Continued from episode 1
1) Oxidation – Reduction reaction
Ø Half reaction = It explicitly shows the electrons involved in a redox reaction.
Ø Oxidation reaction = Half reaction that involves loss of electrons
Ø Reduction reaction = Half reaction that involves gain of electrons
Ø Reducing agent = It donates electrons.
Ø Oxidizing agent = It accepts electrons.
Ø Oxidation number (Oxidation state) = The number of charges the atom would have in a molecule (or an ionic compound) if electrons were transferred completely.
2) Type of redox reaction
Ø Combination reaction = A reaction in which two or more substances combine to form a single product.
Ø Decomposition reaction = The breakdown of a compound into two or more components
Ø Combustion reaction = A reaction in which a substance reacts with oxygen, usually with the release of heat and light to produce a flame.
Ø Displacement reaction = An ion (or atom) in a compound is replaced by an ion (or atom) of another element. It can be divided into three groups: hydrogen displacement, metal displacement, and halogen displacement.
Ø Disproportionation reaction = An element in one oxidation state is simultaneously oxidized and reduced.
3) Concentration of solution
Ø Molarity (M, Molar concentration) = The number of moles of solute per liter of solution (mol/L).
Ø Dilution = The procedure for preparing a less concentrated solution from a more concentrated one
Ø Quantitative analysis = The determination of the amount or concentration of a substance in a sample
4) Gravimetric analysis
Ø Gravimetric analysis = An analytical technique based on the measurement of mass.
5) Acid – Base titration
Ø Titration = A technique as quantitative studies of acid – base neutralization reaction
Ø Standard solution = A solution of accurately known concentration
Ø Equivalence point = The point at which the acid has completely reacted with or been neutralized by base
Ø Indicator = Substances that have distinctly different colors in acidic and basic media
Ø End point = It occurs when indicator changes color.
6) The gas law
Ø The pressure – volume relationship: Boyle’s law = The pressure of a fixed amount of gas at a constant temperature is inversely proportional to the volume of the gas.
Ø The temperature – volume relationship: Charles’s and Gay – Lussac’s law = The volume of a fixed amount of gas maintained at constant pressure is directly proportional to the absolute temperature of gas.
Ø The volume – amount relationship: Avogadro’s law = At constant pressure and temperature, the volume of a gas is directly proportional to the number of moles of the gas present.
7) The ideal gas equation
Ø Ideal gas equation = The relationship among the four variables P, V, T, and n.
Ø Ideal gas = A hypothetical gas whose pressure – volume – temperature behavior can be completely accounted for by the ideal gas equation.
Ø Standard temperature and pressure (STP) = The conditions 0 °C and 1 atm
8) Dalton’s law of partial pressure
Ø Partial pressure = The pressures of the individual gas components in the mixture
Ø Dalton’s law of partial pressure = The total pressure of a mixture of gases is just the sum of the pressures that each gas would exert if it were present alone.
9) Gas diffusion and effusion
Ø Diffusion = The gradual mixing of molecules of one gas with molecules of another by virtue of their kinetic properties
Ø Effusion = The process by which a gas under pressure escapes from one compartment of a container to another by passing through a small opening.
10) Graham ‘s law of diffusion
Ø Graham ‘s law of diffusion = Under the same condition of temperature and pressure, rates of diffusion for gases are inversely proportional to the square roots of their molar masses.
11) Types of energy
Ø Energy = The capacity to do work
Ø Work = Directed energy change resulting from a process
Ø Thermal energy = The energy associated with the random motion of atoms and molecules
Ø Chemical energy = It is stored within the structural units of chemical substances.
Ø Potential energy = Energy available by virtue of an object’s position
12) Energy changes in chemical reaction
Ø Heat = The transfer of thermal energy between two bodies that are at different temperatures
Ø Thermochemistry = The study of heat change in chemical reaction
Ø Surrounding = The rest of the universe outside the system
Ø Open system = It can exchange mass and energy, usually in the form of heat with its surroundings.
Ø Closed system = It allows the transfer of energy (heat) but not mass.
Ø Isolated system = It does not allow the transfer of either mass or energy.
Ø Exothermic process = Any process that gives off heat (transfer thermal energy to the surroundings)
Ø Endothermic process = Heat has to be supplied to the system by the surroundings
Ø Thermodynamics = The scientific study of the interconversion of heat and other kinds of energy
Ø The first law of thermodynamics = Energy can be converted from one form to another but cannot be created or destroyed.
Ø Entropy = A measure of how spread out or dispersed the energy of a system
Ø The second law of thermodynamics = The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process.
Ø The third law of thermodynamics = The entropy of a perfect crystalline substance is zero at the absolute zero of temperature.
Ø Calorimetry = The measurement of heat changes
Ø Specific heat (s) = The amount of heat required to raise the temperature of one gram of the substance by one degree Celsius. [J/g.°C]
Ø Heat capacity (C) = The amount of heat required to raise the temperature of a given quantity of the substance by one degree Celsius. [J/ °C]
15) Properties of waves
Ø Wavelength = The distance between identical points on successive waves
Ø Frequency = The number of waves that pass through a particular point in one second
Ø Amplitude = The vertical distance from the midline of a wave to the peak or trough
16) Photoelectric effect
Ø Photoelectric effect = A phenomenon in which electrons are ejected from the surface of certain metals exposed to light of at least a certain minimum frequency
17) Bohr ‘s theory of hydrogen atom
Ø Emission spectra = Continuous or line spectra of radiation emitted by substances
Ø Line spectra = The light emission only at a specific wavelengths
Ø Ground state (Ground level) = The lowest energy state of a system
Ø Excited state (Excited level) = It is higher in energy than the ground state.
18) Quantum mechanic
Ø Heisenberg uncertainty principle = It is impossible to know simultaneously both the momentum p (defined as mass times velocity) and the position of a particle with certainty.
Ø Electron density = The probability that electron will be found in a particular region of an atom
Ø Atomic orbital = The wave function of an electron in an atom
19) Quantum number
Ø Quantum number = The distribution of electrons in hydrogen and other atoms
· The principal quantum number (n)
· The angular momentum quantum number (l)
· The magnetic quantum number (ml)
· The electron spin quantum number (ms)
20) Electron configuration
Ø Electron configuration = How the electrons are distributed among the various atomic orbitals, in order to understand electronic behavior
Ø Pauli exclusion principle = No two electrons in an atom can have the same set of four quantum numbers.
Ø Paramagnetic substance = It contains net unpaired spins and are attracted by a magnet.
Ø Diamagnetic substance = It does not contain net unpaired spins and is slightly repelled by a magnet.
Ø Hund ‘s rule = The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins.
Chang, R. and Goldsby, K.A. (2014). Chemistry (Eleventh edition). New York: McGraw-Hill Education.
Hornby, A.S. (2010). Oxford Advanced Learner’s Dictionary (Eighth edition). New York: Oxford University Press.