Honors Final Review
Note: this is in no way guaranteed to be comprehensive. Anything from the year is fair game, but the emphasis
will be on more recent material. Bolded items are most important.
Chapter 1: Introductory stuff.
 Qualitative vs quantitative measurements.
 Scientific method
Chapter 2: Atoms, elements, and mixtures.
 Matter is made of up atoms
 Elements, compounds, and molecules: elements are only one type of atom, compounds have
multiple elements, and molecules are atoms that are stuck together.
 States of matter—solid, liquid, and gas—and their properties.
 Physical versus chemical changes.
 Mixtures: homogeneous and heterogeneous.
Chapter 3: The elements and ions.
Dalton’s atomic theory: elements are made of atoms, all atoms of a given element are identical, atoms of
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different elements are different, and reactions don’t alter the atoms—only rearrange them.
Structure of the atom: protons and neutrons in the nucleus, electrons around the outside.
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J.J. Thomson discovered electrons.
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Rutherford shot particles at gold foil and saw some scatter, some go through, meaning that part of the
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atom was hard while most was empty space.
Element symbols: atomic number and atomic mass
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Atomic number = number of protons, defines an element.
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Atomic mass = # protons + # neutrons
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Isotopes have same number of protons, different number of neutrons (same element, but weigh
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different amounts).
Number of electrons changes, leading to charges. Things with a charge are called ions.
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Positively charged atoms/groups of atoms are cations. Negatively charged ones are anions.
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Periodic table: alkali metals, alkaline earth metals, transition metals, metalloids, nonmetals, halogens,
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and noble gases.
Certain groups of elements always form ions with the same charge. See page 74.
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Compounds containing ions conduct electricity if the ions can move.
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Writing formulas for ionic compounds: the charges have to balance out to neutral. Use the
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crossing method with the charges, but reduce if there’s a common factor. For example:
Mg2+ and N3- for Mg3N2
but
Mg2+ and O2- form MgO
Chapter 4: Naming compounds.
Ionic compounds have a metal and at least one non-metal, and are stuck together by charges.
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Covalent compounds have only non-metals, and are held together by bonds.
Anions of a single element use the start of the element name, and end in -ide. Nitrogen with a negative
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charge becomes nitride. Oxygen with a negative charge is oxide.
Ions can be a group of elements all stuck together by covalent bonds, that overall have a charge.
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If the group has an element attached to some number of oxygens, then it ends in -ate, and starts with the
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beginning of the name of the non-oxygen element. For example, NO3- is nitrate, PO43- is phosphate.
You can have a different number of oxygens, which have different naming schemes. The one that ends
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in –ate is the most common one.
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There are eight common complex ions to know: ammonium, carbonate, nitrate, chlorate, sufate,
phosphate, permanganate, and hydroxide.
To name compounds, there are two questions to ask:
1. Is it ionic (is there a metal or NH4+ group)?
No: If there are two elements that are both non-metals, then the name is the name of the first
element, followed by the name of the second element with the –ide ending. Then before each
element, you add a prefix indicating the number of that element (mono, di, tri, tetra, penta, hexa,
hepta, octa). You drop mono if it’s the first element, and you can drop the last “a” from the
prefix if it sounds funny
Example:
P2O5 is diphosphorous pentoxide.
CO2 is carbon dioxide
Yes: If there is a metal (or an NH4+) , then the name is the name of the metal (or “ammonium”)
followed by the name of the anion. Then you ask the second question.
2. Is the metal one that always forms the same charge? If yes, then you’re done. If no, then you
add the charge of the metal in parentheses after the metal’s name.
Example:
NaClO3 is sodium chlorate Fe2O3 is iron(III) oxide
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From the name, you can work back to the formula.
Chapter 5: Measurements and calculations
Measurements have limits in their precision, depending on what’s used to measure them.
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To determine how precise a measurement is, we count the number of significant figures.
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Non-zeros always count.
Zeros at the beginning of a number never count.
Zeros in between other numbers always count
Zeros at the end of the number count if there’s a decimal point.
IMPORTANT: nothing in these rules asks whether the number is before or after a decimal point!
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Conversion factors (1000 mL in 1L) and counted numbers (30 students in class today) are
infinitely precise: sig fig rules don't apply to them.
When doing calculations with significant figures, you use these rules to determine where to round off:
Multiplication and division: the prodcut/quotient should be rounded to the same number of
significant figures as the original number with the fewest.
Addition and subtraction: working left to right in the digits, keep significant figures until you
reach a point where one of the numbers doesn’t have a sig fig.
Unit conversions/dimensional analysis:
To convert units, multiply by a fraction arranged such that the old unit cancels, and the new unit
remains. If you have a unit that’s squared or cubed, you have to square or cube the conversion if
your normal conversion is not squared/cubed.
For example: 1002 cm2 = 12 m2
1m3 = 1000 L ← don't cube this one; you already have the conversion
Temperature conversions follow a different formula:
o
F = (9/5).oC + 32
K = oC + 273.15
Chapter 6: Moles!
We can use weight to count the number of atoms, if we know how much each weighs.
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A mole is defined as 6.022x1023 items (usually atoms or molecules), aka Avogadro’s number.
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1 mole of protons or neutrons weighs 1g, so 1 mole of an element weighs its atomic mass in grams.
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The atomic mass on the periodic table is an average of all the isotopes of that element, weighed by how
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common they are.
To find the weight of 1 mole of a compound, you add the atomic masses, multiplied by the number
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of each type of atom. This is the molecular mass/molecular weight.
The molecular mass is a conversion factor that allows conversion between grams and moles.
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Empirical formula: given the percent of each element in a compound, the formula can be found by
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assuming that you have 100 g total, converting the grams of each element to moles, and dividing
all of them by the smallest one. If there are still decimals, multiply everything by a number that
makes them all into whole numbers.
Chapter 7: Reactions!
Reactants go on the left of the arrow, products on the right.
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Coefficients are added to make the number of each element match up on both sides (balancing).
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The coefficients describe the ratios in which the starting materials react, and the amount of
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products they form.
The state of the compound is added after it in the equation, in parenthesis. In addition to solid (s), liquid
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(l), and gas (g), aqueous—dissolved in water—is also indicated (aq).
The major types of reaction are: precipitation (two aqueous make a solid), acid-base (and H+ is
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transferred from one thing onto another), and oxidation-reduction aka redox (an element changes
charge).
Redox has three special types: synthesis (two or more reactants make 1 compound), decomposition
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(1 compound makes two or more products), and combustion (something reacts with oxygen).
For combustion reactions (reacting with O2, i.e. burning), the products are determined by the
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elements in the starting material. Anything with carbon yields CO2, hydrogen yields H2O, sulfur
yields SO2, and nitrogen leads to NO2
Chapter 9: Yields
Because the reaction equation is a ratio, if we know how much starting material we put in, then we can
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multiply by the ratio from the reaction equation to determine the amount of product that comes out.
This calculation must be done in moles, which can then be converted back to grams.
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If two things are reacting together, the amount of product will be determined entirely by which
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one runs out first—the limiting reagent.
Determination of the limiting reagent must be done in moles, and must account for the coefficients
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in the reaction equation (divide each reactant's moles by its coefficient, then compare).
The yield calculated from the reaction equation is the theoretical yield (take the number from the
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limiting reagent in the last step, and multiply by the product's coefficient).
The actual yield divided by the theoretical yield, times 100, is the percent yield.
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Chapters 8/15: Solutions and Solubility
 Concentration of a solution is the amount of stuff dissolved divided by how much stuff it's dissolved in.
 We'll use molarity (M) as a measure of concentration: Molarity = moles of solute / L of solution
 When making simple dilutions, we can use C1V1 = C2V2.
 In more complicated cases, such as mixing two solutions, we have to find all the moles from every
source, add them, and divide by the final volume.
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Molarity can be used to find moles; this can be used with a reaction equation to find limiting reagents
and theoretical yields.
If a compound dissolves, it's soluble. If not, it's insoluble.
A precipitate is a solid that forms when liquids are mixed. It happens because an insoluble
compound is formed.
To predict when a precipitate forms, use the solubility rules (NASH CHOPS).
By combining an unknown solution with known solutions and observing whether a precipitate forms,
the identity of the unknown can be determined.
Solubility Rules:
Soluble Compounds:
Nitrates, Alkali metals (and NH4+), Sulfates (EXCEPT with Ba2+,Pb2+,Ca2+), Halides (EXCEPT with Ag+, Pb2+)
Insoluble Compounds:
Carbonates, Hydroxides (EXCEPT with Ca2+), Oxides, Phosphates, Sulfides
If there is a conflict (Na2CO3, for example has an “always soluble” and a “never soluble” paired), then the
“soluble” one wins—the compound is soluble.
Chapter 10: Energy and Thermodynamics
* Energy comes in many forms, but for chemical reactions, mostly in the form of heat.
* Heat and temperature are not the same thing.
* Temperature is the average energy of all the molecules in a substance.
* Heat is the total energy.
* We can't measure amount of heat something has, but we can measure the flow of heat—how much was
transferred to something else.
* Total heat absorbed or released is always given the symbol “q”.
* The units of energy, and therefore heat, are Joules.
* Heat given off is given a negative sign, and referred to as “exothermic”.
* Heat taken in is given a positive sign, and referred to as “endothermic”.
For objects:
q = m*c*T
where T = Tf – Ti
and c the heat capacity, which measures the ease with which a material accepts or gives off heat.
c has one value for each material. For liquid water, it is 4.184 J/g*oC
Special object:
q = Ccal*T
Calorimeters—the containers in which you measure the flow of heat—have m and c combined together into one
term—Ccal. This is just used because m and s won't change anyway and might be inconvenient to find
separately.
For reactions:
q = H*moles
H is the heat given off or taken in by one mole of a reaction. It's a way to compare reactions using a
standardized number, like miles per gallon for a car.
The number of moles is determined using the limiting reagent divided by its coefficient. For example:
2 Al + Fe2O3 –> Al2O3 + 2 Fe
If the reaction is done using 5 moles of Al and 4 moles of Fe2O3, plug '2.5' in for moles.
Chapter 11: Electrons
* The Bohr (solar system) model of the atom doesn't work.
* Electrons actually don't have nice neat little orbits.
* Instead, they ‘live’ in orbitals, which are places where we are 90% likely to find an electron.
* Orbitals come in different sizes and shapes.
* The Pauli Exclusion Principle states that each orbital can hold a maximum of two electrons—one spinning in
each direction.
* The electron configuration is a list of orbitals (in the form 2s, 3p, 6p, etc) and how many electrons are in
each as a superscript.
* Electrons get filled in the following order:
* Electrons get filled in the following order:
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
6f
7s
7p
7d
7f
* The idea that electron configurations build up from previous elements is the aufbau principle.
* The configurations can be written in shorthand using a noble gas as a base element from which to start:
[Ne]3s2 means that everything is the same as neon, with two more electrons in a 3s orbital.
* You can use the periodic table to count off configurations using the s, p,d, and f blocks. The s and p
blocks have n=row number, while d is row-1 and f is row-2.
* Hund's Rule deals with cases like p orbitals where you have choices about whether to fill each one first, or
pair electrons up. It says to put electrons in different orbitals first, with the same spin, before pairing any up.
* Valence electrons are those in the highest energy level (n). These are the ones that will react.
* Electrons in d orbitals don't count as valence, except for transition metals.
* Lewis dot structures have the element surrounded by dots—one for each valence electron
* Everything wants to end up with a full octet (8 electrons) in the valence shell. Except hydrogen, which
wants two.
* In ionic compounds, they get to the octet by taking or giving away electons. In covalent compounds,
they get to the octet by sharing (making a bond).
* General procedure for coming up with a structure: count the total valence electrons. Connect everything once
and fill up the octets. Count up the electrons—if there are too many, you need to remove lone pairs and make
more bonds instead. It's also useful to look at how many electrons an atom needs to get to 8—that will tell you
how many bonds it makes under most circumstances (C makes 4, N makes 3, O makes 2, F makes 1).
* Formal charge is determined by counting up the electrons that belong to an atom, and comparing it to
the normal number of valence electrons for that element. Atoms get one electron from each bond.
* When comparing potential structures, lower formal charge is better.
* When necessary, elements in the third row or larger can break the octect rule by making extra bonds
using their d orbitals. They only do this if needed to minimize charges.
* If you have a single element followed by a bunch of oxygens, the single one is in the middle with the others
radiating out. See SO42- for an example.
Chapter 13: Ideal gas laws.
* Various gas laws (Boyle's, Dalton's, Avogadro's, Gay-Lussac's) are all simplifications of one ideal gas law.
* PV= nRT
Where R = 0.08206 L*atm/mol*K so your units better match that.
* 1 atm = 760 mmHg
* If two variables are held constant and two changed, solve for the two variable ones and set them equal
to themselves under the new conditions: P1V1 = P2V2 or V1/T1 = V2/T2.
(a proportion!)
2
2
* van der Waals corrected the gas law to be
(P+ an /V )(V-nb) = nRT
where a and b are constants that you have to look up for each different gas.
* Dalton's Law says that for mixtures of ideal gases, total pressure is the sum of the partial pressures of
each gas. As a consequence of this, the partial pressure of any gas in a mixture is the percent (by moles)
of that gas, times the total pressure:
PTOT = P1 + P2 + P3
and P1 = %1*PTOT
Acids and Bases
* Bronsted defines an acid to be a proton (H+) donor, and a base as a proton acceptor.
* Arrhenius instead defines acids to be things that make H+ in water, and bases to be things that make
OH- in water.
* Acids can usually be recognized by splitting a compound apart in H+ and an anion. If the anion is
recognizable, then the compound is probably an acid.
* Acids have various strengths. A strong acid is defined as any one that completely breaks apart in H+ and an
anion when put in water. The strong acids are: HCl (and HBr, HI), HNO3, H2SO4, and HClO4.
* Strong bases are OH-, O2-, H-, and NH2-.
* If your base is hydroxide, the acid will react with it to make water and an ionic compound.
* pH measures both the strength of an acid or base, and how concentrated it is. 7 is neutral, lower is acidic,
higher is basic.
* pH = -log[H+]
* For bases, calculate the pOH = -log[OH-], then subtract from 14 to get pH (pH + pOH =14).
* Watch out for sneaky bases that have two hydroxides, like calcium hydroxide. Likewise, oxides produce two
hydroxides for each oxide that goes in.
* Phenolphthalein is an indicator which changes color from clear (in acid) to pink (in base). It can help you
figure out when you have neutralized an acid or base.