A1. International System of Units (SI). Base Units and Derived Units

A1. International System of Units (SI). Base units and derived units.

Fractions and multiples of base units.

The International System of Unit is the modern form of the metric system. It is the world’s most widely used system of units.

The reference standards and the definitions of base and derived quantities and units together make up the International

system of units. (A reference standard is the physical description or embodiment of a base unit.)

The units of SI can be divided into 2 subsets, base unit and derived unit.

There are seven base units, and each of these base units are nominally dimensionally independent.

length(m), mass(kg), time(s), temperature(K), mole(mol), (electric current(A), Luminous intensity(cd))

From these seven base units, several other units are derived, which means the derived units can be defined in terms of the base units, for example, liter is derived from m3.

SI prefix can be attached to the names of the base units to express multiples or submultiples of these units. Using prefixes

we can put very large or very small numbers into scientific notation with the same unit.

They are in fractions based on 10, letting them be divisible by 10 makes many calculations much easier.

A2. The conservation laws. Chemical equations.

Law of conservation of mass

In any chemical reaction, the sum of the masses of the reactants always equals the sum of the masses of the products.

Law of conservation of energy

The total energy of the universe is constant and can neither be created nor destroyed, it can only be transformed.

Chemical equations are used to describe reactions. The symbols of the substances involved as reactants, separated by plus

signs, are on one side of an arrow that points to the symbols of the products, also separated by plus signs.

The equation is balanced when all atoms showing in the formulas of the reactants are present in like numbers in the formulas

of the products.(Coefficients, numbers standing in front of formulas, are employed as needed to achieve the correct

balance.)

In both formulas and equations, the numbers used for subscripts and coefficients are generally the smallest whole numbers that show the correct proportions.

A3. The characteristics of states of matter, transition between states.

Matter, anything with mass that occupies space, can exist in three physical states, solid liquid, and gas.

A solid has both a definite volume and a definite shape.

A liquid has a definite volume but no fixed shape.

A gas has no definite volume and no fixed shape.

(Physical changes brought about by heating or cooling convert a given sample of matter into its different states.)

A4. The definitions and types of energy and system.

Energy is the ability to cause change in motion, position, illumination, sound, or chemical composition.

Energy is a conserved quantity, meaning that it cannot be created or destroyed but only converted from one form into

another.

Energy is a scalar quantity because it has no direction in space.

The SI unit of energy is the joule(J), equals 1N applied through 1m, for example.

System is defined as the matter within a definite region of space, and separated from surrounding (the rest of the universe)

by a boundary which may be imaginary.

The possible exchange of work, heat, or matter between the system and surroundings take place across this boundary.

Isolated system; matter and energy don’t cross the boundary.

Closed system; matter don’t cross the boundary.

Open system; matter and energy cross the boundary.

A5. Extensive and intensive properties. The heat of vaporization. Of formation. Of solution.

Properties are called extensive or intensive according to their dependence on sample size.

Extensive properties are directly proportional to the size of the sample, for example mass, volume, length, heat etc.

Intensive properties are independent of the sample’s size, temperature, color, density, pressure, concentration etc.

Heat of vaporization is the energy required to transform a given quantity of a substance into gas.

Values are usually quoted in kJ/mol, although kJ/kg, kcal/mol, cal/g are also possible.

A6. Exothermic, endothermic reactions. The sign of heat. Energy diagrams.

Exothermic reactions are the reactions that continuously release heat from the system to the surroundings, and most (but

not all) spontaneous reactions are exothermic. (ex. Combustion or burning)

(Spontaneous reactions are those that, once arranged or started, continue with no further human intervention.)

Endothermic reactions are the reactions that require a continuous input of heat from the surroundings to the system. (ex. Photosynthesis)

Heat can change an object’s temperature, or its physical state.

Heat, symbolized by Q, is the energy that transfers from one object to another when the two are at different temperature and in some kind of contact.

The SI unit of heat is the joule as it is a form of energy, but also calorie(cal), an older unit of heat, is still used commonly.

1cal is the energy needed to increase the temperature of 1g of water by 1℃, and this is about 4.184 joules.

A7. The enthalpy change of a chemical reaction. The law of hess.

Enthalpy, symbolized by H, is the change of the reaction heat at constant pressure.

Enthalpy change is defined as the enthalpy of the products minus the enthalpy of the reactants.

If the enthalpy change is positive, the reaction is endothermic.

If the enthalpy change is negative, the reaction is exothermic.

The Law of Hess states that if a given reaction can take place in several ways, then the sum of heat change (the enthalpy

change of the reaction) is independent of the order and type of the subreactions, it only depends on the type and state of

the reactants and products.

In other words, only the start and end states matter to the reaction, not the individual steps between.

A8. The change of entropy and free energy, direction of the reactions.

Entropy is the function of measure the disorder, and is central to the second law of thermodynamics, which deals with physical processes and whether they occur spontaneously.

(The second law of thermodynamics states that the total entropy of any isolated thermodynamic system tends to increase over time.)

Gibbs free energy is freely available energy for work in the system.

The change of free energy is defined as ΔG＝ΔH－TΔS, and it predicts the direction of the reaction.

If ΔG is negative, the reaction is exothermic and the reaction will follow.

If ΔG is positive, the reaction is endothermic and the reaction will not follow.

A9. Avogadro’s number. The mole concept, the different types of concentration expressions.

The number of atoms in 12g of the carbon-12 isotope is called Avogadro’s number, and it equals 6.02×1023 atoms.

This many formula units of any pure chemical substance constitutes 1 mole of the substance.

Equal numbers of moles contain identical numbers of formula units.

Percent weight per weight denotes the mass of a substance in a mixture as a percentage of the mass of the entire mixture. (% w/w)

Percent weight per volume (% m/v or % w/v) describes the mass of the solute in g per mL of the resulting solution.

Molarity (M) denotes the number of moles of a given substance per liter of solution. (mol/L)

A10. Components of solutions, their types. The types of solutions, solubility.

A solution is made of a solvent and one or more solutes.

The solvent is the medium into which the other substances are mixed or dissolved.

The solute is anything that is dissolved by the solvent.

A solution can be described as dilute or concentrated according to its ratio of solute to solvent being small or large.

Whether a solution is unsaturated, saturated, or supersaturated depends on its ability to dissolve any more solute at the same temperature.

The solubility

Each substance has a particular solubility in a given solvent at a specified temperature, and this is often expressed as the grams of solute that can be dissolved in 100g of the solvent.

The solubility is the maximum quantity of solute that can dissolve and form a stable solution at the given temperature in the given solvent.

A11. The concentration units.

See A9

A12. The physical properties of water. Water as a solvent.

The higher electronegativity of oxygen over hydrogen and the angularity of the water molecule make it polar, so polar that

hydrogen bonds exist between its molecules.

Hydrogen bonding helps to explain many of water’s unusual thermal properties, such as its relatively high boiling point, its high heats of fusion and vaporization, and its high surface tension.

Water dissolves best those substances whose ions or molecules can strongly attract water molecules.

The hydration (the association of water molecules with dissolved ions or polar molecules) of ions or polar molecules helps them to dissolve in water because water molecules are very polar and have sizable partial charges.

Cations or δ+ sites on polar molecules are surrounded by water molecules whose δ- ends point toward the positively charge.

Anions or δ- sites on polar molecules are surrounded by water molecules whose δ+ ends point toward the positively charge.

A13. The solution process of solid crystalline substances in water. The effect of temperature and pressure of the dissolution process.

When a solid crystalline is dropped into water, their surfaces are instantly bombarded by water molecules, an action that

works to dislodge the ions.

The hydration (the association of water molecules with dissolved ions or polar molecules) of ions or polar molecules helps them to dissolve in water because water molecules are very polar and have sizable partial charges.

Cations or δ+ sites on polar molecules are surrounded by water molecules whose δ- ends point toward the positively charge.

Anions or δ- sites on polar molecules are surrounded by water molecules whose δ+ ends point toward the positively charge.

The solubilities of most solids increase with temperature, because their dissolving is usually endothermic.

The solubilities of some ionic compounds decrease with increasing temperature. (SO42- with metal, Ca(OH)2)

Pressure affects the solubilities only when the solutes are gases.

A.14. Electrolytes, ionization and dissociation. Degree of dissociation, weak and strong electrolytes

Electrolytes is any substance whose solution in water conducts electricity or the solution itself of such a substance.

Dissociation is the separation of preexisting ions from one another as an ionic compound dissolves or melt,

But Ionization is the formation of ions by a chemical reaction.

Strong electrolytes; one that is strongly dissolved or ionized in water. A high percentage ionization. (NaCl, HCl)

Weak electrolytes; one that is weakly ionized in water. A low percentage ionization. (acetylic acid, NH3)

Nonelectrolytes; one that does not dissolve or ionize in water. Essentially zero percentage ionization. (gasoline)

A15. Liquid-vapor equilibria, vapor pressure of pure mater, boiling point, freezing point.

Liquid-vapor equilibrium

The rate of evaporation (the change from liquid state to vapor state) and condensation (the change from vapor state to liquid state) are identical.

The vapor pressure of a pure liquid

The pressure exerted by a vapor that in equilibrium with its liquid state at a given temperature.

A liquid’s vapor pressure can be regarded as the escaping tendency of its molecules.

When the liquid’s vapor pressure equals the pressure of the atmosphere, the liquid boils.

Boiling point

The temperature at which a substance boils when the atmospheric pressure is 760mmHg (1atm)

Freezing point

The temperature at which a substance changes state from liquid to solid.

When considered as the temperature of the reverse change from solid to liquid, it is referred to as the melting point of a crystalline solid.

A16. Solution of gases in water. Henry’s law. The effect of pressure and temperature on the solution of gases.

Temperature, pressure and sometimes reaction with water affect the solubility of a gas in water.

All gases are less soluble in water at higher temperature, because the dissolving of gases in liquids is always exothermic.

Gases are more soluble under higher pressure.

Henry’s Law (Pressure- solubility law) states that gas solubility is directly proportional to gas pressure.

The chemical factor that affect the solubilities of some gases is their ability to react with water. (CO2, SO2, NH3)

A17. Vapor pressure of solutions, Raoult’s law, boiling point elevation, freezing point depression.

Raoult's law states that the vapor pressure of each chemical component in an ideal solution is dependent on the vapor

pressure of the individual component and the mole fraction of the component present in the solution .

Once the components in the solution have reached chemical equilibrium, the total vapor pressure of the solution is:

Psolution = （P1）pure・X1 + （P2）pure・X2 ・・・

and the individual vapor pressure for each component is Pi = (Pi)pure ・ Xi

where ((Pi)pure is the vapor pressure of the pure component, Xi is the mole fraction of the component in solution

Boiling-point elevation is a colligative property (the properties of dilute solutions of non-volatiles solute whose values just

depend on the concentration of solute particles rather than their (solute) individual properties) that states that a solution

will have a higher boiling point than that of a pure solvent after the addition of a dissolved solute.

The change in boiling point can be determined by the equation ΔTB.P.= i ·Kb ·m, where m is the molality of the

solute(mol/kg), i is the Van 't Hoff factor (the number of dissolved particles the solute will create when dissolved), and Kb is

the ebullioscopic constant unique to each solvent.

Freezing-point depression is the difference between the freezing points of a pure solvent and a solution mixed with a solute.

The change in boiling point can be determined by the equation ΔTf = i ·Kf ·m, where Kf is the cryoscopic constant.

A18. Osmosis and dialysis.

Osmosis is the passage of water only, without any solute, from a less concentrated solution (or pure water) to a more concentrated solution when the two solutions are separated by a semipermeable membrane.

The back pressure needed to prevent osmosis is called the osmotic pressure, symbolized by π, and it is directly proportional to the concentration. From PV=nRT → πV=nRT → π=n/V・RT (n/V=mol/L)

Dialysis is the passage through a dialyzing membrane (more permeable than semipermeable) of water and particles in solution, but not of particles that have colloidal size.