21.6 Biological Effects of Radiation

• Increased concerns over the effects of radioisotopes on biological systems have been caused by the increased use of radioisotopes.
• High-energy particles are emitted by all radioactive nuclides.
• Radiation can cause heating, break chemical bonds, or ionize molecule in living cells.
• When radioactive emissions fragment or ionize, it is the most serious biological damage.
• Nuclear decay reactions emit particles with higher energies than ordinary chemical bonds.
• When these particles strike and penetrate matter, they produce ion and molecular fragments that are very reactive.

• Damage to the DNA of cells can be caused by radiation.
• The cells may divide in an uncontrollable manner if the damage is not repaired.

• The lower-energy radiation is nonionizing and the higher-energy radiation is ionizing.

• The energy absorbed from nonionizing radiation is equivalent to heating a sample.
• Although biological systems are sensitive to heat, a large amount of nonionizing radiation is necessary before dangerous levels are reached.
• Ionizing radiation may cause more severe damage by breaking bonds or removing electrons in biological molecule, disrupting their structure and function.

• The hydroxyl radical has an unpaired electron.
• This radical can cause damage to the molecule and disrupt the processes of the body.

• Ionizing radiation can cause damage to a biomolecule by breaking its bonds or creating an H2O+ ion, which reacts with H2O to form a hydroxyl radical, which in turn reacts with the biomolecule.

• The whole body can be harmed by radiation, as well as eggs and sperm.
• The effects are more pronounced in cells that reproduce rapidly, such as the stomach lining, hair follicles, bone marrow, and embryos.
• This is the reason why patients undergoing radiation therapy often feel sick to their stomach, lose hair, and so on, and why particular care must be taken when they are pregnant.

• Different types of radiation have different abilities.
• A sheet of paper or the top layer of skin cells can be used to stop alpha particles.
• Alpha particle sources are not dangerous if outside the body, but they are hazardous if swallowed or breathed in.
• A thin layer of metal stops the particles from passing through a hand or a thin layer of material.
• The radiation can penetrate a thick layer of materials.
• A small amount of high-energy radiation can be seen through a few feet of concrete.

• Lead is one of the dense, high atomic number elements that can be used for shielding.
• Some particles have no tendency to produce ionized particles, while other emissions have a tendency to cause ionized particles.
• Alpha particles have more ionizing power than fast- moving neutrons, b particles, and g rays.

• The radiation's ability to pass through material is shown.
• They are alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha, alpha,

• One of the largest sources of exposure to radiation is from radon gas.
• The half-life of Radon-222 is 3.82 days.
• The radioactive decay series of U-238 is found in trace amounts in soil and rocks.
• The radon gas that is produced slowly escapes from the ground and slowly creeps into homes and other structures above.

• Houses and other buildings are affected by Radon-222 because of the rocks that contain the radioactive substance.

• Depending on where you live, Radon is found in buildings across the country.
• About one in six houses in the US have high levels of the radioactive gas in their air, so it's important to get rid of it.
• Smoking a carton of cigarettes a day is as bad for your health as exposure to radon, which increases one's risk of lung cancer.
• There are two leading causes of lung cancer, one of which is Radon.
• Over 20,000 deaths in the US are believed to be caused by Radon exposure.

• The gas is ionized in a tube by radiation.
• The amount of radiation affects the rate of ionization.
• There are many types of dosimeters.

• Radiation can be measured with devices such as Geiger counters, scintillators, and dosimeters.

• Various units are used to measure radiation.
• The energy and biological effects of radiation are taken into account.

• The number of rems is related to the number of rads with RBE, the number of protons and neutrons is related to the number of rems, and the number of b and g radiation is related to the number of rems.

• The rate of emission from a radioactive source, the amount of energy absorbed from the source, and the amount of damage done by the absorbed radiation are measured in different units.

• The table summarizes the units used for measuring radiation.

• Since the g rays it emits can be focused in small areas where the cancer is located, 1/2 is used in cancer therapy.
• A sample of Co-60 can be used for cancer treatment.

• It has several uses, including self-powered lighting, in which electrons emitted in tritium radioactive decay cause phosphorus to glow.
• The atomic mass of tritium is 3.016 amu.

• The annual radiation exposure for a person in the US is about 600 mrem.
• The bar graph shows the various sources and their relative amounts.

• A short-term, sudden dose of a large amount of radiation can cause a wide range of health effects, from changes in blood chemistry to death.
• The risk of death from short-term exposure to tens of rems of radiation is 50%; a dose of 500 rems is estimated to cause the death of the victim within 30 days.
• It is important to avoid unnecessary exposure to radiation because it can have a cumulative effect on the body during a person's lifetime.
• Table 21.5 shows the health effects of shortterm exposure to radiation.

• Exposure to ionizing radiation is impossible to avoid.
• Cosmic radiation, rocks, medical procedures, consumer products, and even our own atoms are some of the natural sources of background radiation that we are exposed to.
• blocking or shielding the radiation, moving farther from the source, and limiting the time of exposure are some of the ways we can minimize our exposure.

• An atomic nucleus is made up of protons and neutrons.
• The nucleus is held together by a very strong force called the strong nuclear force.
• The total mass of the nucleus is less than the total mass of the nucleons.
• Only a small number of nuclides are stable.
• It is likely that the nucleus with even numbers of protons or neutrons will be stable.
• A graph of number of protons versus number of neutrons shows a narrow band of stability for these nuclides.
• The most stable nuclei have mass numbers around 56 and have the largest binding energy per nucleon.

• The number of protons, number of neutrons, and energy state can be changed.
• Nuclear reactions can involve many different particles.
• The most common are alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, alpha particles, Nuclear reactions are always balanced.
• The total mass and charge remain the same when a nuclear reaction occurs.

• Nuclei with unstable n:p ratios are prone to radioactive decay.
• The most common types of radioactivity are decay, b decay, g emission, positron emission, and electron capture.
• Nuclear reactions often involve electron capture and g rays.
• Some substances go through multiple decays before ending in a stable isotope.
• Each radioisotope has its own characteristic half-life, the time that is required for half of its atoms to decay.
• There is a wide range of half-lives of radioactive substances because of the large differences in stability.

• Many of the substances have been used for medical diagnosis and treatment, as well as determining the age of archaeological and geological objects.

• New atoms can be created by bombarding other atoms with high-speed particles.
• The transmutation reactions can be stable or radioactive.
• A number of artificial elements, including technetium, astatine, and the transuranium elements, have been produced in this way.

• Nuclear power as well as nuclear weapon detonations can be generated through fission, in which a heavy nucleus is split into two or more lighter nuclei.
• A chain reaction can result when the neutrons combine with other heavy nuclei.
• Useful power can be obtained if a nuclear reactor is used.
• Light and heavier nuclei are converted into energy.
• This energy is too expensive to be viable for commercial energy production.

• A radioactive compound can be used to follow reactions, track the distribution of a substance, diagnose and treat medical conditions, and much more.
• Other radioactive substances can be used to control pests, visualize structures, and provide fire warnings.
• Medicine tests and procedures using a wide variety of radioisotopes with relatively short half-lives are performed every year in the US.
• Most of the radioisotopes have half-lives that are short enough to be made on-site at medical facilities.
• Radiation therapy uses high-energy radiation to kill cancer cells.
• External or internal delivery of the radiation used for this treatment is possible.

• We are exposed to a lot of radiation from a variety of sources.
• Living organisms can be affected by this radiation.
• Ionizing radiation can damage a molecule and cause malfunction in cell processes, so it's the most harmful.
• It can cause damage to biological molecules and disrupt processes.
• Radiation is most harmful to rapidly reproducing cells.
• Alpha particles are the least penetrating type of radiation, with the most damaging and the most penetrating.

• Various devices are used to detect and measure radiation.
• We use several units to measure radiation: becquerels or Curies for rates of radioactive decay, gray or rads for energy absorbed, and rems or sieverts for biological effects of radiation.
• Exposure to radiation can cause a wide range of health effects.
• We can minimize the effects of radiation by shielding with dense materials such as lead, moving away from the source, and limiting time of exposure.

• The atom 23 11 Na has a mass of 22.9898 amu.

• Technetium-99 was prepared in 98Mo.
• An excited form of technetium-99 is represented as 99Tc* when Molybdenum-98 combines with a neutron.
• This excited nucleus relaxes to the ground state when it emits a g ray.
• The ground state of 99Tc emits a particle.

• The mass of the atom is 18.

• There are no nuclei in the band of stability.

The OpenStax book is available for free at http://cnx.org/content/col11760/1.9

• A series of decay reactions involving the step-wise emission of a, a, a, a, b, b particles is what proceeds from 90 Th.

• Explain the term half-life with an example.

• A sample of nobelium has a half-life of 55 seconds.

• The half-life of 239Pu is 24,000 y.

• Technetium-99 is often used to assess heart, lung, and liver damage.
• The half-life is 6.0 h.

• A sample of rock had rubidium-87 and strontium-87 in it.

• Glenn Seaborg and his associates found trace amounts of plutonium in natural uranium deposits in 1941.
• They claimed that the source of this 239Pu was the capture of neutrons by 238U nuclei.

• The 4 Be atom decays into a 3 Li atom by electron capture.

• 5 B atom decays into a 4 B atom by loss of a b+ particle or by electron capture.

• It is thought that 26Al (half-life: 7.2 x 105 years) was present in our solar system as it formed, but have since decayed and are now called extinct nuclides.

• Nuclear reactions are fusion and fission.

• The control rods are necessary to operate a nuclear chain reaction safely for the purpose of energy production.
• Explain why both are necessary.

• The tritium atom is 3.01605 amu, and the 1 H is 1.007825 amu.

• The half-life of Technetium-99m is 6.01 hours.

• In order to control growth and metabolism, iodine is released from the thyroid glands.
• If the body is injected with iodine-131, the thyroid can be imaged.
• I-133 is used to treat cancer of the thyroid in larger doses.
• I-131 has a half-life of 8.70 days.

• A product of two numbers is used to express large and small numbers.

• The power of 10 is the same as the number of places the decimal is shifted to.

• Every large and very small number can be marked with the exponential method.

• Add the digit terms of the numbers and adjust the exponential term to convert the digit back to a number between 1 and 10.

• If appropriate, convert the digit term back to a number between 1 and 10 by adjusting the exponential term.

• Add the exponents of the exponential terms to the digit terms.

• Take the digit term of the numerator and divide it by the digit term of the denominator.

• Divide 3.6 x 105 by 6.0 x 10-4.

• The digit term should be square in the usual way.

• The digit term can be cubed in the usual way.

• The power of 10 is evenly divisible by 2 if the exponential term is decreased or increased.
• Take the square root of the digit and divide it by 2.

• He has 525,341 bees.
• It's difficult to determine the exact number of bees because the last three figures of the number are incorrect, for during the time the keeper was counting the bees, some of them died and others hatched.

• The last three figures are unimportant except for the position of the decimal point.
• Their values are not useful in this situation.
• The accuracy of the measurement warrants is more important than the number of significant figures used.

• Significant figures are important in their application to computation.
• The sum or difference should have as many digits to the right of the decimal as the least certain of the numbers used in the computation.

• The product or quotient should not contain more digits than the factor with the least significant figures.

• If there is a number larger than 5 rounded up, the retained digit should be increased by one.
• If the digits that follow are less than 5, keep the retained digit.
• If the retained digit is followed by 5, then round up if it is even or odd.

• The logarithm of a number is the power to which 10 must be raised to equal that number.
• The common logarithm of 100 is 2 because 10 must be raised to the second power to equal 100.
• There are more examples following.

• A number less than 1 has a negative logarithm.

• Operations involving logarithms follow the same rules as operations involving exponents.

• Second-order polynomials are known as mathematical functions of this form.

• Only those with positive values are of any significance, as the real roots of the equations are always real.

• A two-dimensional data plot can be used to represent the relationship between any two properties of a system.

• There are five points in this table: (1,5), (2,10), (3,7), and (4,14).

• The amount of energy needed to heat a pound of water is called the BTU.
• The relationship of the two energy units depends on the temperature at which they are measured.
• 59 degF (15 degC) is the most widely used reference temperature in the United States.
• The conversion factor is provided in the table.

• The curve shows the absorption of water.

• The axis shows the absorption in cm.
• If we divide 1 by the value, we will get the length of the path after the light beam decays by a factor of the base of the natural logarithm.

• Aqueous solutions are available for acids and bases.
• The table shows the properties of acid and base solutions.
• Nominal values can be provided in cases where the manufacturer cites a range of concentrations and densities.

• Specific gravity data is contained in this column.
• Specific gravity is the ratio of density of a substance to the density of pure water.
• Commercial labels often cite specific gravity.

This solution is sometimes referred to as "ammonium hydroxide."

• y is years, d is days, h is hours, m is minutes and s is seconds y, d, days, h, m, and s are the number of years.

• If the temperature is below zero, it will freeze.

• The heat is determined by the properties of objects.

• Solids can't change their shape.
• Liquids do not experience large volume changes as gases do.
• Solids do not change volume.

• A pure substance has a definite composition, while a mixture can have a variety of compositions.
• Both have the same composition.

• Molecules of elements have one type of atom, while compounds have two or more types of atoms.
• Both of them are comprised of two or more atoms.

• Answers will change.

• A molecule consists of two or more atoms.
• They have different types of atoms that change from one substance to the next.

• Oxygen, gasoline, and nitrogen are consumed.
• The main products are carbon dioxide and water.
• Carbon monoxide and nitrogen oxides are not produced in large quantities.

• The value of an intensive property is the same regardless of the amount of matter being considered.

• Mass and volume are proportional to the amount of substance being studied.
• Dividing one extensive property by another will "cancel" the dependence on amount, yielding a ratio that is independent of amount.

• A/yellow mass is 65.14 kilograms, volume is 3.38 L, density is 19.3 kg/L, likely identity is gold.
• B/blue shows mass, volume, density and likely identity of apple.
• C/green is mass, volume, density and likely identity.
• The mass is 3.10 kg, the volume is 3.38 L, the density is 0.920 kg/L, and the likely identity is diamond.

• There are only two figures that are justified.

• The acid's volume is 123 liters.

• The starting materials are green and purple.
• There are two green and two purple spheres.
• This violates the idea that atoms aren't created during a chemical change.

• The number of atoms in a compound always have the same ratio.

• All atoms of a particular element have the same properties.
• The original idea was violated by the concept of isotopes, in which an element has a different mass.
• The second postulate of his atomic theory was changed to state that atoms of the same element must have the same chemical properties.

• The particles reside in the nucleus of the atom.
• Both have the same mass.
• The charged particles are called protons and un charged particles are called neutrons.

• Positive-positive repulsion will cause particles to pass near the nucleus.
• The larger the angle, the closer the particles are to the nucleus.
• If the nucleus is larger, the positive charge will be larger and the expected particles will be larger.

• The neutral 12C atom has six electrons.
• The mass number is 12 and the net charge is zero.

• Other answers are possible if a different element is chosen.

• The symbol is 6Li or 63Li.

• Neon is an example.
• There is no way to make sure that the total of 20.18 amu average atomic mass is accurate.
• We can guess that the abundances are 9% Ne-22, 9% Ne-20, and 9% Ne-21.
• The average mass is 20.18.
• The nature's mix of isotopes is 90.48% Ne-20, 9.25% Ne-22, and 0.21% Ne-21, so our guess amounts have to be slightly adjusted.

• The element oxygen has a symbol called O that represents both the element and one atom of oxygen.
• The subscript 2 in the formula must be used to distinguish the diatomic molecule from the single oxygen atoms.

• The same chemical composition and number of atoms can be found in these molecules.
• They are structural isomers.

• To get the number of moles, divide the mass of compound by the mass of the compound expressed in grams.

• The formula has more oxygen atoms than the other two compounds.
• 1.20 mol of a compound containing a single oxygen atom is equivalent to 0.60 mol of formic acid.

• The mass of 1 molecule is the same as the mass of 6.022 x 1023 molecule, but the units are different.

• The Al2S3 sample contains the greatest mass of Al.

• 1 serving is about 3.113 x 1025 C atoms.

• The least number of molecule is represented by 20.0 g H2O.

• We need to know how many moles of sulfuric acid are dissolved in the solution.

• This equation can be rearranged to fit the given quantities.

• When the same number of elements are represented on the reactant and product sides, an equation is balanced.

• The law of conserve of matter requires that equations be balanced.

• SiC + 2CO, 3.50 kilo of SiO2 x 103 kilo of 1.28 kilo of CO2 is the limiting reactant.

• The amount of acid present should be compared with the amount of Cr.
• The limiting reactant is cr.

• The limit reactant is Na2C2O4.

• There are only four molecules that can be made.

• The amount cannot be weighted by ordinary balances.

• The empirical formula is BH3.
• The formula is B2H6.

• The empirical formula is WCl4.

• For a match and a bonfire, the temperature of 1 gram of burning wood is the same.
• This is an intensive property and depends on the material.
• This is an extensive property and the amount of heat produced depends on the amount of material.
• The amount of wood in a bonfire is more than the amount of heat produced in a match, which is why we can sit around a bonfire to stay warm, but a match would not provide enough heat to keep us from getting cold.

• The heat capacity and specific heat are related to the temperature of the mass of the substance.
• Specific heat is an intensive one, and heat capacity is an extensive one.

• We assume that the density of water is 1.0 g/ cm3 and that it takes as much energy to keep the water at 85 degrees.
• The water is going to be heated.

• The temperature of the coffee will go down.

• The heat produced shows that the reaction is exothermic.

• The twofold increase in the amount of water leads to a two-fold decrease of the temperature change since the mass and heat capacity of the solution is equal to that of the water.

• The heat in the example is produced by 0.0500 mol HCl and 0.0500 mol NaOH.

• B2H6 is the best rocket fuel because it gives off the most heat.

• The spectrum has at least one colored line that is red.

• The energy mol-1 was 1.823 x 105 J mol-1 and red 9.
• The color of (a) is red and (b) is blue.

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• No quantized energy means that the electrons can only have certain energy values; values between those quantized values are not allowed.

• The Bohr model only works for one-electron atoms or strontiums, and both involve a heavy nucleus with electrons moving around it.
• According to classical mechanics, the Rutherford model predicts a miniature solar system with electrons moving about the nucleus in a circular or elliptical pattern that are confined to planes.
• If the requirements of classical theory are ignored, such atoms would be stable, have constant energy and have no visible light.
• If classical theory is applied, the Rutherford atom would emit radiation of continually increasing frequencies, which would cause it to collapse in a short time, contrary to the long-term stability of atoms.
• The atom can emit or absorb radiation when the electron moves around the nucleus, but the model assumes that the electron won't emit radiation while moving about the nucleus.
• The "quantum jumps" will produce a spectrum in agreement with observations.

• Both models have a central positively charged nucleus with electrons moving about the nucleus.
• This quantization is an ad hoc assumption made by Bohr to incorporate quantization into an essentially classical mechanics description of the atom.
• When the electron switches are in a different location, the electrons around the nucleus do not emit or absorb radiation.

• Both (b) and (c) are correct, but (e) is the best answer.

• Normal chemical reactions do not change the nucleus protons.
• The electrons move by themselves.

• Positive charges form when electrons are lost.

• P, I, Cl, and O are nonmetals.
• The metals would form cations.

• NaCl is composed of ionized atoms arranged in a crystal lattice.

• Two of the electrons in the molecule are shared, and the other six are located on the Cl atom.

• There is a sharing of electrons between atoms.
• Two electrons are shared in a single bond, four electrons are shared in a double bond, and six electrons are shared in a triple bond.

• The figure on the left has greater bond energy.
• The form is more stable.

• The longest bonds are the C-C single bonds.

• The 3 level is much smaller.
• An electron is removed from a lower energy level where the attraction is stronger for the electron.
• Energy is required to unpair two electrons.
• The lone electron in the exposed outer energy level is what the second ionization potential requires for Ca.

• The HBeH molecule has only two electrons to bond with the two electrons from the hydrogens and is therefore linear.

• Space is needed for each pair of electrons whether they are in a bond or lone pairs.

• The placement of all electrons is considered by the geometry.
• The bonding-pair geometry is only considered by the structure.

• As long as the polar bonds are compensated.

• There are polar bonds in all of these molecule and ion.

• 5, PCl3 and SeF4 have dipole moments.

• SeS2, CCl2F2, PCl3 and ClNO all have dipole moments.

• The structures are very similar.
• The bond angle is shown in the model mode because each electron group occupies the same amount of space.
• The hydrogens are compressed in the "real" mode when the lone pairs are larger.
• This leads to a smaller angle.

• Both types of bonds contain a maximum of two electrons because of the overlap of atomic orbitals on adjacent atoms.
• s bonds are stronger and result from end-to-end overlap and all single bonds are s bonds, p bonds between the same two atoms are weaker because they result from side-by-side overlap, and multiple bonds contain one or more p bonds.

• The distance with the lowest energy is the average bond distance.
• The positive charges on the two nuclei repel each other at distances less than the bond distance.

• The geometry of bonding orbitals is explained in hybridization.

• Oxidation states P, S, and O.

• 3 hybrid to form three bonds and hold one pair.

• A triple bond consists of two bonds.
• The s bond is stronger than the p bond.

• There are four electrons in a separate orbital and an electron on an oxygen atom.

• s orbitals are end-to-end combinations of atomic orbitals, whereas p orbitals are formed by side-by-side overlap of orbitals.
• The holding of two or more atoms is different.
• Any bonding that has occurred can be destabilizing due to antibonding orbitals.

• An odd number of electrons can never be matched.

• Bonding orbitals have close proximity to more than one nucleus.
• The system is stable because of the interaction between the positively charged nuclei and negatively charged electrons.

• The energy of the system is lowered relative to the energy of the nonbonded electrons when the two bonding electrons arepaired.

• The p orbitals are the last filled in N2 because 2 has s-p mixing.

• A sharper knife has a smaller surface area than a dull knife.
• A sharp knife will cut through material more effectively if it exerts a higher pressure with the same amount of force.

• Lying down forces less pressure on the ice compared to standing up.
• You are less likely to break through thin ice if you exert less pressure.

• With a closed-end manometer, there would be no change since the liquid in the manometer tube would contribute the same pressures in both arms.

• The volume of the bubbles increases as the pressure decreases.

• The curve would be the same shape as before.

• The pressure goes down by a factor of 3.

• The moles of HgO should be determined by using the chemical equation, and the moles of O2 should be determined by using the volume of O2 from the moles of O2, temperature, and pressure.
• The balanced equation can be used to calculate the moles of H2 for the This OpenStax book.
• Determine the amount of CO2 produced and the number of moles.
• Determine the volume of gas from the ideal gas law.

• A gas escapes through a hole into a vacuum.

• A sample of gas has a range of values.
• A molecule can speed up or slow down when it collides with another molecule.
• The average speed of the molecule is constant.

• The He atoms behave as though they were heavier when cooling their velocities.

• Under the conditions in (b), the gas behavior will be like an ideal gas.
• Molecules have high speeds and move through greater distances, but they have shorter contact times and interactions are less likely.
• There are conditions described in (a) and (c).
• Some gases may liquefy.
• Most gases will liquefy.

• At low pressures, the correction factor for intermolecular attractions is more significant, and the effect of the volume of the gas molecule on Z would be a small lowering compressibility.

• The effect of intermolecular attractions would be larger at higher pressures and low temperatures.

• Liquids and Solids are similar in their composition.
• They have similar densities that are much larger than gases.
• Liquids and Solids are not the same in that they do not have a fixed shape.

• The atoms are free to move from one position to another.
• The particles of a liquid are confined to the shape of the vessel in which they are placed.
• A gas expansion without limit will fill the space in which it is placed.

• The attractive forces exceed the energy of the molecule in a liquid or solid.

• The structure is more prevalent in large atoms.
• A second atom can be distorted by the appearance of the dipole in the first atom.
• The electrons of the second atom are attracted to the positive end of the first atom.
• The net result is temporary dipoles that attract one another.

• As the number of electrons increases, the London forces increase.

• Chloroethane has larger dipole interactions because of the Cl-C bond, which leads to a higher boiling point.

• Water has a higher melting point because it has stronger hydrogen bonds.

• The hydrogen bond between two hydrogen fluoride molecule is stronger than that between two water molecule because the electronegativity of F is greater than that of O.
• The hydrogen bond between H and O will be stronger than that between H and F.

• The principle of H-bonding is holding the strands together.
• The H-bonding is between the two.

• The water has strong hydrogen bonding.
• The water is attracted to one another and has a large surface tension, forming a type of "skin" at its surface.

• If placed on the water, this skin can support a bug or paper clip.

• The higher the temperature, the more fluid the liquid is, and the less the intermolecular forces are overcome.

• Intermolecular forces and surface tension are related.

• The energy required to partially overcome intermolecular attractive forces in the solid and cause a phase transition to liquid water is absorbed by the ice.
• Until the ice is melted, the solution remains at 0 degC.
• Until the ice disappears, the amount of water is constant.
• The temperature of the water can go up.

• There is a decrease in the amount of liquid in an open container.

• The higher the temperature, the more energy the molecule of gasoline has to escape from the liquid.

• Water will boil at a lower temperature at 5000 feet because the atmospheric pressure is lower than at sea level.
• A longer time is required to fully cook the egg because of the physical and chemical changes that will be caused by the lower temperature.

• The amount of intermolecular attraction via dispersion forces and the amount of energy required to overcome these forces increases as the number of atoms in this homologous series increases.

• The attractive forces are stronger in CS2 than in CO2 because of the higher boiling point.
• It would be expected that the heat of vaporization would be greater than that of CO2.
• It would seem reasonable to have a value of 28 kJ/mol.
• It is implausible that a value of -8.4 kJ/ mol would indicate a release of energy.

• The heat that is needed to evaporate the liquid is removed from the skin.

• The heat of vaporization is likely to have a larger magnitude since the intermolecular interactions have to be completely overcome.

• The water is a gas at low pressures.
• As the pressure increases, the water becomes a liquid.
• At 40 degC, water is a liquid at pressures higher than 75 torr.
• Water goes from a gas to a solid as the pressure increases.

• It may take several days for ice to break.
• Some ice atoms form gas and escape from the ice crystals.

• The clothes are dry.

• The same amount of energy is required to overcome these intermolecular forces.

• The melting process occurs over a wide temperature range because the attractions of various strengths are overcome at different temperatures.

• This form of iron has a body-centered cubic structure.
• There are eight corners to the cube and one atom in the center.

• There are two holes for each anion in the closest-packed array.
• The numbers of anions and cations are the same if only half the holes are occupied.
• The formula is called CdS.

• Only one hole can be occupied for each anion in the array.
• The formula for thallium is TlI.

• The oxidation number of titanium is +4.

• The ion sizes are 888-609- 888-609- 888-609- 888-609- 888-609- The two can interchange easily.
• The switch of Si4+ for Al3+ usually compensates for the difference in charge.

• A compound cannot vary in composition.
• Heterogeneity is not present in solutions at the molecular level.

• The dissolution process increases the thermal energy of the solution to make up for the difference.

• When the total intermolecular forces between the solute and solvent are stronger than the total in the pure solute and in the pure solvent, heat is released.
• When the totalIMFs in the solution are weaker than the total in the pure solute and in the pure solvent, heat is absorbed.

• The individual ion atoms become strongly solvated when they are dissolved in water with a large dipole moment.
• A nonpolar liquid with a zero dipole moment does not interact with the NaCl crystals.

• 3)3 is a strong electrolyte, so it should dissociation into Fe3+ and NO3.
• The solution is best represented by (z)

• The solubility of gases decreases when a solution is cooled.

• The bonds between molecule are stronger than the bonds between molecule.

• Some regions will have water and oil in them, while other regions will have oil and water in them.

• The boiling point elevations of both solutions are the same.
• Both solutions have the same Osmotic pressure and freezing point.

• H2O has a mole fraction of 0.622.

• Determine the number of moles in the solution.
• The molality is determined by the number of moles and the mass of solvent.

• The assumption that HCl is not ionized is supported by these values.

• The ion and compounds present in the water lower the freezing point of the beef.

• The observed change is equivalent to the theoretical change.

• The particles inoidal dispersions are larger than the solutes of typical solutions.

• On a visual scale, collioids are not different from solutions on a small scale.

• If they are placed in an electrolytic cell, dispersed particles will move towards the electrode that carries a charge opposite to their own charge.
• The charged particles will be coagulated at this electrode.

• The instantaneous rate is the rate of a reaction at any point in time, a period of time that is so short that the concentrations of reactants and products don't change.
• As product just begins to form, the initial rate is the instantaneous rate of reaction.
• The average rate is the average of the instantaneous rates.

• The rate of the reaction increases with higher molarity.
• The rate of the reaction increases when the temperature is higher.
• Smaller magnesium metal pieces will react more quickly than larger pieces because of the reactivity of the surface.

• The process reduces the rate by a factor of 4.

• The reaction is zero order.

• The plot is linear and the reaction is second order.

• The first order is the reaction.

• The reactants may be moving too slowly to have enough energy for the reaction, or the orientation of the molecule may prevent the reaction from happening.

• The minimum amount of energy needed to form an activated complex is called the activation energy.

• We can't predict the effect of changing concentration without knowing the rate equation.

• The rate is increased by a factor of 4.
• The rate-determining step is Step II.
• This reaction is in line with the rate law.
• If you combine steps 1 and 2 with step 3, you can give the appropriate stoichiometry.

• The mode of action for a catalyst is to provide a mechanism by which the reactants can unite more readily by taking a path with a lower reaction energy.
• The rates of the forward and reverse reactions are increased, which leads to a faster achievement of equilibrium.

• They are not used up, which is a characteristic of catalysts.

• The effect of a catalyst is indicated by the lowering of transition state energy.

• They both have the same rate.

• C6H6 dominates over C2H2.
• If the equilibrium rate is suitable, the reaction would be feasible.

• To reach equilibrium, the system will shift toward the reactants.

• There is a situation in (a) and (b).

• H2O is 0.042.

• CO2 is required for equilibrium.

• The change in enthalpy can be used.
• The heat produced in the reaction can be thought of as a product.
• The heat added can be thought of as a reactant if the reaction is endothermic.
• The exothermic reaction to the reactants would be shifted back to the products by additional heat.
• cooling an exothermic reaction causes it to shift to the product side, while cooling an endothermic reaction causes it to shift to the reactants' side.

• It isn't at equilibrium.
• Products and reactants escape from the region of the flame because the system is not confined.

• Reduce the container volume and heat the mixture.

• The OpenStax book is available for free.

• CO no change, H2O no change, and H2O decreases.
• In (b), (c), (d), and (e), the mass of carbon will not change.

• Add NaCl or some other salt to the solution.

• [C]2 [A]2.

• The mass of Ni doesn't change, but the activities of pure crystals are the same.

• The concentration of N2O4 is less than the maximum allowed.

• The error is very small.

• The gases must increase in pressure.

• After H2 is added, some water forms to compensate for the removal of water vapor and as a result of a shift to the left.

• NH4 [O2]7 [NO2]4[H2O]6.

• The conjugate acid of the labels is CA.

• H2S(BA), NH2 (BB), HS-(CB), NH3(CA) 11.
• Amphiprotic species can either gain or lose a protons in a chemical reaction and act as a base or acid.
• H2O is an example.

• The OpenStax book is available for free at 40 degrees.

• The water reacts with the salt to form a weak acid.
• This reaction causes the solution to be basic.

• The oxidation state of sulfur in H2SO4 is greater than the oxidation state of sulfur in H2SO3.

• PH3 is weaker than HI.
• The stronger acid is HBr, which is to the left and below S.

• The larger number of oxygen atoms on the central atom creates a higher oxidation state and leads to a stronger acid.
• The acidity increases in the same way as a salt.

• The basicity of the anions in a series of acids will be different from the acid's acidity.
• As the electronegativity of the central atom increases, the acidity increases.
• I am the least electronegative of the three.
• I am the least electronegative of the three.

• The weak bases of NH2 and PH2 act as strong bases toward H+.
• Weak acids have less basic character.
• The basic anion is found in a periodic group.

• As the oxidation state of the central ion increases, the acid becomes more acidic and the anion less basic.

• When a weak acid or weak base is present, equilibrium calculations are needed.

• The initial concentration of the acid can be assumed constant and equal to the initial value of the total acid concentration if the equilibrium is neglected.

• The equilibrium concentration of H3O+ is dependent on the contribution of water.

• Adding H3O+ ion will lower their concentration by reacting with OH- ion.

• There is a mixture of acids in the solution.
• The weak acid is suppressed by the strong acid and HCO2H is primarily a molecule of HCO2H.
• The HCO2H contributes a small amount of hydronium ion to the solution.
• The strongest acid is the one that is completely ionized.
• The stronger acid determines the concentration of hydronium ion and the weaker acid is ionization by the stronger acid.

• The assumption that it can be neglected is valid because the value is less than 5%.

• The assumption that it can be neglected is valid because the value is less than 5%.

• The assumption that it can be neglected is valid because the value is less than 5%.

• The assumption that it can be neglected is valid because the value is less than 5%.

• The assumption that it can be neglected is valid because the value is less than 5%.

• 3CO2 will increase the concentration of CH3 CO2 which will react with H3O+ and produce CH3CO2 H. H3O+ decreases and CH3CO2H increases.

• 3CO2H will produce CH3 CO2 in the process.
• CH3CO2H decreases and CH - 3 CO2 increases.

• H3O+ increases slightly and CH3 CO2 increases.

• You should use a weak base and salt for buffers with pHs greater than 7.

• The pOH of the buffer is 14.

• The mass of the substance is 205.169 g/mol.

• The solution of a weak base with a strong acid is slightly acidic due to the presence of the conjugate acid.
• Pick an indicator that changes color in the acidic range and brackets the pH at the equivalence point.
• It is a good example.

• Water is the only source of OH- ion in an acid solution.
• The concentration of OH- would be zero if the contribution from water was neglected.

• There is no change.
• There is an activity of 1 for a solid.

• The new temperature must be known about the solubility of silver bromide.
• Some of the solid silver bromide can be dissolved.

• The value is less than 5% and can be ignored.

• The condition is satisfied.

• The condition is satisfied.

• The above value is less than 5%.

• A more exact method, such as successive approximations, must be used.

• The maximum value for ignoring the change is greater than the concentration changes.

• AgI will start to form.

• The electronic andmolecular shapes are the same.

• HgCl2 dissolved.

• The Le Chatelier's principle states that the equilibrium will shift to the reactants' side when added heat appears on the product side.
• Less reagent will be dissolved.
• This situation was found in our case.
• The reaction is cold.

• A reaction can happen without the constant input of energy from an external source.

• The rate of oxidation is very slow.
• Even though plastic is stable, it does not break down over long periods of time.

• There are four initial and four final microstates.

• There is a chance that all the particles will be on the same side.
• The 18 result for the four-particle system has a lower probability.
• The conclusion we can make is that the probability for all the particles to stay in one part of the system will decrease rapidly as the number of particles increases.

• There is one initial state.
• The energy can be contained in pairs A-C, A-D, B-C, or B-D.

• There are four possible states.

• The opposite trend in their entropies would be suggested by the mass of these molecules.
• The trend is a result of the more significant variation of entropy.
• I2 is a solid, Br2 is a liquid, and Cl2 is a gas at room temperature.

• There is a decrease in the number of mobile ion in solution.
• Three moles of gas are lost from reactants to products.
• There is an increase in gas from reactants to products.

• It is assumed that the values for enthalpy and entropy don't change much at the higher temperatures used in the problem.

• At room temperature, the reaction is nonspontaneous.

• As it approaches zero, it becomes less positive.

• There will be a reaction at hotter temperatures.

• The process is cold.

• Under the stated conditions, this is the maximum pressure of the gases.

• The air is saturated with water at 25% humidity.

• Under these conditions, the forward reaction to produce F6P is spontaneously occurring.

• It becomes more negative.

• It becomes more positive.

• Reducing agent: (a) Hg; (b) Al 9.

• The concentration of hydrogen ion is zero.
• It would react with the excess hydroxide ion to produce water if it were produced.

• Current can flow through the circuit with a salt bridge.

• The oxidation-reduction reaction involves active electrodes.
• If metal atoms were to oxidize and go into solution, the electrode would lose mass.
• There isOxidation at the anode.

• The cost of the materials used in the battery, toxicity of the various components, should it be a primary or secondary battery, energy requirements (the "size" of the battery/how long should it last), will a particular battery leak when the new device is some of the considerations.

• The reaction can be interfered with by battery reaction byproducts.
• As long as reagents are supplied, a fuel cell can continue to function.

• It is the same effect as a battery running dead if the term is decreased at low temperatures.

• Both examples involve protection from the elements.
• The metal that oxidizes or reacts is the sacrificial anode.

• In the case of iron and zinc, zinc has a more negative standard reduction potential.
• In the case of iron and copper, iron has a smaller standard reduction potential and so it is corrodes.

• While the reduction potential of lithium would make it capable of protecting the other metals, it would also have a reaction with most substances.
• The metal it is trying to protect would react quickly with other substances, even those that wouldn't oxidize it.
• The anode would need to be replaced frequently if activity like this continued.

• The alkali metals are more reactive than the alkaline earth metals in the same period.

• There are a number of ways in which to distinguish between the two, including a flame test that shows the yellow color of strontium and a comparison of their solubilities in water.

• It's easy to test the solubility of NaCl by heating to 100 degC, since it's 39.12 g 100 mL.
• It is difficult to determine density on a solid, but there is a way to do it, and it is the easiest and least expensive test to perform.

• Tin reacts with acid to make hydrogen gas.

• The bonding is ionic, as indicated by its melting point.
• The liquid in PbCl4 is an unstable liquid at room temperature.

• magnesium can be used in construction even when it comes in contact with a flame because a protective oxide coating is formed, preventing gross oxidation.
• A high-intensity flame will cause its rapid burning if the metal is finely subdivided or present in a thin sheet.

• A solution of hydrofluoric acid would not harm the diamond.

• The nitrogen atoms have s bonds and p bonds that hold them together in the N2 molecule.
• N2 is a very stable molecule because it has three strong bonds.
• The bonding requirement of Phosphorus is fulfilled by forming three s bonds because it does not form p bonds efficiently.

• The nonmetals have more electronegativity than hydrogen.
• The nonmetal has a tendency to attract electrons in the bond to itself, which makes it a better representation of the negative charge.

• There is only one orbital for hydrogen to bond to other atoms.
• Only one two-electron bond can be formed.

• Ammonia is a Lewis base because it accepts both protons and electron pairs.

• The molecule has a bent geometry with an ONO bond angle.

• The ONO bond angle is slightly less than 120 degrees.

• The molecule has an ONO bond angle of 180 degrees.

• The stronger attraction of the oxygen electron results in a stronger attraction of the oxygen for the electrons in the O-H bond.

• The acid strength depends on the relative electronegativity of the central atom as H2SO4 and H2SeO4 have the same oxidation number.
• H2SO4 is the stronger acid as sulfur is more negative than positive.

• Sulfur can form double bonds only at high temperatures, which is not the case for oxygen.

• It is not a salt because it is covalently bonding.
• A salt has ionic bonds.

• The strength of HClO3 is stronger than that of HBrO3.

• The only option that can provide enough driving force to convert La(III) into La is Al.

• The molten iron is denser than CaSiO3 so it can easily be separated.
• The floating slag layer prevents molten iron from being exposed to O2 and oxidizes Fe back to Fe2O3.

• The OpenStax book is free and can be found at http://cnx.org/content/col11760/1.9.

• The complex has no unpaired electrons.
• The complex does not have any geometric isomers, but it does have an optical isomer.

• The structure of the product is affected by the two reactions.
• The way in which that is done is different.
• A bond between the C and the Br can be formed when an existing C-H bond is broken.
• The only bond broken in the hydrocarbons is the p bond, which can be used to form a bond to one of the bromine atoms.

• All orientations of the substituents about the C-C bonds are interchangeable by rotation in unbranched alkanes.
• In the unbranched alkenes, the inability to rotate about the C bond results in different isomers.

• The answer key explains phenomena at the molecular level.

• They are the same compound because they are saturated with an unbranched chain of six carbon atoms.

• The -COOH functional group is found in both of them, since they are both carboxylic acids.
• There are no double or triple bonds in the hydrocarbon chain in a saturated fat acid, whereas there are one or more multiple bonds in the unsaturated fat acid.

• The book is free at http://cnx.org/content/col11760/1.9 14 Si, 15 P, 25 Mn, and 56 Ba 5.

• Nuclear reactions change one type of nucleus into another.
• Nuclear reactions have larger energies than chemical reactions.

• A radioactive element may be emitted from 92 U + 1 H.

• An inner atomic electron can be absorbed in a protons-rich nucleus.

• When an electron falls from a higher energy level to a lower one, the difference in its two energy levels is given off as an X-ray.

• It is most likely to decay by positron emission.
• The n:p ratio is 29 24 for Cr 53, 26 25 for Mn 51, and 33 26 for Fe-59.
• Positron decay occurs when the n:p ratio is low.

• Answer Key n:p ratio is most likely to decay.

• 26 Fe is a stable isotope.

• Half-life is the time required for half the atoms in a sample to decay.
• The half-life is 5770 years for C-14.
• A 10-g sample of C-14 would contain 5 g of C-14 after 5770 years, and a 0.20-g sample of C-14 would contain 0.10 g after 5770 years.

• The rock would be younger than the age calculated in part.
• If Sr was in the rock, the amount produced by radioactive decay would be the same as the initial amount.
• This amount is smaller than the amount used to calculate the age of the rock and the age is proportional to the amount of Sr.

• Since the formation of the earth, no Pu-239 could remain.
• plutonium could not have been formed with the Uranium.

• For fusion to happen, two nuclei must collide.
• The nuclei need high temperatures to have enough energy to overcome the strong repulsion.

• A controlled chain reaction can be achieved if a large amount of a fissionable isotope is present.
• Tubes called fuel rods contain the radioactive isotope.

• A moderator slows the production of nuclear reactions so that they can be absorbed by the fuel.

• The heat from the reactor is carried to an external boiler and turbine where it is converted into electricity.

• The control system consists of control rods that are placed between fuel rods to absorb neutrons and keep the chain reaction at a safe level.

• The function of this component is to protect workers from the radiation produced by the nuclear reactions.

• An external steam generator is used to generate heat from the fission of uranium.
• The steam turns the turbine into a generator.

• When radioactive Ag+ or radioactive Cl- is introduced into the solution containing the stated reaction, it will produce a radioactive precipitate that was previously devoid of radiation.

• Alpha particles have a stronger ionizing potential than x-rays and g-rays and can be stopped by very thin shielding.
• The OpenStax book is available for free at http://cnx.org/content/col11760/1.9 and can damage the cells of the lungs and cause cancer.

Document Outline

• Preface

• Chapter 1. Essential Ideas 1.1. Chemistry in Context* 1.2. Phases and Classification of Matter* 1.3. Physical and Chemical Properties* 1.4. Measurements* 1.5. Measurement Uncertainty, Accuracy, and Precision* 1.6. Mathematical Treatment of Measurement Results* Glossary

• Chapter 2. Atoms, Molecules, and Ions 2.1. Early Ideas in Atomic Theory* 2.2. Evolution of Atomic Theory* 2.3. Atomic Structure and Symbolism* 2.4. Chemical Formulas* 2.5. The Periodic Table* 2.6. Molecular and Ionic Compounds* 2.7. Chemical Nomenclature* Glossary

• Chapter 3. Composition of Substances and Solutions 3.1. Formula Mass and the Mole Concept* 3.2. Determining Empirical and Molecular Formulas* 3.3. Molarity* 3.4. Other Units for Solution Concentrations* Glossary

• Chapter 4. Stoichiometry of Chemical Reactions 4.1. Writing and Balancing Chemical Equations* 4.2. Classifying Chemical Reactions* 4.3. Reaction Stoichiometry* 4.4. Reaction Yields* 4.5. Quantitative Chemical Analysis* Glossary

• Chapter 5. Thermochemistry 5.1. Energy Basics* 5.2. Calorimetry* 5.3. Enthalpy* Glossary

• Chapter 6. Electronic Structure and Periodic Properties of Elements 6.1. Electromagnetic Energy* 6.2. The Bohr Model* 6.3. Development of Quantum Theory* 6.4. Electronic Structure of Atoms (Electron Configurations)* 6.5. Periodic Variations in Element Properties* Glossary

• Chapter 7. Chemical Bonding and Molecular Geometry 7.1. Ionic Bonding* 7.2. Covalent Bonding* 7.3. Lewis Symbols and Structures* 7.4. Formal Charges and Resonance* 7.5. Strengths of Ionic and Covalent Bonds* 7.6. Molecular Structure and Polarity* Glossary

• Chapter 8. Advanced Theories of Covalent Bonding 8.1. Valence Bond Theory* 8.2. Hybrid Atomic Orbitals* 8.3. Multiple Bonds* 8.4. Molecular Orbital Theory* Glossary

• Chapter 9. Gases 9.1. Gas Pressure* 9.2. Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law* 9.3. Stoichiometry of Gaseous Substances, Mixtures, and Reactions* 9.4. Effusion and Diffusion of Gases* 9.5. The Kinetic-Molecular Theory* 9.6. Non-Ideal Gas Behavior* Glossary

• Chapter 10. Liquids and Solids 10.1. Intermolecular Forces* 10.2. Properties of Liquids* 10.3. Phase Transitions* 10.4. Phase Diagrams* 10.5. The Solid State of Matter* 10.6. Lattice Structures in Crystalline Solids* Glossary

• Chapter 11. Solutions and Colloids 11.1. The Dissolution Process* 11.2. Electrolytes* 11.3. Solubility* 11.4. Colligative Properties* 11.5. Colloids* Glossary

• Chapter 12. Kinetics 12.1. Chemical Reaction Rates* 12.2. Factors Affecting Reaction Rates* 12.3. Rate Laws* 12.4. Integrated Rate Laws* 12.5. Collision Theory* 12.6. Reaction Mechanisms* 12.7. Catalysis* Glossary

• Chapter 13. Fundamental Equilibrium Concepts 13.1. Chemical Equilibria* 13.2. Equilibrium Constants* 13.3. Shifting Equilibria: Le Chatelier's Principle* 13.4. Equilibrium Calculations* Glossary

• Chapter 14. Acid-Base Equilibria 14.1. Bronsted-Lowry Acids and Bases* 14.2. pH and pOH* 14.3. Relative Strengths of Acids and Bases* 14.4. Hydrolysis of Salt Solutions* 14.5. Polyprotic Acids* 14.6. Buffers* 14.7. Acid-Base Titrations* Glossary

• Chapter 15. Equilibria of Other Reaction Classes 15.1. Precipitation and Dissolution* 15.2. Lewis Acids and Bases* 15.3. Multiple Equilibria* Glossary

• Chapter 16. Thermodynamics 16.1. Spontaneity* 16.2. Entropy* 16.3. The Second and Third Laws of Thermodynamics* 16.4. Free Energy* Glossary

• Chapter 17. Electrochemistry 17.1. Balancing Oxidation-Reduction Reactions* 17.2. Galvanic Cells* 17.3. Standard Reduction Potentials* 17.4. The Nernst Equation* 17.5. Batteries and Fuel Cells* 17.6. Corrosion* 17.7. Electrolysis* Glossary

• Chapter 18. Representative Metals, Metalloids, and Nonmetals 18.1. Periodicity* 18.2. Occurrence and Preparation of the Representative Metals* 18.3. Structure and General Properties of the Metalloids* 18.4. Structure and General Properties of the Nonmetals* 18.5. Occurrence, Preparation, and Compounds of Hydrogen* 18.6. Occurrence, Preparation, and Properties of Carbonates* 18.7. Occurrence, Preparation, and Properties of Nitrogen* 18.8. Occurrence, Preparation, and Properties of Phosphorus* 18.9. Occurrence, Preparation, and Compounds of Oxygen* 18.10. Occurrence, Preparation, and Properties of Sulfur* 18.11. Occurrence, Preparation, and Properties of Halogens* 18.12. Occurrence, Preparation, and Properties of the Noble Gases* Glossary

• Chapter 19. Transition Metals and Coordination Chemistry 19.1. Occurrence, Preparation, and Properties of Transition Metals and Their Compounds* 19.2. Coordination Chemistry of Transition Metals* 19.3. Spectroscopic and Magnetic Properties of Coordination Compounds* Glossary

• Chapter 20. Organic Chemistry 20.1. Hydrocarbons* 20.2. Alcohols and Ethers* 20.3. Aldehydes, Ketones, Carboxylic Acids, and Esters* 20.4. Amines and Amides* Glossary

• Chapter 21. Nuclear Chemistry 21.1. Nuclear Structure and Stability* 21.2. Nuclear Equations* 21.3. Radioactive Decay* 21.4. Transmutation and Nuclear Energy* 21.5. Uses of Radioisotopes* 21.6. Biological Effects of Radiation* Glossary

• Appendix A. The Periodic Table*

• Appendix B. Essential Mathematics* B.1. Exponential Arithmetic B.2. Significant Figures B.3. The Use of Logarithms and Exponential Numbers B.4. The Solution of Quadratic Equations B.5. Two-Dimensional (x-y) Graphing

• Appendix C. Units and Conversion Factors*

• Appendix D. Fundamental Physical Constants*

• Appendix E. Water Properties*

• Appendix F. Composition of Commercial Acids and Bases*

• Appendix G. Standard Thermodynamic Properties for Selected Substances*

• Appendix H. Ionization Constants of Weak Acids*

• Appendix I. Ionization Constants of Weak Bases*

• Appendix J. Solubility Products*

• Appendix K. Formation Constants for Complex Ions*

• Appendix L. Standard Electrode (Half-Cell) Potentials*

• Appendix M. Half-Lives for Several Radioactive Isotopes*

• Solutions Chapter 1 Chapter 2 Chapter 3 Chapter 4 Chapter 5 Chapter 6 Chapter 7 Chapter 8 Chapter 9 Chapter 10 Chapter 11 Chapter 12 Chapter 13 Chapter 14 Chapter 15 Chapter 16 Chapter 17 Chapter 18 Chapter 19 Chapter 20 Chapter 21 Index

• Chemistry.pdf Blank Page