CHEM 1 CONCEPT PACKET Complete

Transcription

1 CHEM 1 CONCEPT PACKET Complete Written by Jeremy Robinson, Head Instructor Find Out More +Private Instruction +Review Sessions

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3 Calculations Calculations are fundamental to all of chemistry, and certain rules and methods apply to the calculations. Significant Figures and Results Significant Figures are very important in sensitive calculations which are frequent in chemistry. A number is always made up of one or more figures or digits which may or may not be significant and possibly together with a decimal point. Significant Figures The following rules determine whether or not a single figure in a number is significant to the overall number: 1. A number is always stated to the accuracy such that only the last digit may be uncertain or rounded off to. 2. All nonzero figures within a number are significant no matter where they are located in the number. 3. Some of the zero figures within a number are significant. Those that appear before the first nonzero figure are not significant but place the decimal point and those that appear after the first nonzero figure are significant. Another method is to rewrite the number in scientific notation which is always reported with significant figures only. Significant Results Rules behind Significant Figures ensure that combined accuracy of different measurements is preserved in equations: 1. When two numbers are added or subtracted together, the result should be reported with an amount of significant figures out to the lowest amount of significant figures found beyond the decimal point in any one number in the calculation. If no decimal point exists in the calculation, significant figures can be ignored. 2. When two numbers are multiplied or divided together, the result should be reported with an amount of significant figures equal to the lowest amount of significant figures found in any one number in the calculation. Fundamental Units of Measurement and Unit Prefixes Fundamental Units of Measurement The fundamental units of measurement in both Standard International and US Customary for the following quantities: These fundamental units can also be combined together within formulas to produce units with new names. Unit Prefixes The most commonly used prefixes in chemistry are those that involve power multiples of three except for These prefixes are combined on units of measurement to indicate a multiplication factor of a certain power of ten. Temperature Conversions in Fahrenheit, Celsius, and Kelvin Temperatures here in the U.S. are reported on the Fahrenheit scale, while most of the world uses the Celsius scale The value of temperature within most equations must be calibrated with an absolute lowest value of zero. The Kelvin scale has a degree size the same as its counterpart Celsius scale, but calibrated with an absolute lowest value of zero. Conversion Factors Conversion Factors are used to convert one unit measurement into another equivalent unit measurement as follows: Where each of the equal proportion values is determined by the relation between the values of the equivalent units. 3

4 The Atom An atom consists of a collection of a number of particles called the Nucleons that comprise the nucleus of an atom, and a collection of a number of particles called the Electrons in orbitals around this nucleus. Protons and Neutrons together are the Nucleons. The number of Protons determines the element and its nuclear charge, and the number of Nucleons (Protons and Neutrons together) determines the isotope and the atomic mass. The standard form for an atom or ion is: is the number of Protons and Neutrons together, the nuclear mass in amu, the molar mass in g, or the atomic number is the number of Protons, or the Atomic Number that determines the element and the nuclear charge in terms of +e is the number of Neutrons in the nucleus is the number of Electrons in orbitals around the nucleus is the overall charge of the atom which may be either zero (Neutral), positive (Cation), or negative (Anion) A Neutral atom is one that has an overall charge of zero and an equal number of Protons and Electrons ( ) A Cation has an overall positive charge and occurs when a neutral atom loses electrons and has more Protons ( ) An Anion has an overall negative charge and occurs when a neutral atom gains electrons and has more Electrons( ) Moles, Avogadros Number, Atomic Number, Atomic Weight and Molecular Weight The Atomic Number locates the position of the element within the periodic table and its chemical properties. The Atomic Weight and Molecular Weight are the masses of individual or collections of atoms and molecules. Amount in Moles and Avogadros Number The Mole is any amount that contains a standard number known as Avogadros Number of components in the amount: The components will typically be either atoms within a substance or molecules within a substance. Atomic Weight or Atomic Molar Mass The Atomic Weight is the number at the bottom of each and every element on the periodic table and is defined by: Remember mass abundance percentages are calculated as decimals within any equation Sum mass abundance percentage of each isotope commonly is a numerical value of 1 or a percentage value of 100%. Mass of a collection of atoms can be calculated from the Atomic Weight of the atoms by the following: Molecular Weight or Molecular Molar Mass The Molecular Weight is composed of the Atomic Weight from all atoms within the molecule and is defined by: Be careful with the subscript of each particular atom, which can be found according to the following rules: Mass of a collection of molecules can be calculated from the Molecular Weight of the molecules by the following: Atomic Mass Units The mass of individual or small collections of atoms or molecules is extremely small and expressed in Atomic Mass Units. The atomic mass unit is defined by the average mass of a single nucleon or a single component in the Carbon-12 atom. 4

5 Light, Photons, Spectra, and Quantized Energy Light has properties and behavior consistent with a particle or a wave depending on the measurements and situation. Light as a Wave, Wavelength, Light Spectrum, and Relations Light behaves like a wave in its ability to combine in such a way so as to become brighter as addition in Constructive Interference or to become dimmer as subtraction in Destructive Interference. From the interference, it can be determined that light is a wave with speed as the rate at which all points of the wave propagate through space, with wavelength as the distance between equivalent identical points on the wave, with period as the time needed for one full wavelength or one full cycle to pass, and frequency as the number of wavelengths or cycles that pass per second: Where the speed of light is and the Plancks Constant is Light Spectrum Light exists in a continuous spectrum of different wavelengths, frequencies, and energies as follows: The Electromagnetic Energy of Light is determined solely by its Wave Property of Wavelength. The wavelengths of visible light are colored and can be memorized in order of increasing frequency by The wavelengths longer than visible light are harmless and are used for radio broadcasts and telecommunications. The wavelengths shorter than visible light are harmful and are only used in small doses for medical and dental imaging. Light as a Particle, Quantized Energy, and Relations Light behaves like a particle in its ability to carry a finite amount of energy in a quantum packet, and interact with mass particles in collisions such as in the photoelectric effect. Energy carried by light of frequency and wavelength is Where Plancks Constant is the constant ratio for each energy equation Electron Energy Levels in Hydrogen Electron Energy Levels and Wavelength of light either emitted or absorbed from a Hydrogen atom are quantized: Where is initial energy level, is final energy level,, Rydberg Constant Absorption occurs when a photon is absorbed by the atom and the electron energy level is increased as Emission occurs when a photon is emitted by the atom and the electron energy level is decreased as Series are named by one value of as -2.0 ev -4.0 ev -6.0 ev -8.0 ev ev ev ev Lyman Series Balmer Series Paschen Series Paschen Series Wavelengths are longer than visible Balmer Series Wavelength Red for Wavelength Green for Wavelength Blue for Wavelength Violet for Wavelength Ultraviolet for Lyman Series Wavelengths are shorter than visible 5

6 Electrons and Electron Configuration The standard form for an atom or ion is: is the number of Protons and Neutrons together, the nuclear mass in amu, the molar mass in g, or the atomic number is the number of Protons, or the Atomic Number that determines the element and the nuclear charge in terms of +e is the number of Neutrons in the nucleus is the number of Electrons in orbitals around the nucleus is the overall charge of the atom which may be either zero (Neutral), positive (Cation), or negative (Anion) A Neutral atom is one that has an overall charge of zero and an equal number of Protons and Electrons ( ) A Cation has an overall positive charge and occurs when a neutral atom loses electrons and has more Protons ( ) An Anion has an overall negative charge and occurs when a neutral atom gains electrons and has more Electrons( ) Electron Configuration, Valence Electrons, and Core Electrons Electrons within an atom occupy locations known as orbitals. For a ground state atom, the electrons will completely fill a lower energy orbital before the next highest energy orbital. The list of electron orbitals in order of increasing energy is: The superscript of each electron orbital is the number of electrons that will completely fill the orbital. The Periodic Table itself follows the order of the electron configurations as the left two columns (alkali and alkaline) are the orbital being filled, the right six columns (metalloids and nonmetals) are the orbital being filled, the middle ten columns (transition metals) are the orbital being filled, and the lower fourteen columns (rare earth) are the orbital being filled. Electron Configuration To determine the electron configuration of any atom or ion: 1. Calculate the number of electrons in orbitals around the nucleus of the atom or ion. 2. Determine the location of the atom or ion within the periodic table. For a neutral atom, the location will be at an atomic number of which is at the location of the element. For a cation, the location will be at atomic number of which is units to the left of the element. For an anion, the location will be at atomic number of which is units to the right of the element. 3. Compare the location of the atom or ion in the Periodic Table and determine the configuration. Valence Electrons and Core Electrons Each time either the highest orbital is the and it is completely filled with 2 electrons or the highest orbital is any orbital and it is completely filled with 6 electrons, the atom is in a very stable and inert Noble Gas state. The highest completely filled or any completely filled orbital and all lower energy orbitals are the Core Electrons and can optionally be written in the electron configuration as the equivalent Noble Gas in brackets: The total number of electrons in an atom or ion is equal to the atomic number of the atom minus charge The core electrons are the number of electrons that are part of the Noble Gas Core Electrons The valence electrons are the number of electrons that are outside of the Noble Gas Core electrons When an atom loses electrons to become a cation, it will lose only its valence electrons in this order: When an atom gains electrons to become an anion, it will gain only its valence electrons in this order: 6

7 Periodic Properties and Trends A Periodic Property of an element has a certain value according to the location of the element on the Periodic Table. A Periodic Trend is a general name for a certain Property of the elements that has a consistent change as one moves topbottom or left-right across the Periodic Table. The Periodic Table was constructed for grouping together elements that have similar chemical properties and for displaying the Periodic Trends. The common Periodic Properties and Trends are: Electronegativity Electronegativity is a unitless measure of an atoms ability to attract electrons to itself after bonding to other atoms. Increases across the Periodic Table from lower left to upper right, Cs the lowest value 0.7 and F the highest value 4.0 The highest Electronegativity values from higher to lower can be remembered from the mnemonic The difference in the Electronegativity between two atoms will determine the type of bond that they will form Electron Affinity Electron Affinity is the ability of an atom to gain an electron or energy change when a gaseous atom gains an electron. Increases across the Periodic Table from lower left to upper right, Cs the lowest value and F the highest value. Ionization Energy Neutral Atom Ionization Energy is the energy needed to completely remove the outermost valence electron from the atom or ion. Increases across the Periodic Table from lower left to upper right, Cs the lowest value and Nobel Gases the highest value Ionization Energy Isoelectronic Series Ionization Energy is the energy needed to completely remove the outermost valence electron from the atom or ion. Isoelectronic Series is a set of different element atoms that all have the same number of electrons by some being ions. Increases across the Periodic Table with increasing Atomic Number. Atomic Radius Neutral Atom Atomic Radius is the radius and size of a particular element atom due to the outermost valence electron orbitals. Decreases across the Periodic Table from lower left to upper right, Cs the largest value and F the smallest value. Atomic Radius Isoelectronic Series Atomic Radius is the radius and size of a particular element atom due to the outermost valence electron orbitals. Isoelectronic Series is a set of different element atoms that all have the same number of electrons by some being ions. Decreases across the Periodic Table with increasing Atomic Number. Metallic Character Decreases across the Periodic Table from lower left to upper right, Cs the largest value and F the smallest value The diagonal border between the Metals and Nonmetals are the Metalloids B, Si, Ge, As, Sb, Te, At Metals: shiny, malleable into sheets, ductile into wires, good conductor of heat and electricity, higher melting/boiling point, donate bond electrons and form cations in aqueous solutions, oxides are ionic and form basic aqueous solutions. Nonmetals: dull, brittle but may be hard or soft, poor conductor (insulator) of heat and electricity, lower melting/boiling point, gain bond electrons and form anions in aqueous solutions, oxides are covalent and form acidic aqueous solutions. Reactivity Reactivity is the ability and rate to which a given element atom will react and form a bond with another atom. The Reactivity of an element atom is related to the extremity of its Electronegativity value, whether highest or lowest. Increases across the Periodic Table both towards lower left Cs and upper right F away from Nobel Metals Ag, Au, Pt, Pd 7

8 Atomic Structure The players within an atom and some useful information about each Name Symbol Mass (amu) Charge (e) Proton 1 +1 Neutron 1 0 Electron 0-1 Atom A C=Z-E An atom consists of a collection of a certain number of particles called the Nucleons that comprise the nucleus of an atom, and a collection of a certain number of particles called the Electrons that are in certain stable orbitals around this nucleus. Protons and Neutrons together are the Nucleons. The number of Protons determines the element and its nuclear charge, and the number of Nucleons (Protons and Neutrons together) determines the isotope and the atomic mass. The standard form for the number of Protons, Neutrons, and Electrons is the number of Nucleons, or the number of Protons and Neutrons together, the nuclear mass in atomic mass units amu, the molar mass in grams g, or the atomic number of the element and the numerical name of the element is the number of Protons, or the atomic number that determines the element and the nuclear charge in terms of +e is the number of Neutrons in the nucleus is the number of Electrons in orbitals around the nucleus is the overall charge of the atom which may be either zero (Neutral), positive (Cation), or negative (Anion) A Neutral atom is one that has an overall charge of zero and an equal number of Protons and Electrons ( ) A Cation has an overall positive charge and occurs when a neutral atom loses electrons and has more Protons ( ) An Anion has an overall negative charge and occurs when a neutral atom gains electrons and has more Electrons( ) Quantum Numbers Each electron has a set of 4 quantum numbers,,, and, to describe its location and energy on an orbital diagram Symbol Quantum Number Name Allowed Values Number of Possibilities Principle Quantum Number Angular Momentum Quantum Number Orbital Magnetic Quantum Number Spin Magnetic Quantum Number For a given there are a total of different possible allowed combinations of the four quantum numbers For a given there are a total of different possible allowed values of the Orbital Magnetic quantum number For a given there are a total of different possible allowed values of the Orbital Magnetic quantum number The Pauli Exclusion Principle states that only one electron can have a given combination of the four quantum numbers For a ground state atom, the electrons will all exist in the lowest energy combination of the four quantum numbers The energy of an electron is mainly dependent on the Principle Quantum Number, but to a lesser degree the energy is also dependent on the Angular Momentum Quantum Number Hunds Rule states that for a given Angular Momentum Quantum Number (collectively known as degenerate orbitals) a single electron will fill each orbital with the possible values of the Orbital Magnetic Quantum Number and identical signs for the Spin Magnetic Quantum Number before Spin Pairing a second electron having exactly the opposite sign of the Spin Magnetic Quantum Number together in each orbital. Atoms or Molecules with unpaired electrons in any orbitals will be Paramagnetic and be affected by a Magnetic Field Atoms or Molecules with paired electrons in all orbitals will be Diamagnetic and not be affected by a Magnetic Field 8

9 Electron Orbitals Electrons will exist only in certain allowed orbitals around the nucleus, with two electrons able to exist in each orbital, one with a positive spin and one with a negative spin. The allowed subshell orbitals and their shapes are the following Names Shapes Number of Orbitals Number of Electrons s 0 1 Sphere 1 2 p 1 3 Dumbbells 3 6 d 2 4 Cloverleaves, 1 Dumbbell and Donut 5 10 f f subshell Each of the orbitals (boxes) can contain 2 spin paired electrons, one with spin up and one with spin down An subshell is filled through the fourteen column Rare Earth elements Rare Earth elements are found at the bottom of the Periodic Table d subshell Each of the orbitals (boxes) can contain 2 spin paired electrons, one with spin up and one with spin down A subshell is filled through the ten column Transition Metal elements Transition Metal elements are found in the middle of the Periodic Table p subshell Each of the orbitals (boxes) can contain 2 spin paired electrons, one with spin up and one with spin down A subshell is filled through the six column Nonmetal elements Nonmetal elements are found at the far right of the Periodic Table s subshell Each of the orbitals (boxes) can contain 2 spin paired electrons, one with spin up and one with spin down An subshell is filled through the two column Alkali/Alkaline elements Alkali/Alkaline elements are found at the far left of the Periodic Table Effective Charge and Ionization Energy The (First) Ionization Energy is the energy needed to completely remove the outermost electron of an atom or ion. The outermost electron will be the one with the highest Principal Quantum Number, and first in the order: The core electrons shield or repel the outermost valence electrons from the number of protons in the nucleus, so that the amount of attractive charge felt by the valence electrons is lessened to a value known as is the least for the extreme left side of the Periodic Table, Group 1A or the Strong Alkali Metals. The one valence electron is weakly bounded to the attractive nucleus due to high shielding repulsion of the core electrons. These atoms have very low ionization energy, are very good conductors, and readily form ionic bonds with non metals. The ionization energy is so low that these atoms are never found in elemental form in nature and only exist in compounds. is also low for the far left side of the Periodic Table, Group 2A or the Strong Alkaline Metals. The two valence electrons are weakly bounded to the attractive nucleus due to high shielding repulsion of the core electrons. These electrons have low ionization energy, are good conductors, and readily form ionic bonds with non metals. The ionization energy is so low that these atoms are never found in elemental form in nature and only exist in compounds. throughout the center of the Periodic Table, or the Transition Metals. The Transition Metals are the elements whose highest energy electron orbital is the subshell. When a Transition metal is ionized, it loses its valence electrons first followed by its valence electrons since its valence electrons always have a higher Principal Quantum. The outermost valence electrons are weakly bounded to the attractive nucleus due to high shielding repulsion of the core electrons. These electrons have low ionization energy and are good conductors. throughout the upper right of the Periodic Table, or the Nonmetals. The Nonmetals are the elements whose highest energy electron orbital is the subshell. Even the outermost valence electrons are strongly bounded to the attractive nucleus due to low shielding repulsion of the core electrons. These electrons have high ionization energy and are poor conductors. The ionization energy is so high that these elements take electrons to fill the subshell octet. 9

10 Orbital Shapes, Wavefunctions, and Electron Density Electrons and other particles such as protons or more electrons interact with each other in such a way that each electron will tend to be found in only certain regions known as orbitals that represent the regions of highest probability of locating the electron at any given instant in time. It cannot be predicted with certainty where the electron will be located at any given instant in time, but it is 90% likely to be found somewhere within the orbital boundaries. There is even a vanishing small probability the electron will be found further away from the nucleus or anywhere in the Universe. Orbital Orbital Shape Wavefunction Graph Electron Density Graph 1s has 0 Nodes 1s has 0 Nodes 1s 2s has 1 Node 2s has 1 Node 2s 3s has 2 Nodes 3s has 2 Nodes 3s 2p has 1 Node 2p has 1 Node 2p 3p has 2 Nodes 3p has 2 Nodes 3p 10

11 Valence Shell Electron Pair Repulsion (VSEPR) Theory Valence Shell Electron Pair Repulsion Theory describes the electron distribution within atoms, the electron distribution within a molecule after the bonding of atoms, and the overall shape of a molecule after the bonding of atoms. and Atomic Orbitals The orbital has a spherical shape and has a distribution closer to the nucleus of the atom. The orbital has a dumbbell shape and has a distribution further away from the nucleus of the atom. Hybrid Atomic Orbitals The Hybrid Orbitals are created when the single valence electron orbital and number of valence electron orbitals interact and combine together forming number of Hybrid Orbitals. These Hybrid Orbitals will form number of Bond Complexes or Lone Pairs around the atom. The remaining number of valence electron orbitals are available to form number of extra bonds in the Bond Complexes, creating multiple bonds. The value of is called the index of the Hybrid Orbital. The index determines the size of the orbital and the angle associated with the orbital: an orbital has a larger angle ( ) and a smaller radius and a orbital has a smaller angle ( ) and a larger radius. The larger the index of a Hybrid Orbital, the larger the character and the larger the radius associated with the orbital, but the smaller the character and the smaller the angle associated with the orbital. Hybrid Atomic Orbitals The Hybrid Orbitals are created when the single valence electron orbital, number of valence electron orbitals and number of valence orbitals interact and combine together forming number of Hybrid Orbitals. These Hybrid Orbitals will form number of Bond Complexes or Lone Pairs around the atom. The atom will then exceed the Octet Rule, instead reaching the Decatet Rule (10 electrons) or Dodecatet Rule (12 electrons). Atomic Bond Complexes and Lone Pairs Two atoms will combine their valence electron wavefunctions and form bonds in one of two different ways: Valence Electrons of each and every bonded atom will be involved either in a Bond Complex or in a Lone Pair. A Bond Complex is formed by the shared valence electrons between two atoms as their wavefunctions combine A Bond Complex exists between every pair of bonded atoms in a molecule, and the energy holds the molecule together A Lone Pair is formed by two valence electrons that are not shared in bonding as their wavefunctions combine A Lone Pair exists around some atoms in a molecule, but always away from any Bond Complexes of that atom Atom Hybridization Type The Atom Hybidization Type is a value attributed to each atom in a molecule and determines the number of Electron Domains around the particular atom. Each Electron Domain holds either a Bond Complex or a Lone Pair. Hybridized Atoms one Electron Domain at : one single bond. Only Hydrogen exists as this Hybridization Type. Hybridized Atoms two Electron Domains at : two double bonds or triple bond plus one single bond or lone pair Hybridized Atoms three Electron Domains at : one double bond plus two total single bonds or lone pairs Hybridized Atoms four Electron Domains at : four total single bonds or lone pairs Hybridized Atoms five Electron Domains three at and two at : five total single bonds or lone pairs Hybridized Atoms six Electron Domains at : six total single bonds or lone pairs 11

12 Orbital Overlap Interactions to form Atomic Hybrid Orbitals The value of is called the index of the Hybrid Orbital. The index determines the size of the orbital and the angle associated with the orbital: an orbital has a larger angle ( ) and a smaller radius and a orbital has a smaller angle ( ) and a larger radius. The larger the index of a Hybrid Orbital, the larger the character and the larger the radius associated with the orbital, but the smaller the character and the smaller the angle associated with the orbital. It is not necessary for the Hybrid Orbital index to be an integer. It is only an integer for the more symmetrical molecules. The exact value of the index is determined solely by the angle the Hybrid Orbital occupies within the atom. It is possible to predict the influence a Bond Complex or Lone Pair has on the Hybrid Orbital index, radius, and angle Lone Pairs use more of the orbital, pull into a smaller radius, and spread out more around an atom, forcing any Bond Complexes on the same atom to use more of the orbitals, have a larger radius and pull closer together to each other. Hybrid Orbital Overlap Interaction Hybrid Orbital Result Hybrid Angle Geometry Spherical Linear Trigonal Planar Tetrahedral Decreasing Angle with Increasing m Perpendicular 12

13 Electron Domain Geometry and Molecular Geometry Electron Domain Geometry describes the configuration of the Electron Domains around a particular atom. Molecular Geometry describes the configuration of just the Bond Complex Domains around a particular atom and ignores the Lone Pair Domains which always remain closer to the atom and are not noticeable in the overall atomic shape. To determine: 1. Draw the entire Lewis Structure of the molecule including the Lone Pairs on any atoms. 2. Count the number of Electron Domains, which include both Bond Complex Domains and Lone Pair Domains. 3. Determine the number of both Bond Complex Domains and Lone Pair Domains. 4. Use the table to determine the Electron Domain Geometry, Molecular Geometry, Hybridization, and Bond Angle Number of Electron Domains Electron Domain Geometry Hybridization and Bond Angle Bond Complex or Bonding Domains Lone Pair or Nonbonding Domains Molecular Geometry 2 Linear 2 0 Linear 3 Trigonal Planar 3 0 Trigonal Planar 3 Trigonal Planar 2 1 Bent 4 Tetrahedral 4 0 Tetrahedral 4 Tetrahedral 3 1 Trigonal Pyramidal 4 Tetrahedral 2 2 Bent 5 Trigonal Bipyramidal 5 0 Trigonal Bipyramidal 5 Trigonal Bipyramidal 4 1 Seesaw 5 Trigonal Bipyramidal 3 2 T Shaped 5 Trigonal Bipyramidal 2 3 Linear 6 Octahedral 6 0 Octahedral 6 Octahedral 5 1 Square Pyramidal 6 Octahedral 4 2 Square Planar Electronegative Difference, Dipole Moment, and Polar Molecules The difference in the Electronegativity between two atoms will determine the type of bond that they will form Any bond made between two identical atoms will have no electronegative difference as electrons are equally shared. Any bond made between two different atoms will have an electronegative difference as electrons are unequally shared. A bond with an electronegative difference will have a dipole moment arrow pointed from the less electronegative atom towards the more electronegative atom and with a magnitude equal to the electronegative difference within the bond. If all dipole arrows contained within a molecule cancel in magnitude and direction, the molecule will not be a Polar Molecule and will have a net magnetic dipole moment of zero even if some of the bonds within the molecule are polar. If one or more dipole arrows contained within a molecule do not cancel in magnitude and direction, the molecule will be a Polar Molecule and have a net magnetic dipole moment in the direction summation of all bond dipole moments. Some Molecular Geometries will not be a Polar Molecule if all atoms bonded to the central atom are exactly the same: Some Molecular Geometries will always be a Polar Molecule regardless of the atoms bonded to the central atom: Intermolecular Forces Intermolecular Forces hold molecules together in liquid and solid states. The stronger the intermolecular forces, the higher the boiling point and the lower the vapor pressure of the liquid. The types of intermolecular forces are: London Dispersion Forces(All Molecules especially larger ones) induced dipole moments by electron cloud deformities. Dipole Interactions(Ion-Dipole, Dipole-Dipole) attraction between two polar molecules or polar with cations and anions. Ion Interactions (Ion-Ion, Ion Dipole) attraction between two ionic molecules or an ionic molecule with a polar molecule. Hydrogen Bonding (Molecules with bonded directly to,, or ) large electronegative difference with a polar bond. 13

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15 Molecular Bonding Octet Rule is that most atoms seek a minimum energy state of completely filled and orbitals for a total 8 electrons: Group 1A Alkali Metal Atom will lose its one valence electron in forming an aqueous solution or an ionic bond. Group 2A Alkaline Metal Atom will lose its two valence electrons in forming an aqueous solution or an ionic bond. Upper Right Nonmetals will gain other atom electrons in forming an ionic bond, polar covalent bond, or covalent bond. Duet Rule is that Hydrogen and Helium have a orbital only that is filled with just 2 electrons. Lithium (Group 1A) will lose its one valence electron and Beryllium (Group 2A) will lose its two valence electrons to a filled orbital. Hextet Rule is that Group 3A Aluminum, Boron, Indium have 6 electrons in three covalent bonds and an empty orbital. Decatet and Dodecatet Rule is that Phosphorous, Sulfur, Chlorine, Bromine, Iodine use orbital for 10 or 12 electrons. Each atom in a molecule has a Formal Charge assigned to it, a measurement of its tendency to be in the Lewis Structure Lone Pair Electrons is the Number of Electrons involved in a lone pair around the atom, 2 for each lone pair Bond Complex Electrons is the Number of Electrons in each bond, 2 for single bond, 4 for double bond, 6 for triple bond Higher electronegative atoms have a negative Formal Charge while lower electronegative have a positive Formal Charge. Atom Form 1 Form 2 Form 3 Form 4 Form 5 Notes Hydrogen Hydrogen only forms one bond Alkali 1A Li,Na,K,Rb,Cs The Alkali Group 1A always lose one electron in forming ionic bonds for Formal Charge 1+ Alkaline 2A Be,Mg,Ca,Sr,Ba Group 3A B,Al,Ga,In Carbon C Nitrogen N Phosphorous P Oxygen O Sulfur S Halogen 7A F,Cl,Br,I The Alkaline Group 2A always lose two electrons in forming ionic bonds for Formal Charge 2+ The Group 3A is the Hextet Rule with an empty p orbital, but can follow the Octet Rule Carbon follows the Octet Rule with four bonds and readily forms multiple bond complexes Nitrogen follows the Octet Rule and prefers three bonds with one lone pair but may have four bonds Phosphorous is the Octet Rule at three bonds with one lone pair or the Decatet Rule at five bonds Oxygen follows the Octet Rule at two bonds with two lone pairs or one bond with three lone pairs Sulfur follows the Octet Rule at two bonds with two lone pairs, four bonds, or six bonds Fluorine is a terminal atom only with Form 1. Others may exceed the Octet Rule as central atoms 15

16 Resonance Resonance forces molecules towards symmetry of bonds and lone pairs to achieve a lower energy state. Symmetry is created by the interchange of bonds, empty orbitals, and lone pairs or radicals while bonds remain unchanged. Allylic Resonance A lone pair is on an atom directly bonded to a second atom that itself has a bond with a third atom. The lone pair moves between the first and second atoms to become a bond while the already existing bond moves onto the third atom to become a lone pair on the atom. This forms an equivalent resonance structure. Vicinic Resonance A lone pair is on an atom directly bonded to an atom with an empty orbital. The lone pair moves between the two atoms to become a bond. This forms an equivalent resonance structure. Aromatic Resonance A ring is composed of alternating single and double bonds, known as conjugated bonds. Each bond moves over to the adjacent single bond forming a new bond. This forms an equivalent resonance structure. Bond Order and Resonance Bond Order is a measure of the strength of a bond between any two atoms. A Bond Order of 0 means no bond will form, and Bond Order of 1 is a single bond, a Bond Order of 2 is a double bond, and a Bond Order of 3 is a triple bond. Bond Order can also have a fractional value as single, double, and triple bonds average out through resonance. Formal Charge and Resonance Each atom in a molecule has a Formal Charge assigned to it, a measurement of its tendency to be in the Lewis Structure Lone Pair Electrons is the Number of Electrons involved in a lone pair around the atom, 2 for each lone pair Bond Complex Electrons is the Number of Electrons in each bond, 2 for single bond, 4 for double bond, 6 for triple bond Higher electronegative atoms have a negative Formal Charge while lower electronegative have a positive Formal Charge. Exceptions occur when the bonded atoms have a large electronegative difference as in ionic bonds or polar covalent bonds. The higher electronegative atom will gain extra electrons within the bond and have a negative Formal Charge while the lower electronegative atom will lose extra electrons within the bond and have a positive Formal Charge. Common Formal Charge Results Most common results in the Formal Charge of the following atoms: Li, Na, K, Rb, Cs form ionic bonds with nonmetals and lose one valence electron in doing so for a Formal Charge of 1+ Be, Mg, Ca, Sr, Ba form ionic bonds with nonmetals and lose two valence electrons in doing so for a Formal Charge of 2+ B, Al usually form three single bonds for a Formal Charge of 0, rarely form four single bonds for a Formal Charge of 1- H always forms one bond (single) for a Formal Charge of 0 C usually forms four bonds (four single, two single and double, two double, or triple and single) for a Formal Charge of 0 N, P usually forms three bonds (three single, single and double, or triple) with one lone pair for a Formal Charge of 0 N rarely forms four bonds (four single, two single and double, two double, or triple and single) for a Formal Charge of 1- P commonly forms five bonds (five single) for a Formal Charge of 0 O, S usually forms two bonds (two single, or double) with two lone pairs for a Formal Charge of 0 O commonly forms one bond (single) with three lone pairs for a Formal Charge of 1- S commonly forms four bonds (four single) with one lone pair for a Formal Charge of 0 F, Cl, Br, I usually form one bond (single) with three lone pairs for a Formal Charge of 0 Cl, Br, I commonly form three bonds (three single) with two lone pairs for a Formal Charge of 0 F, Cl, Br, I form ionic bonds with Li, Na, K, Rb, Cs, Be, Mg, Ca, Sr, Ba and gain a single electron for a Formal Charge of 1- Formal Charge can also have a fractional value as single, double, and triple bonds average out through resonance. 16

17 Oxidation Number (Oxidation State) of an Atom, Molecule, or Ion Molecules or ions often stabilize with an increase or a decrease of electrons and have a total charge or Oxidation Number. The Oxidation Number of a molecule or an ion is the total oxidation number of all atoms in the molecule or ion: An atom with an unknown Oxidation Number or charge can be found by relating the total molecule Oxidation Number or charge of the molecule and the total known Oxidation Numbers or charges of all other atoms in the molecule or ion. Periodic Atom Oxidation Numbers 1A Group +1 (Alkali Metals ) 2A Group +2 (Alkaline Metals ) 3A Group +3 ( ) 5A Group -3 ( ) or +5 ( ) or +3 ( ) 6A Group -2 ( ) or -2 Peroxide ( ) Metal Atom Oxidation Numbers Chromium(II) or Chromous Chromium(III) or Chromic Cobalt(II) or Cobaltous Cobalt(III) or Cobaltic Iron(II) or Ferrous Iron(III) or Ferric 7A Group -1 unless internal (Halogens ) H +1 ( ) unless bonded with any Metal then -1 ( ) O -2 ( ) unless bonded to another O each -1 ( ) Ammonium, Silver +1 ( ) Cadmium, Nickel, Zinc +2 ( ) Copper(I) or Cuprous Copper(II) or Cupric Tin(II) or Stannous Tin(IV) or Stannic Lead(II) or Plumbous Lead(IV) or Plumbic +1 Oxidation Number Alkali Ammonium Hydrogen Silver +2 Oxidation Number Alkaline Iron Zinc +3 Oxidation Number Iron -1 Oxidation Number Acetate Hypochlorite Chlorite Chlorate Perchlorate Cyanide Halide Hydride Hydroxide Nitrite Nitrate -2 Oxidation Number Carbonate Chromate Oxide Peroxide Sulfide Sulfite Sulfate -3 Oxidation Number Phosphate Oxidation The Oxidation Number increases due to the loss of electrons. The molecule containing the atom that undergoes this reaction is the reducing agent since it reduces the other molecule containing the reduced atom. Reduction The Oxidation Number decreases due to the gain of electrons. The molecule containing the atom that undergoes this reaction is the oxidizing agent since it oxidizes the other molecule containing the oxidized atom. Bonding Ionic bonds occur when electrons are donated from a metal atom making a cation to a nonmetal atom making an anion. Covalent bonds occur when electrons are shared equally between atoms and usually occur between two nonmetals. Metallic bonds occur when electrons are shared among all metallic atoms and all atoms are bonded together in a solid. Lattice Energy Lattice Energy is the energy contained within an ionic bonded solid. Lattice Energy increases with Oxidation Number Pair To compare molecules with equal Oxidation Number Pair, the Lattice Energy in an ionic bonded solid is higher for molecules made with element atoms that are smaller in size or equivalently element atoms that are higher in a group column on the Periodic Table. Ionic bonded solids with a higher Lattice Energy have higher intermolecular forces. Higher intermolecular forces lead to a stronger and harder substance with a higher melting point and a higher boiling point. 17

18 Organic Chemistry Organic chemistry is of carbon which has the unique periodic properties of being a small atom with four valence electrons and can make single, double, or triple bonds to itself or other atoms producing billions of possible compounds. Structure of the molecule is determined by the carbon-carbon bond complexes and the ring shapes of the molecule: Structure Molecular Formula Condensed Structures Suffix Notes Alkane Has only single carbon-carbon bonds but not in a ring Alkene Alkyne Cycloalkane Benzene Has all single except exactly one double carbon-carbon bond Has all single except exactly one triple carbon-carbon bond Has at least three or more ringed single carbon-carbon bonds Has alternating single and double conjugated carbon-carbon bonds Functional Groups of the molecule are common groups that are directly connected to the carbons of the structure: Functional Group Group Formula Condensed Structures Suffix Notes Alcohol Weak Acid by donating from the group Amine Ketone Aldehyde Carboxylic Acid Ester Amide Moderate Base by bonding the nitrogen N lone pair Weak Base by bonding the oxygen O lone pair Weak Base by bonding the oxygen O lone pair Moderate Acid by donating from the group Weak Base by bonding the oxygen O lone pair Moderate Base by bonding the oxygen O lone pair Formulas and Isomers Structural Isomers are molecules with the same molecular formula but atoms bonded together in different combinations Geometric Isomers are molecules with the same molecular formula and combinations but different spatial geometry Cis Isomers have groups on the same side of the atom and Trans Isomers have groups on the opposite side of the atom 18

19 Chemical Formulas and Combustion Analysis The Chemical Formula is the number of atoms in a molecule and Combustion Analysis can find the Chemical Formula. Chemical Formulas and Mass Percent The Chemical Formula describes the number for each type of atom that is involved in bonding to comprise a molecule. Mass Percent The Mass Percent of an atom in a molecule is the percent ratio of the atom mass to the entire molecule mass: Empirical Formula The Empirical Formula is the proportion in lowest integer form for the number of each atom that comprises a molecule. To find the Empirical Formula from the Mass Percent of each atom: 1. Assume 100 grams of the substance and multiply the 100 by each atom mass percent for the mass of each atom. 2. Calculate the number of moles of each atom by dividing the mass with the atomic weight of each atom. 3. If the number of moles for every atom is an integer, the number of moles is its subscript in the Empirical Formula. If one or more of the number of moles for every atom is not an integer, continue to step number Divide each of the number of moles for every atom by the lowest number of moles of an atom found. If all of the numbers for every atom after the division is an integer, the integer is its subscript in the Empirical Formula. If one or more of the numbers for every atom after the division is not an integer, continue to step number Take the values of the numbers for every atom after the division and multiply each by the lowest number needed to make all of them become integers after the multiplication by the number. After the lowest number multiplication to make each an integer, the integer of an atom is its subscript in the Empirical Formula. Molecular Formula The Molecular Formula is the exact value in integer form for the number of each atom that comprises a molecule. To find the Molecular Formula from the Mass Percent of each atom: 1. Calculate the Empirical Formula using the steps as given above. 2. Calculate the molecular weight of the Empirical Formula and divide the given molecular formula weight by the calculated Empirical Formula molecular weight to produce a number. 3. Multiply the subscript of each atom in the Empirical Formula by this number to produce the Molecular Formula. Combustion Analysis Combustion Analysis is a method to calculate the Empirical Formula or Chemical Formula of a hydrocarbon or alcohol. 1. Calculate the moles of carbon dioxide and water by dividing the given masses by molecular weights. 2. Calculate the moles of reactant carbon atoms as equal to moles of carbon dioxide: 3. Calculate the moles of reactant hydrogen atoms as twice the moles of water:. For a hydrocarbon which contains no oxygen, skip to step Calculate the moles of product oxygen atoms 5. Calculate the mass of atmosphere oxygen atoms by subtracting the masses of the products and reactants: 6. Calculate the moles of atmosphere oxygen by dividing the mass of atmosphere oxygen by its molecular weight. 7. Calculate the moles of reactant oxygen by subtraction: 8. Divide each of the number of moles for every atom by the lowest number of moles of an atom found. If all of the numbers for every atom after the division is an integer, the number of an atom is its subscript in the Empirical Formula. If one or more of the numbers for every atom after the division is not an integer, multiply each by the lowest number needed to make all of them integers after the multiplication for the Empirical Formula. 9. Divide the given molecular weight by the calculated Empirical Formula molecular weight to produce a number. 10. Multiply the subscript of each atom in the Empirical Formula by this number to produce the Molecular Formula. 19

20 Ideal Gases and Ideal Gas Law, Real Gases and Van der Waals Equation Gases expand to fill the entire volume of a container since their kinetic energy at a given temperature is much greater than the attractive intermolecular forces that hold the gas molecules together. Gases may be either Ideal or Real. Pressure Pressure is the collective force per area created by the collisions of gas molecules with a container or surface Pressure units include the atmosphere or pressure of air at sea level, millimeters of mercury or or the height that a column of mercury is forced to by a gas, pounds per square inch or force in pounds created on each square inch of the surface, and the Pascal unit or force in newtons created on each square meter of the surface: Partial Pressure is the contribution of a single gas to the total overall pressure of a mixture of two or more gases Partial Pressure may calculated from parts per million as or from parts per billion as Mercury Manometer is a device used to measure the pressure of a gas by its effect on a column of mercury in a tube. Ideal Gas Laws Many gases approximate the Ideal Gas state. A gas must meet two conditions to be ideal and follow the Ideal Gas Laws: The Ideal Gas Laws relate state variables (pressure, volume, and temperature) and amount of gas in moles or molecules. Avogadros Law The proportion of volume occupied by a gas to the amount of gas in moles remains constant Boyles Law The product of the pressure created by a gas and the volume occupied by a gas remains constant Charles Law The proportion of volume occupied by a gas to the temperature of the gas in Kelvin remains constant Ideal Gas Law Moles n Version Ideal Gas Law Density d Version Standard Temperature and Pressure STP Standard Temperature and Pressure is defined as the state where pressure and temperature From the Ideal Gas Law it can be calculated that the molar volume occupied by any ideal gas at STP is Temperature used in Kelvin Real Gas Laws and Van der Waals Equation Real gases do not always exist at low pressures and/or high temperatures necessary for the accuracy of the Ideal Gas Law. At non ideal conditions, the equation that models the state variables of the gas is the Van der Waals Equation: With the term being the increase in pressure correction due to the attractive intermolecular forces, and the term being the decrease in available volume correction from the small yet finite volume of individual molecules. 20

21 Kinetic Molecular Theory Kinetic Molecular Theory is the set of laws describing the motions and behavior for a collection of gas molecules, which may involve an extremely large number of molecules moving in random directions and with random speeds. To handle a large variety of motions, Kinetic Molecular Theory is based on statistical probabilities under the following conditions: 1. The gas molecules themselves must be small and occupy a very small fraction of the total container volume. 2. The gas molecules must have small intermolecular forces and have no attractive or repulsive forces. 3. The gas molecules each move in a straight line but in a random direction and with a random speed. 4. The gas molecules only interact in perfectly elastic collisions such that no energy is lost in any interactions. 5. The gas molecules have an average kinetic energy proportional to the absolute temperature in Kelvin. Molecular Speed Distribution and Root Mean Square Speed Each molecule in a gas will be moving with a random speed; some molecules will be moving slow, a few will be moving very fast, but the majority will be moving around the same middle range of speeds as seen in the following graphs: Fraction of Molecules at each Speed Molecular Speed The above colored coded graphs can be compared to each other according to the following schemes: Absolute Temperature The higher absolute temperature in Kelvin, the faster individual molecules move and the more the curve is stretched Molecular Weight The higher molecular weight in grams, the slower individual molecules move and the more the curve is compressed Root Mean Square Speed The Root Mean Square Speed is the value around which the middle range of speeds is centered in a Molecular Speed Distribution graph and has a value determined by the average amount of Kinetic Energy for all molecules in the gas. Effusion and Grahams Law Effusion is the rate at which gas molecules escape through a small hole in a container. Grahams Law relates the rates of effusion and or the times of effusion and for two gases with different molecular weights and : From Grahams Law it can be determined that the rate of effusion for a gas decreases as the molecular weight increases. Diffusion, Mean Free Path, Collision Frequency, and Pressure Diffusion is the rate at which a gas spreads out into an open volume which decreases as the molecular weight increases. Mean Free Path is the average distance a gas molecule travels between collisions and is inversely related to pressure. Collision Frequency is the rate at which collisions occur in a gas and is directly related to only the pressure of the gas. 21