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化学 C2-C10,C13,C15
Type I: Ionic compounds with metals that always form a cation with the same charge、Type II: Ionic compounds with metals (usually transition metals)that form cations with various charges。
编辑于2023-01-24 00:36:30 江苏省
- 化学 C2-C10,C13,C15
Type I: Ionic compounds with metals that always form a cation with the same charge、Type II: Ionic compounds with metals (usually transition metals)that form cations with various charges。
- Nuclear Power
这是一篇关于Nuclear Power的思维导图，《Nuclear Power》是一本图书，作者是National Research Council; Committee on Future Nuclear Power Develo; Energy Engineering Board
化学 C2-C10,C13,C15
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- 化学 C2-C10,C13,C15
Type I: Ionic compounds with metals that always form a cation with the same charge、Type II: Ionic compounds with metals (usually transition metals)that form cations with various charges。
- Nuclear Power
这是一篇关于Nuclear Power的思维导图，《Nuclear Power》是一本图书，作者是National Research Council; Committee on Future Nuclear Power Develo; Energy Engineering Board
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Chemistry Mindmap
C2:Measurements and Calculations
2.1 Scientific Notation
Scientific notation simply expresses a number as a product of a number between 1 and 10 and the appropriate power of 10.( a x 10^n)
2.2 Units
The units part of a measurement tells us what scale or standard is being used to represent the results of the measurement
In1960 an international agreement set up a comprehensive system of units called the International System (le Système Internationale in French), or SI.
some of the SI Units
Commonly Prefices:
2.3 Measurements of Length,Volume, and Mass
Length: The fundamental SI unit of length is the meter, which is a little longer than a yard (1 meter =39.37 inches). In the metric system fractions of a meter or multiples of a meter can be expressed by powers of 10, as summarized in Table 2.3.
1 inch=2.54 centimeters
Table2.3
Volume:The amount of three-dimensional space occupied by a substance.The fundamental unit of volume in the SI system is based on the volume of a cube that measures 1 meter in each of the three directions--one cubic meter.
Measuring tool: graduated cylinder
Figure 2.2
Table 2.4
Mass: The quantity of matter present in an object. The fundamental SI unit of mass is the kilogram.
Measuring tool: balance
Table 2.5
Other English–metric equivalences are given in Section 2.6.
2.4 Uncertainty in Measurement
The numbers recorded in a measurement (all the certain numbers plus the first uncertain number) are called significant figures.
2.5 Significant Figures
Rules for Counting Significant Figures:
1. Nonzero integers. Nonzero integers always count as significant figures. For example, the number 1457 has four nonzero integers, all of which count as significant figures.
2. Zeros. There are three classes of zeros:
a. Leading zeros are zeros that precede all of the nonzero digits. They never count as significant figures.
For example, in the number 0.0025,the three zeros simply indicate the position of the decimal point. The number has only two significant figures, the 2 and the 5.
b. Captive zeros are zeros that fall between nonzero digits. They always count as significant figures.
For example, the number 1.008 has four significant figures.
c. Trailing zeros are zeros at the right end of the number. They are significant only if the number is written with a decimal point.
The number one hundred written as 100 has only one significant figure, but written as 100., it has three significant figures.
3. Exact numbers. Often calculations involve numbers that were not obtained using measuring devices but were determined by counting: 10 experiments, 3 apples, 8 molecules. Such numbers are called exact numbers. They can be assumed to have an unlimited number of significant figures. Exact numbers can also arise from definitions.
For example, 1 inch is defined as exactly 2.54 centimeters. Thus in the statement 1 in. = 2.54 cm, neither 2.54 nor 1 limits the number of significant figures when it is used in a calculation.
Rounding Off Numbers
rules for rounding off:
1. If the digit to be removed
a. is less than 5, the preceding digit stays the same. For example, 1.33rounds to 1.3.
b. is equal to or greater than 5, the preceding digit is increased by 1. For example, 1.36 rounds to 1.4, and 3.15 rounds to 3.2.
2. In a series of calculations, carry the extra digits through to the final result and then round off.* This means that you should carry all of the digits that show on your calculator until you arrive at the final number (the answer) and then round off, using the procedures in Rule 1.
Determining Significant Figures in Calculations：
1. For multiplication or division, the number of significant figures in the result is the same as that in the measurement with the smallest number of significant figures.
2. For addition or subtraction, the limiting term is the one with the smallest number of decimal places. For example, consider the following sum:
2.6 Problem Solving andDimensional Analysis
2.7 Temperature Conversions:An Approach to ProblemSolving
Familyhrenheit scale / Celsius scale / absolute or Kelvin scale
1. The size of each temperature unit (each degree) is the same for the Celsius and Kelvin scales. This follows from the fact that the difference between the boiling and freezing points of water is 100 units on both of these scales.
2. The Fahrenheit degree is smaller than the Celsius and Kelvin units. Note that on the Fahrenheit scale there are 180 Fahrenheit degrees between the boiling and freezing points of water, as compared with 100 units on the other two scales.
3. The zero points are different on all three scales.
Converting
Celsius and Kelvin: T°C+273=T K
Fahrenheit and Celsius Scales: T°F = 1.80 (T°C)+32
2.8 Density
Density can be defined as the amount of matter present in a given volume of substance
The densities of various common substances are given in Table 2.8
In certain situations, the term specific gravity is used to describe the density of a liquid. Specific gravity is defined as the ratio of the density of a given liquid to the density of water at 4 °C. Because it is a ratio of densities, specific gravity has no units.
C3:Matter
3.1 Matter
Matter has two characteristics: it has mass and it occupies space.
The substances we have just described illustrate the three states of matter: solid, liquid, and gas. The state of a given sample of matter depends on the strength of the forces among the particles contained in the matter; the stronger these forces, the more rigid the matter.
3.2 Physical and Chemical Properties and Changes
Properties
Substances also have physical properties physical properties: odor, color, volume, state (gas, liquid, or solid), density, melting point, and boiling point
We can also describe a pure substance in terms of its chemical properties, which refer to its ability to form new substances.
Changes:
physical changes: A physical change involves a change in one or more physical properties, but no change in the fundamental components that make up the substance. The most common physical changes are changes of state: solid ⇔ liquid ⇔ gas.
chemical changes: A chemical change involves a change in the fundamental components of the substance; a given substance changes into a different substance or substances. Chemical changes are called reactions: silver tarnishes by reacting with substances in the air; a plant forms a leaf by combining various substances from the air and soil; and so on.
3.3 Elements and Compounds
Elements: As we examine the chemical changes of matter, we encounter a series of fundamental substances called elements. Elements cannot be broken down into other substances by chemical means.
Compounds: The atoms of certain elements have special affinities for each other. They bind together in special ways to form compounds, substances that have the same composition no matter where we find them
3.4 Mixtures and PureSubstances
Mixtures: A mixture can be defined as something that has variable composition. Mixtures can be classified as either homogeneous or heterogeneous.
A homogeneous mixture is the same throughout. A homogeneous mixture is also called a solution.
A heterogeneous mixture contains regions that have different properties from those of other regions
Pure Substances: pure substance, on the other hand, will always have the same composition.
3.5 Separation of Mixtures
Distillation: We can separate the water from the minerals by boiling, which changes thewater to steam (gaseous water) and leaves the minerals behind as solids. If wecollect and cool the steam, it condenses to pure water.
Filtration: Suppose we scooped up some sand with our sample of seawater. This sample is a heterogeneous mixture, because it contains an undissolved solid as well as the saltwater solution. We can separate out the sand by simple filtration.
C4:Chemical Foundations: Elements, Atoms, and Ions
4.1 The Elements
4.2 Symbols for the Elements
4.3 Dalton’s Atomic Theory
1. Elements are made of tiny particles called atoms.
2. All atoms of a given element are identical.
3. The atoms of a given element are different from those of any other element.
4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms.
5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.
4.4 Formulas of Compounds
A compound is a distinct substance that is composed of the atoms of two or more elements and always contains exactly the same relative masses of those elements.
The types of atoms and the number of each type in each unit (molecule) of a given compound are conveniently expressed by a chemical formula.
In a chemical formula the atoms are indicated by the element symbols, and the number of each type of atom is indicated by a subscript, a number that appears to the right of and below the symbol for the element.
Rules for Writing Formulas
1. Each atom present is represented by its element symbol.
2. The number of each type of atom is indicated by a subscript written to theright of the element symbol.
3. When only one atom of a given type is present, the subscript 1 is not written.
4.5 The Structure of the Atom
Electrons(at an average distance of about 10^-8)
nucleus(about 10^-13 cm in diameter)
proton
neutron(* no for H)
4.6 Introduction to the ModernConcept of Atomic Structure
Table 4.4
4.7 Isotopes
These atoms are isotopes, or atoms with the same number of protons but different numbers of neutrons. The number of protons in a nucleus is called the atom's atomic number. The sum of the number of neutrons and the number of protons in a given nucleus is called the atom's mass number. Mass number = number of protons + number of neutrons
4.8 Introduction to the PeriodicTable
The elements were first arranged in this way in 1869 by Dmitri Mendeleev, a Russian scientist.
Families of elements with similar chemical properties that lie in the same vertical column on the periodic table are called groups.
the first column of elements (Group 1) has the name alkali metals. The Group 2 elements are called the alkaline earth metals, the Group 7 elements are the halogens, and the elements in Group 8 are called the noble gases. A large collection of elements that spans many vertical columns consists of the transition metals.
Most of the elements are metals. Metals have the following characteristic physical properties:
1. Efficient conduction of heat and electricity
2. Malleability (they can be hammered into thin sheets)
3. Ductility (they can be pulled into wires)
4. A lustrous (shiny) appearance
The relatively small number of elements that appear in the upper-right corner of the periodic table are called nonmetals. Nonmetals generally lack those properties that characterize metals and show much more variation in their properties than metals do. Whereas almost all metals are solids at normal temperatures.
The elements that lie close to the “stair-step” line as shown in blue in often show a mixture of metallic and nonmetallic properties. These elements, which are called metalloids or semimetals, include silicon, germanium, arsenic, antimony, and tellurium.
4.9 Natural States of the Elements
diatomic molecules: molecules made up of two atoms
Only two elements are liquids in their elemental forms at 25 °C: the nonmetal bromine (containing Br2 molecules) and the metal mercury. The metals gallium and cesium almost qualify in this category; they are solids at 25 °C, but both melt at ~30 °C.
Different forms of a given element are called allotropes.
4.10 Ions
We can produce a charged entity, called an ion, by taking a neutral atom and adding or removing one or more electrons.
positive ion, called a cation , is produced when one or more electrons are lost from a neutral atom
When electrons are gained by a neutral atom, an ion with a negative charge is formed. A negatively charged ion is called an anion.
4.11 Compounds That Contain Ions
Many substances contain ions. In fact, whenever a compound forms between a metal and a nonmetal, it can be expected to contain ions. We call these substances ionic compounds.
One fact very important to remember is that a chemical compound must have a net charge of zero. This means that if a compound contains ions, then
1. Both positive ions (cations) and negative ions (anions) must bepresent.
2. The numbers of cations and anions must be such that the netcharge is zero.
C5:Nomenclature
5.1 Naming Compounds
binary compounds—compounds composed of two elements.
1. Compounds that contain a metal and a nonmetal
2. Compounds that contain two nonmetals
5.2 Naming Binary Compounds That Contain a Metal and aNonmetal (Types I and II)
The metal loses one or more electrons to become a cation, and the nonmetal gains one or more electrons to form an anion. The resulting substance is called a binary ionic compound
Type I compounds: The metal present forms only one type of cation.
Type II compounds: The metal present can form two (or more) cations that have different charges.
To help distinguish between Type I and Type II cations, remember that Group 1 and 2 metals are always Type I. On the other hand, transition metals are almost always Type II.
5.3 Naming Binary CompoundsThat Contain Only Nonmetals (Type III)
Type III binary compounds contain only nonmetals.
5.4 Naming Binary Compounds:A Review
Type I: Ionic compounds with metals that always form a cation with the same charge
Type II: Ionic compounds with metals (usually transition metals)that form cations with various charges
Type III: Compounds that contain only nonmetals
5.5 Naming Compounds ThatContain Polyatomic Ions
As their name suggests, polyatomic ions are charged entities composed of several atoms bound together.
Note in Table 5.4 that several series of polyatomic anions exist that contain an atom of a given element and different numbers of oxygen atoms. These anions are called oxyanions.
When there are two members in such a series, the name of the one with the smaller number of oxygen atoms ends in -ite, and the name of the one with the larger number ends in -ate. When more than two oxyanions make up a series, hypo- (less than) and per- (more than) are used as prefixes to name the members of the series with the fewest and the most oxygen atoms, respectively.
5.6 Naming Acids
When dissolved in water, certain molecules produce H+ ions (protons). These substances, which are called acids
5.7 Writing Formulas from Names
C6:Chemical Reactions: An Introduction
6.1 Evidence for a Chemical Reaction
6.2 Chemical Equations
Chemists have learned that a chemical change always involves a rearrangement of the ways in which the atoms are grouped. A chemical change such as this is called a chemical reaction.We represent a chemical reaction by writing a chemical equation in which the chemicals present before the reaction (the reactants) are shown to the left of an arrow and the chemicals formed by the reaction (the products) are shown to the right of an arrow.The arrow indicates the direction of the change and is read as“yields” or “produces”
chemical reaction involves changing the ways the atoms are grouped.
Balancing the chemical equation: It is important to recognize that in a chemical reaction, atoms are neither created nor destroyed. All atoms present in the reactants must be accounted for among the products.
The chemical equation for a reaction provides us with two important types of information:
1. The identities of the reactants and products
2. The relative numbers of each
Physical States:
(s) solid
(l) liquid
(g) gas
(aq) dissolved in water (in aqueous solution)
6.3 Balancing ChemicalEquations
The principle that lies at the heart of the balancing process is that atoms are conserved in a chemical reaction.
The identities (formulas) of the compounds must never be changed in balancing a chemical equation.
The accepted convention is that the “best” balanced equation is the one with the smallest integers (whole numbers). These integers are called the coefficients for the balanced equation.
How to Write and Balance Equations
Step 1 Read the description of the chemical reaction. What are the reactants, the products, and their states? Write the appropriate formulas.
Step 2 Write the unbalanced equation that summarizes the information fromstep 1.
Step 3 Balance the equation by inspection, starting with the most complicatedmolecule. Proceed element by element to determine what coefficients are necessary so that the same number of each type of atom appearson both the reactant side and the product side. Do not change theidentities (formulas) of any of the reactants or products.
Step 4 Check to see that the coefficients used give the same number of each type of atom on both sides of the arrow. (Note that an “atom” may be present in an element, a compound, or an ion.) Also check to see that the coefficients used are the smallest integers that give the balanced equation. This can be done by determining whether all coefficients can be divided by the same integer to give a set of smaller integer coefficients.
C7:Reactions in Aqueous Solutions
7.1 Predicting Whether a Reaction Will Occur
The most common of these driving forces are
1. Formation of a solid
2. Formation of water
3. Transfer of electrons
4. Formation of a gas
7.2 Reactions in Which a Solid Forms
One driving force for a chemical reaction is the formation of a solid, a process called precipitation. The solid that forms is called a precipitate, and the reaction is known as a precipitation reaction.
When each unit of a substance that dissolves in water produces separated ions, the substance is called a strong electrolyte
In this text we will use the term soluble solid to mean a solid that readily dissolves in water; the solid “disappears” as the ions are dispersed in the water.
The terms insoluble solid and slightly soluble solid are taken to mean the same thing: a solid where such a tiny amount dissolves in water that it is undetectable with the naked eye.
7.3 Describing Reactions in Aqueous Solutions
This is called the molecular equation for the reaction; it shows the complete formulas of all reactants and products
the complete ionic equation better represents the actual forms of the reactants and products in solution
Ions such as these, which do not participate directly in a reaction in solution, are called spectator ions
This equation, called the net ionic equation, includes only those components that are directly involved in the reaction.
7.4 Reactions That Form Water:Acids and Bases
It was not until the late 1800s that the essential nature of acids was discovered by Svante Arrhenius, then a Swedish graduate student in physics.
Arrhenius proposed that an acid is a substance that produces H+ ions (protons) when it is dissolved in water.
Because these substances are strong electrolytes that produce H ions, they are called strong acids.
Arrhenius also found that aqueous solutions that exhibit basic behavior always contain hydroxide ions. He defined a base as a substance that produces hydroxide ions (OH-) in water.
Because these hydroxide compounds are strong electrolytes that contain OH- ions, they are called strong bases.
The second product is an ionic compound, which might precipitate or remain dissolved, depending on its solubility. This ionic compound is called a salt.
7.5 Reactions of Metals with Nonmetals(Oxidation–Reduction)
A reaction that involves a transfer of electrons is called an oxidation–reduction reaction.
7.6 Ways to Classify Reactions
the formation of a solid when two solutions are mixed is called precipitation, we call this a precipitation reaction.
In the products these associations are reversed. Because of this double exchange, we sometimes call this reaction a double-exchange reaction or double-displacement reaction. AB+CD -->AD+CB
We classify these reactions as acid–base reactions. You can recognize this as an acid–base reaction because it involves an H+ ion that ends up in the product water.
The process of electron transfer is also called oxidation–reduction. Thus we classify the preceding reaction as an oxidation–reduction reaction.
We can call this a single-replacement reaction in contrast to double-displacement reactions, in which two types of anions are exchanged. We can represent a single replacement as A+BC-->B+AC
7.7 Other Ways to Classify Reactions
When a given compound is formed from simpler materials, we call this a synthesis (or combination) reaction.
Many chemical reactions that involve oxygen produce energy (heat) so rapidly that a flame results. Such reactions are called combustion reactions.
In many cases a compound can be broken down into simpler compounds or all the way to the component elements. This is usually accomplished by heating or by the application of an electric current. Such reactions are called decomposition reactions
C8:Chemical Composition
8.1 Counting by Weighing
8.2 Atomic Masses: Counting Atoms by Weighing
To avoid using terms like 1023 when describing the mass of an atom, scientists have defined a much smaller unit, of mass called the atomic mass unit, which is abbreviated amu. In termsof grams, 1 amu=1.66*10^-24g
Therefore, just as with the nonidentical jelly beans, we need to use an average mass for the carbon atoms. The average atomic mass for carbon atoms is 12.01 amu.
8.3 The Mole
The mole (abbreviated mol) can be defined as the number equal to the number of carbon atoms in 12.01 grams of carbon.
Techniques for counting atoms very precisely have been used to determine this number to be 6.022x10^23. This number is called Avogadro's number. One mole of something consists of 6.022x10^23 units of that substance.(This definition of the mole is slightly different from the SI definition but is used because it is easier to understand at this point)
8.4 Learning to Solve Problems
'Where Are We Going?'
'How Do We Get There?'
Reality Check
8.5 Molar Mass
The molar mass* of any substance is the mass (in grams) of 1 mole of the substance. The molar mass is obtained by summing the masses of the component atoms.
8.6 Percent Composition of Compounds
The mass fraction for each element is calculated as follows: Mass fraction for a given element = mass of the element present in 1 mole of compound / mass of 1 mole of compound
The mass fraction is converted to mass percent by multiplying by 100%.
The mass percent (sometimes called the weight percent)
8.7 Formulas of Compounds
The formula of a compound that expresses the smallest whole-number ratio of the atoms present is called the empirical formula or simplest formula.
The actual formula of a compound—the one that gives the composition of the molecules that are present—is called the molecular formula.
8.8 Calculation of Empirical Formulas
8.9 Calculation of Molecular Formulas
C9:Chemical Quantities
9.1 Information Given by Chemical Equations
It is important to recognize that the coefficients in a balanced equation give us the relative numbers of molecules.
9.2 Mole–Mole Relationships
we will develop a more convenient procedure, which uses conversion factors, or mole ratios, based on the balanced chemical equation.
9.3 Mass Calculations
The process of using a chemical equation to calculate the relative masses of reactants and products involved in a reaction is called stoichiometry. Chemists say that the balanced equation for a chemical reaction describes the stoichiometry of the reaction.
9.4 The Concept of Limiting Reactants
The reactant that runs out first and thus limits the amounts of products that can form is called the limiting reactant (limiting reagent).
9.5 Calculations Involving a Limiting Reactant
9.6 Percent Yield
Products stop forming when one reactant runs out. The amount of product calculated in this way is called the theoretical yield of that product. The actual yield of product, which is the amount of product actually obtained, is often compared to the theoretical yield. This comparison, usually expressed as a percentage, is called the percent yield. (Actual yield / Theoretical yield) x 100% = percent yield
C10:Energy
10.1 The Nature of Energy
For our purposes we will define energy as the ability to do work or produce heat.
One of the most important characteristics of energy is that it is conserved. The law of conservation of energy states that energy can be converted from one form to another but can be neither created nor destroyed.
Energy can be classified as either potential or kinetic energy
Potential energy is energy due to position or composition.
The kinetic energy of an object is energy due to the motion of the object and depends on the mass of the object m and its velocity v: KE=1/2 (mv^2 )
This brings us to a very important idea, the state function. A state function is a property of the system that changes independently of its pathway.
10.2 Temperature and Heat
Temperature is a measure of the random motions of the components of a substance.
Heat can be defined as a flow of energy due to a temperature difference.
10.3 Exothermic and Endothermic Processes
The system is the part of the universe on which we wish to focus attention
the surroundings include everything else in the universe
When a process results in the evolution of heat, it is said to be exothermic; that is, energy flows out of the system.
Processes that absorb energy from the surroundings are said to be endothermic. When the heat flow moves into a system, the process is endothermic
10.4 Thermodynamics
The study of energy is called thermodynamics.
The law of conservation of energy is often called the first law of thermodynamics and is stated as follows: The energy of the universe is constant.
The internal energy, E, of a system can be defined most precisely as the sum of the kinetic and potential energies of all “particles” in the system.
10.5 Measuring Energy Changes
In the metric system the calorie is defined as the amount of energy (heat) required to raise the temperature of one gram of water by one Celsius degree.
The “calorie” with which you are probably familiar is used to measure the energy content of food and is actually a kilocalorie (1000 calories), written with a capital C (Calorie) to distinguish it from the calorie used in chemistry.
The joule (an SI unit) can be most conveniently defined in terms of the calorie: 1 calorie = 4.184 joules
The amount of energy required to change the temperature of one gram of a substance by one Celsius degree is called its specific heat capacity or, more commonly, its specific heat.
Energy (heat) required (Q)=Specific heat capacity (s) x Mass (m) in grams of sample x Change in temperature (delta T) in °C
10.6 Thermochemistry (Enthalpy)
For a reaction occurring under conditions of constant pressure, the change in enthalpy (delta H) is equal to the energy that flows as heat
A calorimeter is a device used to determine the heat associated with a chemical reaction.
10.7 Hess’s Law
One of the most important characteristics of enthalpy is that it is a state function.
Consequently, in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps. This principle is known as Hess's law.
two characteristics of H for a reaction:
1. If a reaction is reversed, the sign of H is also reversed.
2. The magnitude of H is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of H is multiplied by the same integer.
10.8 Quality Versus Quantity of Energy
10.9 Energy and Our World
By the process of photosynthesis, plants store energy that can be claimed by burning the plants themselves or the decay products that have been converted over millions of years to fossil fuels.
Petroleum is a thick, dark liquid composed mostly of compounds called hydrocarbons that contain carbon and hydrogen.
Natural gas, usually associated with petroleum deposits, consists mostly of methane, but it also contains significant amounts of ethane, propane, and butane.
Coal was formed from the remains of plants that were buried and subjected to high pressure and heat over long periods of time.
In a way, the atmosphere acts like the glass of a greenhouse, which is transparent to visible light but absorbs infrared radiation, thus raising the temperature inside the building.greenhouse effect
There are several potential energy sources: the sun (solar), nuclear processes (fission and fusion), biomass (plants), and synthetic fuels.
10.10 Energy as a Driving Force
two very important driving forces:
Energy spread
Energy spread means that in a given process, concentrated energy is dispersed widely.
Matter spread
Matter spread means exactly what it says: the molecules of a substance are spread out and occupy a larger volume
Entropy is a function we have invented to keep track of the natural tendency for the components of the universe to become disordered—entropy (designated by the letter S) is a measure of disorder or randomness.
This idea leads to a very important conclusion that is summarized in the second law of thermodynamics: The entropy of the universe is always increasing.
A spontaneous process is one that occurs in nature without outside intervention—it happens “on its own.”
C13:Gases
13.1 Pressure
A device that measures atmospheric pressure, the barometer, was invented in 1643 by an Italian scientist named Evangelista Torricelli (1608– 1647), who had been a student of the famous astronomer Galileo.
The unit mm Hg (millimeters of mercury) is often called the torr in honor of Torricelli. A related unit for pressure is the standard atmosphere.
1 standard atmosphere = 1.000 atm = 760.0 mm Hg = 760.0 torr = 101,325 Pa
The SI unit for pressure is the pascal (abbreviated Pa).
A unit of pressure that is employed in the engineering sciences and that we use for measuring tire pressure is pounds per square inch, abbreviated psi.
1.000 atm = 14.69 psi
13.2 Pressure and Volume: Boyle's Law
For Boyle’s law to hold, the amount of gas (moles) must not be changed. The temperature must also be constant.-- pressure times volume equals a constant. PV = k--k=1.41 x 10^3 (in Hg) x in.^3.
13.3 Volume and Temperature: Charles's Law
All of the lines extrapolate to zero volume at the same temperature: -273 °C. This suggests that -273 °C is the lowest possible temperature, becausea negative volume is physically impossible. In fact, experiments haveshown that matter cannot be cooled to temperatures lower than -273 °C. Therefore, this temperature is defined as absolute zero on the Kelvin scale.
The direct proportionality between volume and temperature (in kelvins) is represented by the equation known as Charles's law: V = bT (Charles's law applies only when both the amount of gas (moles) and the pressure are constant.)
13.4 Volume and Moles:Avogadro's Law
In words, this equation means that for a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas. This relationship is called Avogadro's law after the Italian scientist Amadeo Avogadro, who first postulated it in 1811. V=an
13.5 The Ideal Gas Law
where R is the combined proportionality constant and is called the universal gas constant
the ideal gas law: PV = nRT
gas that obeys this equation is said to behave ideally. That is, this equation defines the behavior of an ideal gas. Most gases obey this equation closely at pressures of approximately 1 atm or lower, when the temperature is approximately 0 °C or higher.
P1V1/T1 = P2V2/T2 is often called the combined gas law equation. It holds when the amount of gas (moles) is held constant.
13.6 Dalton’s Law of Partial Pressures
For a mixture of gases in a container, the total pressure exerted is the sum of the partial pressures of the gases present. The partial pressure of a gas is the pressure that the gas would exert if it were alone in the container. This statement, known as Dalton's law of partial pressures Ptotal = P1 + P2 + P3 = ntotala(RT/V)
The fact that the pressure exerted by an ideal gas is affected by the number of gas particles and is independent of the nature of the gas particles tells us two important things about ideal gases:
1. The volume of the individual gas particle (atom or molecule) mustnot be very important.
2. The forces among the particles must not be very important.
13.7 Laws and Models: A Review
13.8 The Kinetic Molecular Theory of Gases
A relatively simple model that attempts to explain the behavior of an ideal gas is the kinetic molecular theory. This model is based on speculations about the behavior of the individual particles (atoms or molecules) in a gas.
13.9 The Implications of the Kinetic Molecular Theory
In fact, the Kelvin temperature of a gas is directly proportional to the average kinetic energy of the gas particles
13.10 Gas Stoichiometry
This volume of 22.4 L is called the molar volume of an ideal gas.
The conditions 0 °C and 1 atm are called standard temperature and pressure (abbreviated STP). Properties of gases are often given under these conditions. Remember, the molar volume of an ideal gas is 22.4 L at STP. That is, 22.4 L contains 1 mole of an ideal gas at STP.
C15:Solutions
15.1 Solubility
15.2 Solution Composition: An Introduction
When a solution contains as much solute as will dissolve at that temperature, we say it is saturated.
A solution that has not reached the limit of solute that will dissolve in it is said to be unsaturated.
A relatively large amount of solute is dissolved in a concentrated solution (strong coffee is concentrated). A relatively small amount of solute is dissolved in a dilute solution (weak coffee is dilute).
15.3 Solution Composition: Mass Percent
One common way of describing a solution’s composition is mass percent (sometimes called weight percent), which expresses the mass of solute present in a given mass of solution. Mass percent = mass of solute / mass of solution x 100%
15.4 Solution Composition: Molarity
We define the concentration of a solution as the amount of solute in a given volume of solution. The most commonly used expression of concentration is molarity (M). Molarity describes the amount of solute in moles and the volume of the solution in liters. Molarity is the number of moles of solute per volume of solution in liters. M = molarity = moles of solute / liters of solution = mol / L
15.5 Dilution
The process of adding more solvent to a solution is called dilution
Moles of solute after dilution = moles of solute before dilution
15.6 Stoichiometry of Solution Reactions
15.7 Neutralization Reactions
An acid–base reaction is often called a neutralization reaction. When just enough strong base is added to react exactly with the strong acid in a solution, we say the acid has been neutralized. One product of this reaction is always water.
15.8 Solution Composition: Normality
One equivalent of an acid is the amount of that acid that can furnish 1 mole of H+ ions.
One equivalent of a base is defined as the amount of that base that can furnish 1 mole of OH- ions.
The equivalent weight of an acid or a base is the mass in grams of 1 equivalent (equiv) of that acid or base
Normality (N) is defined as the number of equivalents of solute per liter of solution. Normality = N = number of equivalents /1 liter of solution = equivalents / liter = equiv / L
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