Gas laws describe relationships between molecules in the gas phase. Gas phase particles can be thought of as individual particles moving randomly in an enclosed container. There are four main laws that govern the behavior of gases.
Charles’s law describes the relationship between volume (V) and temperature (T) for a gas at constant pressure (P). Volume and temperature are directly related.
As V ↑, then T↑; as V↓, then T↓
If a system experiences a change in volume or temperature, a proportional relationship between the initial and final conditions is given by
V1/T1 = V2/T2
Boyle’s law describes the relationship between volume (V) and pressure (P) for a gas at a constant temperature (T). Volume and pressure are inversely related.
As V ↑, then P↓; as V,↓ then TP↑
If a system experiences a change in volume or pressure, a proportional relationship between the initial and final conditions is given by
P1V1 = P2V2
Avogadro’s law describes the relationship between volume (V) and moles of gas (n) for a gas at constant temperature (T) and pressure (P). Volume and moles of gas are directly related.
As V ↑, then n↑; as V↓, then n↓
If a system experiences a change in volume or moles of gas, a proportional relationship between the initial and final conditions is given by
V1/n1 = V2/n2
The combined gas law incorporates all of the relationships in the previous three laws to describe the relationship between volume (V), pressure (P), and temperature (T) for a given gas sample (constant moles).
If a system experiences a change in V, P, or T parameters, a proportional relationship between the initial and final conditions is given by
P1V1/T1 = P2V2/T2
Endothermic vs. Exothermic
Chemical reactions transfer heat as molecular bonds break and form between atoms. The amount of heat that is transferred depends on the strength of the bond.
Exothermic reactions are chemical reactions that release heat, which means that the reactants are higher energy than the products. When the products form, the extra energy is transferred to the environment as heat energy. The reaction coordinate diagram shows the general progression of an exothermic reaction.
Endothermic reactions are chemical reactions that absorb heat, which means that the reactants are lower energy than the products. When the products form, the extra energy is transferred from the environment as heat energy. The reaction coordinate diagram shows the general progression of an endothermic reaction.
History of the Periodic Table
The periodic table is foundational to all chemical science. It organizes the elements, pure atomic substances, in a way that creates periodic trends and groups with similar properties.
Discovery of Elements
The discovery of phosphorus by Robert Boyle in 1680 is believed to be the first published finding of a pure element. Over a century later in 1789, Antoine Lavoisier defined an element as a substance that does not break down to a simpler substance in a chemical reaction. Lavoisier named oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc, and sulfur and classified metal and nonmetal elements.
Classification and Organization
Dimitri Mendeleev, a Russian chemist, is widely acknowledged as the father of the modern periodic table. In 1869 Mendeleev arranged the known 63 elements by atomic weight. This new organization left spaces for undiscovered elements to be added in the future. In the late 1890s, the discovery of new inert gases by William Ramsay and Lord Rayleigh prompted the addition of an additional noble gases group (He, Ne, Ar, Kr, Xe, Rn) to the periodic table.
Rare Earth Elements
In 1905 a Swiss chemist named Alfred Werner solved another part of the periodic puzzle. Rare earth elements (lanthanides and actinides) were difficult to place in the original table because their number and size were still largely unknown. Werner expanded the periodic table by adding new columns for the lanthanide rare-earth elements. Actinide rare-earth metals were added several decades later after experiments by chemist Glenn Seaborg and physicist Edwin McMillan, both Americans.
The most recent additions to the periodic table include four new elements added in 2015 (nihonium-113, moscovium-115, tennessine-117, and oganesson-118). New elements are created by fusing atomic nuclei to form superheavy atoms. These artificial elements are very unstable and may only exist for a fraction of a second before they undergo radioactive decay.
Types of Bonds
Intramolecular bonds are chemical bonds that connect atoms within a molecule. Intramolecular bonds are stronger than the intermolecular bonds between molecules.
There are three main types of intramolecular bonds:
- Covalent bond – valence electrons between adjacent atoms are shared
- Single covalent bond – one pair of electrons are shared, as in methane (CH4)
- Double covalent bond – two pairs of electrons are shared, as in diatomic oxygen (O2)
- Triple covalent bond – three pairs of electrons are shared, as in diatomic nitrogen (N2)
- Ionic bond – valence electrons between adjacent atoms are transferred
- Creates a strong electrostatic attraction between the positively charged cation (electron donor) and the negatively charged anion (electron acceptor)
- Examples include salts like sodium chloride (NaCl), potassium iodide (KI), and magnesium oxide (MgO)
- Metallic bond – valence electrons are delocalized throughout the metal
- Electrostatic forces between the positively charged metal nuclei and the negatively charged electrons hold adjacent metal nuclei in place
- Conduction of electric current from a flow of charged electrons through the metal
- Examples of metallic bonding include pieces of solid metal like Silver (Ag), Copper (Cu), and Sodium (Na).
Kinetic Molecular Theory – Arrhenius Equation
Kinetic molecular theory explains how the behavior of particles in the gas phase are related to the observable quantities of pressure, volume, and temperature. Kinetic molecular theory makes a number of assumptions about gas molecules.
Assumptions of Kinetic Molecular Theory
- Gas particles are in constant, random motion, and they move in straight lines until they collide with other particles or the sides of the container.
- Gas particles are inert and do not experience any attractive or repellent forces.
- Collisions between gas particles are completely elastic.
- The size of the gas particles is negligible compared to the volume of the container.
- The average kinetic energy of the gas particles depends solely on temperature. Therefore there is no molecular motion in a system at a temperature of 0 K.
The Arrhenius equation uses the assumptions of kinetic molecular theory to calculate the reaction rate constants for chemical reactions.
The Arrhenius equation is expressed by
k = Ae-Ea/RT
k ⇒ reaction rate constant
A ⇒ pre-exponential factor
Ea ⇒ activation energy (J/ mol)
R ⇒ universal gas constant; 8.314 J/mol K
T ⇒ temperature (K)
Since the pre-exponential factor is a constant and RT is kinetic energy, the reaction rate is inversely proportional to the ratio of activation energy to kinetic energy.
Near room temperature (293 K), reaction rate doubles with approximately a 10-degree rise in temperature. Use the Arrhenius equation as a guide to show that raising the temperature from 290 to 300 can cause doubling for a chemical reaction with an activation energy of 65000 j/mol.
- Start with the Arrhenius equation:
k = Ae-Ea/RT
- Summarize the equation to focus on the dependent terms:
k ~ e-Ea/RT
- Substitute the relevant values for each temperature and compare the results:
kT=290 ~ e-Ea/RT
= e-65000/(8.31 x 290) = 1.932 x 10-12 = k @ 290 K
kT = 300 ~ e-Ea/RT
= e-65000/(8.31 x 300) = 4.749 x 10-12 = k @ 300 K
The rate at the higher temperature is about 2.4 times the rate at the lower temperature.
Nomenclature – Acids
Naming conventions for acids depend on the number of atoms and the identity of the conjugate base.
- If the acid is diatomic, use the prefix “hydro-” and the suffix “-ic acid.”
- If the acid is polyatomic AND the conjugate base anion ends in “-ate,” use the base name and the suffix “-ic acid.”
- If the acid is polyatomic AND the conjugate base anion ends in “-ite,” use the base name and the suffix “-ous acid.
Molarity is a measure of concentration expressed by units of moles of solute per liter of solution (mol/L or M). Molar concentration conveys the number of solute molecules in the solution. Moles can be converted to grams using molar mass.
Find the molarity, rounded to the nearest integer, of a solution that contains 361 g glucose (C6H12O6) dissolved in 200 mL of water.
- Calculate the moles of glucose in 1.4 g glucose:
molar mass C = 12.010 g/mol
molar mass H = 1.007 g/mol
molar mass O = 15.999 g/mol
molar mass glucose = 6(12.010 g/mol) + 12(1.007 g/mol) + 6(15.999 g/mol)
= 180.128 g/mol
? mol glucose = 361 g glucose (1 mol glucose/180.128 g)
= 2.004 mol glucose
2. Divide by solvent volume in liters:
200 mL = 0.2 L
2.004 mol glucose/0.2 L water = 10 M
The solution is about 10 M glucose in water.
Enthalpy is a thermodynamic principle that describes the total capacity of a system to release heat. The SI unit for enthalpy is joule per kilogram (J/Kg), which is the energy released as heat per kilogram of substance. In equations, such as the Gibb’s free energy equation, enthalpy is denoted by H and a change in enthalpy is denoted by ΔH.
ΔG = ΔH – TΔS
The Gibbs free energy equation relates the change in enthalpy to the total free energy change of a system. In addition to enthalpy, the total change in free energy also depends on the temperature (T) multiplied by the change in entropy (ΔS), another thermodynamic principle.
If a system undergoes a positive change in enthalpy (ΔH > 0), then the system has gained heat energy from the environment.
If a system undergoes a negative change in enthalpy (ΔH < 0), then the system has lost heat energy to the environment.