Chemistry Honors
Semester 1 Study Guide
Good luck studying guys. –Sandeep :)
Introduction to Chemistry: Ch1
Key Terms:
Chemistry
Substance
Model
Mass
Weight
Conclusion
Control
Dependent Variable
Experiment
Hypothesis
Independent Variable
Qualitative Data
Quantitative Data
Scientific Law
Scientific Method
Applied Research
the study of matter and the changes it undergoes
matter that has a definite and uniform composition, (also called a chemical)
visual/verbal/mathematical explanation of experimental data
measurement that reflects amount of matter
measure of matter and the effect of Earth’s gravitational pull on that matter
judgment based on information obtained
a standard for comparison
value that changes in response to change in independent variable (often y-axis)
set of controlled observations that test the hypothesis
tentative explanation for what has been observed
value that you intend to change (often x-axis)
information that describes a physical characteristic; relating to 5 senses
numerical description; tells how much, how long, how fast, etc.
relationship in nature that is supported by many experiments
organized approach to doing an experiment (see 1.3, Scientific method)
research undertaken to solve a specific problem
Pure Research
Theory
done to gain knowledge for the sake on knowledge itself
explanation of a natural phenomenon based on many observations and investigations over time
Notes
Ozone Layer
When O2, oxygen gas, is exposed to ultraviolet radiation, O3, ozone, is formed. In the past, Chlorofluorocarbons (CFC’s)
were used in refrigerators and air conditioners, as well as plastic foams and propellants in aerosol cans. Scientists
discovered that CFC’s had gone into the atmosphere and bonded
with the O2, preventing it from being formed into ozone and leading
to the lower levels of ozone we have in our atmosphere today.
1.3, Scientific Method
The scientific method is a systematic approach to experimentation.
First, make an observation: when ever I throw something, it comes
down. Next, ask a question: if I throw this ball up, it will come down?
Then, make a hypothesis, which would be what you think will
happen: if I throw this ball up, it will come down. Predict what will
happen: a ball will fall back down. Test experiment: it worked!
Finally, make more predictions and test them out: A rock will also
fall. However, if you were wrong, go back and revise your
hypothesis, then test out the experiment again.
Chemistry Honors Study Guide| Notes
2
Analyzing Data: Ch2
Key Terms:
Base unit
Density
defined unit based on object in physical world: second, meter, kilogram, mol, etc.
amount of mass per volume (formula below)
Derived Unit
Kilogram
unit that is a combination of base units: volume, density, speed, etc.
SI unit for mass
Kelvin
Liter
Meter
Second
Conversion Factor
Dimensional Analysis
Scientific Notation
Accuracy
Precision
SI unit for temperature, at 0 Kelvin, molecular motion stops (formula below)
measure of volume, 1 L= 1 dm3
SI unit for length
SI unit for time
ratio of equivalent values having different units.
systematic approach to problem solving that uses conversion factors.
used to express any number as a number between 1-10 times 10 to a power.
how close a measurement is to an accepted value.
how close a series of measurements are to each other
Error
Significant Figure
difference between experimental value and accepted value
include all known digits plus 1 more. (explained below)
Graph
visual display of data
Formulas:
Density =
Mass
----------------
Volume
Error = experimental value – accepted value
Error
Percent Error =
----------------------------
Accepted value
y2 - y 1
Slope =
----------------------------
x2 - x1
Chemistry Honors Study Guide| Notes
3
Analyzing Data: Ch2 –cont.
Notes
The SI (Systeme Internationale d’Unites) is a French system that updated the metric system. It is used between
scientists of different nations in most of the world.
A good guide to dimensional analysis can be found at
http://www4.ncsu.edu/unity/lockers/users/f/felder/public/kenny/papers/units.html.
Accuracy vs. Precision
Graphs
Pie charts show parts of a whole. Bar graphs show how a factor varies with time, location, or temperature.
Information about significant digits can be
found at the very end, in the chemistry
skills section.
Chemistry Honors Study Guide| Notes
4
Matter—Properties and Change: Ch3
Key Terms:
states of matter
solid
liquid
gas
different physical forms of all matter that exists on earth
definite shape and definite volume
flows, has constant volume, take shape of container
flows to conform to shape of container and fills entire volume
vapor
physical property
gaseous state of a substance that is sold or liquid at room temperature
prop. that can be observed/ measured without changing the sample’s composition
chemical property
extensive property
the ability of a substance to combine with or change into one or more other s
dependent on the amount of a substance; length, volume, mass
intensive property
independent of the amount; density, scent
physical change
alters a substance without changing its composition; cutting, crushing
chemical change
process that involves one or more substance changing in to new ones; chemical reactions
phase change
transition of mass from one state to another
law of conservation of mass
mass is neither created nor destroyed during a chemical reaction
mixture
combination of two or more pure substance, both retain individual chem. properties
heterogeneous mixture
mixture in which individual substances remain distinct; salad dressing, pulpy juice
homogeneous mixture
mixture that has constant composition throughout
solution
same as homogenous mixture
filtration
uses porous barriers to spate solids from liquids
distillation
uses different boiling points of substances to separate them
crystallization
results in the formation of solid crystals from a mixture; rock candy
sublimation
separating by changing solid substances directly into vapors
chromatography
separates components of mixture based on the ability of each to travel across a surface(marker lab)
element
pure substance that cannot be separated into simpler substances by physical or chemical means
periodic table
ment organizes elements into grid of horizontal rows (periods) and vertical columns (families)
compound
law of definite proportions
law of multiple proportions
two or more different elements that are combined chemically
compounds are always composed of the same elements in the same proportion by mass
when two compounds are made of the same elements, they will have whole number ratios
percent by mass
percentage
ratio of the mass of each element to the total mass of the compound expressed as a
Formulas:
Mass element
Percent by Mass =
----------------------------
X 100
Mass compound
Chemistry Honors Study Guide| Notes
5
Matter—Properties and Change: Ch3 –cont.
Notes
Law of Definite Proportions
The law stating that a pure substance, e.g. H2O, will
always have the same percent by weight, e.g. 11.2%
H and 88.8% O. In other words, oxygen will always
make up about 88% of any amount of water.
Law of Multiple Proportions
When two or more elements form more than one
compound, the ratio of the weights of one element
that combine with a given weight of another
element in the different compounds is a ratio of
small whole numbers. For example, carbon and
oxygen combine in carbon dioxide (CO2) and carbon
monoxide (CO). A sample of carbon dioxide
containing 1 gram of carbon contains 2.66 grams of oxygen; a sample of carbon monoxide containing 1 gram of carbon
contains 1.33 grams of oxygen. The ratio of the two weights of oxygen (2.66:1.33) is exactly 2:1. So therefore there are
twice as many oxygen atoms in a molecule of carbon dioxide compared to carbon monoxide.
The law of Multiple Proportions is definitely one that you would want to do practice problems in the book for.
Chemistry Honors Study Guide| Notes
6
The Structure of the Atom: Ch4
Key Terms:
Atom
Cathode Ray
Electron
Nucleus
smallest particle of an element that still retains its properties
Radiation that travels through a cathode ray; these were used in TVs.
negatively charged particles that travel around the nucleus
small, dense, positively charged center of an atom, contains protons and neutrons
Neutron
Proton
neutral particle in the nucleus that has a mass nearly equal to the proton
particle in the nucleus with a charge of 1+
Atomic Mass
Atomic Mass Unit
uses porous barriers to spate solids from liquids
uses different boiling points of substances to separate them
Atomic Number
Isotope
Mass Number
number of protons in an atom of an element
two atoms are isotopes if they have the same number of protons but different numbers of neutrons
sum of the number of protons and number of neutrons.\
Chapter 4 section 4 doesn’t seem like one that we studied. However, it may be a good idea to look it it.
Notes
Democritus- 400BC
Though that the world was made of tiny individual particles
Atoms are solid, indivisible, and indestructible
Different kinds of atoms have different sizes and shapes
It is these size and shapes that give atoms different properties
Aristotle- 350BC
Rejected Democritus. He denied the existence of atoms and was so influential that this denial went largely
unchallenged for the next 2000 years or so.
Matter is made of earth fire, air, water.
Dalton- 1800
Matter is composed of extremely small particles called atoms
Atoms are indivisible and indestructible
All atoms of a given element are identical in mass and properties
Compounds are formed by a combination of two or more different kinds of atoms
A chemical reaction is a rearrangement of atoms
J.J. Thompson
Discovered electron
Robert Millikan
Calculated charge of electron by making an atom of oil drop and measure the charge it took to control its
descent.
Chemistry Honors Study Guide| Notes
7
The Structure of the Atom: Ch4 –cont.
The Plum Pudding model was commonly accepted, until Rutherford came along and did his gold foil experiment. He shot
electrons at thin gold foil, expecting them to go through. However, some actually bounced back, which led him to
discover the nucleus, which he said contained positively charged protons.
Soon after, James Chadwick discovered the neutron.
The atomic number is the number of protons, which will usually also be the
number of electrons. Because the amount of neutrons in atoms of the
same element may vary, the atomic mass is actually a weighted (no pun
intended) average of the different weights.
We can also write this in formation in the format:
29
Writing an element as carbon-14 indicates that the atomic mass is
14. Knowing this, you can subtract to find that there are 8 neutrons
In this isotope.
The atomic mass unit is about the weight of a
Proton or neutron, but it was calculated using
A carbon-12 atoms.
In this format, 63 is the amount of
protons plus neutrons. 29 is the amount
of protons. This tells you that there are
29 protons, 29 electrons, and 34
neutrons.
If you’re confused, I did 63-29= 34
Chemistry Honors Study Guide| Notes
8
Electrons in Atoms: Ch 5
Key Terms:
Amplitude
Wavelength
Frequency
Atomic emission spectrum
wave’s height from origin to crest/trough
shortest distance between equivalent points on continuous wave; crest-crest/trough-trough
number of waves that pass a given point per second
set of frequencies of the electromagnetic waves emitted by atoms of that element
Electromagnetic radiation
Electromagnetic spectrum
form of energy that exhibits wavelike behavior as it travels through space
includes all forms of electromagnetic radiation
Photoelectric effect
Photon
the strange effect that not all light can eject electrons from a metal
massless particle that carries a quantum of energy
Planck’s constant
Quantum
Atomic orbital
Energy sublevel
Ground state
Principal energy level
Principal quantum number
shows that the energy of radiating increases as the radiation frequency increases
minimum amount of energy that can be gained or lost
3D region around the nucleus which describes the electron’s probable location
“slots” for pairs of electrons, as n increases, there are more sublevels Book pg 153
lowest allowable energy state of an atom is called its ground state
each major energy level of n
Book pg 153
indicates the relative size and energy of atomic orbitals
Book pg 153
Quantum mechanical model
Quantum number
Electron configuration
Electron-dot structure
Valence electron
model in which electrons are treated as waves
the value of n that specifies orbital
Book pg 153
arrangement of electrons in an atom
elements symbol surrounded by a number of dots equal to that elements valence electrons
electrons in the outermost orbitals
Formulas:
c = 3.00 x 108
h = 6.626 x 1034
λ
ν
(m/s)
(j·s)
(m)*
(Hz, 1/s)*
speed of light
Planck’s constant
has to be in meters for formulas, so
wavelength Wavelength
don’t forget to convert!
frequency Frequency is measured in Hertz. 1 Hertz is 1 wave per second.
c = λν
Ephoton = hν
Electromagnetic Wave Relationship
Energy of a Photon
λ=
Particle Electromagnetic-Wage Relationship* m is the mass of the particle.
h
mν
Chemistry Honors Study Guide| Notes
9
Electrons in Atoms: Ch 5
Notes
The black line is called the origin. In real life, this
wave is moving. The amount of times each crest
(called peak in this picture) passes you, that is
one Hz.
Make sure to understand that high energy waves have high
frequencies (they are fast) and short wavelengths. Low energy
waves have low wavelengths and low frequencies (they are
slow)
De Broglie equation
All moving particles have wave characteristics.
Heisenberg Uncertainty Principle
Its impossible to know both the location and speed of a particle. (the more you know about one, the less you know
about the other)
Hund’s Rule
Single electrons with the same spin must occupy each equal energy orbital before additional electrons with opposite
spins can occupy the same orbitals.
Aufbau Principle
Each electron occupies the lowest energy orbital available.
Pauli Exclusion Principle
A maximum of two electrons can occupy a single atomic orbital.
Hey guys, I didn’t cover a couple things here because the book does a
great job already. These topics are: Electron Configurations, Orbital
Diagrams, and Dot Diagrams.
Chemistry Honors Study Guide| /Notes
10
The Periodic Table and Periodic Law: Ch 6
Key Terms:
Periodic law
there is a periodic repetition of chemical and physical properties of elements
Group
columns on the period table, also called families. These share traits.
Period
rows on the periodic table
Representative Element
elements in groups 1,2,13,18
Transition element
Metal
elements in groups 3-12
shiny, solid at room temperature, ductile, malleable
Alkali metal
Alkaline earth metal
group 1, extremely reactive
group 2, highly reactive
Transition metal
an element in groups 3-12
Inner transition metal
t he lanthanide and actinide series that is set off usually at the bottom of the graph
Lanthanide series
f block elements from period 6 (starting with lanthanum)
Actinide Series
f block elements from period 7 (starting with actinium)
Nonmetal
elements that are generally gasses or brittle, dull-looking solids
Halogen
group 17, very reactive
Noble gas
group 18, extremely unreactive
Metalloid
Electronegativity
Ion
Ionization energy
Octet rule
elements that have physical and chemical properties of both metals and nonmetals
ability of an element to attract electrons to it.
atom or bonded group of atoms that has a positive or negative charge
energy required to remove electron from gaseous atom
atoms tend to gain, lose, or share electrons in order to acquire a full set of 8 valence electrons
Notes
Lavoiser, Newlands, Meyer and Mendeleev all worked to create the periodic table. Mendeleev is generally given
the most credit for organizing the elements into a series of rows of elements with similar characteristics. Mosley
was able to then add more and slightly reorganize the table.
Chemistry Honors Study Guide| Notes
11
The Periodic Table and Periodic Law: Ch 6
Notes
Periodic Trends
As you go from top to bottom on the
periodic table, obviously your element will
be bigger. However, when you go from left
to right, since the electrons in that row are
in the same energy level regardless, the
more protons in the atom, the harder they
pull the electrons in close.
Also, when comparing ions, if the amount of
electrons are equal, then whichever has less
protons will be larger. If the amount of
protons are the same, then the one with
more electrons is bigger.
The easiest way for me to remember this
one is that Fluorine is the most
electronegative element. As you get further
away from it, you get less electronegative.
Chemistry Honors Study Guide| Notes
12
Ionic Compounds and Metals: Ch 7
Key Terms:
Anion
Cation
negatively charged ion
positively charged ion (Why is Willy the wildcat a cation...Because he’s pawsative :)
Chemical Bond
Crystal Lattice
force that holds two atoms together
3D arrangement of particles in which each positive ion is surrounded by negative ions
Electrolyte
Ionic bond
ionic compound whose aqueous solution conducts an electric current
bond in which one atom “takes” the others electron, often between a metal and nonmetal; NaCl
Ionic compound
Lattice energy
compound with ionic bond
energy required to separate 1 mol of the ions of an ionic compound
Formula unit
chemical formula for an ionic compound (similar to a molecule)
Monatomic ion
one atom ion; Mg2+, Br
Oxidation number
Oxyanion
Polyatiomic ion
Alloy
Delocalized electron
Electron sea model
Metallic bond
charge of a monatomic ion
polyatomic ion composed of an element, usually a nonmetal, bonded to one/more oxygen atoms
ions made up of more than one atom
mixture of elements that has metallic properties
small, dense, positively charged center of an atom, contains protons and neutrons
all the metal atoms in a metallic solid lend their valence electrons to form a “sea” of electrons
attraction of a metallic a metallic cation for delocalized electrons
-
Notes
To the left is a picture of the electron sea
model.
Naming Ionic Compounds has been placed
at the “Chemistry Skills” section at the end
of this document.
Chemistry Honors Study Guide| /Notes
13
Covalent Bonding: Ch 8
Key Terms:
Covalent bond
Molecule
Lewis Structure
Sigma Bond
bond in which atoms share valence electrons
formed when two or more atoms bond covalently
Take a dot diagram, replace bonds with lines (page 242-244)
single covalent bond (σ)
Pi Bond
double covalent bond (π)
Endothermic reaction
requires energy, generally becoming cold
Exothermic reaction
Oxyacid
releases energy, generally in the form of heat
an acid that contains both a hydrogen atom and an oxyanion
Structural formula
similar to Lewis structure, but without dots. Does not show lone pairs.
Resonance
condition that occurs when there is more than one way to draw the Lewis structure
Coordinate covalent bond
forms when one atom donates both the electrons to be shared with an atom or ion that
needs two electrons to form a stable electron arrangement.
VSEPR model
Valence Shell Electron Pair Repulsion, used to determine 3D shape of a molecule
Resonance
a process that carbon atoms undergo in which atomic orbitals mix and form new shapes
Polar Covalent Bond
polar covalent bond
Notes
Covalently bonded atoms.
Sigma bonds are single bonds, and are direct
between the two carbon atoms. However, when
atoms form double or even triple bonds, they have
to form them where there is not already one
present. So the pi bonds, which signify double/triple
bonds go around.
Chemistry Honors Study Guide| Notes
14
Covalent Bonding: Ch 8
Notes
All the three figures at the top are valid. To show this, we
use dotted lines as it shows on the bottom. However, you
can ignore the numbers (1-,2/3-); they are unnecessary for
the purposes of Chem Honors.
Above is a little chart the differences between the Lewis Structure in 2D and the 3D representation. I believe we actually
need to know the differences between the different Tetrahedral figures, but don’t quote me on that.
Electronegativity Difference
Bond Character
>1.7
Mostly ionic
0.4-1.7
Polar covalent
<0.4
Mostly covalent
0
Nonpolar covalent
This chart here is how you determine how to categorize the atoms. The teachers don’t expect you to memorize the
values, but think about it like this: the closer to each other two elements are, the more they will share electrons. This is
why metals and nonmetals form mostly ionic bonds with each other- they don’t share very well. Elements like oxygen
will form pretty much perfect nonpolar covalent bonds with themselves.
Chemistry Honors Study Guide| Notes
15
Covalent Bonding: Ch 9
Key Terms:
Chemical reaction
Reactants
Products
Chemical equation
the process by which the atoms of one or more substances are rearranged into new ones
starting substances
substances formed
statement that uses chemical formation to show identities and amounts of substances
Coefficient
Synthesis
number in front of product or reactant
a chemical reaction in which two or more substance react to produce a single product
Combustion
Decomposition
releases energy, generally in the form of heat
an acid that contains both a hydrogen atom and an oxyanion
Single-replacement
similar to Lewis structure, but without dots. Does not show lone pairs.
Double-replacement condition that occurs when
Precipitate
solid produced during chemical reaction
Aqueous solution
one or more solutes dissolved in solution; saltwater
Solute
substance dissolved into the solvent; salt
Solvent
substance solute is dissolved into; generally water
Complete ionic equation
equation that shows all the particles in a solution as ions
Spectator ion
Net ionic equation
ions that do not participate in a reaction
ionic equations that include only particles that participate in reaction
Notes
I’m assuming you guys know how to balance a chemical equation…but just in case, check page 287 in your
book for a step by step guide.
Types of Reactions
Synthesis
Combustion
Decomposition
Single-replacement
Double-replacement
Chemistry Honors Study Guide| Notes
16
Total/Net Ionic equations
Example Expirement
Molecular Equation: CaCO3(s) + 2HCl(aq)
CaCl2(aq) + H2O(l) + CO2(g)
Total Ionic Equation: CaCO3(s) + 2H+(aq) + 2Cl-(aq)
Net Ionic Equation: CaCO3(s) + 2H+(aq)
Ca2+(aq) + 2Cl-(aq) + H2O(l) + CO2(g)
Ca2+(aq) + H2O(l) + CO2(g)
If a reactant or product is aqueous, split it up in your total ionic equation. If not aqueous (if its solid, gas, liquid) then
leave them as is. If there is a precipitate (a solid forms on the right side that wasn’t there on the left) then generally your
net Ionic equation will show the formation of that solid. Ignore the previous sentence if it confuses you.
Solubillity Rules
1. All common compounds of Group I and ammonium ions are soluble.
2. All nitrates, acetates, and chlorates are soluble.
3. All binary compounds of the halogens (other than F) with metals are soluble, except those of Ag,
Hg(I), and Pb. Pb halides are soluble in hot water.)
4. All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I).
The latter three are slightly soluble.
5. Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.
6. Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium,
and ammonium.
Here are some basic rules about solubility. Remember to check with the ones that your teacher gave you!
Chemistry Honors Study Guide| Total/Net Ionic equations
17
The Mole: Ch 10
Key Terms:
Mole
SI base unit to measure amount of substance. A mol of something is 6.02x10 23 of that thing.
Avogadro’s Number
6.02x1023 (originally calculated as the amount of carbon atoms in exactly 12 g of carbon-12)
Molar mass
Percent composition
the mass in grams of one mole of any pure substance; the molar mass of water is 18g
percent by mass of each element in a compound is the percent composition
Empirical Formula
Molecular Formula
formula with the smallest whole-number ratio of the elements
specifies the number of atoms of each element in one molecule of that substance
Hydrate
compound that has a specific number of water molecules bound to its atoms
Notes
If you have trouble understanding mols, think about it like this: if I had a mol of cars, I would have 6.02x1023
cars. If there was a mol of pages in this study guide, there would be 6.02x1023 pages.
If I had a mol of carbon, I would have 6.02x1023 atoms. If I had a mol of H2O, I would have 6.02x1023
molecules of water, 6.02x1023 atoms of oxygen. As for hydrogen, I actually have 2x6.02x1023, or 1.2x1024
atoms of hydrogen. Here’s why: for every 1 molecule of oxygen, you have 2 atoms of
hydrogen and 1 atom of oxygen. Hence the doubling.
To convert from grams to mols, divide the grams by the molar mass
To convert from mols to atoms/molecules, multiply the mols by 6.02x1023
To convert from mols to grams, multiply the grams by the molar mass
To convert from atoms/molecules to mols, divide the atoms by 6.02x1023
For percent composition and hydrates, my explanation
would not be any shorter than the book’s. Read pgs.342353. Sounds like a lot, but most of it is math. However, it is
important math :)
Chemistry Honors Study Guide| Notes
18
Stoichiometry: Ch 11
Key Terms:
Stoichiometry
the study of quantitative relationships between the amounts of reactants and products
Mole Ratio
Limiting reactant
ratio between the numbers of moles of any two of the substances in a balanced chemical equation
limits amount of product created
Excess reactant
Theoretical yield
Actual yield
reactants that are not completely used up in the reaction
maximum amount of product that can be produced from a given amount of reactant
amount of product poruduced when chemical reaction is carried out
Percent yield
ration of the actual yield to the theoretical yield
Notes
Stochiometry
The thing about stoichiometry is that it is just a series of grams to mole conversions. However, its easy
to get confused by the steps.
Lets take the above equation. See how you need 2 groups of 2 hydrogen atoms and 1 group of 2 oxygen
atoms? That leads us to a total of 4 atoms of Hydrogen and 2 Oxygen. This is a 2:1 ratio. So if I had 1 mol
of O2, I would need 2 mol of H2. The coefficients in front of the letters is just a ratio of mols. So if I
had 2.5 mol O2, how many mol H2O would I make? 2x2.5= 5! I would make 5 mol H2O.
o
A little trick for remembering how to use stoichiometric ratios goes as follows:





1) Make sure you have your reactants/product in moles
2) You would like to convert mol of one substance, lets call it A, to another, called B.
2) Take your moles of substance A, and divide it by the coefficient of substance A in the
equation.
3) Then, multiply this new number by the coefficient of the substance B.
4)The number you get is your mols of substance B
Example problems for this can be found on pages 375-377
Chemistry Honors Study Guide| Notes
19
Limiting Reactants
This is easy to understand if you think about it in other terms first.
Each car takes exactly 1 car body and 4 tires to make. In this example, we have 48 tires. 48/4=12. We have enough tires
for 12 cars. We also have 8 car bodies so we have enough to make 8 cars. Therefore, we can only make 8 cars, no matter
how many tires we have.
Lets take this problem: A 2.00 g sample of ammonia is mixed with 4.00 g of oxygen. Which is the limiting
reactant and how much excess reactant remains after the reaction has stopped?
First, write the chemical equation:
4 NH3(g) + 5 O2(g)
4 NO(g) + 6 H2O(g)
Then, plug in the two values to see which substance makes the most of the product. Because the question did not
specify, we can just pick a product; lets use NO.
Above is the conversion the way that the book would do it. They calculated the mols of NO created twice, using the 2 g
of ammonia the first time and the 4 grams of oxygen the second, then converted it in grams. We have enough ammonia
for 3.53 g NO. We have enough oxygen for 3 g NO. Therefore, at most we can make 3 g NO. Oxygen is the limiting
reactant.
For the second part of the problem, we have to do some sneaky math. We know that oxygen is the limiting reactant; so
why not just convert mol of oxygen into grams of ammonia? Doing this conversion (divide by molar mass O2, divide by 5,
multiply by 4, multiply by molar mass NH3) gives you the amount of NH3 that was used. So if 2 g were given to us, and we
used 1.7, how much is left? .3 g NH3
Theoretical/Actual Yield
Its important to remember that when you do experiments in real life, results aren’t always perfect. The theoretical yield
is the amount that the math says you’ll get; the actual yield is what happens in a lab setting. Your percent yield is
therefore a percentage of how well your lab compared to a perfect one. For example, I plugged in my 4 grams of Oxygen
into my chemical formula (this is an imaginary formula, don’t look for it) and found my theoretical yield to be 10 grams
of Magnesium. However, when I did the lab, I only got 6 grams of magnesium. 6/10=.6….my percent yield was 60%.
Chemistry Honors Study Guide| Notes
20
States of Matter: Ch 12
Key Terms:
Elastic collision
on in which no energy is lost; think if it like pool balls as opposed to playdough
Temperature
Diffusion
measure of the average kinetic energy
describes movement of one material through another
Effusion
Pressure
Barometer
describes how a gas escapes through a small hole
force per unit area; kPa, atm, mmHg, torr
instrument that measures atmospheric pressure
Pascal
Atmosphere
Dispersion force
Dipole-dipole force
force of one newton per square meter, unit for pressure
average air pressure at 0 Celsius and sea level, unit for pressure
weak forces that result from temporary shifts in electron density
attractions between oppositely charged regions of polar molecules
Hydrogen bond
Viscosity
Surface tension
Surfactant
dipole-dipole bond between hydrogen and N, O, or F. stronger than regular dipole-dipole
measure of a resistance to flow (the difference between pouring water and syrup)
energy required to increase the surface area by a given amount
compounds that lower the surface tension of water
Crystalline solid
Unit cell
Allotrope
Amorphous solid
Melting point
solid whose atoms, ions, or molecules are arranged in an orderly geometric structure
smallest arrangement of atoms in a crystal lattice that has the same structure as the whole
an element that exists in different forms at the same state
one in which the particles are not arranged in a regular, repeating pattern
temperature at which the forces holding a crystal lattice together break
Vaporization
Evaporation
process of a liquid changing to a gas
vaporization that occurs only at the surface of a liquid
Vapor pressure
pressure exerted above the liquid
Boiling point
temperature at which the vapor pressure of a liquid equals the atmospheric pressure
Freezing point
temperature at which a liquid is converted into a crystalline solid
Condensation
process by which a gas becomes a liquid
Deposition
process by which a substance changes from a gas to a sold without being a liquid
Phase diagram
graph of pressure
versus temperature, shown
to the right
Triple point
point on a phase
diagram at which all three
phases can coexist
Critical point
point above which
a substance cannot exist as
a liquid
Chemistry Honors Study Guide| Notes
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Stoichiometry: Ch 12- cont.
Notes
Kinetic-molecular theory
1. Gases consist of large numbers of molecules (or atoms, in the case of the noble gases) that are in
continuous, random motion
2. The volume of all the molecules of the gas is negligible compared to the total volume in which the gas is
contained
3. Attractive and repulsive forces between gas molecules is negligible
4. The average kinetic energy of the molecules does not change with time (as long as the temperature of
the gas remains constant). Energy can be transferred between molecules during collisions (but the
collisions are perfectly elastic)
5. The average kinetic energy of the molecules is proportional to absolute temperature. At any given
temperature, the molecules of all gases have the same average kinetic energy. In other words, if I have
two gas samples, both at the same temperature, then the average kinetic energy for the collection of gas
molecules in one sample is equal to the average kinetic energy for the collection of gas molecules in the
other sample.
Graham’s law of effusion
The two M’s are the molar masses of the two gasses.
Dalton’s law of partial pressures
PressureTotal = Pressure1 + Pressure2 ... Pressuren
Dalton’s law of partial pressures basically states that if you know the total pressure, then you can subtract the
different parts. For example, the total pressure on my head from a combination of nitrogen and oxygen is 1.5
atm. The nitrogen has a pressure of 1.25 atm. What is the pressure of Oxygen? 1.5-1.25= .25
Unit
Number Equivalent to 1 atm
kPa
101.3
Atm
1
mmHg
760
Torr
760
Psi
14.7
Bar
1.01
(mm Hg is the same thing as Torr. We don’t use Psi or Bar very often.)
Dispersion forces are weaker than dipole-dipole forces which are weaker than hydrogen bonds. If you struggle
with understanding what dispersion forces are, remember that the electrons in an atom are moving around randomly.
Dispersion forces are the attractions that occur when it just so happens that the electrons in molecules are more
numerous on one side than another.
Only crystalline solids have a melting point. Amorphous solids actually have more of a range of temperatures
where they melt. Also,
Chemistry Honors Study Guide| Notes
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Gases: Ch 13
Key Terms:
Absolute zero
Molar volume
Avogadro’s principle
particles
zero on the Kelvin scale, at which temperature all molecular motion stops
volume that 1 mole of a gas occupies at STP, 22.4 L
equal volumes of gas at the same temperature and pressure contain equal numbers of
Formulas:
Boyles Law
Charles Law
The ideal gas constant is represented by
the symbol R. Depending on the unit of
pressure, there are different values of R.
Gay-Lussac’s law
atm
.0821
kPa
8.314
Combined Gas Law
mmHg 62.4
Personally, I just remember the R value
for atm and just convert. But ,since we
get to use a cheat sheet, this would
probably be a good idea to write down.
Ideal gas law
STP is 0 Celsius, 1 atm
Notes:
Also, don’t forget to use Kelvin instead
of Celsius of Fahrenheit when
performing these calculations.
o
As the temperature goes up, the pressure goes up
o
As the temperature goes down, the pressure goes down
temperature and pressure are………… directly proportional
o
As the temperature goes up, the volume goes up
pressure and volume are…………………..inversely proportional
o
As the temperature goes down, the volume goes down
o
As the volume goes down, the pressure goes up
o
As the volume goes up, the pressure goes down
volume and temperature are…………….directly proportional
Chemistry Honors Study Guide| Notes
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Gases: Ch 13
Notes:
Most gasses are more or less ideal at most conditions. However, they are furthest from being ideal gasses at
high pressures and low temperatures.
Polar gasses (water vapor) do not behave ideally
Gas Stoichiometry
Gas stoichiometry is not that much different than regular stoichiometry. The main difference here is that you
often have to use conversions. The key thing that keeps you from gas stoichiometry being just regular stoichiometry is
the fact that you often are given gas as a volume. Use the gas laws to find out moles, then its back to basics.
This is a shortcut, but you have to make sure that the problems you are given tell you to assume that pressure
and temperature remain constant. This means that you can actually just take the volume in L that you were given and
treat them like mols, without converting. Example on page 461.
I’ll admit, Gas stoichiometry isn’t that easy. It is highly advisable to check out the
information on pages 461 to 464 in your book.
Chemistry Honors Study Guide| Notes
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Chemistry Honors Study Guide| Notes
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Significant Digits
Digits from 1-9 are always significant.
Zeros between two other significant digits are always significant
One or more additional zeros to the right of both the decimal place and another significant digit are significant.
Zeros used solely for spacing the decimal point (placeholders) are not significant.
Examples of Significant Digits
EXAMPLES
# OF SIG. DIG.
453 kg
3
5057 L
4
5.00
3
0.007
1
COMMENT
All non-zero digits are always
significant.
Zeros between 2 significant
digits are significant.
Additional zeros to the right of
decimal and a significant
digits are significant.
Placeholders are not
significant.
Adding and Subtracting
RULE: When adding or subtracting your answer can only show as many decimal places as the
measurement having the fewest number of decimal places.
Exampled: When we add 3.76 g + 14.83 g + 2.1 g = 20.69 g
We look to the original problem to see the number of decimal places shown in each of the original
measurements. 2.1 shows the least number of decimal places. We must round our answer, 20.69, to one decimal
place (the tenth place). Our final answer is 20.7 g
Chemistry Honors Study Guide| Notes
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Significant Digits –cont.
Multiplying and Dividing
RULE: When multiplying or dividing, your answer may only show as many significant digits as the
multiplied or divided measurement showing the least number of significant digits.
Example: When multiplying 22.37 cm x 3.10 cm x 85.75 cm = 5946.50525 cm3
We look to the original problem and check the number of significant digits in each of the original
measurements:
22.37 shows 4 significant digits.
3.10 shows 3 significant digits.
85.75 shows 4 significant digits.
Our answer can only show 3 significant digits because that is the least number of significant digits in the
original problem.
5946.50525 shows 9 significant digits, we must round to the tens place in order to show only 3 significant
digits. Our final answer becomes 5950 cm3.
Chemistry Honors Study Guide| Notes
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Naming Chemical Compounds
Binary Molecular Compounds
Binary molecular compounds are composed of two types of nonmetals. The nonmetals are normally ordered
with the element leftmost on the periodic table first. If both elements are in the same column, then the element
lower on the periodic table is first. The order is the same for the formula and the name.
Because the nonmetals can combine in many ways, prefixes are used to express how many atoms of each
element are in the atom. Memorize the following list of prefixes.
1 mono*
2 di
3 tri
4 tetra
5
penta
6 hexa
7 hepta
8 octa
9 nona
10 deca
The prefix for one is starred because its use is optional for the second element. If there is only one atom of the
first element, no prefix is used.
Examples:
Carbon Monoxide
Oxygen Difluoride
Tetranitrogen Decaoxide
Arsenic Tribromide
Boron Trichloride
Iodine Heptafluoride
CO
OF2
N4O10
AsBr3
BCl3
IF7
Binary Molecular Compounds
Binary ionic compounds may contain more that two elements but are binary because they contain two ions.
First, take your cation (it will usually be a metal from the left). Make no changes; this will be the first part.
Then, take your negative part (usually an element from the right side) add it on to the end. Drop the ending, add
an –ide. If your working with a polyatomic ion, just stick it in as is. They are already “conjugated.”
Remember, Alkali metals always have a charge of +1. Alkali earth metals always have a +2 charge. In addition
to these two groups, aluminum is always Al3+, zinc is Zn2+, and silver is Ag+.
Examples:
sodium carbonate
cobalt(II) nitrate
tin(IV) sulfide
Na2CO3
Co(NO3)2
SnS2
Chemistry Honors Study Guide| Notes
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Binary Molecular Compounds
Since acids are substances that release H+ in water, it is traditional to write the hydrogen atom first in the
formula. The names of these acids are based on the anion the acid came from. (Hydrogen acts as a cation, H+.
Although acids are molecular compounds, they react with water to form ions.) If the anion has an ate ending,
the ate is changed to ic and the word acid added. If the anion has an ite ending, the ite is changed to ous and is
followed by acid.
Examples:
hydroiodic acid
carbonic acid
sulfurous acid
perchloric acid
HI
H2CO3
H2SO3
HClO4
(that’s H with a capital i after it)
Chemistry Honors Study Guide| Notes
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