CHM1045 Exam 1 Study Guide (ch.1-3)
CHM1045 Exam 1 Study Guide (ch.1-3) CHM1045
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CHAPTER 1. Chemistry 1.1 Chemistry: A Science for the Twenty-First Century Chemistry is the study of matter and the changes it undergoes • also called the central science 1.2 Chemistry: The Study of Chemistry How a chemist thinks: to look at the macroscopic world (the things we can see, touch and measure directly) and visualize the particles and events of the microscopic world (we cannot experience without modern technology and our imaginations) 1.3 Chemistry: The Scientific Method Scientific method: a systematic approach to research • Step 1: OBERVATION o carefully define the problem o deals with events in the macroscopic world § NOTE: atoms and molecules are in the microscopic world • Step 2: REPRESENTATION o performing experiments, making careful observations, and recording information, or data, about the system o representation is a scientific shorthand for describing an experiment in symbols and chemical equations § Types of data obtained in a research study: • Qualitative: consisting of general observations about the system, and • Quantitative: comprising numbers obtained by various measurements of the system • Step 3: INTERPRETATION o after the data has been recorded, the next step is to attempt to explain the observed phenomenon § Hypothesis: a tentative explanation for a set of observations. • formed based on the recorded data Law: is a concise verbal or mathematical statement of a relationship between phenomena that is always the same under the same conditions. • After a large amount of data has been collected, it is important to summarize the information in a concise way (as a law) Theory: is a unifying principle that explains a body of facts and/or those laws that are based on them. • Hypotheses that survive many experimental tests of their validity may evolve into theories • Theories are constantly being tested • If a theory is disproved by experiment, then it must be discarded or modified so that it becomes consistent with experimental observations. • Proving or disproving a theory can take years, even centuries KEEP IN MIND: • Scientific progress is seldom made in step-by-step fashion. • Sometimes a law precedes a theory; sometimes it is the other way around. • Two scientists may start working on a project with exactly the same objective, but will end up taking different approaches 1.4 Chemistry: Classifications of Matter Matter: is anything that occupies space and has mass. • What we can see and touch (such as water, earth, and trees), as well as things we cannot (such as air) • everything in the universe has a “chemical” connection Substances and Mixtures Substance: is a form of matter that has a definite (constant) composition and distinct properties o Ex. water, ammonia, table sugar (sucrose), gold, and oxygen. o differ from one another in composition o can be identified by their appearance, smell, taste, and other properties. Types of Substances: • Element: is a substance that cannot be separated into simpler substances by chemical means • 118 elements have been positively identified. • Most of them occur naturally on Earth Atoms of most elements can interact with one another to form compounds • Compound: a substance composed of atoms of two or more elements chemically united in fixed proportions. • Unlike mixtures, compounds can be separated only by chemical means into their pure components. Mixture: is a combination of two or more substances in which the substances retain their distinct identities o any mixture can be created and then separated by physical means into pure components without changing the identities of the components o Ex. are air, soft drinks, milk, and cement. o do not have constant composition • ex. samples of air collected in different cities would probably differ in composition because of differences in altitude, pollution, and so on Types of Mixtures: • homogeneous mixture: the composition of the mixture is the same throughout. o ex. When a spoonful of sugar dissolves in water • heterogeneous mixture: the composition is not uniform o ex. If sand is mixed with iron filings, however, the sand grains and the iron filings remain separate 1.5 Chemistry: The Three States of Matter Atomic Theory of Matter * macroscopic properties (easily visible) of each physical state are caused by the microscopic properties to their left 1.6 Chemistry: Physical and Chemical Properties of Matter I. Categories: Substances are identified by their properties as well as by their composition Physical property: any property of a substance that can be measured and observed without changing the composition or identity • ex. we can measure the melting point of ice by heating a block of ice and recording the temperature at which the ice is converted to water. Water differs from ice only in appearance, not in composition, so this is a physical change; • ex. we can freeze the water to recover the original ice. Therefore, the melting point of a substance is a physical property Chemical property: any property of a substance that cannot be studied without converting the substance into some other substance • ex. Burning II. Subcategories: all measurable properties of matter fall into one of two additional categories Extensive property: depends on how much matter is being considered • ex. mass: the quantity of matter in a given sample of a substance • values of the same extensive property can be added together • ex. volume: length cubed • value of an extensive quantity depends on the amount of matter Intensive property: does not depend on how much matter is being considered • ex. density: the mass of an object divided by its volume • ex. temperature • unlike mass, length, and volume, temperature and other intensive properties are not additive • ex. Suppose that we have two beakers of water at the same temperature. If we combine them to make a single quantity of water in a larger beaker, the temperature of the larger quantity of water will be the same as it was in two separate beakers 1.7 Chemistry: Measurement macroscopic properties: can be determined directly vs microscopic properties: on the atomic or molecular scale, must be determined by an indirect method I. SI Units Measurement : Base unit • Length - meter (me) • Mass - gram (g) • Volume - liter (L) • Time - second (s) • Frequency - hertz (Hz) • Bigger è smaller : multiply • Smaller è bigger : divide II Mass and Weight Mass is a measure of the amount of matter in an object • The SI unit of mass is the kilogram (kg) vs. Weight: is the force that gravity exerts on an object III. Volume • The SI unit of length is the meter (m), and the SI-derived unit for volume is the cubic meter (m ) • Another common unit of volume is the liter (L) o Liter: is the volume occupied by one cubic decimeter. 3 § One liter of volume = 1000 milliliters (mL) or 1000 cm § and one milliliter is equal to one cubic centimeter IV. Density or • because density is an intensive property and does not depend on the quantity of mass present, for a given substance the ratio of mass to volume always remains the same • V increases as m does • usually decreases with temperature • The SI-derived unit for density is the kilogram per cubic meter (kg/m ) Densities of Some Substances at 25°C 3 Substance Density (g/cm ) Air* 0.001 Ethanol 0.79 Water 1.00 Graphite 2.2 Table salt 2.2 Aluminum 2.70 Diamond 3.5 Iron 7.9 Lead 11.3 Mercury 13.6 Gold 19.3 Osmium † 22.6 *Measured at 1 atmosphere. † Osmium (Os) is the densest ele ment known. V. Temperature Scales The Fahrenheit scale: • most commonly used scale in the United States outside the laboratory • normal freezing and boiling points of water to be exactly 32°F and 212°F The Celsius scale: • divides the range between the freezing point (0°C) and boiling point (100°C) of water into 100 degrees. The Kelvin scale: • is the SI base unit of temperature • it is the absolute temperature scale: zero on the Kelvin scale, denoted by 0 K, is the lowest temperature that can be attained theoretically Note that there are 100 divisions, or 100 degrees, between the freezing point and the boiling point of water on the Celsius scale, and there are 180 divisions, or 180 degrees, between the same two temperature limits on the Fahrenheit scale. Th e Celsius scale was formerly called the centigrade scale. Note that the Kelvin scale does not have the degree sign. Also, temperatures expressed in kelvins can never be negative. • F à C o The size of a degree on the Fahrenheit scale is only 100/180, or 5/9, of a degree on the Celsius scale. To convert degrees Fahrenheit to degrees Celsius, • C à F o The following equation is used to convert degrees Celsius to degrees Fahrenheit: • C à K o Both the Celsius and the Kelvin scales have units of equal magnitude; that is, one degree Celsius is equivalent to one kelvin. Experimental studies have shown that absolute zero on the Kelvin scale is equivalent to −273.15°C on the Celsius scale 1.8 Chemistry: Handling Number • Significant figures: are the meaningful digits in a measured or calculated quantity 1. any digit that is not zero is significant 2. zeros between nonzero digits are significant 3. zeros to the left of the first nonzero digit are not significant 4. if a number is greater than 1, then all the zeros written to the right of the decimal point count as significant figures 5. for numbers that do not contain decimal points, the trailing zeros (that is, zeros after the last nonzero digit) may or may not be significant § addition and subtraction: • the answer cannot have more digits to the right of the decimal point than either of the original numbers § multiplication and division: • the number of significant figures in the final product or quotient is determined by the original number that has the smallest number of sig figs 1.9 Chemistry: Dimensional Analysis in Solving Problems Dimensional analysis: (also called the factor-label method) is the procedure we use to convert between units in solving chemistry problems § based on the relationship between different units that express the same physical quantity. § in general, to apply dimensional analysis we use the relationship § and the units cancel as follows: Remember that the unit we want appears in the numerator and the unit we want to cancel appears in the denominator. Problem Solving Example: Example: Example: 1.10 Chemistry: Real-World Problem Solving: Information, Assumptions, and Simplifications Chapter 1 ‘Chemistry’ summary I. Key Equations: Equation for density Converting °F to °C Converting °C to °F Converting °C to K II. Summary: 1. The study of chemistry involves 3 basic steps: § Observation: refers to measurements in the macroscopic world; § Representation: involves the use of shorthand notation symbols and equations for communication; § Interpretation: are based on atoms and molecules, which belong to the microscopic world. 2. The scientific method: is a systematic approach to research that begins with the gathering of information through observation and measurements. § In the process, hypotheses, laws, and theories are devised and tested. 3. Chemists study matter and the changes it undergoes. § Substances: o have unique physical properties that can be observed without changing their identity and unique chemical properties that, when they are demonstrated, do change the identity of the substances. o can exist in three states: solid, liquid, and gas. (the interconversion between these states can be effected by changing the temperature) § Elements: The simplest substances in chemistry § Compounds: are formed by the chemical combination of atoms of different elements in fixed proportions. § Mixtures: o can be separated into pure components by physical means § homogeneous (uniform) § heterogeneous (not uniform) 4. SI units are used to express physical quantities in all sciences, including chemistry 5. Numbers expressed in scientific notation have the form N × 10 n § where N is between 1 and 10, and n is a positive or negative integer § this helps us handle very large and very small quantities. Chapter 2: Atoms, Molecules, and Ions 2.1 Atoms, Molecules, and ions: the atomic theory Dalton’s hypotheses: about the nature of matter The Atomic Theory: I. First hypotheses: • states: elements are composed of extremely small particles called atoms. II. Second hypotheses: • states: the atoms of one element are different from the atoms of all other elements o All atoms of a given element are identical, having the same size, mass, and chemical properties. o Dalton made no attempt to describe the structure or composition of atoms (he had no idea what an atom is really like) o But he did realize that the different properties shown by elements III. Third hypotheses: • states: to form a certain compound (composed of one or more element) , we need not only atoms of the right kinds of elements, but specific numbers of these atoms as well • in any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction o Law of definite proportions: (integer) § different samples of the same compound always contain its constituent elements in the same proportion by mass o Law of multiple proportions: (simple faction) § If two elements can combine to form more than one type of compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers. § different compounds made up of the same elements differ in the number of atoms of each kind that combine IV. Fourth Hypotheses: • A chemical reaction involves only the separation, combination, or rearrangement of atoms; no creation or destruction o Law of conservation of mass: § matter can be neither created nor destroyed o Because matter is made of atoms that are unchanged in a chemical reaction, it follows that mass must be conserved as well According to Dalton's atomic theory, atoms of the same element are identical, but atoms of one element are different from atoms of other elements. (b) Compound formed from atoms of elements X and Y. In this case, the ratio of the atoms of element X to the atoms of element Y is 2:1. Note : that a chemical reaction results only in the rearrangement of atoms, not in their destruction or creation. 2.2 Atoms, Molecules, and Ions: The Structure of the atom The study of radiation: Def. the emission and transmission of energy through space in the form of waves § Helped gain information to understand atomic structure o Cathode ray: § Instrument used to investigate atoms § It is a glass tube from which most of the air has been evacuated § Cathode: • a negatively charged plate that emits an invisible ray, when the two metal plates are connected to a high-voltage source § Anode: • A positively charged plate that the cathode ray is drawn to. • There it passes through a hole and continues traveling to the other end of the tube • When the ray strikes the specially coated surface, it produces a strong fluorescence (bright light) Radioactivity: Def. the spontaneous breakdown of an atom by emission of particles and/or radiation. § any element that spontaneously emits radiation is said to be radioactive. § 3 types of rays are produced by the decay (breakdown) of radioactive substances: o Alpha (α) rays § Helium ions with a positive charge of +2. § consist of positively charged particles, called α particles § are deflected by the positively charged plate o Beta (β) rays § consist of negatively charged particles, called β particles § electrons § are deflected by the negatively charged plate o gamma (γ) rays § consists of high-energy rays called § no charge and are not affected by an external field Atom: Def. the basic unit of an element that can enter into chemical combination § To maintain electric neutrality, an atom must contain an equal number of positive and negative charges I. The Electron § Def. a subatomic particle that has a very low mass and carries a single negative electric charge § Rotating the nucleus −19 § Charge: −1.6022 × 10 C II. The Proton § Def. a subatomic particle having a single positive electric charge § The mass of a proton is about 1840 times that of an electron § In the nucleus: the central core of an atom o The mass of a nucleus constitutes most of 13e mass of the entire atom, but the nucleus occupies only about 1/10 of the volume of the atom o atomic (and molecular) dimensions are in picometer (pm) III. The Neutron § Def. a subatomic particle that bears no net electric charge. § mass is slightly greater than a proton's § in the nucleus Mass and Charge of Subatomic Particles Charge Particle Mass (g) Coulomb Charge Unit −28 −19 Electron* 9.10938 × 10 −1.6022 × 10 −1 Proton 1.67262 × 10 −24 +1.6022 × 10 −19 +1 −24 Neutron 1.67493 × 10 0 * coulomb: is the equivalent of one ampere-second // An electric current of A represents 1 C of unit electric charge carriers flowing past a specific point in 1 s 2.3 Atoms, Moles, and Ions: Atomic Number, Mass Number, and Isotopes All atoms can be identified by the number of protons and neutrons they contain I. Atomic number (Z) Def. the number of protons in the nucleus of each atom of an element • in a neutral atom the number of protons is equal to the number of electrons, so the atomic number also indicates the number of electrons present in the atom • the chemical identity of an atom can be determined solely from its atomic number II. Mass number (A) Def. is the total number of neutrons and protons present in the nucleus of an atom of an element. • all atomic nuclei contain both protons and neutrons o except for the most common form of hydrogen, which has one proton and no neutrons Note: that all three quantities (atomic number, number of neutrons, and mass number) must be positive integers, or whole nu.bers III. Isotopes Def. atoms that have the same atomic number but different mass numbers • isotopes of the same element have similar chemistries (forming the same types of compounds and displaying similar reactivities) Ex. 2.4 Atoms, Molecules, and Ions: The Periodic Table Periodic table: Def. a chart in which elements having similar chemical and physical properties are grouped together • elements are arranged by the atomic number • Periods: § horizontal rows • Groups (or families): § Vertical columns § similarities in their chemical properties • The elements can be divided into 3 categories: I. Metal: • Elements that are good conductors of heat and electricity and have the tendency to form positive ions in ionic compounds II. Nonmetal: • Elements that are usually poor conductors of heat and electricity III. Metalloid: • An element with properties intermediate between those of metals and nonmetals • Special elements: § alkali metals: o The Group 1A elements (Li, Na, K, Rb, Cs, and Fr) § alkaline earth metals: o the Group 2A elements (Be, Mg, Ca, Sr, Ba, and Ra) § halogens: o Elements in Group 7A (F, Cl, Br, I, and At) § noble gases: or rare gases o elements in Group 8A (He, Ne, Ar, Kr, Xe, and Rn) 2.5 Atoms, Molecules, and Ions: Molecules and Ions Molecules: Def. is an aggregate of at least two atoms in a definite arrangement held together by chemical forces (also called chemical bonds) § may contain atoms of the same element or atoms of two or more elements joined in a fixed ratio, in accordance with the law of definite proportions § not necessarily a compound (made up of two or more elements) § diatomic molecule: o contains only two atoms § polyatomic molecules: o containing more than two atoms o They can be atoms of the same element, as in ozone (O ), wh3ch is made up of three atoms of oxygen, or they can be combinations of two or more different elements. Ions: Def. an atom or a group of atoms that has a net positive or negative charge § The number of positively charged protons in the nucleus of an atom remains the same during ordinary chemical changes (called chemical reactions), but negatively charged electrons may be lost or gained § With very few exceptions, metals tend to form cations and nonmetals form anions § ionic compound: o Any neutral compound containing cations and anions § Cation: • an ion with a net positive charge • the results of the loss of one or more electrons from a neutral atom § Anion: • ion whose net charge is negative • due to an increase in the number of electrons § monatomic ion: o an ion that contains only one atom 2.6 Atoms, Molecules, and Ions: Chemical Formulas Chemical formulas: Def. An expression showing the chemical composition of a compound in terms of the symbols for the atoms of the elements involved • By composition we mean not only the elements present but also the ratios in which the atoms are combined. Two types of formulas: I. Molecular formulas: Def. An expression showing the exact numbers of atoms of each element in a molecule in the smallest unit of substance o Allotropes: § Two or more forms of the same element that differ significantly in chemical and physical properties § Ex. oxygen (O )2and ozone (O ) 3re allotropes of oxygen § Molecular models: o Two standard types of molecular models: 1. Ball-and-stick model: • Pros: show the three-dimensional arrangement of atoms clearly • Cons: not proportional • the atoms are wooden or plastic balls with holes in them • Sticks or springs = represent chemical bonds • The angles they form between atoms approximate the bond angles in actual molecules • (with the exception of the H atom) the balls are all the same size and each type of atom is represented by a specific color. 2. Space-filling model: • Pros: more accurate because they show the variation in atomic size • Cons: more time consuming time-consuming to put together and they do not show the three-dimensional positions of atoms very well • atoms are represented by truncated balls held together by snap fasteners, so that the bonds are not visible • balls are proportional in size to atoms § Building a Molecular model: 1. Structural formula: o Def. shows how atoms are bonded to one another in a molecule o First step of building a molecular model is to write the structural formula. o A line connecting the two atomic symbols represents a chemical bond. II. Empirical Formulas: Def. tells us which elements are present and the simplest whole number ratio of their atoms, but not necessarily the actual number of atoms in a given molecule § The simplest chemical formulas § they are written by reducing the subscripts in the molecular formulas to the smallest possible whole numbers § Molecular formulas are the true formulas of molecules § If we know the molecular formula, we also know the empirical formula, but the reverse is not true Example of difference: 2.7 Atoms, Molecules, and Ions: Naming Compounds Organic compounds: Def. contain carbon (usually in combination with elements such as hydrogen, oxygen, nitrogen, and sulfur) vs. Inorganic compounds: Def. All other compounds (however, some carbon-containing compounds groups are considered to be inorganic compounds) 4 categories of inorganic compounds: I. Ionic compounds: • ionic compounds are made up of cations (positive ions) and anions (negative ions) • many ionic compounds are binary compounds • binary compounds: o compounds formed from just two elements o the first element named is the metal cation, followed by the nonmetallic anion • ternary compounds: o compounds consisting of three elements • Certain metals (especially the transition metals) can form more than one type of cation 2+ 3+ o Ex. Iron can form two cations: Fe and Fe . • Stock System: Mn : MnO manganese(II) oxide Mn : Mn O manganese(III) oxide 4+ 2 3 Mn : MnO 2 manganese(IV) oxide Keep in mind that the Roman numerals refer to the charges on the metal cations. II. Molecular compounds (binary molecular): • Unlike ionic compounds, molecular compounds contain discrete molecular units • They are usually composed of nonmetallic elements • Many molecular compounds are binary compounds • Naming binary molecular compounds is similar to naming binary ionic compounds o We place the name of the first element in the formula first, and the second element is named by adding -ide to the root of the element name Ex. CO carbon monoxide CO 2 carbon dioxide SO 2 sulfur dioxide SO 3 sulfur trioxide NO 2 nitrogen dioxide N O dinitrogen tetroxide 2 4 Greek Prefixes Used in Naming Molecular Compounds: Prefix Meaning mono- 1 di- 2 tri- 3 tetra- 4 penta- 5 hexa- 6 hepta- 7 octa- 8 nona- 9 deca- 10 • The following guidelines are helpful in naming compounds with prefixes: o The prefix “mono-” may be omitted for the first element. For example, PCl is 3 named phosphorus trichloride, not monophosphorus trichloride o For oxides, the ending “a” in the prefix is sometimes omitted. For example, N O 2 4 may be called dinitrogen tetroxide rather than dinitrogen tetraoxide. • Exceptions to the use of Greek prefixes are molecular compounds containing hydrogen B 2 6 diborane CH methane 4 SiH 4 silane NH 3 ammonia PH 3 phosphine H 2 water H 2 hydrogen sulfide ü Summary of the steps for naming ionic and binary molecular compounds: III. Acids and bases: Acids: Def. a substance that yields hydrogen ions (H ) when dissolved in water. (H is equivalent to one proton, and is often referred to that way.) • Formulas for acids contain one or more hydrogen atoms as well as an anionic group • Anions whose names end in “-ide” form acids with a “hydro-” prefix and an “-ic” ending • Note: acids all exist as molecular compounds in the gas phase • the name assigned to the compound depends on its physical state • Oxoacids: o are acids that contain hydrogen, oxygen, and another element (the central element) o formulas written with the H first, followed by the central element, then O • Starting with our reference oxoacids whose names all end with “-ic, we use the following rules to name these compounds: 1. Addition of one O atom to the “-ic” acid: The acid is called “per … -ic” acid. Thus, adding an O atom to HClO changes chloric acid to perchloric acid, 3 HClO .4 2. Removal of one O atom from the “-ic” acid: The acid is called “-ous” acid. Thus, nitric acid, HNO ,3becomes nitrous acid, HNO . 2 3. Removal of two O atoms from the “-ic” acid: The acid is called “hypo … - ous” acid. Thus, when HBrO is 3onverted to HBrO, the acid is called hypobromous acid. • The rules for naming oxoanions (anions of oxoacids) 1. When all the H ions are removed from the “-ic” acid, the anion's name ends with “-ate.” For example, the anion CO 32− derived from H 2O is3called carbonate. 2. When all the H ions are removed from the “-ous” acid, the anion's name ends with “-ite.” Thus, the anion ClO 2erived from HClO is c2lled chlorite. 3. The names of anions in which one or more but not all the hydrogen ions have been removed must indicate the number of H ions present Bases: Def. as a substance that yields hydroxide ions (OH ) when dissolved in water ex. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH) 2 barium hydroxide V. Hydrates: Def. compounds that have a specific number of water molecules attached to them. Example: § in its normal state, each unit of copper(II) sulfate has five water molecules associated with it. § The systematic name for this compound is copper(II) sulfate pentahydrate, and its formula is written as CuSO 4 5H 2 § The water molecules can be driven off by heating § When this occurs, the resulting compound is CuSO , w4ich is sometimes called anhydrous copper(II) sulfate; “anhydrous” means that the compound no longer has water molecules associated Chapter 2 ‘Atoms, Molecules, and ions’ summary I. Key Equations: II. Summary: 1. Modern chemistry began with Dalton's atomic theory of matter • all matter is composed of tiny, indivisible particles called atoms • all atoms of the same element are identical • compounds contain atoms of different elements combined in whole- number ratios • atoms are neither created nor destroyed in chemical reactions (the law of conservation of mass) 2. law of definite proportions: • Atoms of constituent elements in a particular compound are always combined in the same proportions by mass law of multiple proportions: • When two elements can combine to form more than one type of compound, the masses of one element that combine with a fixed mass of the other element are in a ratio of small whole numbers 3. An atom consists of a very dense central nucleus containing protons and neutrons, with electrons moving about the nucleus at a relatively large distance from it. 4. Protons are positively charged neutrons have no charge electrons are negatively charged • Protons and neutrons have roughly the same mass, which is about 1840 times greater than the mass of an electron. 5. atomic number: of an element is the number of protons in the nucleus of an atom of the element; it determines the identity of an element mass number: is the sum of the number of protons and the number of neutrons in the nucleus 6. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. 7. Chemical formulas combine the symbols for the constituent elements with whole-number subscripts to show the type and number of atoms contained in the smallest unit of a compound. 8. The molecular formula conveys the specific number and type of atoms combined in each molecule of a compound. The empirical formula shows the simplest ratios of the atoms combined in a molecule. 9. Chemical compounds are either • molecular compounds (in which the smallest units are discrete, individual molecules) • ionic compounds, which are made of cations and anions. 10. The names of many inorganic compounds can be deduced from a set of simple rules. The formulas can be written from the names of the compounds. 11. Organic compounds contain carbon and elements like hydrogen, oxygen, and nitrogen. • Hydrocarbon is the simplest type of organic compound. CHAPTER 3. MASS RELATIONSHIPS IN CHEMICAL REACTIONS 3.1 MASS RELATIONSHIPS IN CHEMICAL REACTIONS: Atomic mass atomic mass: (sometimes called atomic weight) is the mass of the atom in atomic mass units (amu) One atomic mass unit is defined as a mass exactly equal to one-twelfth the mass of one carbon-12 atom Dalton: What one atomic mass unit is called average atomic mass: • when we measure the atomic mass of an element, we must generally settle for the average mass of the naturally occurring mixture of isotopes • ex. o the natural abundances of carbon-12 and carbon-13 are 98.90 percent and 1.10 percent, respectively. o The atomic mass of carbon-13 has been determined to be 13.00335 amu. o Thus, the average atomic mass of carbon can be calculated as follows: 3.2 Avogadro’s Number and the Molar Mass of an Element Mole (mol) Def. is the amount of a substance that contains as many elementary entities (atoms, molecules, or other particles) • as there are atoms in exactly 12 g (or 0.012 kg) of the carbon-12 isotope • The adjective formed from the noun “mole” is “molar.” Avogadro's number (N )A 6.022 × 10 ; the number of particles in a mole Molar mass (M) Def. the mass (in grams or kilograms) of one mole of atoms, molecules, or other particles • In calculations, the units of molar mass are g/mol or kg/mol. Ex. • the mass of one carbon-12 atom is given by Ø We can use the preceding result to determine the relationship between atomic mass units and grams. o Because the mass of every carbon-12 atom is exactly 12 amu, the number of atomic mass units equivalent to 1 gram is o Thus, and • Conversion o The notions of Avogadro's number and molar mass enable us to carry out conversions between mass and moles of atoms and between moles and number of atoms 3.3 Molecular Mass Molecular mass (sometimes called molecular weight) Def. is the sum of the atomic masses (in amu) in the molecule. • For example, the molecular mass of 2 O is or • In general, we need to multiply the atomic mass of each element by the number of atoms of that element present in the molecule and sum over all the elements 3.4 The Mass Spectrometer Mass spectrometry: Def. The most direct and most accurate method for determining atomic and molecular masses 3.5 Percent Composition of Compounds Percent composition by mass: Def. is the percent by mass of each element in a compound. • Percent composition is obtained by: o dividing the mass of each element in 1 mole of the compound by the molar mass of the compound o and multiplying by 100 percent • where n is the number of moles of the element in 1 mole of the compound • For example: o in 1 mole of hydrogen peroxide (H O ) 2h2re are 2 moles of H atoms and 2 moles of O atoms o The molar masses of H O , 2,2and O are 34.02 g, 1.008 g, and 16.00 g, respectively o Therefore, the percent composition of H O i2 c2lculated as follows: v Procedure for calculating the empirical formula of a compound from its percent compositions: § The procedure used in the example can be reversed if necessary o Given the percent composition by mass of a compound, we can determine the empirical formula of the compound o Because we are dealing with percentages and the sum of all the percentages is 100 percent à it is convenient to assume that we started with 100 g of a compound ß 3.6 Experimental Determination of Empirical Formulas The percent composition allows you to identify compounds experimentally: • Steps 1. The chemical analysis tells us the number of grams of each element present in a given amount of a compound 2. convert the quantities in grams to number of moles of each element 3. using the method given in the example below we find the empirical formula of the compound. • Ex. oThe masses of CO an2 of H O p2oduced can be determined by measuring the increase in mass of the CO 2nd H O2absorbers, respectively. oSuppose that in one experiment the combustion of 11.5 g of ethanol produced 22.0 g of CO 2nd 13.5 g of H O2 oWe can calculate the mass of carbon and hydrogen in the original 11.5-g sample of ethanol as follows: o Thus, 11.5 g of ethanol contains 6.00 g of carbon and 1.51 g of hydrogen. The remainder must be oxygen, whose mass is o o The number of moles of each element present in 11.5 g of ethanol is o The formula of ethanol is therefore C 0.50 1.5 0.25 o Because the number of atoms must be an integer, we divide the subscripts by 0.25, the smallest subscript, and obtain for the empirical formula Answer: C H2O6 Determination of Molecular Formulas: o The formula calculated from percent composition by mass is always the empirical formula because the subscripts in the formula are always reduced to the smallest whole numbers. o To calculate the actual, molecular formula we must know the approximate molar mass of the compound in addition to its empirical formula. o Knowing that the molar mass of a compound must be an integral multiple of the molar mass of its empirical formula, we can use the molar mass to find the molecular formula 3.7 Chemical Reactions and Chemical Equations Chemical reaction: Def. a process in which a substance (or substances) is changed into one or more new substances. Chemical equation: Def. uses chemical symbols to show what happens during a chemical reaction. In this section, we will learn how to write chemical equations and balance them. I. Writing Chemical Equations: Example: o “plus” sign means “reacts with” o arrow means “to yield.” o interpretation: “Molecular hydrogen reacts with molecular oxygen to yield water.” o The reaction is assumed to proceed from left to right as the arrow indicates. • Balancing: § It is not complete because there are twice as many oxygen atoms on the left side of the arrow (two) as on the right side (one) o To conform with the law of conservation of mass, there must be the same number of each type of atom on both sides of the arrow o Must have as many atoms after the reaction ends as we did before it started. § We can balance the example by placing the appropriate coefficient o (2 in this case) in front o2 H and 2 O: NOTE: When the coefficient is 1, as in the case of O2, it is not shown. • 3 Ways of reading a balanced equation: • Three ways of representing the combustion of hydrogen. In accordance with the law of conservation of mass, the number of each type of atom must be the same on both sides of the equation. 1. “two hydrogen molecules can combine or react with one oxygen molecule to form two water molecules” § the ratio of the number of molecules is equal to the ratio of the number of moles 2. “2 moles of hydrogen molecules react with 1 mole of oxygen molecules to produce 2 moles of water molecules.” 3. “4.04 g of H react with 32.00 g of O to give 36.04 g of H O.” 2 2 2 § the mass of a mole of each of these substances Reactants: Def. the starting materials in a chemical reaction. Ex. refer to H 2nd O 2 § the reactants are written on the left Product: Def. the substance formed as a result of a chemical reaction Ex. Water § the products on the right of the arrow Steps to Balancing Chemical Equations: 1. Identify all reactants and products and write their correct formulas on the left side and right side of the equation 2. Try different coefficients to make the number of atoms of each element the same on both sides of the equation. • We can change the coefficients (the numbers preceding the formulas) but not the subscripts (the numbers within formulas). o Changing the subscripts would change the identity of the substance. 3. First: o look for elements that appear only once on each side of the equation with the same number of atoms on each side: The formulas containing these elements must have the same coefficient. Therefore, there is no need to adjust the coefficients of these elements at this point. Next: o look for elements that appear only once on each side of the equation but in unequal numbers of atoms. Balance these elements. Finally, balance elements that appear in two or more formulas on the same side of the equation. 4. Check your balanced equation to be sure that you have the same total number of each type of atoms on both sides of the equation arrow. 3.8 Amounts of Reactants and Products Stoichiometry Def. is the quantitative study of reactants and products in a chemical reaction. mole method Def. to simply that the stoichiometric coefficients in a chemical equation can be interpreted as the number of moles of each substance. o For example, industrially ammonia is synthesized from hydrogen and nitrogen as follows: o The stoichiometric coefficients show that one molecule of N reac2s with three molecules of H t2 form two molecules of NH . It 3ollows that the relative numbers of moles are the same as the relative number of molecules: o Thus, this equation can also be read as “1 mole of N ga2 combines with 3 moles of H 2as to form 2 moles of NH gas.3 In stoichiometric calculations, we say that three moles of H 2re equivalent to two moles of NH , th3t is, o where the symbol ≏ means “stoichiometrically equivalent to” or simply “equivalent to.” This relationship enables us to write the conversion factors o Similarly, we have 1 mol N ≏ 2 mol NH and 13mol N ≏ 3 mol2H . 2 • Let's consider a simple example in which 6.0 moles of H reac2 completely with N to 2 form NH .3 o To calculate the amount of NH pr3duced in moles, we use the conversion factor that has H 2n the denominator and write • Now suppose 16.0 g of H react2completely with N to form 2H . 3 o How many grams of NH will b3 formed? o To do this calculation, we note that the link between H and NH is the mole 2 3 ratio from the balanced equation. o So we need to first convert grams of H to2moles of H , the2 to moles of NH , 3 and finally to grams of NH . 3he conversion steps are o First, we convert 16.0 g of H to 2umber of moles of H , using2the molar mass of H as the conversion factor: 2 Next, we calculate the number of moles of NH produ3ed. Finally, we calculate the mass of NH pro3uced in grams using the molar mass of NH as3the conversion factor These three separate calculations can be combined in a single step as follows: Similarly, we can calculate the mass in grams of N cons2med in this reaction. The conversion steps are By using the relationship 1 mol mol H ,2we write The general approach for solving stoichiometry problems is summarized next. 1.Write a balanced equation for the reaction. 2.Convert the given amount of the reactant (in grams or other units) to number of moles. 3.Use the mole ratio from the balanced equation to calculate the number of moles of product formed. 4.Convert the moles of product to grams (or other units) of product.
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