Final Exam Study Guide
Final Exam Study Guide Chem 1010
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FINAL EXAM STUDY GUIDE Fall Quarter Chemistry 1010 Chapter I: SECTION 1.1 States of Matter Chemistry: The study of matter and its properties, the changes that matter undergoes, and the energy associated with the changes. Matter: Anything that has mass and volume. Composition of Matter: The types and amount of simpler substances that make up matter. Matter occurs in three physical forms called states: o Solid: Has a fixed shape; does not conform to its container; atomic particles lie next to each other in a regular, threedimensional array. o Liquid: Has a varying shape; conforms to the container shape but only to the extent of the liquid’s volume; atomic particles lie close together but move randomly amongst each other. o Gas: Has a varying shape; conforms to the container shape and fills the entire container; does not have a surface; atomic particles have large distances between them and move randomly throughout the container. The Properties of Matter Properties: The characteristics that give each substance its unique identity. Physical Properties: Characteristics a substance shows by itself, without interacting with another substance (ex: melting point, electrical conductivity, density). o Ex: Water (solid state) Water (liquid state) A physical change occurs when a substance alters its physical properties, not composition. Chemical Properties: Characteristics a substance shows as it changes into or interacts with another substance. A physical change caused by heating can be reversed by cooling. This is not true for a chemical change. The distinction between chemical and physical change is defined by composition. Importance of Energy in the Study of Matter Energy: The ability to do work. The total energy an object possesses is the sum of its potential energy and its kinetic energy. Potential Energy: The energy due to the position of the object relative to other objects. Kinetic Energy: The energy due to the motion of the object. When energy is converted from one form to another, it is conserved, not destroyed. Situations of lower energy are favored over situations of higher energy. Electrostatic forces (interactions); like charges repel each other and opposites attract. The chemical potential energy of a substance results from the relative positions and the attractions and repulsions among its particles. Chapter II: SECTION 2.6 Elements: A First Look at the Periodic Table There are 118 known elements today. The original periodic table was listed by increasing atomic mass, but today it is arranged by atomic number. Each element has a box that contains its atomic number, atomic symbol, and atomic mass. The periodic table consists of periods (horizontal rows) and groups (vertical columns). o Each period has a number 17, and each group has a number from 18 and either the letter A or B. The eight A groups contain the maingroup elements. The ten B groups contain the transition elements. Two horizontal series of inner transition elements (lanthanides and actinides) fit between the elements in group 3B(3) and group 4B(4). Classifying the elements: o Metals lie in the large lowerleft portion of the table. Shiny solids at room temperature (Mercury is the only liquid). Conduct heat and electricity well. Malleable and ductile. o Nonmetals lie in the smaller upperright portion of the periodic table. Generally gases or dull, brittle solids at room temperature (bromine is the only liquid). Conduct heat and electricity poorly. o Metalloids have properties between those of metals and nonmetals. Elements in a group have similar chemical properties and elements in a period have different chemical properties. Group 1A(1) except for hydrogen, consists of alkali metals. Group 2A(2) consists of the alkaline earth metals. The halogens, Group 7A(17) are highly reactive nonmetals. Noble gases, Group 8A(18) are highly unreactive nonmetals. SECTION 2.7 Compounds: Introduction to Bonding The overwhelming majority of elements occur in compounds combined with other elements. The noble gases (helium, neon, argon, krypton, xenon, and radon) occur in the air as separate atoms. Oxygen, nitrogen, and sulfur also occur in their most common elemental form as the molecules O , N , 2 2 and S 8 Elements combine in two general ways: o Transferring electrons from one element to another to form ionic compounds. o Sharing electrons between atoms of different elements to form covalent bonds. Chemical bonds: The forces that hold the atoms together in a compound. Ions: Charged particles that form when an atom gains or loses one or more electrons. The simplest type of ionic compound is a binary ionic compound; one composed of two elements, and typically forms when a metal reacts with a nonmetal. o Each metal atom loses one or more electrons and becomes a cation, a positively charged ion. o Each nonmetal atom gains one or more electrons lost by the metal atom and becomes an anion, a negatively charged ion. A cation or anion derived from a single atom is called a monatomic ion. All binary ionic compounds are solid arrays of oppositely charged ions. Coulomb’s Law: o Ions with higher charges attract (or repel) each other more strongly than ions with lower charges. o Smaller ions attract (or repel) each other more strongly than larger ions, because the charges are closer to each other. Ionic compounds are neutral because they contain equal numbers of positive and negative charges. Nonmetal atoms gain electrons to form ions with the same number of electrons as in an atom of the nearest noble gas. Metals lose electrons: elements in Group 1A(1) lose one electron, elements in Group 2A(2) lose two electrons, and elements in Group 3A(3) lose three electrons. Nonmetals gain electrons: elements in Group 7A(17) gain one electron, elements in Group 6A(16) gain two electrons, and elements in Group 5A(15) gain three electrons. Atoms of different elements share electrons to form the molecules of a covalents compound. Most covalent substances consist of molecules–and there are no molecules in an ionic compound. The nature of the particles attracting each other in covalent and in ionic substances is fundamentally different. Many ionic compounds contain polyatomic ions, which consist of two or more atoms bonded covalently and have a net positive or negative charge. Chapter IV: SECTION 4.1 Solution Concentration and the Role of Water as a Solvent Uneven charge distribution creates a polar bond, one with partially charged “poles.” Bent molecular shape: bond angle of 104.5˚ Solvated: Surrounded closely by solvent molecules. A substance that conducts a current when dissolved in water is an electrolyte. Ionic compounds are strong electrolytes because they dissociate completely in water and conduct a large current. Aqueous solutions do not conduct an electric current, and these substances are called nonelectrolytes. Concentration: The quantity of solute dissolved in a given quantity of solution. Molarity (M): Moles of solute per liter of solution. Molarity = moles of solute/liters of solution Solving dilution problems: M dil =dilount (mol) = M concx Vconc SECTION 4.2 Writing Ionic Equations for Aqueous Ionic Reactions Molecular equation shows all the reactants and products as if they were intact, undissociated compounds. Total ionic equation shows all soluble ionic substances as they actually exist in solution, where they are dissociated into ions. Ions that appear unchanged on both sides of the equation are called spectator ions. + 2– Example: 2Ag (aq) + CrO (aq) 4 KCrO (aq) AgC4O (s) 4 Molecular: 2AgC H O 2(a3) 2 K CrO (a2) A4CrO (s) + 2KC H 4 (aq) 2 3 2 + – – 2– + Total Ionic: 2Ag (aq) + 2C H O 2aq3 +22K (aq) + CrO (aq) Ag Cr4 (s) + 2K (aq2 + 24 H O (aq) 2 3 2 Net Ionic: 2Ag (aq) + CrO 42–(aq) Ag Cr2 (s)4 SECTION 4.3 Precipitation Reactions In a precipitation reaction, two soluble ionic compounds react to form an insoluble product, a precipitate. When the ions exchange partners, it is called a doubledisplacement reaction. Solving stoichiometry problems for any reaction that takes place in a solution: 1. Balance the equation. 2. Find the amount (mol) of one substance using its molar mass (for a pure substance) or the volume and molarity (for a substance in solution). 3. Relate that amount to the stoichiometrically equivalent amount of another substance. 4. Convert the desired units. SECTION 4.4 AcidBase Reactions Acidbase reactions (where an acid reacts with a base) involve water as a reactant or product. An acid is a substance that produces H ions when dissolved in water. + HX H (aq) + X (aq) A base is a substance that produces OH ions when dissolved in water. + MOH M (aq) + OH (aq) Acidic solutions arise when certain covalent Hcontaining molecules dissociate into ions in water. Strong acids AND strong bases dissociate completely into ions when placed in water. Weak acids AND weak bases dissociate very little into ions. Strong acids, such as HNO and 3 SO , an2 we4k acids, such as HF and H PO , have one o3 mor4 H atoms as part of their structure. Strong bases have either OH or O as part of their structure. Weak bases, such as ammonia (NH ) do not 3ontain OH ions, but they all have an electron pair on N. + The key event in aqueous reactions between a strong acid and a strong base is that an H ion from the acid and an OH ion from the base form a water molecule. Acidbase reactions occur through the electrostatic attraction of ions and their removal from solution as the product. The ionic compound that results from the reaction of an acid and base is called a salt. Acidbase reactions are metathesis (doubledisplacement) reactions. An acid is a molecule (or ion) that donates a proton. A base is a molecule (or ion) that accepts a proton. H O ion acts as the acid and donates a proton to OH ion, which acts as the base and accepts it. 3 Titration: The known concentration of one solution is used to determine the unknown concentration of another. Equivalence point: Occurs when the amount (mol) of H ions in the original volume of acid has reacted with the same amount (mol) of OH ions from the buret. End point: Occurs when a tiny excess of OH ions changes the indicator permanently to its basic color. The amount of base needed to reach the end point is the same as the amount needed to reach the equivalence point. Oxidationreduction (redox) reaction: Net movement of electrons from one reactant to another. Ionic compounds: transfer of electrons. Covalent compounds: shift (sharing) of electrons. Oxidation is the loss of electrons. Reduction is the gain of electrons. Example: Formation of MgO Oxidation (electron loss by Mg): Mg Mg + 2e – Reduction (electron gain by O )2 ½ O +22e O 2– The oxidizing agent is the species doing the oxidizing (causing electron loss). The reducing agent is the species doing the reducing (causing electron gain). In MgO, O 2xidizes Mg by taking electrons that Mg gives up. Mg reduces O 2y providing the electrons that O t2kes, so Mg is the reducing agent, and O is2 the oxidizing agent. The oxidizing agent is reduced, the reducing agent is oxidized. Oxidation numbers: 1. For Group 1(A): O.N. = +1 in all compounds 2. For Group 2(A): O.N. = +2 in all compounds 3. For hydrogen: O.N. = +1 in combination with nonmetals = –1 in combination with metals and boron 4. For fluorine: O.N. = –1 in all compounds 5. For oxygen: O.N. = –1 in all peroxides = –2 in all other compounds (except with F). 6. For Group 7(A): O.N. = –1 in combination with metals, nonmetals (except O), and other halogens lower in the group Transferred electrons are never free because the reducing agent loses electrons and the oxidizing agent gains them simultaneously. Atoms occur as an element on one side of an equation and as part of a compound on the other. In a combination reaction, two or more reactants form a compound. X + Y Z Metal and nonmetal form an ionic compound. The metal is the reducing agent and the nonmetal is the oxidizing agent. Two nonmetals form a covalent compound. Nearly every nonmetal reacts with O t2 form a covalent oxide. In a decomposition redox reaction, a compound forms two or more products, at least one of which is an element. Z X + Y In doubledisplacement (metathesis) reactions, atoms of two compounds exchange places. AB + CD AD + CB In solution, singledisplacement reactions occur when an atom of one element displaces the ion of another. The most reactive metals displace H f2om liquid water. Group 1(A) metals and Ca, Sr, and Ba from Group 2(A) displace H f2om water. Chapter VI: Thermodynamics: The study of energy and its transformations. Thermochemistry: A branch of thermodynamics that deals with heat in chemical and physical change. System: The part of the universe we are focusing on (in a reaction). Surroundings: Everything else (in a reaction). Internal Energy (E): Sum of all potential and kinetic energy in a system. ∆E is the difference between internal energy after the change (E finaland before the change (E initial Final state minus the initial state: ∆E = E finalE initial productsEreactants A change in the energy of a system must be accompanied by an equal and opposite change in the energy of the surroundings. By releasing some energy in a transfer to the surroundings: E < E so ∆E < 0 final initial By absorbing some energy in a transfer from the surroundings: E final Einitial ∆E > 0 Heat: Thermal energy (symbolized by q) is the energy transferred as a result of a difference in temperature. Work: (symbolized by w) the energy transferred when an object is moved by a force. The total change in a system’s internal energy is the sum of the energy transferred as heat and/or work: ∆E = q + w Energy transferred to the system is positive because the system ends up with more energy. Energy transferred from the system is negative because the system ends up with less energy. Heat flowing out of a system: heat is released so q is negative and ∆E is negative. Heat flowing into a system: heat is absorbed so q is positive and ∆E is positive. Law of Conservation of Energy – First Law of Thermodynamics: The total energy of the universe is constant. ∆E universe∆E system– ∆Esurroundings Joule (J): The SI unit of energy 2 2 1 J = 1 kg • m /s Calorie (cal): Quantity of energy needed to raise the temperature of 1 g of water by 1 ˚C. 1 cal = 4.184 J or 1 J = 1/4.184 cal = 0.2390 cal British thermal unit: Quantity of energy required to raise the temperature of 1 lb of water by 1 ˚F. 1 BTU = 1055 J ∆E does not depend on how the change takes place, but only on the difference between the final and initial states. Pressurevolume work (PV work): The mechanical work done when the volume of the system changes in the presence of an external pressure (P). w = –P∆V At constant pressure, enthalpy (H) is defined as the internal energy plus the product of the pressure and volume. H = E + PV Change in enthalpy (∆H): The change in internal energy plus the product of the pressure, which is constant, and the change in volume (∆V). ∆H = ∆E + P∆V Exothermic and endothermic process: ∆H = H final HinitialHproducts reactants Exothermic: Releases heat and results in a decrease in the enthalpy of a system: H productsHreactants ∆H < 0 Endothermic: Absorbs hear and results in an increase in the enthalpy of a system: H productsHreactants ∆H > 0 q/∆T = constant Heat capacity: [refer to above equation] the quantity of heat required to change its temperature by 1 K. Specific heat capacity: The quantity of heat required to change the temperature of 1 gram of a substance or material by 1 K. Specific heat capacity (c) = q/mass x ∆T Molar heat capacity: The quantity of heat required to change the temperature of 1 mole of a substance by 1 K. Molar heat capacity (C) = q/amount (mol) x ∆T Calorimeter: A device used to measure the heat released (or absorbed) by a physical or chemical process. Finding the specific heat capacity of a solid: c = c x mass x ∆T /mass x ∆T solid H2O H2O H2O solid solid A thermochemical equation is a balanced equation that includes the enthalpy change of the reaction (∆H). Hess’s Law: The enthalpy change if an overall process is the sum of the enthalpy changes of its individual steps: ∆H = ∆H + ∆H … + ∆H overall 1 2 n Standard states: o For a gas, the standard state is 1 atm* and ideal behavior o For a substance in an aqueous solution, the standard state is 1 M concentration. o For a pure substance, the standard state is usually the most stable form of the substance at 1 atm and the temperature of interest (the temperature is usually 25 ˚C (298 K). Chapter VII: SECTION 7.1 The Wave Nature of Light The wave properties of electromagnetic radiation are described by three variables and one constant: o Frequency: The number of complete waves, or cycles, per second. Measured in Hertz (Hz). o Wavelength: The distance the wave travels in one cycle. o Speed: The speed of a wave is the distance it moves per unit of time (meters per second). Speed of light: c = v x λ o Amplitude: Height of the wave crest; related to the intensity of the radiation or its brightness. Electromagnetic Spectrum Light of a single wavelength is called monochromatic; light of many wavelengths is called polychromatic. The Classical Distinction Between Energy and Matter Refraction: When a light wave strikes the boundary between two media (ex: air and water) at an angle other than 90º, the change in speed causes a change in direction, causing the wave to continue at a different angle. Dispersion: When white light enters through a prism, it is separated into its component colors because each incoming wave is refracted at a different angle. Diffraction: When a light wave bends around an object, it bends around both edges of the slit and forms a semicircular wave on the other side of the opening. Interference (constructive and destructive): The interaction of emerging circular waves due to waves of light passing through two adjacent slits. When the amplitudes are added together (in phase) to form a brighter region, the interference is constructive. When the crests coincide with the troughs (out of phase) the interference is destructive. The Particle Nature of Light The Quantum Theory: Each change in an atom’s energy occurs when the atom absorbs or emits one or more “packets” or definite amounts of energy. Max Planck developed a formula to fit this phenomenon. ∆E = hv h = 6.62606874 x 10 joules/second SECTION 7.2 Line Spectra and the Rydberg Equation Line Spectrum: When light from electrically excited gaseous atoms passes through a slit and is refracted by a prism, it creates a series of fine lines at specific frequencies separated by black spaces. Rydberg’s equation predicts the position and wavelength of any line in a given series: R = 6.626 x 10 34 The Bohr Model of the Hydrogen Atom The H atom has only certain energy levels called stationary states. Each state I associated with a fixed circular orbit of the electron around the nucleus. The higher the energy level, the farther the orbit is from the nucleus. The atom does not radiate energy while in one of its stationary states. The atom changes to another stationary state only by absorbing or emitting a photon. E photon ∆E atom= Efinal Einitialhv The quantum number n is a positive integer (1…2…3…) o The lower the n value, the smaller the radius of the orbit, and the lower the energy level. When the electron is in the first orbit (n = 1), it is closest to the nucleus and in it’s lowest energy level, called the ground state. If the electron is in any orbit farther from the nucleus, the atom is in an excited state. If an H atom absorbs a photon whose energy equals the difference between lower and higher energy levels, the electron moves to the outer orbit. If an H atom in a higher energy level returns to a lower energy level, the atom emits a photon whose energy equals the difference between two levels. Since an atom’s energy is not continuous, but rather has certain states, an atomic spectrum is not continuous. The Bohr model failed to predict the spectrum of any other atom because: o It is a oneelectron model. o Electrons do not have defined, fixed orbits. o The energy of an atom occurs in discrete levels and it changes when the atom absorbs or emits a photon of specific energy. The Energy Levels of the Hydrogen Atom Finding the difference between two energy levels: ∆E = E finalE initial 18 x 10 J (1/n 2final1/n2initial SECTION 7.3 The Wave Nature of Electrons and the Particle Nature of Photons Matter and energy are alternate forms of the same entity. If energy is particlelike, matter is wavelike. deBroglie Wavelength: Matter behaves as though it moves in a wave. Results on the atomic scale show electrons moving in waves and photons having momentum. Both matter and energy show both behaviors. o This dual character of matter and energy is known as the waveparticle duality. Uncertainty Principle: It is impossible to know simultaneously the position and momentum of a particle. SECTION 7.4 The Atomic Orbital and the Probable Location of the Electron Quantum Mechanics: Examines the wave nature of objects on an atomic scale. Schrödinger Equation: Ψ is called a wave function; a mathematical representation of the electrons matterwave in three dimensions. The H symbol is called the Hamiltonian operator, representing a set of mathematical operations that, when carried out with a particular Ψ, yields one of the allowed energy states of the atom. 2 Ψ is called the probability density, a measure of the probability of finding an electron in some tiny volume of the atom. Electron Density Diagram: The value for Ψ for a given volume is shown with dots: the greater the density of the dots, the higher the probability of finding an electron in that volume. The electron probability density decreases with distance from the nucleus. Radial Probability Distribution Plot: The volume of each layer increases faster than its density of dots decreases. Probability Contour: We cannot assign a definite volume to an atom. Therefore, we must use the probability contour (ex. an atom with a 90% probability contour means the electron is somewhere within that volume 90% of the time). Quantum Numbers of an Atomic Orbital Principle Quantum Number: (n) a positive integer (1, 2, 3 and so forth). o Indicates the relative size of the orbital, and the relative distance from the nucleus. o The higher the n value, the higher the energy level. Angular Momentum Quantum Number: (l) is an integer from 0 to n – 1. o Related to orbital shape. o N =1 l = 0 (1 value) N = 2 l = 0, 1 (two values) N = 3 l = 0, 1, 2, 3 (three values) Magnetic Quantum Number: (m) is an lnteger from –l through 0 to +l o Prescribes the threedimensional orientation of the orbital in space around the nucleus. o An orbital with l = 0 can only have m = l. o An orbital with 1 = 1 can have m = l1, 0, +1. o The number of m valles = 2l + 1 = the number of orbitals given for l. o The total number of m vall s for a given n value = n = the total number of orbitals in that energy level. Quantum Numbers and Energy Levels The atom’s energy levels (shells) are given by the n value. The atom’s levels are divided into sublevels that are given by the l value. L = 0 is an s sublevel. L = 1 is a p sublevel. L = 2 is a d sublevel. L = 3 is an f sublevel. Orbitals: n l Sublevel Name Possible m lalues # of Orbitals 3 2 3d 2, 1, 0, +1, +2 5 2 0 2s 0 1 5 1 5p 1, 0, +1 3 4 3 4f 3, 2, 1, 0, +1, +2, +3 7 Shapes of Atomic Orbitals The s Orbital: l = 0 has a spherical shape, with the nucleus at it’s center. o The 1s orbital holds the electron in the H atom’s ground state. The electron probability is highest at the nucleus. o The 2s orbital has two regions of higher electron density. The radial probability distribution is higher than that of the closer one because the sum of Ψ for it is taken over a much larger volume. Between the two regions is a spherical node. o The 3s orbital has three regions of high electron density and two nodes. The highest radial probability is at the greatest distance from the nucleus. The p Orbital: An orbital with an l = 1is a p orbital. o Two regions of high probability, one on either side of the nucleus. o The nucleus lies at the nodal plane of this dumbbellshaped orbital. o The maximum value of l is n – 1, only levels with n = 2 or higher have a p orbital. o p orbitals have different spatial orientations; three possible m values of 1, 0, +1. l The d Orbital: An orbital with l = 2 is called a d orbital. o M values for l = 2 are 2, 1, 0, +1, +2. o Two mutually perpendicular nodal planes between them and the nucleus at the junction of the lobes. o A d orbital must have a principal quantum number of n = 3 or higher, so 3d is the lowest energy d sublevel. Orbitals with l = 3 are f orbitals and have a principal quantum number of at least n = 4. o Each f orbital has a complex, multi lobed shape with several nodal planes. Chapter VIII: SECTION 8.1 The ElectronSpin Quantum Number Spin Quantum Number: (m ) has swo possible values; 1/2 or +1/2. o Each electron in an atom is described by a et of four quantum numbers: the first three describe its orbital, and the fourth describes its spin. The Exclusion Principle No two electrons in the same atom can have the same four quantum numbers. An atomic orbital can hold a maximum of two electrons, which must have opposing spins. Electrostatic Effects and EnergyLevel Splitting The energy of an orbital in a manyelectron atom depends mostly on its n value (size) and to a lesser extent on its l value (shape). This energy difference arises from three factors: nuclear attraction, electron repulsions, and orbital shape. Their interplay lead to two phenomena–shielding and penetration, which occur in all atoms except hydrogen. A higher nuclear charge increases nucleuselectron attractions and thus, lowers sublevel energy (stabilizes atom). Repulsions counteract the nuclear attraction by making the electron easier to remove. Shielding: Reduces the full nuclear charge to an effective nuclear charge (z ): eff nuclear charge an electron actually experiences, and this lower nuclear charge makes it easier for an electron to be removed. Penetration: o Increases the nuclear attraction. o Decreases shielding. Order of sublevel energies: S < P < D < F SECTION 8.2 The Quantum Mechanical Model and the Periodic Table Aufbau Principle: We start at the beginning of the periodic table and add one proton to the nucleus and one electron to the lowest energy sublevel available. The electron configuration diagram consists of the principal energy level (n value), the letter designation of the sublevel (l value) and the number of electrons in the sublevel, written as a superscript: nl . # An orbital diagram consists of a box for each orbital in a given energy level, grouped by a sublevel, with an arrow representing an electron and its spin. 2 o Helium: He (z=2) 1s o Lithium: Li (z=3) 1s 2s 2 1 o Beryllium: Be (z=4) 1s 2s 2 2 2 2 1 o Boron: B (z=5) 1s 2s 2p o Carbon: C (z=6) 1s 2s 2p 2 2 2 2 3 o Nitrogen: N (z=7) 1s 2s 2p o Oxygen: O (z=8) 1s 2s 2p2 2 4 2 2 5 o Fluorine: F (z=9) 1s 2s 2p o Neon: N (z=10) 1s 2s 2p2 2 6 Elements in the same group have similar outer electron configurations and similar patterns of reactivity. 2 5 In Group 7A(17), F and Cl have the outer electron configuration ns np , as do the halogens. All are reactive nonmetals that occur as diatomic molecules and all form ionic compounds with metals. 2 In Group 8A(18), He has the electron configuration ns , and all the other elements in the group have the outer configuration ns np – all members are very unreactive monatomic gases. Sublevels are filled in order of increasing energy, which leads to outer electron configurations that recur periodically, which leads to chemical properties that recur periodically. Period 4 contains the first series of transition elements. o The 3d sublevel is filled in period 4, but the 4s sublevel is filled first. o The 4s orbital is slightly lower in energy than the 3d and fills first. o Halffilled and filled sublevels are unexpectedly stable. Inner (core) electrons: They fill all the lower energy levels of an atom. Outer electrons: Those in the highest energy level (highest n value). Farthest from the nucleus. Valence electrons: For main group elements, the valence electrons are the outer electrons. For transition elements, in addition to the outer ns electrons, the (n–1)d electrons are also valence electrons, though the metals Fe (z=26) through (z=30) may use only a few, if any of their d electrons in bonding. The A number equals the number of outer electrons. The period number is the n value of the highest energy level. The n value squared (n ) is the number of orbitals, and the 2n is the maximum number of electrons (or elements). In a transition series, the filling order is 5s, then 4d, and then 5p. Inner transition elements: The period 6 inner transition series, called the lanthanides, occurs after lanthanum and in it the 4f orbitals are filled. o Period 7 inner transition series, called the actinides, occurs after actinium and in it the 5f orbitals are filled. SECTION 8.3 Trends in Three Atomic Properties Atomic size is defined in terms of how close one atom lies next to another. We measure the distance between atomic nuclei in a sample of an element and divide that distance in half. Metallic radius: Onehalf the shortest distance between nuclei of adjacent, individual atoms in a crystal of the element. Covalent radius: Onehalf the shortest distance between nuclei of bonded atoms. As the principal quantum number (n) increases, the atomic size increases. As the effective nuclear charge increases (Zeffincreases, atomic size decreases. Atomic radius increases down a group. Atomic radius generally decreases across a period. Down a transition group, n increases but only a small size increase occurs between 4 and 5, and no increase occurs between 5 and 6. Across a transition series, atomic size shrinks through the first two or three elements because of the increasing nuclear charge. But from then on, size remains relatively constant. A transition series affects atomic size in neighboring main groups. A major size decrease occurs from Group 2A(2) to Group 3A(13). Ionization energy: The energy required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions. o Atoms with a low IE tend to form cations during reactions. o Atoms with high IE (except noble gases) tend to form anions. Ionization energy generally decreases down a group. Ionization energy generally increases across a period. Trends in Electron Affinity Electron affinity: The energy change (kj/mol) accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms or ions. The first electron affinity refers to the formation of 1 mol of monovalent gaseous anions. Down a group, a smooth decrease is expected because atomic size increases. But only Group 1A(1) exhibits this behavior. Reactive nonmetals have high IEs and highly negative (exothermic) EAs. o Lose electrons with difficulty but attract them strongly. o In their ionic compounds, they form negative ions. Reactive metals have low IEs and slightly negative EAs. Lose electrons easily but attract them weakly. In their ionic compounds, they form positive ions. Members of the noble gases have very high IEs and slightly positive (endothermic) EAs. They tend not to lose or gain electrons. SECTION 8.4 Atomic Properties and Chemical Reactivity For most elements, the Agroup number is the highest oxidation number. The exceptions are O and F. For nonmetals and metalloids, the Agroup number minus 8 is the lowest oxidation number (always negative) of any element in the group. With their low IEs and small EAs, the members of Groups 1A(1) and 2A(2) lose electrons readily, so they are strong reducing agents and become oxidized. With their high IEs and large EAs, nonmetals in group 6A(16) and 7A(17) gain electrons readily, so they are strong oxidizing agents and become reduced. Most maingroup metals transfer electrons to oxygen, so their oxides are ionic. In water, these oxides act as bases. Nonmetals share electrons with oxygen, so nonmetal oxides are covalent. They react with water to form acids. Some metals and many metalloids form oxides that are amphoteric: they can act as acids or bases in water. As elements become more metallic down a group, their oxides become more basic. As the elements become less metallic across a period, their oxides become more acidic. When elements at either end of a period form ions, they attain a filled outer level–a noble gas configuration. Isoelectronic: Same electronic ion charge. Elements in Groups 1A(1) and 2A(2) lose electrons and become isoelectronic with the previous noble gas (ex: the Na ion becomes isoelectronic with neon). Elements in Groups 6A(16) and 7A(17) become isoelectronic with the next noble gas (ex: Br ion is isoelectronic with Krypton (Kr)). Cations occur when an electron is removed. Anions occur when an electron is added. Transition metal ions rarely attain a noble gas configuration. A transition element typically forms more than one cation by losing all of its ns and some of its (n–1)d electrons. A species with unpaired electrons exhibits paramagnetism; it is attracted by an external field. A species with paired electrons exhibits diamagnetism; it is not attracted by the field. Ionic radius: A measure of the size of an ion and is obtained from the distance between the nuclei of adjacent ions in a crystalline ionic compound. Cations are smaller than parent atoms. Anions are larger than parent atoms. Down a group, ionic size increases because n increases. From last cation to first anion, a great jump in size occurs. Chapter IX: SECTION 9.1 Atomic Properties and Chemical Bonds Bonding lowers the potential energy between positive and negative particles. There is (in general) a gradation from more metallic elements to more nonmetallic elements across a period and up a group. Metal with nonmetal: electron transfer and ionic bonding. We observe electron transfer and ionic bonding between atoms with large differences in their tendencies to lose or gain electrons. Nonmetal with nonmetal: electron sharing and covalent bonding. When two atoms differ little, or not at all, in their tendencies to lose or gain electrons, we observe electron sharing and covalent bonding, which occurs most commonly between nonmetals. The shared electron pair is typically localized between the two atoms, linking them in a covalent bond of a particular length and strength. Metal with metal: electron pooling and metallic bonding. In the simplest model of metallic bonding, the enormous number of atoms in a sample of a metal pool their valance electrons into a sea of electrons that “flow” between and around each metalion core, thereby attracting and holding them together. Electrons in metallic bonding are delocalized, moving freely throughout the entire piece of metal. Lewis Symbols and the Octet Rule Lewis electrondot symbol: the element symbol represents the nucleus and inner electrons, and dots around the symbol represent the valence electrons. For a metal, the total number of dots is the number of electrons an atom loses to for a cation. For a nonmetal, the number of unpaired dots equals either the number of electrons an atom gains to form an anion or the number it shares to form covalent bonds. Octet rule: When atoms bond, they lose, gain, or share electrons to attain a filled outer level of eight electrons. SECTION 9.2 The Ionic Bonding Molecule The transfer of electrons from metal atoms to nonmetal atoms to form ions that attract each other and form a solid compound. In ionic bonding, the total number of electrons lost by the metal atom(s) equals the total number of electrons gained by the nonmetal atom(s). Why Ionic Compounds Form: The Importance of Lattice Energy Energy is actually absorbed during electron transfer. The electron transfer process: o The first ionization energy is the energy absorbed when 1 mol of gaseous atoms loses 1 mol of valence electrons. o The first electron affinity (EA 1 is the energy released when 1 mol of gaseous atoms gains 1 mol of electrons. If the overall reaction releases energy, there must be some step that is exothermic enough to outweigh the endothermic steps. This step involves the strong attraction between pairs of oppositely charged ions. Even more energy is released when the separate gaseous ions coalesce into a crystalline solid because each ion attracts several oppositely charged ions. The lattice energy (∆H˙ lattices the enthalpy change that accompanies the reverse of the previous equation–1 mol of ionic solid separating into gaseous ions. In a BornHaber cycle, a series of steps from elements to ionic solid for which all the enthalpies are known except the lattice energy. The BornHaber cycle shows that the energy required for elements to form ions is supplied by the attraction among the ions in the solid. And the “take home lesson” is that ionic solids exist only because the lattice energy far exceeds the total energy needed to form the ions. Electrostatic energy decreases between cations and anions and thus, lattice energy should decrease as well. How the Model Explains the Properties of Ionic Compounds A typical ionic compound is hard, rigid, and brittle. Ionic compounds typically do not conduct electricity in the solid state but do conduct when melted or dissolved. We expect ionic compounds to have high melting points and much higher boiling points. Ion pairs: gaseous ionic molecules, rather than individual ions. Ionic compounds are solid arrays of ions and no separate molecules exist. SECTION 9.3 The Formation of a Covalent Bond Sharing electrons is the main way that atoms interact. A covalent bond arises from the balance between the nuclei attracting the electrons and electrons and nuclei repelling each other. Formation of a covalent bond always results in greater electron density between the nuclei. The shared bonding pair is represented by a pair of dots or a line. An outerlevel electron bonding pair that is not involved in bonding is called a lone apir, or unshared pair. Properties of a Covalent Bond: Order, Energy, and Length The bond order is the number of electron pairs being shared by a given pair of atoms. A single bond is the most common bond and consists of one bonding pair of electrons: a single bond has a bond order of 1. Multiple bonds usually involve C, O, and/or N atoms. A double bond consists of two bonding electron pairs, four electrons shared between two atoms, so the bond order is 2. A triple bond consists of three shared pairs: two atoms share six electrons, so the bond order is 3. The strength of a covalent bond depends on the magnitude of the attraction between the nuclei and shared electrons. The bond energy, called bond enthalpy, is the energy needed to overcome this attraction and is defined as the standard enthalpy change for breaking the bond in 1 mol of gaseous molecules. The same quantity of energy absorbed to break the bond is released when the bond forms. Bond formation is an exothermic process, so the sign of its enthalpy change is always negative. A(g) + B(g) A–B(g) ∆H˙ bond formingBE A–B (always < 0) Stronger bonds have a larger BE because they are lower in energy (have a deeper energy well). Weaker bonds have a smaller BE because they are higher in energy (have a shallower energy well). Bond length is the distance between the nuclei of two bonded atoms. The order, energy, and length of a covalent bond are interrelated. For a given pair of atoms, a higher bond order results in a smaller bond length and a higher bond energy. Thus, a shorter bond is a stronger bond. How the Model Explains the Properties of Covalent Substances Strong bonding forces hold the atoms together within the molecule and weak intermolecular forces act between separate molecules in the sample. Network covalent solids are held together by covalent bonds between atoms throughout the sample, and their properties do reflect the strength of covalent bonds. Most covalent bonds are poor electrical conductors, because their electrons are localized as either shared or unshared pairs, and no ions are present. SECTION 9.5 Between the Extremes: Electronegativity and Bond Polarity Electronegativity: the relative ability of a bonded atom to attract shared electrons. Linus Pauling reasoned that if one element attracts the shared electron pair more strongly than the second element in a compound (that is, if the first element is electronegative), then the electrons will spend more time closer to the first element, and this unequal sharing makes the first element end of the bond partially negative and the second element partially positive. Electronegativity is inversely related to atomic size. Down a main group, electronegativity decreases as size increases. Across a period of maingroup elements, electronegativity increases. Nonmetals are more electronegative than metals. The most electronegative element is fluorine. The more electronegative atom in a bond is assigned all the shared electrons; the less electronegative atom is assigned none. Each atom in a bond is assigned all of its unshared electrons. The oxidation number is given by: – – – O.N. = # of valence e – (# of shared e + # of unshared e ) Bond Polarity and Partial Ionic Character Polar covalent bond: whenever atoms having different electronegativities form a bond and the bonding pair is shared unequally. Nonpolar covalent bond: where atoms are identical, and the bonding pair is shared equally. The electronegativity difference (∆EN) is the difference between the EN values of the bonded atoms, is directly related to the bond’s polarity. Partial ionic character: a greater ∆EN results in larger partial charges and higher potential ionic character. Two approaches that quantify ionic character: o ∆EN range: this approach divides bonds into mostly ionic, polar covalent, mostly covalent, and nonpolar covalent based on range of ∆EN values. o Percent ionic character: this approach is based on the behavior of a gaseous diatomic molecule in an electric field. Percent ionic character generally increases with ∆EN. Electron sharing occurs to some extent in every bond. The Graduation in Bonding Across a Period A metal and a nonmetal–elements from the left and right sides of the periodic table have a relatively large ∆EN and typically form an ionic compound. The relative strengths of the bonds in reactants and products determine whether heat is released or absorbed in a chemical reaction. Kinetic energy: Molecules’ movements through space and their rotations and vibrations. Potential energy: Phase changes and changes in the attraction between vibrating atoms. A certain quantity of heat is absorbed (∆H˚ > 0) to break the reactant bonds and form separate atoms. A different quantity of heat is then released (∆H˚ < 0) when the atoms form product bonds. Chapter X: SECTION 10.1 Depicting Molecules and Ions with Lewis Structures Lewis structure (or Lewis Formula) shows symbols for the atoms, the bonding electron pairs as lines, and the lone electron pairs that fill each atom’s outer level (valence shell) as pairs of dots. Applying the Octet Rule to Write Lewis Structures The four steps for writing Lewis structures for species with only single bonds: o Place the atoms relative to each other. o Determine the total number of valence electrons. o Draw a single bond from each surrounding atom to the central atom, and subtract from the total for each bond to find the number of e remaining. – o Distribute the remaining electrons in pairs so that each atom ends up with 8e . Hydrogen atoms form one bond, carbon atoms form four bonds, nitrogen atoms form three bonds, oxygen atoms for two bonds, and surrounding halogens form one bond; fluorine is always a surrounding atom. Steps for writing Lewis Structures for species with multiple bonds: o Same steps for single bonds. o And then if a central atom does not end up with an octet, form one or more multiple bonds. Resonance: Delocalized ElectronPair Bonding Resonance Structures (Forms): Have the same relative placement of atoms but different locations of bonding and lone electron pairs. An average of the resonance structures is called a resonance hybrid. Electronpair delocalization: In a single, double, or triple bond, each electron pair is localized between the bonded atoms. In a resonance hybrid, two of the electron pairs (one bonding and one lone pair) are delocalized: their density is spread over a few adjacent atoms. Formal Charge: Selecting the More Important Resonance Structure Formal charge: The charge an atom would have if the bonding electrons were shared equally. Formal charges must sum to the actual charge on the species. Smaller formal charges (positive or negative) are preferable to larger ones. The same nonzero formal charges on adjacent atoms are not preferred. A more negative formal charge should reside on a more electronegative atom. For a formal charge, bonding electrons are shared equally by the atoms, so each atom has half of them: – – – Formal charge = valence e – (lone pair e + ½ bonding e ) For an oxidation number, bonding electrons are transferred completely to the more electronegative atom (as if the bonding were pure ionic): Oxidation number = valence e – (lone pair e + bonding e )– Electron deficient: They have fewer than eight electrons around the central atom. Most molecules have a central atom from an oddnumbered group, such as N (Group 5A(15)) or Cl (Group 7A(17)). These are called free radicals, species that contain a lone (unpaired) electron, which makes them paramagnetic and extremely reactive. Many molecules (and ions) have more than eight valence electrons around the central atom. That atom expands its valence shell to form more bonds, which releases energy. Expanded valence shells occur only with nonmetals from period 3 or higher because they have d orbitals available. SECTION 10.2 Valenceshell Electronpair Repulsion (VSEPR) Theory Valenceshell electronpair repulsion theory: to minimize repulsions, each group of valence electrons around a central atom is located as far as possible from the others. Only electron groups around the central atom affect shape. The molecular shape is the three dimensional arrangement of nuclei joined by the bonding groups. The electrongroup arrangement is defined by the bonding and nonbonding electron groups around the central atom, but the molecular shape is defined by the relative positions of the nuclei, which are connected by bonding groups only. The same electrongroup arrangement can give ride to different molecular shapes. AX Em nesignation, where m and n are integers, A is the central atom, X is the surrounding atom, and E is a nonbonding valenceelectron group (usually a lone pair). The bond angle is the angle formed by the nuclei of two surrounding atoms with the nucleus of the central atom at the vertex. Electrongroup Arrangements and Molecular Shapes Using VSEPR theory to determine molecular shape: o Write the Lewis structure from the molecular formula. o Assign an electrongroup arrangement by counting all electron groups. o Predict the ideal bond angle from the electrongroup arrangement and the effect of and deviation caused by the lone pairs or double bonds. o Draw and name the m
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