Big idea 1: Atoms & Elements (Annie Jho)

Molecules & Elements

Elements, Molecules, and Mixtures

All matter can be broken down into pure substances and mixtures. Pure substances include elements and molecules. Although they are both composed of uniform parts, elements are made up of one type of atom and molecules are composed of two or more (different) types of atoms bonded together, but still with same ratio of average masses.


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Chemical Analysis

Atomic Weight

One can determine how many moles (more information about the mole will be addressed later), is in a given mass by utilizing the average atomic mass, or the atomic weight. Atomic weight can be first calculated by multiplying the weight of each isotopes by their corresponding frequency (or percentage) in nature.


Usually, atomic weight is written on the periodic table under the chemical symbol until the 3rd decimal place.

Finding the Empirical Formla

When one has the measured mass of the given amount of substance, then one can divide the given mass by the atomic weight (average atomic mass) in order to find out how many moles (bunch of particles) is in that given amount.

By repeating the calculation for all measured mass of each element in the given amount, one can find all the mole numbers for corresponding elements. If one divides the amount of each element in moles by the smallest amount, then one can find the rough ratio for each elements in that compound or molecule, which will be the empirical formula. Sometimes the ratio may not be in whole numbers. In that case one MUST make it to whole number ratio by multiplying common factor to all elements.

Sometimes, the questions will have them not as a numerical mass measurement, but as percentage. Do not panic -- in that case, just think as having 100g of the sample. That will make all the percentages to be just the mass. For example, if the question said a compound has 73% mercury (Hg) and 27% chlorine (Cl). Simply by hypothetically saying that we only have 100g of that compound, one can state that there is 73g of Hg and 27g of Cl. Then it’s much easier to figure out the empirical formula.

Finding the Molecular Formula

ALSO! Since the mass is measured, one can divide the given mass by the molecular mass of the empirical formula in order to find the molecular formula.

Structural Formula

        In addition to empirical and molecular formulae that represents the simplest and the actual ratio or the number of the compound, there is another formula called structural. It visually shows the structure of the compound, as well as represent (in form of letter formula) how the molecule or the compound is arranged.

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The Mole

The Mole

        Mole is basically a specific “unit” (technically not) used when dealing and calculating incredibly small and great number of atom or particles. Similar to how we call “12” a “dozen,” (ex. Dozen of pencils, dozen of doughnuts, etc.) a “mole” is “6.02 x 1023” atoms/particles/ions/whatever. That specific and weird 6.02 x 1023 number has its special name called Avogadro’s Number. Why do we use it? It makes life much easier when one deals with the massive number of atoms or particles in their chemistry calculation (and you will be thankful of this number later on). Well the ‘mole’ is the Avogadro’s number, but when one writes it in calculations like ‘1 mole of Carbon’, then one must write it as ‘1 mol (of) C’.

Particles to Mass (vice versa)

        The mole serves as an important conversion tool between molecular mass of the compound/molecule, and its number of particles. First of all, the atomic mass unit is always (already) calculated and written based on 1 mol for the element. In other words, the atomic weight on periodic table is the mass of 1 mol of that element. It is almost impossible to calculate with the mass of a single atom for that element, so scientists use mole to calculate in a bigger and proper scale. This means when one is trying to find the molar mass of 1 mol of CO2, one must add 1 mol of C and 2 moles of O to find 1 mol of CO2. So 1 mol of C will be 12.01 grams, again with 6.02 x 1023 Carbon atoms in that 12 grams, and 2 moles of O will be 32 grams, which results in 1 mol of CO2 with 44.01 grams.

        If one wants to find the number of molecules, not moles, of -- let’s say -- water in 9.01 grams, one can use the idea of molecular mass and mole (Avogadro’s number) to find it.

In order to determine the mass from the number of molecules or number of moles is very similar, just the opposite direction. This will be dealt in depth in stoichiometry, but this was the basic mass to particles conversion using the idea that the mole is the bridge.

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Coulomb’s Law

Coulomb’s Law

        In the simple or basic atoms, there are neutrons and protons (usually in equal amount depending on their atomic number) clumped in the center, and also equal number of electrons are floating around the nucleus. While neutron has neutral charge (meaning no charge), the protons have positive charge and the electrons have negative charge. And these opposite charges attract each other: this was Coulomb’s first law. Coulomb discovered that with opposite charges, the particles will attract each other. And another part he discovered was that like charges -- positive with positive and negative with negative charges -- will repel each other. This will make protons in the nucleus to repel each other and become unstable, but the neutrons hold it all together by its nuclear charge.

        The actual Coulomb's law equation shows the relationship between the force and the charge and the distance (or radius) between the charged particles.

The Coulomb’s law tells us that the force between charged particles is proportional to the magnitude of the charges of each particles. It also tells us that it is inversely proportional to the radius, or the distance between two. So if any one of the two charges get larger, then the force will be greater as well; if the radius between two particles get greater (meaning further away), then the force will be smaller. However, as one can see, it is radius squared, so is radius and the charge increased by the same (multiplication) factor, then radius will have greater impact.

Ionization Energy

        Coulomb’s law, especially in chemistry, can be used to predict ionization energy. One can actually measure the ionization energy with photoelectron spectroscopy (PES) and predict the shell structure as well. Keeping in mind that ionization energy is essentially energy required to “take out” an electron from its shell, one can determine how many electrons are located in which level shell by looking at this data.

The closer the electron is to the nucleus, the smaller the radius will be, and according to the Coulomb’s law, smaller radius means greater force (or ionization energy). Therefore, 1s2, the closest shell, will have the electron that requires most energy compared to other electrons of that atom: 3206 eV for Argon, 870 eV for Neon, and about 22.5 eV for Helium.

        Although electron configuration will be discussed in depth later, PES analysis utilizes electron configuration as shown in both pictures. Usually, PES shows the number of electrons in peaks with relatively different in height (twice as high or six times high). Again, same in the data image on the right side, the peak (electrons) in the greater energy domain is the closest one to the nucleus. Since the picture on the right has two peaks, one at 6.26 MJ/mol and another at 0.52 MJ/mol, one can know that the element has until 2s shell. Now, if you look at the 6.26 MJ/mol peak, it is approximately twice the size of the 0.52 MJ/mol peak. Thus, the element will have total of 3 electrons and an electron configuration of 1s22s1, occupying two electron shell levels.

        This is how PES works:

        Furthermore, Coulomb’s law explains why it is easier to take away further electrons (electrons in higher energy level or shell) than closer ones. This is also due to the radius: electrons further away are much easier to take away than electrons located nearby.

Photoelectric Effect

        The photoelectric effect is the emission of electrons or other free carriers when light shines on a material. Electrons emitted in this manner can be called photoelectrons, which is used in PES.

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Electron Configuration & Quantum Mechanical Model


        Electron configuration primarily shows the distribution of electrons in atoms or ions. A good way to find that out is by looking at the different ionization energy. There are 4 types of orbitals in order:

  1. s-orbital
  2. p-orbital
  3. d-orbital
  4. f-orbital

S orbital has 1 subshell, p has 3, d has 5, and f has 7. The position and the orientation can also be different depending on other factors of the electron configuration such as the followings:

As N increases, the orbital becomes larger, and electrons stay longer farther from the nucleus. L (angular momentum quantum number) shows the variety of the orbital shapes, and it is related to the principal quantum number. Ml shows the variety of distinguishes the orbitals available within a subshell. Spin quantum number simply shows the direction of the electron spin.


To give you simple visualization of the orbitals, refer to the picture below.

There are 3 rules that helps one define how the electrons fill in the orbitals within the atomic structure.

  1. Hund’s rule
  1. Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

  1. The Aufbau Principle
  1. In the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels.

  1. The Pauli Exclusion Principle
  1. Two or more identical fermions (particles with half-integer spin) cannot occupy the same quantum state within a quantum system simultaneously.

In the periodic table, the orbitals are located as shown.

For example, Carbon, which has atomic number of 6 (meaning 6 protons AND 6 electrons) will have 1s22s22p2 as it’s electron configuration. You write the number of electrons occupied for that shell as a small number written as a superscript on that orbital’s letter.

Different Electron Configurations

  1. Regular type
  1. The regular ones start from the beginning, 1s, no matter how far the element is from the beginning. This will include every electron configuration in the middle until that specific element.
  1. Ex. Calcium → 1s22s22p63s23p64s2
  1. Condensed type→ you should know how to write this!
  1. The condensed form is a simplified, or condensed, form of writing the electron configuration. Starting from the closest previous noble gas, you write it with [Noble gas]_____ (and the rest of the electron configuration starting from that noble gas until the element).
  1. Ex. Calcium → since Argon is the closest, past noble gas, we start from [Ar]. [Ar]4s2.
  1. (Not a type but exceptions from regular rules) Anomalous Electron Configuration
  1. This is because when scientists observed them in lab, they found out that this form was much more stable than the expected electron configuration.

First Ionization Energy

        The “first” ionization energy is the amount of energy needed to take out the first, outermost valence electron. The “second” ionization energy will be the energy required to take away the second outermost valence electron, and the “third” is third outermost valence electron, and so on.

This first ionization energy will differ with different elements because each elements have different number of valence electrons and protons in the nucleus. Compared to H, He will have much higher ionization energy because there’s more protons, but He will also require greater ionization energy compared to Li because the Li will have 1 valence electron on the second shell and the radius will increase.

And by looking at the amount that the first, second, third and further consecutive ionization energy increases, it may be possible to spot a huge leap in the ionization energy, which will indicate the number of valence electrons. Eventually one will be able to identify the element by the sudden increase in ionization energy.

There is specific trend for this first ionization energy throughout the periodic table in increasing atomic number.

Shielding Effect

        Also known as screening effect, shielding effect is when the inner electrons shield the outer electrons from getting close to the nucleus, and at the same time electrons repel and make the atom to be bigger.


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Atomic Radius

        Atomic radius decreases from left to right within a period. This is caused by the increase in the number of protons and electrons across a period. One proton has a greater effect than one electron; thus, electrons are pulled towards the nucleus, resulting in a smaller radius(Nuclear Effective Charge increases as we go across the period, resulting with a higher attraction between the nucleus and the surrounding electron shield). Atomic radius increases from top to bottom within a group. This is caused by electron shielding and the increase of energy levels.

Image result for atomic radius trend

Ionization Energy

        The ionization energy of the elements within a period generally increases from left to right. This is due to valence shell stability. Decreases from top to bottom. This is due to electron shielding.

The noble gases possess very high ionization energies because of their full valence shells as indicated in the graph. Note that helium has the highest ionization energy of all the elements.

Ionic Radius

        Ionic radius will increase for anions, which gains an electron to become an ion. The addition of another electron will repel other electrons to spread out wide, and thus become bigger in radius. However, the ionic radius for cations will be smaller, because they will lose an electron. That will allow protons to pull closer, and sometimes ionization might lead to one less shell (or level).


         The electronegativity increases as we go across the period because if the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one.

The electronegativity decreases as we go down a group because the atomic number increases down the group, increasing the distance between the valence electrons and nucleus(a greater atomic radius).

For the transition metals, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.

Exceptions include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values.

Electron Affinity

        Electron affinity increases from left to right within a period. This is caused by the decrease in atomic radius.Electron affinity decreases from top to bottom within a group. This is caused by the increase in atomic radius.

Melting Point

        Metals generally possess a high melting point.Most non-metals possess low melting points.

The non-metal carbon possesses the highest boiling point of all the elements. The semi-metal boron also possesses a high melting point. Temperature is directly proportional to energy, a high bond dissociation energy correlates to a high temperature.

Image result for melting point trend

Boiling Point

        Again, temperature is directly proportional to energy, a high bond dissociation energy correlates to a high temperature.

Image result for melting point trend

Metallic Character

        From right to left across a period, metallic character increases because the attraction between valence electron and the nucleus is weaker, enabling an easier loss of electrons. Metallic characteristics increase down a group. Electron shielding causes the atomic radius to increase thus the outer electrons ionizes more readily than electrons in smaller atoms.

Metallic characteristics decrease as we go right across a table because of the decrease in radius(Zeff) of the atom allows the outer electrons to ionize more readily.

Increases as we go down a group since electron shielding causes the atomic radius to increase, thus the outer electrons ionize more readily than electrons in smaller atoms.

Nonmetallic Character

        Increases going from left to right across the periodic table. Metallic character relates to the ability to lose electrons, and nonmetallic character relates to the ability to gain electrons.

Image result for nonmetallic character trend

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Atomic Models

John Dalton’s Atomic Theory

  1. All matter is composed of tiny, indivisible particle called atoms, which cannot be destroyed or created.
  2. Each element has atoms that are identical to each other in all of their properties, and these properties are different from the properties of all atoms.
  3. Chemical reactions are simple rearrangements of atoms from one combination to another in small whole-number ratios.

Dalton proposed the “Billiard ball” model, which states that all matter is composed of composed of small, spherical particles. However, this was rejected when later evidence proved that even the small atoms are composed of smaller subatomic particles.

J.J. Thomson’s Plum-pudding model

        J.J. Thomson came up with the idea of having charges embedded with Dalton’s Billiard Balls. He also used cathode ray experiment to discover the existence of the negatively charged electron.


Here, Millikan did not propose any further, different model, but he discovered the mass of the electron through oil drop experiment. He determined that the mass of the electron: 1/1840 the mass of a hydrogen atom, and is 9.11 x 10-28 g.

Rutherford’s Nuclear Model

        Ernest Rutherford conducted the gold foil experiment to test whether atoms were made out of totally empty space, or not. He discovered the existence of the neutron after the gold foil experiment, which was set up like the picture below.

The alpha particles from was shot through a thin gold foil, and as Rutherford expected, most of the rays just passed through the gold foil. He thought the atoms would be mostly hollow with tiny electrons and protons, but surprisingly some rays were deflected. He then concluded that there is a clustered matter at the center of the atom which was dense and big enough to deflect the light ray: that was the discovery of nucleus. At the end, he made 3 conclusions:

  1. The nucleus is small.
  2. The nucleus is dense.
  3. The nucleus is positively charged.

Niels Bohr’s Bohr Model

        Bohr model is like the current day textbook image of atoms with the nucleus in the center and rings of electron shell around the nucleus. Bohr discovered that electrons exist in several distinct layers or levels, and that the electrons travel around nucleus like planets travel around sun. He also found out that electrons can jump between levels with energy being added or released.

Bohr also suggested that the electrons can only revolve in certain orbits, or at certain energy levels (ie, the energy levels are quantized).

Heisenberg Uncertainty Principle

        Heisenberg stated that it is impossible to know the exact momentum & location in space of an electron. If one can know the exact momentum, than they cannot know the another, position. And the reverse is also true.

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Mass Spectroscopy

How it works

        Mass spectroscopy works by utilizing the difference in mass of each isotope for that element. The heaviest isotope (with most number of neutrons) will bend the least, and the lightest isotope (with least number of neutrons) will bend the most, and the detector measures the abundance of each isotope.

When the mass spectroscopy measures the relative abundance, the data looks like this:

Which shows numerous bars on various atomic weight for isotopes. And by multiplying the mass with its corresponding abundance, one can find the atomic weight as addressed in Atomic weight section in Chemical Analysis.

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Light and Matter


        In electromagnetic radiation (radiant energy), there are many portions that includes visible light, infrared, and ultraviolet radiation. Infrared lights are important when observing the vibrations of the molecules, and identifying which bonds exist between atoms. Through ultraviolet or visible light, one can see the electronic structure: where the electrons are found, and what they are doing.

One can also look at solutions using light and measure the absorbance to calculate and determine the concentration of the solute in the solution. Beer’s law states that absorbance is proportional to the concentration and the path length. Sometimes scientists can measure the concentration from transmittance.

        According to the Planck’s equation, the radiation with high energy will have high frequency, and therefore small wavelength. The radiation with low energy will have low frequency, but will have wide wavelength.

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Symbolic Representations & Conservation Laws

Conservation Laws

  1. Conservation of mass
  2. Conservation of energy
  3. Conservation of atoms

Three conservation laws above all states that mass, energy, and atoms respectively are neither created nor destroyed. Conservation of matter can be used to calculate the products from reactants or vice versa. One can also look at the macroscopic changes, such as movements of particles. Physical and chemical changes will still follow the law of conservation of mass as well, and an example of physical change would be change in water phases and chemical change will be hydrogen combustion like hydrogen bomb.

Conservation of energy does not come out often in chemistry but mostly in physics -- from potential to kinetic energy. So you do not have to worry much about this law.

However, conservation of atoms is very significant in chemistry because that is the fundamental law behind writing chemical equations. This will be even more important in stoichiometry with all of the complex calculations, but this law is the reason why we MUST write the coefficients in chemical equations. If there was 2 H atoms (maybe in form of H2 gas) in the reactant side, there must be 2 H atoms still of the product side.

Chemical Symbols

        The change of state of matter in chemical equation before and after the reaction must be written in the chemical equation. Gas is noted as (g), liquid is (l), solid is (s), and aqueous precipitate is written as (aq).

It is also important to be clear whether it is a new formation or breaking of chemical bond, or whether it is just changes in the number of molecules or elements.


Also, the input and output of heat or any energy formed from the reaction (light, heat, sound, etc.) must be noted by “+E” or “+heat” on whichever side the energy was involved in the reaction.

        One way arrow means that the reaction is one wayed, irreversible, and is only spontaneous for one direction. Two sided arrow represents that the reaction is reversible depending on different factors such as number of moles on either side, temperature, pressure, and so forth.

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Big Idea 2: Structure & Properties of Matter (Heumll Wang)

Bonding and Intermolecular Forces

Intramolecular force: forces that hold atoms together within a molecule

Intermolecular force: forces that hold molecules together

Types of Intramolecular Forces:

  1. Ionic Bonding: non-metal + metal bonding. There’s usually a big difference in electronegativity between the atoms so electron gets transferred from one atom to another.
  1. Examples: NaCl, CaF2

-> Sodium has 1 valence electron, whereas chloride has 7 valence electrons. Sodium (non-metal) and chloride (metal) will form an ionic bond with chloride stealing an electron from sodium. At the end both reach octet and they all are happy.

  1. Covalent Bonding: nonmetal + nonmetal
  1. Examples: CH4, H2, O2, H2O, CO2 
  2. Atoms share electrons

Answer: Yes and No. It all depends on atoms’ electronegativity

Oxygen and hydrogen do not evenly share their electrons. Oxygen is much more electronegative than hydrogen, so it will hold electrons for much longer period of time, making oxygen slightly negative and hydrogens slightly positive.

Covalent bonding wouldn’t form charges when atoms have similar (if not same) electronegativity. Carbon and hydrogen have about the same electronegativity, so they evenly share their electrons.

Types of Intermolecular Forces:

  1. London dispersion force: weakest intermolecular force, present whenever molecules bond with each other

  1. Dipole-dipole force: attraction between positive portion of molecule and negative portion of molecule

- In the diagram above HCl molecules are having dipole-dipole force as Cl is partially negative and hydrogen is partially positive

  1. Hydrogen bond: strongest intermolecular force, force that occurs when hydrogen is attracted to fluorine, oxygen, or nitrogen.

- In the diagram above, hydrogen (partially positive) of a water molecule is attracted to oxygen (partially negative) of another water molecule, creating hydrogen bond between the water molecules.

Big Idea 3: Chemical Reactions (Brian Sim)

Molecular, Ion & Net Ionic Equations

Regardless of whether you’re in an AP Chemistry class or a normal Chemistry class, you’re going to be dealing with molecular equations almost all of the time. What is a molecular equation? A molecular equation is merely a reaction written out in the form of words. For example, a reaction between lead(II) nitrate and potassium iodide. (*If you are unfamiliar with the elements, it’s highly advised to study the elements of the periodic table. Memorizing information about commonly asked elements and molecules, which includes their name, symbol, and charge, will greatly help you for the AP course and test.)

The reaction between lead(II) nitrate and potassium iodide will result in the following equation:

Pb(NO3)2 + KI → PbI2 + KNO3

Lead(II) nitrate - Pb(NO3)2 - and potassium iodide - KI - are expressed in an equation that represents the entire reaction. However, this is not the correct molecular equation. Why? Because the equation above is not properly balanced. In order to solve chemistry questions correctly, it’s essential to make sure that the molecular equation that you end up with is balanced. Balanced, in terms of chemistry, means that there must be an equal amount of atoms on both sides of the equation. If you notice above, there are two NO3 molecules on the left side while there is only one NO3 molecule on the other. There are, also, two molecules of I on the right side side while there is only one I molecule on the left side. Hence, this equation is not balanced. In order to balance the equation, you must add more molecules on either side to balance the equation. As a result, you’ll end up with:

Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)

(*Notice the symbols (aq) and (s) in the equation. These represent the current state of the molecules. (aq) stands for aqueous and (s) stands for solid. Other symbols are (l) which is liquid and (g) which is gas. Writing down these states will be a great habit to develop and will ensure your teacher’s satisfaction. But note: if you don’t know the states, don’t write it. You may get it wrong and lose points depending on the situation, whether it’s in class or on the test.)

With that, we’re done with the molecular equation. We’ve successfully written out the reaction and balanced the equation so that there was no unevenness on either side. Now we go into the ionic equation. This is when memorizing the charges helps! In an ionic equation, what you’re doing is basically taking the molecular equation and splitting it based on the ions in the equation. So the ionic equation for the same example will be:

Pb2+ (aq) + 2(NO3)- (aq) + 2K+ (aq) + 2I- (aq) → PbI2 (s) + 2K+ (aq) + 2(NO3)- (aq)

As you can see, all of the ions in the aqueous form was split apart to show each individual ion. The solid PbI2 was left as it was because it is a solid and does not want to be split into ions. The ionic equation is fairly long and may scare some students. That’s why we have the net ionic equation, which is the ionic equation without the spectator ions. Spectator ions are ions that have gone through zero change throughout the reaction. Looking at the ionic equation above, we can see that the two (NO3)- ions and two K+ ions are on both sides of the equation, meaning that they were not influential to the reaction. Removing them will give us the net ionic equation:

Pb2+ (aq) + 2I- (aq) → PbI2 (s)

By removing the spectator ions, we are left with the influential ions in the reaction and the final net ionic equation.


Stoichiometry takes what was described before about molecular equations and uses them to calculate amounts of reactants and products. Stoichiometry is one of many topics you initially learn and is vital for a significant portion of chemistry. Early into the year of an AP Chemistry course, most stoichiometry problems will have you calculate the expected product. You will be taught more difficult and complex ways that you can apply stoichiometry in chemistry (which includes finding the limiting reactant/reagent, percent yield, molar mass of gases, and titration), but in here, only the basics of stoichiometry will be gone over. (*Stoichiometry is an important topic in chemistry, and it’s confusing to many initially. A horrible combination for students. Because of this, it’s highly advised to seek help from a teacher if you don’t understand the concept.)

In a common stoichiometry problem, you will be performing calculations similar to what you learn in algebra about unit conversions. Take, for example, the reaction between methane and oxygen. The balanced chemical (molecular) equation will be: (*It must be balanced!!!!)

CH4 + 2O2 → CO2 + 2H2O

A simple problem will give you the amount of either a reactant or product. For this example, let’s say that there’s 28.3 grams of CH4, and we want to find the amount of H2O from that amount of CH4. To start off this stoichiometry problem, we must first convert the grams of CH4 into moles of CH4 (you will be converting values to moles for most stoichiometry problems). To get the moles, we multiply with the following conversion rate:

So for this example, we would multiply 28.3 grams of CH4 by 1/16.05 to convert our grams into moles. It would look like this:

Now that we have the moles of CH4, we have to convert it to moles of H2O. How do we do this? You look at the molar ratio, which is the ratio between the moles of the two molecules of interest. You can determine the molar ratio based on the balanced chemical equation, which is determined by the coefficients from the equation. (*This is why you must correctly balance the equation. Otherwise, you’ll end up with the wrong molar ratio and get the problem wrong.) From the equation, we can see that for one mole of CH4, we get two moles of H2O. So we multiply the appropriate conversion ratio and then multiply that by the molar mass of H2O to determine the amount of H2O by the end of this reaction. The whole stoichiometry will look like this:

Through stoichiometry, we’ve found out that with 28.3 grams of CH4, 63.5 grams of H2O will be formed.

Synthesis & Decomposition Reactions

The concept of synthesis and decomposition reactions is fairly simple. However this simple concept can be used and applied to the harder sections of AP Chemistry so keep that in mind.

A synthesis reaction refers to the combination of multiple of atoms or molecules to form a new compound. For example:

C + O2 → CO2

A decomposition reaction is the exact opposite of a synthesis reaction as it will break apart a larger compound into smaller atoms or molecules. For example:

H2O → 2H + O

Neutralization Reactions

A neutralization reaction can also be called an acid-base reaction. Acids and bases are opposites; so when an acidic solution meets an equally strong basic solution, they neutralize, hence, a neutralization reaction. Acidic solutions have a pH level lower than 7. The acid becomes stronger as its pH value decreases. Basic solutions tend to have a pH level higher than 7. The strength of the base increases as its pH value increases.

One key distinction between the two is how they deal with a proton. When referring to a proton, in terms of acid and base, a proton is a hydrogen atom or H+. An acid is known as a proton donor or proton giver, while a base is known as a proton receiver. So in a reaction between an acid and a base, the acid will give up its H+ ions to the base. The acid’s counterpart (the atom/molecule without the H+ ion located on the other side) is now referred to as the conjugate base while the base’s counterpart is referred to as the conjugate acid since it received H+ ions from the original acid. Through an equation, it looks like this:

HA + B A- + BH+

Acid + Base Conj. Base + Conj. Acid

An example of a neutralization or acid-base reaction is the reaction between acetic acid and water:


So in this reaction, we can see that the acetic acid (CH3COOH) lost its H+ ion and became CH3COO- on the products’ side. Since it lost a H+ ion or gave its H+ ion, CH3COOH can be shown to be the acid while water, in this case, in the base. However, water, H2O, is not always the base of every reaction. In some situations, water can also be acidic. Therefore, water is known to be amphoteric, meaning that it can either be an acid or a base.

An acid’s strength can be determined by its dissociation. HCl is an extremely strong acid. It will always dissociate completely, meaning that its contents will always split into H+ and Cl- ions. Weaker acids will not be able to dissociate completely, with some barely being able to dissociate at all. Also, as the acid becomes strong, the lower its potential to be a base becomes. So basically, acids and bases will be going in “opposite directions” of each other in every sense.

Redox Reactions

Redox reactions refers to both oxidation and reduction reactions. The combination of both results in high energy being released. Oxidation is the action of losing electrons, while reduction is the action of gaining electrons. At first glance, this may be hard to remember which is why there is an acronym for this: OIL RIG. The “O” stands for oxidation and “IL” stands for is losing. The “R” stands for reduction and “IG” stands for is gaining. What are they gaining or losing? They’re gaining or losing electrons, which carry high amounts of energy. So the electron always goes from an oxidized substance to a reduced substance (*This concept is very important as you will see it again in electrochemistry). Once you find out which substance is oxidized or reduced, you can, then, create half-reaction equations. Basically, you can write out two equations: one for the oxidized substance and the other for the reduced substance. The oxidation half-reaction always has the electrons on the products’ side, and the reduction half-reaction always has the electrons on the reactants’ side.

It’s sometimes hard to determine which one is which: Which is the reduced and which is the oxidized? To determine this, we look at the oxidation number. The following is a set of rules to determine the oxidation number. (*Don’t worry about this too much as you will be learning about redox reaction at a much deeper level in class.)

  1. Free elements are zero.
  2. For ions it is the charge.
  3. Oxygen is -2 (peroxides -1)

        Hydrogen is +1 (with nonmentals) -1 (with metals)

  1. In neutral compounds the sum is zero (or charge) in ions

By following these rules, you will be able to determine which molecule/element is oxidized or reduced and create your half-reactions from there.

Chemical Change Evidence

Chemical reactions occur everyday in our lives and in our bodies. A vital process, cellular respiration, is an example of a chemical reaction. However, how do we know that chemical reactions that are microscopic truly exist? That’s when we look for evidence. Chemical reactions goes through change that can be seen through one or more of five different characteristics:

  1. Odor
  2. Color
  3. Gas
  4. Heat
  5. Precipitate

Seeing something that would go through and exhibit a change from one of the five forms of evidence provided above would be evidence that a chemical change has occurred.

Endothermic & Exothermic Reactions

A change in heat due to a reaction can be classified as either an endothermic reaction or an exothermic reaction. In an exothermic reaction, heat is released to the surroundings, while in an endothermic reaction, heat is absorbed from the surroundings. So in an exothermic reaction, there is more energy within the reactants before they undergo a chemical reaction. After the reaction, since heat is lost to the surroundings, the resulting energy is lower than its initial start. An endothermic reaction is the opposite of that. The initial energy is lower compared to the final energy, since heat (a form of energy) is absorbed from the reaction from the surroundings.

The graph below on the left side shows an energy diagram of an exothermic reaction. The graph below on the right side shows an energy diagram of an endothermic reaction.


Electrochemistry is a very big part of AP Chemistry. In electrochemistry, we are looking the flow of electricity through redox reactions. To see this, we have one of two electrochemical cells: galvanic cells which generate an electrical current without human interference or an electrolytic cell which requires a separate electrical current. The usage of one of the two will be specified in the question that you will be solving in the future. Back onto the topic of redox reactions, electrochemistry focuses mostly on that. (*So if you’re having trouble with redox reactions, it’s highly advised to first understand that before moving onto electrochemistry.) There are two different settings/cells for the redox reactions to occur in. One is the cathode, and the other is the anode. The reduction reaction will always occur at the cathode and the oxidation reaction will always occur at the anode. An easy way to memorize this is through the acronym: AOCR. The “A” stands for anode, and the “O” stands for oxidation. The “C” stands for cathode, and the “R” stands for reduction.

To determine which half-reaction would occur at the anode or cathode, you would look at the reduction potential table. Look for your two half-reactions, and look at the given reduction potential values along with them. The larger value is placed at the cathode, and the smaller value is placed at the anode. (*Remember that the cathode is where reduction occurs!) Now we can find the standard potential of the cell or the energy generated. To do this, subtract the smaller reduction potential value from the larger one. Or just remember that it’s always Ecathode - Eanode.

In electrochemistry, it’s also important to know how the galvanic cell system is set up and how it works. The two metals are placed at their respective half-cells, the anode or cathode. You would then link the two metals together with a voltmeter to measure the standard potential. It’s important to note that the energy (electrons) travel from anode to cathode. Electrons always travel from anode to cathode and not the other way around. Another important part to this system is the salt bridge. The salt bridge contains a compound that can split into a cation and anion. Without the salt bridge, the entire system will fail since equilibrium cannot be maintained without it. Also note that the metal located at the anode would start to decay and lose mass, while the metal at the cathode would gain mass and grow larger. (*Electrochemistry is considered to be a hard topic for students, but if you take your time, you may turn out to like this topic the most out of all other chemistry topics!)

A galvanic cell set-up

Big Idea 4: Kinetics (Jun-young Lee)

Factors that affect reaction rates

Collision Theory: the collision theory states that when suitable particles of the reactant hit with each other, only a certain percentage of the collisions cause any noticeable or significant chemical change. More collisions cause more successful collisions and increase the rate of reaction

In Collision Theory…

Types of Chemical Reaction Rates

The Rate Constant K


Reaction Mechanisms: a mechanism represents the sequence of bond-making and bond-breaking steps that occurs during the conversion of reactants to products

Catalysis: Increase in rate of a chemical reaction due to the participation of a substance. Catalysts function by lowering the activation energy of an elementary step in a reaction mechanism, and by providing a new and faster reaction mechanism

Big Idea 5: Thermodynamics (Jongwon)


  1. Energy is the capacity to do work or produce heat

Law of Conservation of Energy

  1. Energy can neither be created nor destroyed, but can be converted from one form to another (thermal, mechanical, electrical, light, or chemical).
  1. Potential Energy
  1. Energy due to position or composition
  2. In AP Chemistry, potential energy will most likely refer to chemical potential energy, the energy stored in chemical bonds.
  1. Kinetic Energy
  1. Energy due to motion of an object
  1. When conducting labs in class, there will always be a small error in theoretical and experimental values. Because energy cannot be created or destroyed, the sources of error will almost always include loss in energy to the surrounding atmosphere.


  1. Temperature is the average kinetic energy of molecules.
  1. Each individual molecules will have different speeds, and the mean average of the distribution (AKA Maxwell-Boltzmann Distribution) of molecule speed will determine the temperature.
  1. Celsius to Kelvin conversion: K = o C + 273.15
  1. The AP exam also accepts +273, but it is best to take the extra step just to make sure your calculations are accurate.
  2. At 0 K, there is no motion of molecules, often referred to as absolute zero.

Heat Exchange

  1. Heat is the transfer of energy.
  2. When there is a warmer body and a colder body coming in contact, the kinetic energy of molecules will be transferred from hot to cold, and never vice versa.
  1. When left isolated with enough time, the two bodies will reach “thermal equilibrium”, where both bodies will have the same average kinetic energy, temperature.
  1. Specific Heat Capacity ()
  1. Is the energy required to raise the temperature of a given substance with a given amount.
  1. Lower the heat capacity value is, the easier it is to raise or lower the temperature of the substance.
  2. Higher the heat capacity value is, the harder it is to raise or lower the temperature of the substance.
  1. Every substance have different specific heat capacities, and will always be given on the exam if necessary
  1. The heat capacity of H2O is very common on the test, so it is handy to memorize; (or 4.18).


Phase Changes

  1. When some substances are given or taken away of energy, they will go through phase changes.
  1. There are 3 common states of matter: solid, liquid, and gas. Occasionally plasma is also introduced.
  2. Below is a graph of phase change from solid, liquid, to gas, of an ordinary substance such as water.
  1. The stagnation between the states show that there is an amount of energy required to change from one state to another.
  2. The energy required to melt a solid is referred to as “molar heat of fusion”, and the energy required to vaporize a liquid is referred to as “molar heat of vaporization”, both in units of J/mol.

Enthalpy of Reaction

  1. Enthalpy  is the total amount of energy in the contents of a reaction.
  1. We cannot measure the total amount of energy in a system, but can only measure the changes in enthalpy (ΔH).
  1. As seen in the AP Chemistry formula sheet, under thermodynamics/electrochemistry, change is defined as sum of products - sum of reactants.
  1. A reaction is endothermic if there is more consumption of energy, taking energy away from the surrounding. (
  2. A reaction is exothermic if there is more release of energy, giving energy to the surrounding.
  1. Don’t confuse with surrounding and the reaction system. Think of the reaction and the surrounding as 2 seperate bodies
  1. A chemical reaction will have an activation energy, notated in Ea, where it is the initial energy required to get the reaction started.
  1. The Ea hump can be lowered with a catalyst, which speeds up the reaction, but is not consumed.
  1. On the right, it would be an exothermic reaction, because there are less energy of products than reactants, giving away energy to the surrounding.
  2. The enthalpy of a chemical reaction can be reversed when the reaction is reversed.
  1. For example,

C(s)+O(g) → CO2(g) , ΔH  = -393.5 kJ

 However, if the reaction is flipped like so,

CO2(g) → C(s)+O(g) , ΔH  = 393.5 kJ


  1. Change in entropy (ΔS) is the change in order or chaos/randomness in a system.
  1. The more orderly, the lower the entropy.
  2. The more chaotic, the higher the entropy.
  3. By the 2nd law of thermodynamics, entropy cannot decrease, but can remain the same or increase in the universal scale.
  1. Determinants of entropy
  1. States of matter
  1. From least to most random; solid, liquid, then gas.
  1. Stoichiometry
  1. If there are more moles in the product than in the reactant, entropy has increased.
  1. Gas Volume
  1. Increase in gas volume results in more dispersion, thus increasing entropy.
  1. Gas Temperature
  1. According to the kinetic molecular theory, the energy is more spread out as gas temperature increases, thus increasing entropy.

Spontaneous Processes

  1. A reaction is spontaneous if it occurs without being driven by outside forces, and non-spontaneous if it requires outside forces to initiate.
  2. Spontaneous if:
  1. ΔH < 0,  and ΔS > 0
  2. Products are favored.
  1. Non-spontaneous if:
  1. ΔH > 0,  and ΔS < 0
  2. Reactants are favored.

Gibbs Free Energy

  1. Gibbs free energy (ΔG) is the energy available to do work.
  1. Spontaneous if ΔG < 0
  2. Non-spontaneous if ΔG > 0
  3. At equilibrium if ΔG = 0
  1.                                  ΔG = ΔH - TΔS
  1. The equation above is one of the most used equation in thermochemistry
  1. ΔG = change in gibbs free energy (J)
  2. ΔH = change in enthalpy (J)
  3. T = temperature (K)
  4. ΔS = change in entropy (J/K)

ΔH < 0

ΔH > 0

ΔS > 0


Spontaneous at high temperatures

ΔS < 0

Spontaneous at low temperatures


Big Idea 6: Equilibrium (Dawon Song)

Any bond or intermolecular attraction that can be formed can be broken. These two processes are in a dynamic competition, sensitive to initial conditions and external perturbations. 

Reversible Reactions

Irreversible Reaction: Reaction goes from left to right with reactants forming products

Reversible Reaction: Reactants become products, products become reactants and eventually hits equilibrium (sum amount doesn’t change)

Chemical Equilibrium: A dynamic reversible state in which rates of forward and reverse reactions are the same.

The Reaction Quotient

Reaction Quotient Q: The ratio of the concentration of products to reactions at any given time before equilibrium

Q<K means reaction moves to the right (reactants>products)

Q>K means reaction moves to the left (reactants<products)

Q=K means equilibrium


Green: Reactants

Red: Products

The Equilibrium Constant

K: The ratio of products to reactants in a reversible reaction (only occurs when reaction reaches equilibrium)

Use the value of K in a reversible reaction to determine which chemicals will have very large versus very small concentrations at equilibrium

K>>1 means the reaction moves to the right (product-favored)

K<<1 means the reaction moves to the left (reactant-favored)

Le Chatelier's Principle


Equilibrium Disturbances

Disturbance to a reversible reaction at equilibrium affect the equilibrium constant K and the reaction quotient Q.

The reaction shifts to bring Q and K back into equality

Acid-Base Equilibrium

Acid-base chemistry is a reversible reaction. It will eventually achieve equilibrium. We can calculate the equilibrium constant K (conc. of products over conc. of reactants)

Acid: Donate a proton

Base: Accept a proton

Acid-base reactions have conjugate pairs

Water autoionization: Water acts as both acid and base

Calculate the K value to find or and

pH = -log

pOH = -log

Acid Ionization Constant : the equilibrium constant for the acid

Base Ionization Constant : the equilibrium constant for the base

Neutralization Reaction: Combine acids and bases together, measure this using titration and titration curves

Titration: the slow addition of one solution of a known concentration (called a titrant) to a known volume of another solution of unknown concentration until the reaction reaches neutralization, which is often indicated by a color change.

At the equivalence point, the concentration of titrant (known solution) and the concentration of analyte (unknown solution) are the same.

The of an acid can be determined from the pH at the half equivalence point if the equivalence point is known.

pH & Buffers

Buffer solutions maintain pH in a solution. A buffer solution is made up of a weak acid and its conjugate base. As strong acids or bases are added, the pH remains stable. A good buffer solution has a pKa value equivalent to the pH and equal amounts of the weak acid and the conjugate base.


Dissolution: In a solution, we dissolve a solute in a solvent. The solid solute in the solvent has its bonds broken and the ions are dissolved in the solvent.

Factors that change solubility include

Substances that dissolve readily in water

Ksp>1 means the reaction is going towards the right (more ions)

Ksp<1 means the reaction is going towards the left (back to solid)

Free Energy & the Equilibrium Constant

Free Energy : The amount of free energy available in the products and the reactants and the reaction occuring as we move from reactants to products.

 = -RT(lnK)

*R is the gas constant