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.
A19. The importance of osmolarity of body fluids, isotonic, hypertonic, hypotonic solutions.
The permeability of blood capillaries changes temporarily when a person experiences shock, and macromolecules leave the
blood. Their departure results in the loss of water, too, and the blood volume decrease.
Isotonic solutions; two solutions of equal osmolarity
Hypertonic solutions; one with a higher osmolarity than another.
Hypotonic solutions; one with a lower osmolarity than another.
When red blood cells are put into hypertonic solution, they loose fluid volume, and shrink (this process is called crenation)
In the case of hypotonic solution, the red blood cell is swollen by extra fluids from solution, and the cell burst. (hemolysis)
So, it is important that only isotonic solutions, or those that are nearly so, should be administered in large quantities
intravenously.
A20. The description of reaction rate 1; The transition state and collision theory.
Transition state theory
Transitional state is temporary the highest energy state of the intermediate.
(the energy needed to achieve transitional state is the activation energy)
In the transitional state, the breaking of old bonds and formation of new bonds are being in progress.
The transitional state theory can explain the role of the temperature.
If the temperature is increased, the energy level of the reactants increases, therefore the actual activation energy is
smaller, the reaction is faster.
The transitional state theory can also explain the role of the catalyst.
The catalyst stabilizes the transitional state, and lowers activation energy, therefore the actual activation energy is
smaller, the reaction is faster.
Collision theory
A theory about the rates of chemical reactions that postulates collisions between reacting particles.
A21. The description of reaction rate 2 ; The factors that affect the reaction rate, law, kinetic constant, activation energy.
(Reaction rate is the number of product-forming collisions that occur each second in each unit of volume of a reacting
mixture. （mol/L・s）
Kinetics helps to deal with the reaction rate and mechanism in a chemical reaction.
Kinetic theory is a set of postulates about the nature of an ideal gas,
(1)that is consist of a large number of very small particles in constant, random motion
(2)that in their collisions the particles lose no frictional energy.
(3)that between collisions the particles neither attract nor repel each other.
(4)that the motions and collisions of the particles obey all the laws of motion.)
The main factors that influence the reaction rate include
The physical state of the reactants.
The concentrations of the reactants.
The temperature at which the reaction occurs.
Whether or not any catalysts are present in the reaction.
Reaction rate can be described,
If the reaction is a simple conversion of A into B. (A→B; mono-molecular reaction)
V=k [A] where v=reaction rate, k=rate constant, [A]=concentration of reactant A (mol/L)
The rate is proportional with the concentration of A.
In the case of bimolecular reactions (A+B→C, 2A→C, 2A+B→C)
V=k [A][B], v=k [A]2, v=k [A]2[B]
The sum of all the exponents of the reactants involved in the rate equation is the order of reaction that tells how much the
reaction rate is dependent on the concentration of the reactants.
Zero order reaction: v=K
First order reaction: v=k [A]
Second order reaction: v=k [A][B] or v=k [A]2
Higher order reaction: v=k [A]2[B]3 (2A+3B→C)
B1. Subatomic particles, their number in neutral atoms.
An atom is made of three subatomic particles, the electrons, the protons, and the neutrons.
All three subatomic particles have masses, the proton and the neutron have the atomic mass unit of 1, the electron has that
of 1/1836.
(We can ignore electron masses when working with the masses of whole atoms.)
The proton has a charge of 1+, the electron has a charge 1-, and the neutron has no charge.
An atom’s protons and neutrons make up its atomic nucleus which has all of the atom’s mass and all of its positive charge.
The sum of an atom’s neutrons and protons is its mass number.
Positive charge on its atomic nucleus, same as the number of protons, is called its atomic number.
Because all atoms are electrically neutral, the number of the electrons equals the atomic number.
B2. Types of chemical formulas, atomic mass, atomic weight, formula weight, molecular formula
A chemical formula (=molecular formula) is a concise way of expressing information about the atoms that constitute a
particular chemical compound. For molecular compounds, it identifies each constituent element by its chemical symbol and
indicates the number of atoms of each element found in each discrete molecule of that compound. If a molecule contains
more than one atom of a particular element, this quantity is indicated using a subscript after the chemical symbol
For ionic compounds and other non-molecular substances, the subscripts indicate the ratio of elements in the empirical
formula.
Molecular and structural formula
Relative atomic mass (=atomic weight, average atomic mass) is the average of the atomic masses of all the chemical
element’s isotopes as found in a particular environment.
Atomic weight (=atomic mass) is the average mass of the atoms of the various isotopes of any given element as they occur
in their natural proportions.
Formula weight (=formula mass) of a compound is the sum of the atomic masses of all the atoms present in one formula unit.
Molecular weight (=molecular mass) is the formula mass particularly for molecular compounds.
B3. Atomic number, mass number, atomic symbols, isotopes.
Atomic number is the number of protons charged positively on its atomic nucleus, also equals the number of electrons.
Mass number is the sum of an atom’s neutrons and protons.
A chemical symbol is an abbreviation or short representation of the name of a chemical
element. Natural elements all have symbols of one or two letters; some man-made elements have
temporary symbols of three letters.
Chemical symbols are listed in the periodic table and are used as shorthand and in chemical Equations.
Because chemical symbols are often derived from the Latin or Greek name of the element, they may not bear much similarity
to the common English name, e.g., Na for sodium (Latin natrium) and Au for gold (Latin aurum).
Chemical symbols may also be changed to show if one particular isotope of an atom that is
specified, as well as to show other attributes such as ionization and oxidation state of a chemical compound.
Isotopes have some numbers of protons but different numbers of neutrons.
The isotopes of the same element have same electron configurations, and so an element’s isotopes have identical chemical
properties.
B4. The quantum theory, quantum numbers and their meaning.
The quantum theory
The electrons shift to higher levels when the correct amount of energy is absorbed, and the emit this energy as a quantum
when they drop back again to lower levels.
The quantum numbers
①The principal quantum number
This number, symbolized by n, corresponds to the distance between the electron and the nucleus.
The average distance increase with n, and quantum states with different principal quantum numbers are said to belong to
different shells.
②The orbital (= azimuthal, angular) quantum number
This number is symbolized by l (l= 0, 1, ・・・, n-1), and it specifies the shape of an atomic orbital, subshells.
The subshell with l=0 is called s orbital, l= 1 a p orbital, l= 2 a d orbital and l= 3 an f orbital.
③The magnetic quantum number (me = -l, -l+1, ・・0・・, l-1, l ) means energy shift.
④The spin quantum number (ms = -1/2 or 1/2 ) means the spin direction of an electron.
B5. The electron configuration of atoms. Ground state, excited state.
The electron configuration is the arrangement of electrons in an atom, molecule, or other physical structure.
The electron configuration is described by using the quantum numbers.
In the case of Phosphorus (atomic number 15), is as follows: 1s2 2s2 2p6 3s2 3p3.
Ground state
All electrons of an atom are in the lowest energy states available.
Exited state
Higher energy states than ground state.
B6. Orbital filling. Pauli’s exclusion principle and Hund’s rule.
Orbital filling
To place the electrons one by one into the available orbitals, starting with the one of lowest energy, the 1s orbital.
According to the Pauli exclusion principle, each orbital can hold 2 electrons, but where 2 or more orbitals are available at the
same subshell, electrons spread out (Hund’s rule)
The order of increasing energy of the subshells can be constructed by going through downward-leftward diagonals of the
table below, going from the topmost diagonals to the bottom.
The first (topmost) diagonal goes through 1s; the second diagonal goes through 2s; the third goes through 2p and 3s; the
fourth goes through 3p and 4s; the fifth goes through 3d, 4p, and 5s; and so on. In general, a subshell that is not "s" is
always followed by a "lower" subshell of the next shell; e.g. 2p is followed by 3s; 3d is followed by 4p, which is followed by 5s,
4f is followed by 5d, which is followed by 6p, and then 7s.
A d subshell that is half-filled or full (ie 5 or 10 electrons) is more stable than the s subshell of the next shell. This is the case
because it takes less energy to maintain an electron in a half-filled d subshell than a filled s subshell. For instance, copper
(atomic number 29) has a configuration of [Ar]4s1 3d10, not [Ar]4s2 3d9 as one would expect by the Aufbau principle.
B7. The periodic table, the aufbau method (the relationship between the electrons structure and
the periodic table)
Because many properties of the elements are periodic functions of atomic numbers, the elementsfall naturally into vertical
groups or families in the periodic table.
The atoms of a group of representative elements, those in group ⅠA－ⅦA, have the same number of outside shell
electrons, a number corresponding to the group number itself.
Of the group 0 elements, all except helium have outside shells with 8 electrons, helium has 2.
Several of the groups among representative elements also have names, except for hydrogen, the elements in group ⅠA
(IUPAC 1) are called the alkali metals.
The elements in group ⅡA (IUPAC 2) are called the alkaline earth metals.
The elements in group ⅦA (IUPAC 17) are called the halogens.
The elements in group 0 (IUPAC 18) are called the noble gases.
Other groups of representative elements are named simply after the first member, the boron family (group ⅣA), the
nitrogen family (group ⅤA), and the oxygen family (group ⅥA).
The horizontal rows of the periodic table are called periods.
In the long periods, there are transition and inner transition elements which involve the systematic filling of inner d or f
orbitals.
The nonmetallic elements are in the upper right-hand corner of the table, and the metals make up the rest of table.
At the border between metals and nonmetals occur the metalloids, which have both metallic and nonmetallic properties.
The aufbau method (= aufbau principle?)
-is used to determine the electron configuration of an atom.
According to the principle, electrons fill orbitals starting at the lowest available energy states before filling higher states.
The number of electrons that can occupy each orbital is limited by the Pauli exclusion principle.
If multiple orbitals of the same energy are available, Hund’s rule says that unoccupied orbitals will be filled before
occupied orbitals are reused.
B8. Types of elements in the periodic table.
Representative elements, transition and inner transition elements →see B7.
Metals, nonmetals and metalloids →see B7.
B9. Atomic size, ionization energies and their trends in the periodic table.
In the periodic table, atomic size increases downward in the same group and right to left in thesame period.
Ionization energies increase upward in the same group of the periodic table, and left to right across a period. (Ionization
energy of an atom is the energy required to strip an electron.)
B10. Types of ions. Ionic radius.
(An ion is an atom or group of atoms that normally are electrically neutral and achieve their status as an ion by loss or
addition of one or more electrons)
Anion is a negatively charged ion which has more electrons than its atom.
Cation is a positively charged ion which has fewer electrons than its atom.
Monatomic ion is an ion consisting of a single atom.
Polyatomic ion is an ion consisting of multiple atoms, for example OH-, NH4+)
The ionic radius
Cations are always smaller than their parent atoms.
Anions are always larger than their parent atoms.
B11. The ionic bond
The ionic bond is a type of chemical bond based on electrostatic forces between two oppositely-charged ions.
The ionic bond is similar in strength to the covalent bond, but is stronger than the hydrogen bond.
B12. The covalent bond
The covalent bond is a type of chemical bond characterized by the sharing of pairs of electrons between atoms.
Shared electron pairs are unequally shared when bonded atoms have different electronegativities.
The covalent bond includes σ-bond (single bond), π-bond (double bond) and triple bonds.
The dative bond (= the coordinate bond) is a special type of covalent bond for which both shared electrons come from one
atom.
B13. The shapes of molecules
The shape of molecule is forced by the repulsions of electron clouds of valence-shell electron pairs which stay out of each
other’s way as much as possible. (The VSEPR theory)
Thus molecular shapes are expected for different numbers of pairs of valence-shell electrons.
When there are four pairs, the result is a tetrahedral geometry or close to it.
(tetrahedral ; pyramid shaped with four triangular faces and four corners)
If the number of shared electron pairs is 4 (no lone pairs), all bond angles are 109.5°. (CH4)
(NH3)
If the number of shared electron pairs is 3 (1 lone pair), the shape is called trigonal pyramidal, and bond angle is 107.3°.
If the number of shared electron pairs is 2 (2 lone pairs), it has nonlinear shape, and bond angle is 104.5°. (H2O)
In the case of less electron pairs than 4,
If there are 2 electron pairs, the shape is linear with bond angle of 180°. (BeCl2)
If there are 3 electron pairs, the shape is planar triangular whose axes are in a plane and point to the corners of a regular
triangle, so the bond angle is 120°. (BCl3)
In the case of more electron pairs than 4,
If there are 5 electron pairs, the shape is trigonal bipyramidal consisting of 2 three-sided pyramids joined by sharing a
common face, the triangular plane through the center. (PCl5)
If there are 6 electron pairs, the shape is octahedral which is an eight-sided figure with six corners consisting of two square
pyramids that share a common square base. (SF6)
B14. Electronegativity and its trend in the periodic table. Polarization of chemical bond.
Electronegativity is the ability of an atom to attract electrons in a chemical bond.
Atoms with similar electronegativities will share an electron with each other and form a covalent bond.
If one atom pulls slightly harder than the other, the bond will be polar.
If the difference is too great, the electron will be permanently transferred to one atom and an ionic bond will form.
Electronegativities increase upward each group and increase across the periods (left to right).
When bonded atoms have different electronegativities, the bond is polar and polar bonds have opposite partial charges at
either end.
Even though molecules are electrically neutral, they can still be polar.
Polar molecules are the molecules that have not only individual polar bond caused by electronegativities differences
between the joined atoms, but also have the angular structure (non-symmetrical). Examples, H-F, H2O, NH3
Non polar molecules are the molecules that have only non polar bonds, or have symmetrical structures even if each bond is
polar. Examples, H2, CCl4
B15. The polarity of compounds, weak intermolecular forces.
Polarity → see B14
Intermolecular forces
①Hydrogen bond
Hydrogen bond is a force of attraction between the partial positive charge (δ+) on H in the polar bonds of the H-O, H-N, or
H-F systems to the partial negative charge (δ-) of another O, N, or F.
②Dipole-dipole interaction
The attraction between permanentδ+ andδ- site on separate polar molecules.
③Dipole-induced dipole interaction
The attraction between the permanent dipole molecule and temporary dipole molecule.
④Induced dipole or Van der Waals interaction
The attraction between the hydrophobic molecules.
B16. Dative bond and its significance in living organisms.
The dative bond (= the coordinate bond) is a special type of covalent bond for which both shared electrons come from one
atom.
The metal ion which carries out the coordinate bond to other atoms combines with the side chain of amino acid, such as
histidine, glutamic acid, and cysteine, and is fixed to the inside of protein. Such protein is called metal content protein. Many
things which contain metal ion in the portion which causes the enzyme reaction, and many of enzymes perform the enzyme
reaction through metal ion. On the other hand, iron ions exist in hemoglobin and oxygen conveyance is performed because
oxygen combines with the iron ion.
B17. The hydrogen bonding and its significance in living organisms.
Examples of the hydrogen bond in living organisms
In proteins, the hydrogen bonds help to make secondary, tertiary, and quarterly structures.
In DNA, hydrogen bonds are made between the complimentary base pairs (A-T, G-C).
B18. Lewis structures. Single and multiple bonds.
Lewis structures are also called electron-dot structures that valence shell electrons are shown as dots placed around the
atomic symbol.
When atoms form a molecule by sharing electrons to achieve noble gas configuration, the shared electrons pair is shown
between the 2 atomic symbols.
Single bond is a bond with 1 shared electrons pair. (CH4)
Double bonds are bonds with 2 shared electrons pairs. (H2C=CH2)
Triple bonds are bonds with 3 shared electrons pairs. (N2)
B19. The energetic and kinetic description of equilibria. Le Cahtelier’s principle.
Dynamic equilibrium
A situation in which 2 opposing events occur at identical rates so that no net change happens.
To represent a dynamic equilibrium, we use an equation with double arrows.
Le Chatelier’s principle
If a system in equilibrium is upset by a stress, such as the addition or removal of heat, the system shifts in whichever
direction most directly absorbs the stress and restores equilibrium.
B20. The mass action law.
The rate of a chemical reaction is directly proportional to the product of the effective concentrations of each participating
molecule.
Also see A21.
B21. The dissociation of water, the pH concept.
B22. The ionization of weak electrolytes, dissociation constant, ways of pH calculation.