AP Chemistry/The Basics
You should remember everything here from your high-school level chemistry class.
- 1 Units and Measurement
- 2 States of Matter
- 3 History of Chemistry
- 4 Modern Atomic Theory
- 5 Electrons
- 6 VSEPR Theory
- 7 Hybrid Orbitals
- 8 Sigma and Pi Bonds
- 9 Resonance
- 10 The Periodic Table
- 11 The Quantum Numbers
- 12 Oxidation Numbers
- 13 Naming Compounds
Units and Measurement
- Fahrenheit is not used on the AP exam. Celsius (°C) and Kelvin (K) are used. Pure water freezes at 0° Celsius (273K) and boils at 100°C (373K).
- Digits 1 through 9 are significant, and so are zeroes in between them. For example, the number 209 has three significant figures.
- Zeroes to the right of all other digits are only significant if there is a decimal point written. 290 has 2 sig figs, 290. has three, and 290.0 has four.
Measured vs. Exact Numbers
Exact numbers are either defined numbers or result from a count. A dozen is defined as 12 objects. A pound is defined as 16 ounces. Measured numbers always have a limited number of significant digits. They are certain up to a limited number of digits. A mass reported as 12 grams is implied to be known to the nearest gram and not to the tenth of a gram.
The Mole and Avogadro's Number
12 grams of Carbon-12 contain exactly one mole (6.02 * 10^23) of molecules. This is a measured number known as Avogadro's number. It is easy to convert between atomic mass, grams, and particles using Avogadro's number.
Multiplying and Significant Figures
Multiplying measured numbers in Chemistry is not like multiplying in math. 5 * 92 equals 460 in math class, but it equals 500 in chemistry. This is because the 5 only has one significant figure, so the answer has to be rounded to one sig fig. If 5.0 and 92 were multiplied, on the other hand, the answer would be 460 in both subjects.
Adding and Significant Figures
- First, align all the numbers vertically, as if you were going to add them. DO NOT WRITE IN EXTRA ZEROS AS PLACEHOLDERS.
- Round to the smallest place that contains a digit in every number.
Example: 210 + 370. + 539
210 370. 539
1119 ≈ 1120
States of Matter
- Solid (s) - definite shape and volume. Vibrates in place, but does not flow.
- Fluids - take the shape of their container.
- Liquid (l) - definite volume
- Gas (g) - variable volume (compressible)
History of Chemistry
- Democritus - philosopher who made the idea of atoms.
- Antoine Lavoisier - discovered Law of Conservation of Mass, which states that mass does not appear or disappear in chemical reactions, only rearrange.
- John Dalton - first scientist to scientifically describe atomic theory.
- Matter is made from indestructible particles called atoms.
- Atoms of the same element are the same.
- Compounds are two or more atoms bonded together.
- Chemical reactions are the rearrangement of atoms.
- J.J. Thomson - discovers the electron.
- Robert Millikan - discovers the mass and charge of electrons.
- "Raisin Pudding model" - atoms are like pudding, with electrons as raisins.
- Ernest Rutherford - through his gold foil experiment, discovers the nucleus. Since most of the alpha particles he shot through the gold were not deflected, he concluded that most of an atom is empty space.
Modern Atomic Theory
Atoms are made up of protons, neutrons, and electrons. Protons and neutrons weigh approximately 1 AMU, and electrons have a negligible mass. Elements are determined by the number of protons in the atom, known as the atomic number. The number of neutrons varies, creating different isotopes of different mass. The atomic mass of an atom is the sum of its protons and neutrons, both of which are found in the nucleus.
Electrons are arranged into shells that surround the atom. Each shell has 1-4 subshells, which themselves have 1-7 orbitals, each of which holds two electrons.
|2||s, p||1 + 3 = 4|
|3||s, p, d||1 + 3 + 5 = 9|
|4||s, p, d, f||1 + 3 + 5 + 7 = 16|
Filling Electron Shells
- Aufbau principle - fill the lowest energy subshells first, in accordance with the following image:
- Hund's rule - fill each orbital in a subshell with one electron before putting a second electron in any of those orbitals.
Writing Electron Configurations
E.g. sodium = 1s22s22p63s1
Electrons in a compound will try to move as far apart from each other as possible. Bonded pairs repel more strongly than unbonded pairs.
The moving apart of electron pairs requires the hybridization of orbitals. These hybrids range from sp to sp3d2, having two to six pairs.
Sigma and Pi Bonds
- Sigma bond - forms in all compounds
- Pi bonds - one or more are formed per extra electron pair that is shared among two atoms. These bonds are weaker than sigma bonds.
Sometimes, there is more than one "correct" way to draw a substance. In reality, the structure of the substance is an average of the drawn variations.
The Periodic Table
You should already be familiar with this. Each row is called a period, and each column is a group or family. Nonmetals and metals are separated by a jagged line on the right side. (Hydrogen is also a non-metal). Elements that border the line are called metalloids, and share characteristics with both metals and nonmetals.
The Quantum Numbers
These four numbers are used to describe the location of an electron in an atom.
|Principal Quantum Number|
|Angular Momentum Quantum Number|
|Magnetic Quantum Number|
|Spin Quantum Number|
Principal Quantum Number (n)
Determines the shell the electron is in. The shell is the main component that determines the energy of the electron (higher n corresponds to higher energy), as well as nuclear distance (higher n means further from the nucleus). The row that an element is placed on the periodic table tells how many shells there will be. Helium (n = 1), neon (n = 2), argon (n = 3), etc.
Angular Momentum Quantum Number (l)
Also known as azimuthal quantum number. Determines the subshell the electron is in. Each subshell has a unique shape and a letter name. The s orbital is shaped like a sphere and occurs when l = 0. The p orbitals (there are three) are shaped like teardrops and occur when l = 1. The d orbitals (there are five) occur when l = 2. The f orbitals (there are seven) occur when l = 3. (By the way, when l = 4, the orbitals are "g orbitals", but they (and the l = 5 "h orbitals") can safely be ignored in general chemistry.)
This number also gives information as to what the angular node of an orbital is. A node is defined as a point on a standing wave where the wave has minimal amplitude. When applied to chemistry this is the point of zero-displacement and thus where no electrons are found. In turn angular node means the planar or conical surface in which no electrons are found or where there is no electron density.
Here are pictures of the orbitals. Keep in mind that they do not show the actual path of the electrons, due to the Heisenberg Uncertainty Principle. Instead, they show the area where the electron is most likely to occur (say, 90% of the probability). The two colors represent the two different spin numbers (the choice is arbitrary).
|S orbital →|
|P orbitals →|
|D orbitals →|
|F orbitals →|
Magnetic Quantum Number (ml)
Determines the orbital in which the electron lies. For example, there are three p orbitals in shell n = 2: the magnetic quantum number determines which one of these orbitals the electrons reside in. The different orbitals are oriented at different angles around the nucleus. See how each p orbital has the same general shape, but they point in different directions around the nucleus.
Spin Quantum Number (ms)
Determines the spin on the electron.
Let's examine the quantum numbers of electrons from a magnesium atom. Remember that each list of numbers corresponds to (n, l, ml, ms).
|Two s electrons:||(1, 0, 0, +½)||(1, 0, 0, -½)|
|Two s electrons:||(2, 0, 0, +½)||(2, 0, 0, -½)|
|Six p electrons:||(2, 1, -1, +½)||(2, 1, -1, -½)||(2, 1, 0, +½)||(2, 1, 0, -½)||(2, 1, 1, +½)||(2, 1, 1, -½)|
|Two s electrons:||(3, 0, 0, +½)||(3, 0, 0, -½)|
The Periodic Table
Notice a pattern on the periodic table. Different areas, or blocks, have different types of electrons. The two columns on the left make the s-block. The six columns on the right make the p-block. The large area in the middle (transition metals) makes the d-block. The bottom portion makes the f-block. Each row introduces a new shell (aka energy level). Basically, the row tells you how many shells of electrons there will be, and the column tells you which subshells will occur (and which shells they occur in). The value of ml can be determined by some of the rules we will learn in the next chapter. The value of ms doesn't really matter as long as there are no repeating values in the same orbital.
|To see the pattern better, take a look at this periodic table.|
Oxidation numbers are a way of keeping track of electrons and making sure that components of a compound match by the correct ratios.
- Pure elements have an oxidation number of zero.
- Ions, monoatomic or polyatomic have oxidation numbers equal to their charge.
- The sum of the oxidation numbers in covalent and ionic compounds must equal zero.
- Bonded Group 1 metals are +1, Group 2 are +2, and halogens are -1, unless bonded with oxygen.
- Bonded oxygen is -2 unless it is in a hydroxide (OH), where it is -1, or with fluorine, where it is positive.
- Bonded hydrogen is -1 when bonded with a metal and +1 when bonded with a nonmetal
(First element's name) (Second element's name + ide) e.g. Sodium Chloride.
Hydro(nonmetal+ic) acid. E.g. Hydrobromic acid (HBr)
(First element's name) (polyatomic ion's name) e.g. Sodium Hydroxide (NaOH). Note that there is an exception - the ammonium ion (NH4+) can replace the first element.
In the following order: (hydrogen)(a nonmetal)(oxygen)
If the ion ends in -ate, the acid will be named (nonmetal)ic acid. Example: H2SO4 contains a sulfate ion. It is called sulfuric acid.
If the ion ends in -ite, the acid is named (nonmetal)ous acid
Some elements, especially transition metals, can have many oxidation numbers. As a result, the positive oxidation number has to be written in, using Roman numerals. For example, CuO is Copper (II) oxide and Cu2O is Copper (I) oxide.