Chemistry - Semester 2
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Chapter Summaries: Select the chapter you wish to view by clicking below. (Some summaries taken from Chemistry The Central Science, Brown)
Section 7 - Electronic Structure
Section 9 - Chemical Equilibrium
Useful Links in the Understanding of Chemistry
Section
6-
Behavior of Gases
Textbook Chapters
Objectives
Relate the various units of pressure,
volume, and temperature to one another.
Use the Ideal Gas Law to determine the effect of a change in conditions upon a particular variable
Use the Ideal Gas Law to solve for on variable, given the values of the others.
Use the Ideal Gas Law to calculate the molecular mass of a gas, knowing the mass of a given volume (or the density) at known P and T
Use the Ideal Gas Law
to calculate the density of a gas at a given temperature and
pressure.
Relate the volumes of
gases, measured at the same temperature and pressure, involved
in a chemical reaction.
Use Dalton's Law to
obtain partial pressures of gases in mixtures.
Relate the mole fraction
of a gas to its partial pressure.
Use Graham's Law to
relate the molecular masses of two gases to their rates of effusion.
Calculate the average speed of molecules of a particular gas at a given temperature.
Basic Concepts
This chapter has been concerned with the gaseous state. To describe the state of condition of a gas, it is necessary to specify four variables: pressure, temperature, volume, and quantity of gas. Volume is usually measured n liter (L), and temperature in the Kelvin scale. Pressure is expressed in SI units as pascals (Pa), or more commonly in millimeters of mercury (mm Hg). one standard atmosphere pressure equals 760 mm Hg. A barometer is often used to measure the atmospheric pressure.
The ideal gas equation, PV = nRT, is the equation of state for an ideal gas. Most gases at pressures of about 1 atm and temperature of 30K and above obey the ideal gas law reasonably well. We can use the ideal gas equation to calculate variations in one variable when one or of the more of the others are changed. For example, for a constant quantity of gas at constant temperature, the pressure of the gas is inversely proportional to the volume (Boyle's law). Similarly, for a constant quantity of gas at constant pressure, the volume of the gas is directly proportional to temperature (Charles's law). In gas mixtures, the total pressure is the sum of the partial pressures that each gas would exert if it were present alone under the same conditions. (Dalton's law of partial pressures). We also covered in detail the following gas laws:
In all applications of the ideal gas law we must remember to convert temperatures to the Kelvin temperature scale.
It is important to use the ideal gas equation to solve problems involving gases as reactants or products in chemical reactions. From the gas density, r, or the gas mass, g, it is possible to calculate the molecular mass of the gas.
The kinetic molecular theory accounts for the properties of an ideal gas in terms of a set of assumptions about the nature of gases. These assumptions are:
molecules are in ceaseless, chaotic motion
the volume of the gas molecule is negligible in relation to the volume of the their container
the gas molecules have no attractive forces for one another
the average kinetic energy of the gas is proportional to absolute temperature
Sample Problems
Sample problems can be found in the gas law links above.
Section
7-
Electronic Structure
Objectives
Relate the wavelength
of a spectral line to the change in energy in an atom.
Use the Bohr theory
to calculate the energy of an electron in a given principal energy
level of the hydrogen atom, or the difference in energy between two
levels.
Determine the number of
electrons that may be accommodated by any given principal energy level or
sublevel.
Given the atomic number
of an element, write the electron configuration of its isolated
gaseous atom in the ground state.
Given or having written the electron
configuration of an atom, draw its orbital diagram and
electron dot notation.
Give the four quantum
numbers corresponding to each of the various electrons in an atom.
Given the position of an
element in the Periodic Table, write its outer electron
configuration.
Use the Periodic Table
to predict relative values of atomic radius and ionization energy.
Basic Concepts
Radiant
energy mover through a vacuum at the "speed of light," c =
3.00 x 108
m/s: it has wavelike characteristics that allow it to be
described in terms of
wavelength (l)
and
frequency (n):
these ate interrelated: c = ln. The dispersion
of radiation into its component wavelengths produces a
spectrum.
If all wavelengths are present, the spectrum is sad to be continuous, if only
certain wavelengths are present, it is called a
line
spectrum.
The quantum theory describes the minimum amount of radian energy that an object can gain or lose, E = hn; this smallest quantity is called a quantum. A quantum of radiant light is called a photon. The quantum theory was used to explain the photoelectric effect and the line spectrum of the hydrogen atom. The absorption or emission light by an atom, which produce its line spectrum, correspond to energy changes of electrons within the atom; the energy of the electron in an atom is quantized.
Electrons exhibit wave properties and can be described by a wavelength, l = h/mv. Discovery of the wave properties of the electron led to the Werner Heisenberg's uncertainty principle, which indicates that the position and momentum of a particle can determined simultaneously with only a limited accuracy.
The idea described above culminated in our
current model of the electronic structure atoms in which we speak of the
probability of the electron being found at a
particular point in space.
This mathematical model was first developed by Erwin
Schrodinger. Although the position of the electrons are not defined, except in this average
sense, their energies are precisely known. Each allowed energy state of
the electron in an atom corresponds to a particular set of values for three
quantum
numbers. Each of the allowed energy state is termed an
orbital.
An orbital is described by a combination of an integer and letters,
corresponding to the three values for the quantum numbers. The
principle
quantum number (n) is indicated by the integers 1, 2, 3 ...
This quantum number relates most directly to the size and energy of an
orbital. The
azimuthal quantum number
(l) is indicated by the letters s, p, d, f and so on. corresponding to values of
l of 0, 1, 2, 3, ... The l quantum number defines the shape of the
orbital. The
magnetic quantum number
(ml) describes the orientation of the orbital in space.
Restrictions in the values of the three quantum numbers gives rise tot he following allowed subshells:
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
5s, 5p, 5d, 5f
6s, 6p, 6d
7s, 7p
There is one orbital in an s subshell, three in a p subshell, five in a d subshell, and seven in an f subshell.
Representations the electron structure found in an atom can be shown with electron configuration, orbital notation, or electron dot notation.
The periodic function of the elements is controlled by the electronic structure of valence shell electrons in the atoms and can be demonstrated by ionization energy and electron affinity.
Sample Problems
Section
8-
Chemical Bonding
Objectives
Basic Concepts
Chemical bonding deals with the interactions between atoms that lead to the formation of molecules. Ionic bonding results from the complete transfer of electrons from one atom to another, with formation of a 3-dimensional lattice of charged particles. The stabilities of ionic substances result from the powerful electrostatic attractive force between an ion and all the surrounding ions of opposite charge. These interaction are measured by the lattice energy.
Covalent bonding results from the sharing of electrons between atoms. The rules that govern this sharing are based on the stability of the rare-gas electron configuration (the octet rule). We can represent shared electron-pair structure of molecules by means of Lewis structures, which show the sharing of electron pairs between atoms The sharing of one pair of electrons produces a single bond, the sharing of two pairs of electrons between atoms produces double and triple bonds.
It is important to recognize that even is covalent bonding, electrons may not be shared equally between two atoms. Electronegativity is a measure of the ability of an atom to compete with other atoms for the electrons shared between them. Highly electronegative elements strongly attract electrons. The difference is electronegativities of bonded atoms is used to determine the polarity of a bond.
The bond types we concentrated on are listed below:
Ionic bonds
Covalent bonds
Polar covalent bonds
Coordinate covalent bonds
Metallic bonds
Explanations of each bond type can be found a the chemical bonding page
Sample Problems
Section
9-
Chemical Equilibrium
Objectives
Given the balanced
equation for a reaction involving one or more gases, write the
corresponding expression for the equilibrium constant Kc (Ke).
Given the equation for a
reaction, calculate Kc knowing the equilibrium
concentration of all species.
Given the equation for a
reaction, calculate Kc knowing the original
concentration of all species and the equilibrium concentration of
one.
Given the value of Kc
and all the original concentrations, calculate the equilibrium
concentration.
Given the value of Kc
and all but one equilibrium concentration, calculate the remaining equilibrium
concentration.
Given the value of Kc,
predict the direction in which a chemical system will move to reach equilibrium
and determine the equilibrium concentrations.
Using Le Chatelier's Principle, predict the effect of adding (or removing) a reactant
or product,
Using Le Chatelier's Principle, predict the effect of a changing the volume, or changing the temperature upon the composition of an equilibrium system.
Basic Concepts
If the reactants and products of a reaction are kept in contact, a chemical reaction can achieve a state of dynamic balance in which forward and reverse reactions are occurring at equal rates. This condition is known as chemical equilibrium. A system at equilibrium does not change with time, For such a system, the ratio of products to reactants, each concentration raised to the power corresponding to the coefficient in the balanced equation, is called the equilibrium constant, K : K = [products]/[reactants]. The equilibrium constant changes with temperature but is not affected by changes in relative concentrations of any reacting substance or by pressure or the presence of a catalysts. In heterogeneous equilibria, the concentrations of pure solids or liquids are absent from the equilibrium expression.
LeChatelier's principle indicates that if we disturb a system that is at equilibrium, the equilibrium will shift to minimize the disturbing influence. The effects of adding (or removing) reactants or products, and of changing pressure, volume, or temperature can be predicted using the principle. The equilibrium constant changes value with changes in temperature. Catalysts affect the speed at which equilibrium is reached but do not
affect the equilibrium constant.
Sample Problems
Section
10-
Solutes in Water
Objectives
Given the formula of a
substance, decide whether it is likely to be an electrolyte of a
nonelectrolyte in water solutions.
Predict the relative
solubilties of different solutes in water.
Predict the effect on
solubility of a change in temperature or pressure.
Write a balanced equation
for the formation of a solution or precipitate.
Given or having derived the
equation for a precipitation reaction, relate the amounts of reactants and
products.
Use the limiting laws
pertaining to boiling point elevation and freezing point depression for
solution of nonelectrolytes to relate the boiling point (freezing point),
molality, and Kb (Kf).
Use the limiting laws
pertaining to boiling point elevation and freezing point depression for
solution of nonelectrolytes to relate the boiling point (freezing point),
molality, mode of ionization (or value of i), and Kb (Kf).
Use the limiting laws (Skill
6) as modified for solutions of electrolytes to relate the boiling point
(freezing point), molality, and Kb (Kf).
Use the limiting laws (Skill
6) as modified for solutions of electrolytes to relate: the boiling point
(freezing point), molality, mode of ionization (or value of i), and Kb
(Kf).
Basic Concepts
Solutions are homogeneous mixtures of atoms, ions, or molecules. The relative amounts of solute and solvent in a solution can be described qualitatively (dilute or concentrated) or quantitatively (weight percentage, molarity, molality, normality, and mole fraction).
The extent to which a solute will dissolve in a particular solvent depends on the relative magnitudes of solute-solute, solvent-solvent, and solvent-solvent attractive forces as well as on the changes in disorder accompanying mixing. The rule "like dissolves like" was found to be useful in rationalizing solubilities. It is possible to change the solubility of a solute by changing temperature and pressure. If the solution process is endothermic, an increase in temperature promotes solubility. With a gas, an increase in pressure promotes solubility. These effects can be understood in terms of LeChatelier's principle.
Substances that exist in solutions as ions are called electrolytes. Those substances that are completely ionized are called strong electrolytes. Reactions occur between electrolytes if an insoluble substance, age, or a nonelectrolyte can form. Net ionic equations focus attention on the particular species that actually undergo some change during the reaction.
The presence of a solute in a solvent lowers the vapor pressure and the freezing point and increases the boiling point of the solvent. These changes are termed colligative properties. The magnitude of the change depends on the total concentrations of solute particles in solution and not on their characteristics. Osmotic pressure is the pressure that must be applied to a solution to prevent the transfer of solvent molecules from a pure solvent through a semi permeable membrane. It is a colligative property because it is proportional to the total concentration of solute particles. Colligative properties can be used to calculate the molecular mass of nonvolatile nonelectrolytes.
Sample Problems
Section
11-
Acids and Bases
Objectives
Given one of the concentrations
[H+] or [OH-] for water or a water
solution, calculate the other concentration.
Given either [H+]
or [OH-], calculate the pH.
Write an equation for the dissociation
of a strong acid or strong base in water solution.
Write an equation for the reversible
for the dissociation of a weak acid in water solution; and the
reaction of a weak base with water.
Predict whether a given
ionic compound will give an acidic, basic, or neutral
solution. Write an equation to explain your prediction.
Write an equation for the
reaction of an acid with a base. Describe the solution
that results as being acidic, basic, or neutral.
Use titration data
for an acid-base reaction to determine the concentration of an
acid or base in water solution and the molecular mass of an acid or base.
Select an acid-base
indicator appropriate to a given acid-base titration.
Classify a given species in
a reaction as an acid or base, according to the models of Arrhenius; Bronsted
and Lowry; and Lewis. Indicate the conjugate acid-base
pairs.
Basic Concepts
In this chapter we have considered the general properties of acidic and basic solutions, with emphasis on water as the solvent. We have seen that an acid solution is created when a substance reacts with water in such a way as to increase the concentration of solvated hydrogen ions, which are represented as H+(aq) or H3O+(aq). The concentration of H+(aq) is often expressed on the pH scale: pH = -log [H+]. Solutions of pH less than 7 are acidic; those with pH greater than 7 are basic.
Water spontaneously ionizes to a slight degree (autoionization), forming H+(aq) and OH-(aq). The extent of ionization is expressed by the Water Dissociation constant for water: Kw = [H+][OH-] = 1.0 x 10-14. This relationship describes not only pure water, but aqueous solutions as well. Because the concentration of water is effectively constant in dilute solutions, [H2O] is omitted from this equilibrium constant expression as well as from others associated with reactions in aqueous solutions.
Through most of this chapter we have relied on the Bronsted-Lowery theory of acids and bases. According to this theory, an acid is a proton (H+) donor, a base a proton acceptor. Reaction of an acid with water results in the formation of H+ (aq) and the conjugate base of the acid. Strong acids have conjugate bases that are weaker than H2O. Such acids are strong electrolytes, ionizing completely in solution. The common strong acids are HCl, HNO3, HClO3, and H2SO4. Weak acids are substances for which the reaction with waster is incomplete, and an equilibrium is established. The extent to which the reaction proceeds is expressed by the acid-dissociation constant, Ka. Polyprotic acids are acids such as H2SO3 that have more than one ionizable proton.
The stronger the acid the weaker its conjugate base; the weaker the acid the stronger its conjugate base. This qualitative observation is expressed quantitatively by the expression Ka x Kb = Kw (where Ka and Kb are dissociation constants for conjugate acid-base pairs.).
The acid-base properties of salts can be ascribed to the behavior of their respective cations and anions. The reaction of ions with water with a resultant change in pH is called hydrolysis. The cations of strong bases (alkali metal ions and alkaline earth metal ions) and the anions of strong acids do not undergo hydrolysis.
In the Lewis theory of acids and bases, the emphasis is on the shared electron pairs rather than on the protons. An acid is defined as an electron pair acceptor; a base as an electron pair donor.
Sample Problems