comes out of soln faster
4. 11.4→ The Vapor Pressures of Solutions
a. The presence of a nonvolatile solute lowers the vapor pressure of a solvent
i.
Has to be nonvolatile because if not it contributes to the vapor pressure and makes it more
complicated
ii.
Works because it in essence dilutes the concentration of the solvent, which lowers the rate at
which the solvent vaporizes, thus lowering the vapor pressure of the solution
b. Raoult’s Law
i.
Explains an “ideal solution”
ii.
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iii.
iv.
When solutions dissociate:
1. Count as the number of moles of ions they dissociate into. Eg. NaCl should be counted
as two moles, because it dissociates into one mole of Na+ ions and one mole of Cl- ions
c. Nonideal solutions
i.
A solution that follows Raoult’s law is called ideal
1. Solute solute, solvent solvent, solute solvent interactions about equal
ii.
When solute solvent interactions are stronger than solute solute or solvent solvent interactions
iii.
1.
2. Largest deviation in middle because that is where there are the most solute solvent
interactions
When solute solvent interactions are weaker than solute solute or solvent solvent interactions
1.
5. 11.5→ Boiling Point Elevation and Freezing Point Depression
a. Boiling Point Elevation
i.
ii.
Liquid vapor pressure is lower, meaning it takes a higher T to make the liquid vp the same as the
atmospheric vp
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b. Freezing Point Depression
i.
Solid and liquid vapor pressures both change, freezing point decreases
6. Terms
a.
b.
c.
d.
ii.
Molarity: moles solute/liters solution
Mass percent: (Mass solute/mass solution)(100%)
Molality: Moles solute/kg solvent
Mole fraction: the ratio of moles of a given component in a mixture to the total number of moles in the
mixture
e. Normality: the number of equivalents of a substance dissolved in a liter of solution
f. Enthalpy of solution: Change in energy associated with the formation of the solution
g. Enthalpy of hydration:the amount of energy released when a mole of the ion dissolves in a large amount of
water forming an infinite dilute solution in the process
h. Hydrophobic: water fearing
i. Hydrophilic: water loving
j. Colligative Property: properties that depend upon the concentration of solute molecules or ions, but not
upon the identity of the solute
Chapter 12 Notes: Chemical Kinetics
1.
12.1→ Reaction Rates
a. For a reactant A
i.
ii.
Put a negative sign in front. Rates should always be positive.
b. Rates must work with the stoichiometry of the reaction. For example, if a product to reactant ratio in the
chemical equation is 1:1, these two substances must have the same rate over the same time period
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i.
c. Rates are not constant over the course of a reaction
i.
Also possible to use a derivative (instantaneous rate of change)
d. Given in M/s
e. Reaction Rate
i.
2. 12.2→ Rate Laws: An Introduction
a. If we choose conditions where the reverse reaction can be ignored, the reaction rate will depend only on
the concentrations of the reactants
b.
i.
Note: x and y MUST be experimentally determined
1. Usually: 0, 1, 2 (zero, first, second order)
ii.
c. Types of Rate Laws
i.
Differential Rate Law (rate law)
1. Expresses how rate depends on concentration
ii.
Integrated Rate Law
1. Expresses how concentrations depend on time
3. 12.3→ Determining the Form of the Rate Law
a. See how a change in initial concentration affects the rate in order to determine the reactant orders
b. Then use any of the given experiments to solve for K
c. Overall Reaction Order
i.
Reactant orders added up
4. 12.4→ The Integrated Rate Law
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a.
i.
Integrated rate law is ALWAYS a straight line--look for which of these is a straight line to solve for
the reaction order
b.
c. Types of Reactions
i.
Zero Order: when a substance such as a metal surface or an enzyme is required for the reaction to
occur
ii.
First Order: NUCLEAR
d. Half Life of a First Order Reaction
i.
ii.
Does NOT depend on concentration
1. For both zero and second order reactions, it does
e. Pseudo First Order Rate Laws
i.
Have such large amounts of all but one reaction that they functionally remain the same and the
rate law is determined using only the one reactant that is not in extreme excess
5. 12.5→ Reaction Mechanisms
a. Most reactions occur in a series of steps called the reaction mechanism
i.
Sum of elementary steps must give the overall balanced equation for the reaction
ii.
The mechanism must agree with the experimentally determined rate law
b. A balanced chemical equation for a reaction does not give any direct information about the reaction
mechanism
c. Intermediates: species that are neither reactants nor products but that are formed and consumed during
the reaction sequence
d. Elementary steps
i.
A reaction whose rate law can be written from its molecularity (uni, bi, ter…)
ii.
The slowest step is the rate determining step. We use this step to write the rate law
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1.
Harder when this step is not first because intermediates can go back to reactants
iii.
Mechanism cannot be proved absolutely
6. 12.6→ A Model for Chemical Kinetics
a. Two requirements must be satisfied for reactants to collide successfully
i.
The collision must involve enough energy to product the reaction; that is, the collision energy
must equal or exceed the activation energy
ii.
The relative orientation of the reactants must allow formation of any new bonds necessary to
produce products
b.
i.
1.
Plotting ln K v. 1/T gives straight line with slope -Ea/R
c. Graphs
7.
i.
12.7→ Catalysis
a. A catalyst is a substance that speeds up a reaction without being consumed itself
i.
Provides new reaction pathway with lower activation energy
b. Homogenous/Heterogenous
i.
Homogenous: in same phase as reacting molecules
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ii.
Heterogenous: in a different phase than molecules (usually a solid)
8. Terms
a. Chemical Kinetics: the area of chemistry that concerns reaction rates
b. Reaction Rate: the change in concentration of a reactant or product per unit time
c. Rate Constant: a coefficient of proportionality relating the rate of a chemical reaction at a given
temperature to the concentration of reactant
d. Differential Rate Law (rate law): Expresses how rate depends on concentration
e. Integrated Rate Law: Expresses how concentrations depend on time
f. Order: the positive or negative exponent, determined by experiment, of the reactant concentration in a rate
law
g. Initial Rate: instantaneous rate at t=0
h. Overall Reaction Order: sum of reactant orders
i. Half life: time required for a reactant to reach half of its original concentration (t1/2)
j. Reaction Mechanism: a series of elementary steps
k. Intermediates: species that are neither reactants nor products but that are formed and consumed during
the reaction sequence
l. Elementary step: A reaction whose rate law can be written from its molecularity
m. Molecularity: number of species that must collide to produce the reaction indicated by that step
n. Rate determining step: the slowest elementary step of a reaction (the step that determines the rate of the
reaction)
o. Activation energy: the minimum quantity of energy that the reacting species must possess in order to
undergo a specified reaction
p. Frequency factor (A): combination of collision frequency and steric factor
q. Steric factor: the factor (always less than 1) that reflects the fraction of collisions with orientations that can
produce a chemical reaction
r. Catalyst: a substance that speeds up a reaction without being consumed itself
Chapter 13 Notes (Chemical Equilibrium)
1.
13.1→ The Equilibrium Condition
a. Equilibrium
i.
Dynamic condition
ii.
No net change
1. Forward and reverse rates are equal
iii.
Reaction rates depend on concentrations of reactants
b. At equilibrium, ΔG=0
i.
2. 13.2→ The Equilibrium Constant
a. Keq
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i.
where a, b, c, d represent the coefficients on their
respective substance in the balanced chemical equation
ii.
NO UNITS
iii.
Given for certain T, same for any amounts at that T
iv.
AT EQUILIBRIUM
1. Not at equilibrium, this equation gives Q
b. Conclusions about the equilibrium expression
i.
If reaction is reversed, K→ 1/K
ii.
If reaction is multiplied by a constant A, K→ KA
iii.
If two reactions are added together, K=(K1)(K2)
c. Equilibrium position
i.
Each set of equilibrium concentrations
3. 13.3→ Equilibrium Expressions Involving Pressures
a. KP
i.
ii.
1.
4. 13.4→ Heterogenous Equilibria
a. Involve more than one phase
b. If pure solids or liquids are included in chemical reaction, they ARE NOT included in the equilibrium
expression
i.
Since they are constant at a given T (since you calculate concentration from density)
5. 13.5→ Applications of the Equilibrium Constant
a. The Extent of a Reaction
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i.
ii.
iii.
>> or << by orders of magnitude
iv.
Size of K and time to reach equilibrium are NOT directly related
b. Reaction Quotient
i.
Used to determine which way the reaction is moving
ii.
6. 13.6→ Solving Equilibrium Problems
a. Problem Solving Strategy
i.
Write balanced equation
ii.
Write equilibrium expression
iii.
List initial conditions
iv.
Calculate Q to determine direction of rxn
v. Use ICE to find new concentrations
1.
2. Then use these expressions to find X by plugging in with equilibrium expression
a. REMEMBER: cannot have negative molarity, positive root is right
vi.
Check calculated equilibrium concentrations by making sure they give the correct K
b. Treating Systems That Have Small Equilibrium Constants
i.
Use steps 1-5 given above, then use 5% trick
ii.
Assume x is small enough to be negligible on one side, pretend it is
iii.
Solve for X
iv.
Prove 5%!!! (do assumption-actual/assumption)
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1. Usually works if assumption is greater than 1000 times KC
7. 13.7→ Le Chatelier's Principle
a. If a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction
that reduces that change
b. The Effect of a Change in Concentration
i.
If a component is added, equilibrium shifts to reduce the amount of that component
ii.
If a component is reduced, equilibrium shifts to increase the amount of that component
c. The Effect of a Change in Pressure
i.
Addition of an inert gas
1. Does not affect equilibrium (no effect on partial pressures)
ii.
Reduction in Volume
1. System reacts by reducing moles of gas present
d. The Effect of a Change in Temperature
i.
K changes with T
ii.
Think of a change in T as either heat as either a product (exothermic) or a reactant (endothermic)
1. T goes down→ heat removed
a. exothermic→ shift towards products
b. endothermic→ shift towards reactants
2. T goes up→ heat added
a. exothermic→ shift towards reactants
b. endothermic→ shift towards products
8. Terms
a. Equilibrium: the state in which both reactants and products are present in concentrations which have no
further tendency to change with time. Usually, this state results when the forward reaction proceeds at the
same rate as the reverse reaction
b. Law of mass action: law stating that the rate of any chemical reaction is proportional to the product of the
masses of the reacting substances, with each mass raised to a power equal to the coefficient that occurs in
the chemical equation
c. Equilibrium constant: the value obtained when equilibrium concentrations of the chemical species are
substituted in the equilibrium expression
d. Equilibrium Expression: the expression (from the law of mass action) obtained by multiplying the product
concentrations and dividing by the multiplied reactant concentrations, with each concentration raised to a
power represented by the coef cient in the balanced equation
e. Equilibrium Position: Each set of equilibrium concentrations
f. Homogenous Equilibrium: an equilibrium involving reactants/products of only one phase
g. Heterogenous Equilibrium: an equilibrium involving reactants and/or products in more than one phase
h. Reaction Quotient: a function of the activities or concentrations of the chemical species involved in a
chemical reaction
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i.
Le Chatelier’s Principle: If a change is imposed on a system at equilibrium, the position of the equilibrium
will shift in a direction that reduces that change
Chapter 14 Notes (Chemical Equilibrium)
1.
14.1→ The Nature of Acids and Bases
a. Arrhenius Acids and Bases
i.
Arrhenius Acid: produces H+ ion in aqueous solution
ii.
Arrhenius Base: Produces OH- ion in aqueous solution
iii.
Problems:
1. Only applies to aqueous solutions
2. Not all bases have OHb. Bronsted Lowry Model
i.
Bronsted Lowry Acid: proton donor
ii.
Bronsted Lowry Base: proton acceptor
iii.
Works in non-aqueous solutions
c. Protons
i.
proton=H+=H3O+ (hydronium)
d. Conjugate Acids and Bases
i.
e. Ka and Kb
i.
1.
Acid dissociation constant
ii.
2. 14.2→ Acid Strength
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a.
i.
Strong acid is one where this equilibrium (Ka) lies far to the right (>1)
1. Assume 100% dissociation in calculations
ii.
Weak acid is one where this equilibrium (Ka) lies far to the left (<<1)
b. Strong Acids
i.
HCl, HBr, HI, HNO3, HClO4, H2SO4
c. A strong acid yields a weak conjugate base, a weak acid yields a strong conjugate base
d.
e. Polyprotic Acids
i.
An acid with multiple protons
ii.
iii.
f.
Keeps getting deprotonated, but K decreases for each successive reaction
1. This is because it already has a - charge so it is harder to remove another proton
2. USUALLY second and third are irrelevant, except in the case of H2SO4
Water as an Acid and a Base
i.
Water is amphoteric: it can behave as either an acid or a base
ii.
Autoionization
1.
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iii.
KW=1.0x10-14 at 25 ℃
a. Thus, in neutral water at this T [H3O+]=[OH-]=1.0x10-7
3. 14.3→ The pH scale
1.
a.
b.
i.
As pH goes up, acidity goes down
c. REMEMBER: WEIRD SIG FIGS
i.
Only numbers after the decimal point count as sig figs
4. 14.4→ Calculating the pH of Strong Acid Solutions
a. Assume 100% dissociation (and if small number add 1.0x10-7 to account for H3O+ already in
neutral water (if at
b. NO NEED FOR ICE
25 ℃)
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5. 14.5→ Calculating the pH of Weak Acid Solutions
a. Use ICE (and probably 5% approximation)
b. Solving Weak Acid Equilibrium Problems
i.
Write equation
ii.
(determine Q to determine direction of reaction)
iii.
ICE
iv.
Write the equilibrium expression
v. Solve the equation for x
1. Check 5% if used
vi.
Figure out concentrations
c. Percent Dissociation
i.
Amount dissociated/initial concentration (100%)
6. 14.6→ Bases
a. Strong Bases
i.
NaOH, KOH, metal hydroxide what is soluble
ii.
100% dissociation
b. Not all bases contain OH
i.
NH3 and many other molecules with a lone pair on N also act as bases
c. Weak Bases
i.
Not 100% dissociation
7. 14.7→ Polyprotic Acids
a. When doing successive equilibria make sure to include the initial concentration of H3O+
8. 14.8→ Acid Base Properties of Salts
a. Salts that Produce Neutral Solutions
i.
Cation: from strong base
ii.
Anion: from monoprotic strong acid
b. Salts that Produce Basic Solutions
i.
Cation: from strong base
ii.
Anion: conjugate base of weak acid (stronger Kb)
1.
c. Salts that Produce Acidic Solutions
i.
Cation: conjugate acid of a weak base
ii.
1.
Anion: from monoprotic strong acid
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d. Checking Ks
i.
Cation: from conjugate acid of weak base
ii.
Anion: from conjugate base of weak acid
iii.
CHECK Ks, lower K of weak acid or base means higher K
of conjugate→ higher is the one it will be
e. Process
i.
What ions
ii.
What is K
iii.
What is equilibrium rxn
iv.
ICE
v. math
9. 14.9→ The Effect of Structure on Acid Base Properties
a. Bond Polarity
i.
As the HA bond gets more polar (as the electronegativity of A increases), the acid gets stronger
(because then the H is more 𝛅+ and is thus more easily removed)
1. NOTE: pH is not determined by strength, but rather by strength AND concentration
ii.
b. Oxyacids
i.
As # of O goes up, acid strength increases
ii.
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iii.
1.
Because the electrons get pulled away from the H, and because the - charge on the
conjugate base is delocalized, making it more stable and thus weaker
10. 14.10→ Acid Base Properties of Oxides
a. Covalent: usually gas, acidic
b. Ionic: usually solid, basic
i.
Metal hydroxides
11. 14.11→ The Lewis Acid Base Model
a. Focuses on electron pairs, not protons
i.
Acid: electron pair acceptor, uses lone pair to form bond
ii.
Base: electron pair donor
b.
12. 14.12→ Strategy for Solving Acid Base Problems: A Summary
a. Remember ICE
b. Write everything down
c. Think through the chemistry first
13. Terms
a. Arrhenius Concept: a concept postulating that acids produce hydrogen ions in aqueous solution, while
bases produce hydroxide ions
b. Bronsted Lowry Model: a model proposing that an acid is a proton donor and a base is a proton acceptor
c. Conjugate Acid: the species formed when a proton is added to a base
d. Conjugate Base: the species formed when a proton is removed from an acid
e. Conjugate Acid Base Pair: two species related to each other by the accepting/donating of a single proton
f. Acid Dissociation Constant: Ka
g. Strong Acid: an acid that completely dissociates in aqueous solution to form H3O+ and its conjugate base
h. Weak Acid: an acid that does not completely dissociate in aqueous solution
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i.
j.
k.
l.
m.
n.
o.
p.
q.
r.
Oxyacid: an acid that contains O
Amphoteric: a substance that acts both as an acid and a base
Autoionization: the transfer of a proton from a molecule to another molecule of the same substance
Polyprotic Acid: an acid with multiple acidic protons
Monoprotic Acid: an acid with a single acidic proton
pH Scale: a way to represent a solution’s acidity
Major Species: components present in relatively large amounts
Percent Dissociation: Amount dissociated/initial concentration (100%)
Strong Base: a base that completely dissociates in aqueous solution to form OH- and its conjugate acid
Weak Base: a base that does not completely dissociate in aqueous solution
Chapter 15: Acid Base Equilibria
1.
15.1→ Solutions of Acids or Bases Containing a Common Ion
a. Common ion effect
i.
Addition of an ion already involved in the equilibrium, which in turn shifts the equilibrium
ii.
For weak acid: conjugate base→ pH ↑
iii.
For weak base: conjugate acid→ pH ↓
b. Equilibrium Calculations
i.
Add common ion to ICE table, then solve again
2. 15.2→ Buffered Solutions
a. A solution that resists a change in pH
b. Buffer Composition
i.
Weak acid + conjugate base OR weak base + conjugate acid in significant quantities together in
solution
c.
d.
e. pH
i.
Determined by the ratio of [HA] to [A-]
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f.
ii.
Henderson Hasselbalch Equation (DO NOT USE)
i.
3. 15.3→ Buffering Capacity
a. pH is determined by ratio of [A-] to [HA], capacity of buffered solutions is determined by the magnitudes of
[HA] and [A-] and the amount of each present
b. If you want a buffer for a certain pH range, use a buffer with a pKa within 1 of that pH
c. Ratio of [HA] to [A-] should be between .1 and 10 (within one order of magnitude), but a buffer with more
total concentration will have a higher buffer capacity
d. Problem Solving Strategy
i.
Determine chemical reaction(s) in play
ii.
Do stoichiometry for anything with a huge K→ MOLES
(use chart)
1.
iii.
Set up equilibrium reaction
iv.
ICE→ CONCENTRATION
v. math
4. 15.4→ Titrations and pH curves
a. Titration nuts and bolts
i.
Want 4 sig figs
b. Strong Acid Strong Base Titrations
i.
ii.
iii.
iv.
If reactants are in stoichiometric quantities, the reaction will be a neutralization reaction, pH=7
1. EQUIVALENCE POINT AT 7
Use stoichiometry to find amount of H3O+, then determine pH (use mole chart)
Graph:
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1.
a. Strong acid with strong base
b.
2.
a. Strong base with strong acid
c. Weak Acid Strong Base Titrations
i.
ii.
iii.
This solution will essentially fully react to form a weakly basic solution (huge K)
1. Equivalence point around 8-9
2. NOTICE: buffer zone
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3.
iv.
Bigger original rise in weaker acids, bigger rise around equivalence point for stronger
acids
d. Weak Base Strong Acid Titrations
i.
This solution will essentially fully react to form a weakly acidic solution (huge K)
ii.
1.
1. Equivalence point around 3-4
2. NOTICE: buffer zone
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iii.
e. Polyprotic Acids
i.
1.
Multiple equivalence points
5. 15.5→ Acid Base Indicators
a. Ways to determine equivalence point
i.
pH meter
ii.
Acid base indicator which marks the end point of a titration by changing color. The equivalence
point and the end point are not the same, but are assumed to be very close together. This is why it
is vital to choose a good indicator.
b. Indicators
i.
Indicators change color when they lose or gain a proton, which happens depending on the
amount of H3O+ available or available to be furnished by the solution
ii.
For acids, choose indicator that changes around 8-9; for bases choose one that changes around
3-4
iii.
When you get close to the equivalence point, pH changes VERY quickly
1. Be really careful not to put in too much, should be a very faint change
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iv.
1.
Usually use phenolphthalein
v.
6. Terms
a. Common Ion: ion already involved in equilibrium
b. Common Ion Effect: Addition of an ion already involved in the equilibrium, which in turn shifts the
equilibrium
c. Buffered Solution: the solution which resists changes in pH in the case of addition of small amount of acid
or base
d. Buffering Capacity: a measure of the efficiency of a buffer in resisting changes in pH; how much acid/base
can be added to a buffer before it is no longer a buffer
e. pH Curve: plot showing mL titrant added vs. pH
f. Acid base indicator: marks end point of acid base titration by changing color
g. End point: point in titration when the visual indicator changes color
h. Equivalence Point: point in titration when moles titrant added=moles unknown present in solution
i. Titration: a quantitative determination of the amount of acid/base by the (very nearly) stoichiometric
reaction of that acid or base with a known volume of known concentration of a strong base or acid
j. Titrant: added from burette to Erlenmeyer flask containing an unknown acid or base
Chapter 16: Solubility and Complex Ion Equilibria
1.
16.1→ Solubility Equilibria and the Solubility Product
a. We can think of solubility as an equilibrium, with solubility product constant Ksp
i.
Only includes ions, since pure phases are not included
b. Equilibrium expression only applies when there is solid left in the solution!
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c. Solve problems the same way as a normal equilibrium problem
d. Relative Solubilities
i.
Can only compare Ksp if the stoichiometry is the same, otherwise solve for s and compare those
e. Common Ion Effect
i.
The presence of a common ion greatly decreases s
ii.
Use 5% rule
f. pH and solubility
i.
Salts that are basic are more soluble in acidic solution than in water because the anion can react
with the solution, thus removing products and causing a shift towards the right by Le Chatelier's
Principle
ii.
Strong Base:
1. Add strong base: solubility decreases by Le Chatelier because you are increasing the
concentration of OH- in solution
2. Add strong acid: H3O+ reacts with OH-, so [OH-] decreases, equilibrium shifts towards
products and solubility increases
iii.
Weak Base
1. Add strong base: OH- reacts with weak acid to form more of the basic anion, [anion]
goes up,equilibrium shifts towards reactants and solubility decreases
2. Add strong acid: H3O+ reacts with basic anion to form weak acid, [anion] goes down,
equilibrium shifts towards products and solubility increases
2. 16.2→ Precipitation and Qualitative Analysis
a. Ion Product (Q)
i.
b. Selective Precipitation
i.
ii.
Separating metal ions by adding something that reacts to make some of the metals form a
precipitate (different levels of solubility)
1. Use Ksp to find saturation point of each, then use one that causes the one with lower
solubility to precipitate but does not yet saturate the solution of the one with the higher
solubility
iii.
If one is >99% dissolved at saturation point of other cation, that is a success
c. Qualitative Analysis
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i.
ii.
We can use selective precipitation to qualitatively analyze a mixture of cations to determine what
metal ions are present
1.
Order you go in because don’t want to precipitate everything in first step (then it is not
separating solution)
2. Step 2: many cations form insoluble sulfides, but in acidic conditions, only the most
insoluble sulfides precipitate (because the concentration of S2- is very low)
iii.
3. 16.3→ Equilibria Involving Complex Ions
a. Complex ion is a charged species consisting of a metal ion surrounding by ligands
i.
Ligand is a Lewis base that donates an electron pair to a metal ion in solution
1. OH-, NH3, CN-
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ii.
Number of ligands is coordination number
iii.
Makes it more soluble
b. Use compounded equilibria (and multiply Ks) to solve
i.
Might have to use quadratic
c. Used to help separate post qualitative analysis
i.
d. Amphoteric metal hydroxide
i.
More soluble in acidic solution because lower pH means a lower concentration of OHii.
Some metal hydroxides form complex ions with hydroxide and will be more soluble in basic
solution
4. Terms
a. Solubility product constant: Ksp
b. Molar Solubility (s): moles dissolved/L soln
c. Selective Precipitation: a method of separating metal ions from an aqueous mixture by using a reagent
whose anion forms a precipitate with only one or a few of the ions in the mixture
d. Qualitative Analysis: the use of selective precipitation to qualitatively analyze a mixture of cations to
determine what metal ions are present
e. Complex Ion: a charged species consisting of a metal ion surrounding by ligands
f. Ligand: a Lewis base that donates an electron pair to a metal ion in solution
Chapter 17 (Spontaneity, Entropy, and Free Energy)
1.
17.1→ Spontaneous Processes and Entropy
a. Some reactions are spontaneous, others are not
i.
Something being spontaneous DOESN’T mean it is fast
b. Entropy (S)
i.
The driving force for a spontaneous process is an increase in the entropy of the universe
ii.
The universe tends towards an increase in entropy
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iii.
Entropy depends on the microstates available
1. A more disordered state has higher probability, because there are more microstates
that lead to it happening
iv.
State function
c. Positional entropy
i.
phase changes
1. ΔS (+ or -)
Ending state
Starting state
s
s
ii.
l
-
g
-
l
g
+
+
+
-
Solution formation
1. Positive entropy change because ions move about
freely→ more disorder
iii.
Solid w weak dipole
1. Entropy LOWER than expected because a weak dipole means one orientation is
preferred, # microstates decreases
2. 17.2→ Entropy and the Second Law of Thermodynamics
a. In any spontaneous process, there is always an increase in the entropy of the universe
i.
Entropy increases with the number of energetically equivalent ways to arrange the system
(microstates)
b. ΔSuniv=ΔSsys+ΔSsurr
3. 17.3→ The Effect of Temperature on Spontaneity
a. Entropy changes in the surroundings are primarily determined by heat flow
i.
Exothermic processes release energy in the form of heat, and thus increase the entropy of the
surroundings
1. More energy means more available microstates, more disorder
ii.
The impact of the transfer of a given quantity of energy as heat to or from the surroundings will be
greater at lower T
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b.
i.
The quantity of heat (at constant P) is qrev, which represents reversible heat
1. At constant pressure, q=ΔH
2.
a.
ΔH in this case represents ΔHsys
c. Molecular entropy
i.
S=kB lnW
1. W= number of microstates
2. kB=Boltzmann’s constant=1.3806 x 10-23 J/K
d.
4. 17.4→ Free Energy
a.
b.
c.
i.
If ΔG<0 reaction is spontaneous
ii.
NOTE: ΔH usually in kJ/mol, ΔS usually in J/mol, ΔG should be given in kJ
5. 17.5→ Entropy Changes in Chemical Reactions
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a. In general, when a reaction involves gaseous molecules, the change in positional entropy is dominated by
the relative numbers of molecules of gaseous reactants and products
i.
If # moles of gas increases, ΔS>0
ii.
If # moles of gas decreases, ΔS<0
b. Third Law of Thermodynamics
i.
The entropy of a perfect crystal at 0K is 0
c. Standard Entropy
i.
Normally calculated at 298 K
d.
6. 17.6→ Free Energy and Chemical Reactions
a.
ΔG° is the standard free energy change
i.
b.
the change in free energy that will occur if the reactants in their standard states are converted to
the products in their standard states
ΔG°f is the standard free energy of formation
i.
The change in free energy that accompanies the formation of one mole of a substance from its
constituent elements with all products and reactants in their standard states
ii.
0 for an element in its standard state
iii.
A lower ΔG°f means the compound is more stable
c. What makes a reaction favorable?
i.
More disorder→ increases entropy of system
ii.
Exothermic→ increases entropy of surroundings
d. ΔG°rxn
i.
7. Terms
a. Spontaneous: occurs without outside intervention
b. Entropy: a measure of molecular randomness or disorder
c. Microstate: a specific microscopic configuration of a thermodynamic system that the system may occupy
with a certain probability in the course of its thermal fluctuations
d. Positional probability: a type of probability that depends on the number of arrangements in space that yield
a particular state
e. Free energy: a thermodynamic quantity equivalent to the capacity of a system to do work
f. Standard Free Energy Change: the change in free energy that will occur if the reactants in their standard
states are converted to the products in their standard states
g. Standard Free Energy of Formation: The change in free energy that accompanies the formation of one mole
of a substance from its constituent elements with all products and reactants in their standard states
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h. Reversible Heat: a cyclical process, carried out by a hypothetical pathway, which leaves the universe
exactly the same as it was before. The system must be at equilibrium at every point for reversible heat to
work.
Chapter 18 (Electrochemistry)
18.1→ Balancing Oxidation Reduction Reactions
a. Half Reduction Method
i.
Write oxidation and reduction reactions
ii.
Balance
1. All elements except H and O
2. Balance O using H2O
3. Balance H using H+
4. Balance charge using electrons
iii.
Make number of free electrons in each reaction equal by multiplying by a constant (integer)
iv.
Add half reactions, cancel identical species
v. Check and make sure that charges are balanced
vi.
IF IN BASIC SOLUTION:
1. Add OH- to both sides, cancel H2Os created by this step if necessary
2. 18.2→ Galvanic Cells
a. Must separate the two half reactions in order to utilise the energy of the moving electrons
b. Need a salt bridge or porous disk in the middle, because if not a negative charge accumulates in the
cathode, while a positive charge accumulates in the anode. At this point, no more electrons will move over
to the cathode. A salt bridge fixes this issue by allowing ions into the cathode and anode to help neutralize
the charges of each. A porous disk serves the same purpose.
1.
c.
d. Cell Potential (Ɛ°cell)
i.
ii.
The driving force of the electrons in a galvanic cell
Can be measured using a voltmeter or potentiometer
1. Potentiometer is more accurate, because in voltmeter energy is lost to friction
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e. Units
i.
3. 18.3→ Standard Reduction Potentials
a. We need to standardize cell potential because we can measure the voltage of an entire cell, but not of the
half reactions separately
i.
Standardized with
2H++2e- → H2
b. Standard reduction potentials
i.
Reduction potential
ii.
1M
iii.
All gases at 1.00 atm
c. To find oxidation potential, do -reduction
d. Combining half reactions
i.
The half reaction with the largest positive potential will run as written (reduction), and the other
half reaction will run in reverse. The net potential of the cell will be the difference between the
two.
1.
ii.
Half reactions are NOT multiplied by integers to balance. Standard reduction potential is an
intensive property.
e. Line Notation
i.
Anode on left, cathode on right
ii.
When reactant is aqueous or gas, insert inert conductor as electrode (usually Pt)
f. Complete Description of a Galvanic Cell
i.
The cell potential is always positive for a galvanic cell
ii.
The direction of electrons flow is obtained using the half reactions and determining the cell
potential
iii.
Designate anode and cathode
iv.
Determine all present substances
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4. 18.4→ Cell Potential, Electrical Work, and Free Energy
a. Work and cell potential
i.
Maximum work is (charge)(E°cell)
1. Note charge=nF
ii.
In any real, spontaneous process some energy is always waster--the actual work realized is
always less than the calculated maximum
b. Faraday’s Constant (F)
i.
96,485 C/mol ec. Free energy and Cell Potential
i.
5. 18.5→ Dependence of Cell Potential on Concentration
a. Concentration Cells
i.
A cell in which both compartments have the same components but at different concentrations
b. The Nernst Equation
i.
ii.
1.
When Q=K, the system is at equilibrium, and Ecell=0. A dead battery is not actually dead,
it has just reached equilibrium
At room temperature:
1.
c. Ion Selective Electrodes
i.
Electrode that is sensitive to the concentration of a particular ion
ii.
Eg: pH meter
d. Calculation of Equilibrium Constants of Redox Reactions
i.
ii.
Use the fact that Ecell=0 to determine fomula at temperatures
other than 25℃
e.
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6. 18.6→ Batteries
a. A battery is a group of galvanic cells connected in series, where the potentials of the individual cells add to
give the total battery potential
7. 18.7→ Corrosion
a. The process of returning metals to their natural state--the ores from which they were originally obtained
b. Corrosion of Iron: rust
i.
8. 18.8→ Electrolysis
a. Electrolysis involved forcing a current through a cell to produce a chemical change for which the cell
potential is negative
b. Faraday’s Law of Electrolysis
i.
The mass of element deposited on an electrode in an electrolytic cell is proportional to the total
charge passed through the cell
c. Electrolysis of Water
i.
Add salt to allow current to flow
ii.
iii.
1.
In neutral water (low concentrations of hydronium and hydroxide), Ecell=-1.23 V
iv.
d. Electrolysis of Mixtures of Ions
i.
Can separate these mixtures of metal ions by plating out metals one at a time (one with highest
Ecell will plate out first
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9. Terms
a. Electrochemistry: the study of the interchange of chemical and electrical energy
b. Salt bridge: a U-tube containing an electrolyte that connects the two compartments of a galvanic cell,
allowing ion ow with-out extensive mixing of the different solutions
c. Porous Disk: a disk inside a tube that connects the two compartments of a galvanic cell, allowing ion ow
with-out extensive mixing of the different solutions
d. Galvanic Cell: a device in which chemical energy is changed to electrical energy
e. Anode: the electrode where the oxidation reaction occurs in a galvanic cell
f. Cathode: the electrode where the reduction reaction occurs in a galvanic cell
g. Cell Potential: the driving force in a galvanic cell that pulls electrons from the reducing agent in one
compartment to the oxidizing agent in the other
h. Potential Difference: the amount of work energy required to move an electric charge from one point to
another, given in V
i. Concentration Cell: A cell in which both compartments have the same components but at different
concentrations
j. Battery: a group of galvanic cells connected in series, where the potentials of the individual cells add to
give the total battery potential
k. Electrolysis: a process that involves forcing a current through a cell to cause a nonspontaneous chemical
reaction to occur
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