Chemistry Is the Study of Chemicals and Their Properties, and the Reactions They Undergo

Chemistry is the study of chemicals and their properties, and the reactions they undergo.
Matter has mass and volume:

Matter

Pure Substances or Mixtures

Compound or ElementHomogenous or Heterogenous

(2+ elements) (unique #p+) (uniform) (non-uniform)

Physical properties/changes involve changes in state, color, etc. that do not involve a change in the identity of a chemical
Chemical properties/changes describe a chemicals potential to react, and the changes in identity.
Evidence of a chemical reaction: color change, gas evolution, precipitate formation, release of energy (heat)
Density is mass per unit of volume (g/ml)
Mass is the amount of matter; weight is the force of gravity on an object
SI units (Systeme Internationale); the rest of our units are derived from these basic units

length / l / meter / m
mass / m / kilogram / kg
time / t / second / t
temperature / T / kelvin / K
amount of substance / n / mole / mol
electric current / I / ampere / A
luminosity / Iv / candela / cd

Derived Units

force / newton / N / - / m·kg·s-2
pressure, stress / pascal / Pa / N/m2 / m-1·kg·s-2
energy, work, quantity of heat / joule / J / N·m / m2·kg·s-2

Conversions: the key to any conversion is to determine first what you have and then what you need or want. To set up the problem, simply multiply the data you are given by the conversion factor. The units of the conversion factor are what you need over what you have.
Metric table

Mega / M / 106 / 1,000,000
Kilo / K / 103 / 1,000
Base / - / 100 / 1
Deci / d / 10-1 / 0.1
Centi / c / 10-2 / 0.01
Milli / m / 10-3 / 0.001
Micro /  / 10-6 / 0.000 001

Energy is the capacity to do work and comes in 6 forms: chemical, electrical, light, mechanical, sound, heat

Heat is the exchange of energy from high to low; temperature is the measure of how hot or cold something is

Law of Conservation of Energy: energy is conserved
Heat of Fusion and Heat of Vaporization are at the melting and boiling points of a compound respectively
Specific Heat is the amount of heat needed to raise one gram of a compound by 1 degree Celsius or Kelvin (cp=q/m.T)
Significant figure rules:
Non-zero digits are always significant.
Any zeros between two significant digits are significant.
A final zero or trailing zeros in the decimal portion ONLY are significant.
0.0001 has only 1 significant figure
0.00010 has 2 significant figures
1.0001 has 5 significant figures
200 has only 1 significant figure
200. has 3 significant figures
When quantities are added or subtracted, the number of decimal places in the answer is equal to the number of decimal places in the quantity with the smallest number of decimal places
In multiplication and division, the result may have no more significant figures than the factor with the fewest number of significant figures.
Accuracy is how close data is to the true value; Precision is how close the data in repetitive trials are
Endothermic reactions absorb heat to proceed; Exothermic reactions release heat
Reactant + heat  Productdecrease in temperature of surroundings
Reactant  Product + heatincrease in temperature of surroundings
System: the entire group or sample set; Surroundings: the environment outside of the system
Law of Conservation of Mass: Mass is conserved

Democritus: Atoms are the smallest fraction of a molecule

Three Laws of Atoms
Law of Definite Proportions
Law of Multiple Proportions
Law of Conservation of Mass
Dalton’s Atomic Theory:
All matter is composed of atoms that cannot be subdivided, created or destroyed
Atoms of an element have the same properties
Atoms of different elements have different properties
Compounds are formed when different atoms combine in simple, whole numbers
In a chemical reaction, atoms combine, separate or rearrange, but are not created, destroyed or changed

Atoms are made up of subatomic particles called protons, neutrons and electrons
J.J. Thomson set up a crookes tube with a anodic and cathodic ends
When electricity was applied to the tube, a beam was emitted from the cathodic (-) plate
He tested the tube further by applying an electrical field to the tube using paddles
He concluded that the particles in the tube were negatively charged and had mass
electrons, e, e-, -1.602 x 10-19C; mass = 9.109 x 10-31kg
J.J. Thomson’s Plum Pudding Model
Robert Milikan calculated the charge of an electron
Ernest Rutherford conducted experiments to test the Thomson model
He directed alpha particles through a thin gold foil and measured them with a film
There must be a dense region with positive charges (protons) surrounded by the electrons; nucleus
Rutherford concluded that the nucleus had all of the positive charge and most of the mass, but only a fraction of the volume
Most of the volume of an atom is empty space
The Atomic Number is the number of protons in the nucleus and is unique for each element
The Atomic Mass Number is the number of Protons + Neutrons
Atomic Mass Number = Atomic Number + Number of Neutrons
Different elements can have the same mass number, but different atomic number
You cannot use mass number to identify an element
The structure of the atom can be written using the element symbol, atomic number and mass number
An Isotope has the same atomic number but different mass number (different number of neutrons)
Isotopes can be identified with the symbol, mass and atomic number
3/2He or Helium-3
The Rutherford model described electrons as a cloud of negatively charged particles surrounding the nucleus
If electrons are negatively charged, and the nucleus is positively charged, why don’t the electrons collide with the nucleus? Opposite charges attract?
Niels Bohr was a quantam physicist who described energy levels of an electron
The Bohr Model is probably familar as the "planetary model" of the atom, the figure is used as a symbol for atomic energy
In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun
Louis deBroglie – responsible for connecting the wave theory of particles to the modern atomic theory
Schrodinger Wave Theory: wave-particle duality; electrons exhibit both particle and wave like properties
Plank stated that energy is absorbed and released in discrete units called quanta

An orbital is a region of space that an electron is most likely to occur
S-orbitals: are spherical in shape and there is only 1; there are on each principal energy level
P-orbitals: there are 3 types of p-orbitals starting on the 2nd level
D-orbitals: there are 5 types starting on the 3rd level
F-orbitals: there are 7 types starting on the 4th level
Think of orbitals as different types of rooms, each room can only hold 2 electrons

Floor #
(principal energy level) / Type of Orbitals / Number of Orbitals / Maximum Number of Electrons
1 / S / 1 / 2
2 / S, P / 4 / 8
3 / S, P, D / 9 / 18
4 / S, P, D, F / 16 / 32
n / n types / n2 orbitals / 2n2 electrons

Electron Hotel

4th Floor / 


 / F orbital types
D orbital types
P orbital types
S orbital types / 14 electrons +
10 electrons +
6 electrons +
2 electrons =
32 electrons total
3rd Floor / 

 / D orbital types
P orbital types
S orbital types / 10 electrons +
6 electrons +
2 electrons =
18 electrons total
2nd Floor / 
 / P orbital types
S orbital types / 6 electrons +
2 electrons =
8 electrons total
1st Floor /  / S oribital types / 2 electrons

Quantum numbers are unique numbers assigned to each individual electron which indicate the location and the energy of an electron

The principal quantum number (n = 1, 2, 3, 4 ...) is the principal energy level
Hund’s Rule: electrons entering an orbital-type (sublevel) will half-fill the orbitals in the sublevel before they fill it completely
Pauli’s Exclusion Principle: no two electrons can exist in the exact same state; with the same quantum number
Aufbau’ Principle: electrons occupy the lowest possible energy level
An electron configuration is the arrangement of electrons in an atom
The diagonal rule: fill each orbital type according to the direction of the arrows

6s6p6d

5s5p5d5f

4s4p4d4f

3s3p3d

2s2p

Henry Mosely – periodic law; elements are arranged in the periodic table by their atomic number = # of protons = # of electrons
The most modern form of the periodic table was introduced by Mendeleev.
Elements are arranged in columns (group or families) and rows (periods)
Families have similar characteristics; they are said to have periodicity
New elements can be predicted based on these trends

Main group elements:
Group 1: Alkali Earth Metals
Group 2: Alkaline Earth Metals
Group 17: Halogens
Group 18: Noble Gases
Groups 3-12: Transition Metals
Lanthanides and Actinides are the “f” orbital type elements
Metals are on the left; non-metals are on the right; metalloids are along the stairstep
Trends

Ionization Energy, Ionic Radius and Electronegativity

Atomic Radius

Ionic Compounds are formed from a metal and non-metal; the difference in electronegativity is greater than 2.1
Most ionic compounds are solid at room temperature because of the strength of the bond
An ionic compound contains repeating units or formula units bound together in a crystal lattice
Lattice energy is the amount of energy needed to form or break an ionic bond
Ionic compounds dissociate in water to form ions; these ions conduct electricity
Electrons are exchanged in order to form a compound
Ionic compounds are named with the cation (“+”) and then the anion (“-“)
The anion is the name of the element + “-ide”
If the anion is a polyatomic ion, simply name the ion
Ions and parent atoms have different properties

Octet rule; rule of 8: all compounds want to be more like a noble gas, they want an electron configuration like a noble gas (8 valence electrons).
Therefore, they will lose/gain electrons to be stable

Covalent Compounds are formed from a metal to metal or a non-metal and non-metal; the difference in electronegativity is less than 2.1
Polar covalent bonds are between 0.5 and 2.1
Non-polar covalent bonds are less than 0.5
Electrons are shared in order to form a compound
As bond length increases, bond energy decreases
A triple bond is the shortest and strongest of all the bonds
Bond energy is the amount of energy needed for form or break a covalent bond
Covalent compounds are named with the less electronegative atom first followed by the more electronegative atom. Prefixes are used to indicate ratios of atoms

Lewis dot structures illustrate the configuration of the valence (outer shell) electrons
The shape of a bond can predict the overall polarity of a compound (VSEPR or Valence Shell Electron Pair Repulsion)
If there is a dipole, or an uneven shift of charge (remember that a compound cannot have an overall charge, so there is an unequal sharing of the + and – charges), then the compound is polar
Polar dissolved polar and vice-versa

The Mole is a counting unit; 6.022 x 1023 molecules, particles or atoms
The atomic mass number, atomic weight/mass or molecular mass is the mass of 1 mole of atoms (by the way, you do know where to find the molecular mass? Try the periodic chart!)
In Stoichiometry, the key to solving the quandry is the mole
A balanced chemical equation indicates in simple, small, whole numbers the ratio of compounds to each other
If you are given the mass or the volume of only one of the compounds in a chemical reaction, the remaining information can be calculated or derived using the principles of stoichiometry
Here’s the deal: in a perfect world and on a perfect planet, compounds will come together an react in simple whole numbers. We, however, do not exist in said perfect world. Therefore, we must rely on the simple, small, whole number ratio of compounds
Example:

6CO2 + 6H2O + energy  C6H12O6 + 6O2

(by the way, this is photosynthesis; look familiar in reverse?)

Okay, let’s read this reaction. Six moles of carbon dioxide and 6 moles of water and energy (in the form of light) react to form 1 mole of glucose and 6 moles of oxygen
The ratio of these compounds is 6:6:1:6
So, if I start out with 51.4 grams of CO2 divided by the molecular mass which is 44 g/mol I get a total of 1.17 moles of CO2 to start with.
So then, the ratio looks something like this: 1.17 : 1.17 : 0.195 : 1.17
How, you might ask? Well, it’s called a ratio, and we divide. 6:1 is the same thing as saying 1 : 1/6
Anyway, that’s how you figure out exactly how many moles of each compound you have. Then from there, you can multiply by the molecular mass to get actual mass
Yield is the amount of product that you would expect to recover from a specified amount of reactant. Theoretical because we perform the experiment on paper first. The limiting factor is the reactant that theoretically produces the least amount of product.

Reaction types:
Synthesis: make a new compound from elements or simple substances
Decomposition: break a complex compound into smaller or elemental parts
Combustion: a hydrocarbon with oxygen as a fuel burn or combust to form CO2 and H2O
Single replacement: only one element or polyatomic ion is replaced, exchanged or displaced. The order of this reaction depends upon the activity chart
Double replacement: two elements or polyatomic species are replaced; the product must include a solid; this can be determined by the solubility guidelines

Molar Heat Capacity (C) is the heat required to raise 1 mol of a substance by 1 degree K or C. C = q / n .T ; and
Specific heat (cp) is the heat required to raise 1 gram of a substance by 1 degree K or C. cp = q / m .T; so,
C = cp. M (where M is the molar mass in g / mol)

Thermodynamics is the study of energy and energy is the capacity to do work.

The 1st Law of Thermodynamics is also the Law of Conservation of Energy: Energy cannot be created or destroyed, it is conserved in a chemical reaction.

Enthalpy (H) is the amount of heat absorbed or released in a chemical reaction and is measured in kJ / mol.

Enthalpy of formation Hof is the enthalpy to form a compound from its elements at standard thermodynamic conditions; 25oC and 1 atm.
Therefore, an element has an enthalpy of formation of “0”.

Calorimetry is the experimental measure of an enthalpy change and is measured in kJ (calorie or kilocalorie is another unit)

In an Exothermic Reaction, the change in enthalpy from the reactants to the product, results in a negative change in enthalpy. The potential energy for the products is less than the potential energy for the reactants.

In this reaction, the energy is on the product side (A + B  C + 22 kJ)

Hess’s Law states that the amount of heat released or absorbed does not depend on the number of steps; the sum of all changes in enthalpy equals to the net change in enthalpy.

In an Endothermic Reaction, the change in enthalpy from the reactants to the product, results in a positive change in enthalpy. The potential energy for the products is greater than the potential energy for the reactants.
In this reaction, the energy is on the reactant side (A + B + 30 kJ  C)

Entropy (S) is the measure of disorder in a system and is measured in J / K.

The following events lead to an increase in disorder (entropy):
Increase in temperature;
Phase change from solid to liquid to gas;
More products than reactants;
Simpler products than reactants;
Substances that are put into solution;
Solutions that become more dilute (molecules have more space);
Gases that the pressure decreases (increases the volume and space).

S = Sproducts - Sproducts
Molar entropy (So) is the entropy of one mole of a compound.

Gibb’s Energy (G), also called Gibb’s Free Energy, is the energy in a system available for work and is measured in kJ / mol.

The amount of Gibb’s energy predicts a spontaneous reaction which is a reaction that occurs without any additional energy. A negative G is spontaneous and a positive G is not spontaneous;

If, H =and S =thenG= and the reaction is

-+-spontaneous

--?spontaneous, low T

++?spontaneous, high T

+-+not spontaneous

The factors that affect Gibb’s energy are enthalpy, entropy and temperature

G = H - TS; and
G = Gproducts - Greactants

The 2nd Law of Thermodynamics states that in a spontaneous reaction, entropy is not conserved and must always increase. All spontaneous reactions contribute a portion of energy to perform useful work (Gibb’s Energy).
As a reaction proceeds, G decreases until it equals “0” which is the point of equilibrium.

The states of matter are solid, liquid, gas and sometimes plasma (which is not a naturally occurring state)
Cohesion is an attraction for particles that a liquid has.
Adhesion is an attractive force for particles of solid surfaces
Capillary action is the motion of a liquid up a small surface and is accomplished by adhesion of liquid molecules to the surface of the glass as well as cohesion between the liquid molecules.
Surface tension is the fore that acts of the surface of a liquid and tends to minimize the area of the surface. Why?
1st of all, cohesive forces bring the molecules of a liquid together so that they stay in contact;
2nd, under the surface of the liquid, these cohesive forces are pulling equally in all directions;
3rd, only on the surface, the molecules are being pulled sideways and downward creating surface tension.
It takes energy to increase the surface area of a liquid because this energy must oppose the net forces pulling the molecules; conversely, a liquid decrease energy as the surface area decreases. This tendency toward decreasing the surface area is called surface tension. A high surface tension means that a lot of energy is needed to break the surface.
Application: surface tension is used in laundry. When you put dirty clothes in the washing machine and no laundry detergent, the dirt on your clothes cannot penetrate the surface tension of the water and so it stays on your clothes. When you add the detergent, the soap decreases the surface tension by disrupting hydrogen bonds and therefore the dirt can be carried away by the water!!! Although, it also takes effort to load the washing machine!
Gases do not have the same type of intermolecular forces, because they are farther apart and the attractive forces are minimized. That is why a gas will fill the space available.
Now let’s talk about the physical changes in the states of matter. Refer to the diagram below. The direction of the arrow indicates the change from one state to another by the addition of release of energy. As energy is absorbed, the state changes from solid to liquid; liquid to gas; or solid to gas. As energy is released the state changes in the opposite direction. The name of each process is labeled.

Gas