chapter 1

The Atomic Theory Daltons atomic theory Elements are composed of extremely small particles called as atoms. All atoms of a given element are identical Having the same size, mass and chemical properties The atoms of one element are different from the atoms of all other elements Compounds are composed of atoms of more than one element. A chemical rxtn involves only the separation, combination or rearrangement of atoms. It doesnt result in their creation or destruction Matter cannot be destroyed/created. Law of Conservation of Mass Atom is made up of smaller particles which are called subatomic particlesElectron (-1) Atoms consist of a nucleus (+ve charge) and surrounded by an electron cloud (-ve charge) Component of Atom Atom as the basic unit of an element that can enters into chemical combination Proton, (+1) Neutron (0) electron e e e e nucleus 4 atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m The protons and neutrons are packed in an extremely small nucleus. Electron are shown as clouds around the nucleus no. of proton = no of electronAtoms are neutral5 An ion is an atom, or group of atoms, that has a net positive or negative charge. cation ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na+ 11 protons 10 electrons Cl 17 protons 17 electrons Cl- 17 protons 18 electrons p > e e > p In neutral, proton = electron Each elements is characterized by its proton number. Each element has a differ proton number E.g: atomic number N is 7 N has 7 proton & 7 electron = neutral atom Therefore, atom with 7 protons is always nitrogen, N. Atomic Number, (Z) Atom No. (Z)the number of proton in the nucleus of each atom of an element Mass Number, (A) Mass No. (A)the total amount of neutrons & protons present in the nucleus of an atom of an element Mass number (A) = number of protons + number of neutrons =atomic number (Z) + number of neutrons X A Z Mass NumberAtomic Number Element Symbol 8 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons How many protons, neutrons, and electrons are inC 14 6 ? How many protons, neutrons, and electrons are inC 11 6 ? Example: Isotopes have the same number of electron- they have the the identical chemical properties They have different neutron number have different physical properties (melting point, boling point, density) Isotopes Isotopesatoms that have the same atomic number but different mass number 10 The Isotopes of Hydrogen deuterium hydrogentritium Electron Configuration Show how e- are filled in the orbital Its describe the arrangement of e- in an atom 1s1 principal quantum number n angular momentum quantum number l number of electrons in the orbital or subshell Orbital diagram H 1s The Pauli Exclusion Principle Pauli Principle State that no 2 e- in the same atom can have the same 4 quantum number (n, l, ml, ms) an orbital can hold a max of 2 e- & they must has opposite spins If there are 3 e- in the same orbital; e1: n = 1, l = 0, ml = 0, ms = + (permissible) e2: n = 1, l = 0, ml = 0, ms = - (permissible) e3: n = 1, l = 0, ml = 0, ms = + (not permissible) 2s1s As Li (3 e-) n = 1, l = 0, ml = 0, ms = + n = 1, l = 0, ml = 0, ms = - n = 1, l = 0, ml = 0, ms = + Aufbau Principle Aufbau German word means building up Aufbau Principle State that e- in an atom should be filled in the orbitals in the order of increasing energy level e- should occupy the orbital with the lowest energy first before it enters the one with higher energy E.g:Li has 3 e- Thefirst 2 e- should be placed into 1s orbital (lowest energy level) Follow by 3rd e- in the 2s orbital 1s 2s1s Therefore, electron configuration of Li is= 1s2 2s1 Figure 2: Distribution of energy levels in an atom Energy level diagram to follow the order of higher energy of the orbital, refer to energy diagram 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s Figure 1: Order for fill the e- in the orbital Hunds RuleHunds Rule State thate- will occupy all orbitals of the same energy level singlywith parallel spin before they become paired 2s E.g: N (7 e-) 2p 1s Therefore, electron configuration of N is= 1s2 2s2 2p3 What is the electron configuration of Mg? Mg12 electrons 1s < 2s < 2p < 3s < 3p < 4s1s22s22p63s22 + 2 + 6 + 2 = 12 electrons Abbreviated as [Ne]3s2 What are the possible quantum numbers for the last (outermost) electron in Cl? Cl17 electrons1s < 2s < 2p < 3s < 3p < 4s1s22s22p63s23p52 + 2 + 6 + 2 + 5 = 17 electrons Last electron added to 3p orbital n = 3l = 1ml = -1, 0, or +1ms = or - Electronic conf: Electronic conf: Electron configuration for Cation and Anion Na (11 e-) = 1s2 2s2 2p6 3s1 Na+= 1s2 2s2 2p6 or[Ne] Ca (20 e-) = 1s2 2s2 2p6 3s2 3p6 4s2 Ca2+ = 1s2 2s2 2p6 3s2 3p6 or [Ar] F= 1s22s22p5F-= 1s22s22p6or[Ne] H= 1s1H-= 1s2 or [He] Electron configuration of the transition elements Transition elements are the element with the half-filledd orbitals. These elements are found in the fourth row of periodic table known as the d-block element. The 3d orbitals are filled after 4s orbital is fully occupied by e-. Example: Titanium, Ti: (Z = 22) 2s2p 1s3s3p 3d 4s Electron conf:1s2 2s2 2p6 3s2 3p6 4s2 3d2 [Ar] 4s2 3d2 or The Anomalous Electronic Configuration of Chromium and Copper Chromium , Cr and Copper, Cu = the element in d-block It have irregularities for this element 1s2 2s2 2p6 3s2 3p6 4s2 3d4 [Ar] 4s2 3d4 Cr : 24 electrons or Expected to be based on the rulesBUT To achieve the stability, one of the e- from the 4s orbital occupies one of the 3d orbitals to have a half-filled orbital 1s2 2s2 2p6 3s2 3p6 4s1 3d5 [Ar] 4s1 3d5 Cr : 24 electrons or the actualCu : 29 electrons 1s2 2s2 2p6 3s2 3p6 4s2 3d9 1s2 2s2 2p6 3s2 3p6 4s1 3d10 turn as [Ar] 4s1 3d10or Its said to have fully-filled orbital 24 25 Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n 1)d orbitals. Fe:[Ar]4s23d6 Fe2+:[Ar]4s03d6 or [Ar]3d6 Fe3+:[Ar]4s03d5 or [Ar]3d5 Mn:[Ar]4s23d5 Mn2+:[Ar]4s03d5 or [Ar]3d5 Development of Periodic Table CHAPTER 1.2 CHEMICAL BONDING Introduction Ionic Bonding Covalent Bond Physical Properties Of Ionic Compound Physical Properties Of Covalent Compound Lewis Dot Structure Resonance Octet Rule Exceptions To The Octet Rule Content Atoms are made up of smaller sub-particles: electron, protons & neutrons Proton + neutron are in small dense nucleus of atom Electron are arranged in orbitals outside the nucleus according to Aufbau principle, Pauli Exclusion principle & Hunds rule An atom can gains or loses e- to become ion OR share e- with other atom to become molecule INTRODUCTION Atoms/ions are joined together by 2 main types of chemical bonding Ionic bond Covalent bond Why do atoms combine?? Combine to form a stable e- configuration similar to the unreactive and chemically stable noble gas The force that hold on the atoms/ions are called CHEMICAL BOND Chemical bonds form whenAttractive forces between the positive charged nucleus & negative charged electron are countered balance by repulsive forces among the nucleus and electrons themselves. 2 types of chemical bonds: 1) Ionic bond Formation bonds results of transfer of electron 2) Covalent bond Formation of chemical bonds as results of sharing of electron When atoms interact to form a chemical bond, only their outer region are in contact Valence electrons Lewis dot symbol is used to show how many valence electron is belonged to an atom used in Lewis Structure to demonstrate the formation of bond Lewis dot symbol Consist of the symbol of an atom & one dot for each valence electron in an atom of the element 33 Valence electrons are the outer shell electrons of an atom.The valence electrons are the electrons that participate in chemical bonding. 1A1ns1 2A 2ns2 3A 3ns2np1 4A 4ns2np2 5A 5ns2np3 6A 6ns2np4 7A 7ns2np5 Group# of valence e-e- configuration To draw the Lewis dot symbol atomic symbol is drawn with dots or crosses surrounding E.g:Oxygen (Z= 8)- 1s2 2s2 2p4 valence electron Lewis dot symbol for Oxygen; . .O . ... 35 Lewis Dot Symbols for the Representative Elements & Noble Gases Define as = the electrostatic force that holds ions together in an ionic compound IONIC BOND Ionic bond form when an atoms: transferring (losing/gaining) e- from atom to another atoms Atom that donates its valence electron will be +ve charged (CATION) Atom that accept electron will be ve charged (ANION) The opposite charged ions are held together by strong electrostatic attraction Atoms that donate their valence e- are very electropositive element metal of group 1A & 2A Atoms that accept e- are very electronegative elements nonmetals of group 6A & 7A E.g: reaction between Li and F to form LiF 38 Li+F Li+F - Li Li++e- e-+ FF - F - Li++Li+F - LiF By using Lewis structures and curve arrows, describe the transfer of electron(s) during the formation of the following ionic compounds. 1. Na2O 2. CaF2 Try This: Ionic compound are have high melting and boiling point The ions are stable ( have stable e- configuration) & the ionic bonds are strong Soluble in water but insoluble in organic solvents as benzene Can conduct electricity in liquid form or aqueous form only Because of the present of charged particles that can movePhysical Properties of Ionic Compound Means = is a chemical bond in which 2 or more electrons are shared by 2 atoms Covalent compound = compounds that contain only covalent bond COVALENT BOND A group of atoms are being held together Electron in a shared pair is attracted to nuclei of both atoms Covalent bond usually formed by non-metal atoms having the same or nearly same electronegativity In covalent bond, 43 Why should two atoms share electrons? FF+ 7e-7e- FF 8e-8e- Lewis structure of F2 FF FF lone pairslone pairs lone pairslone pairs single covalent bond single covalent bond Pair of valence electron that are not involve in covalent bond formation Lone pair/unpaired electron The formation of these molecules illustrate the OCTET RULE An atom other than hydrogen tends to form bonds until it is surrounded by 8 valence electrons Atoms can form different types of covalent bonds: Single bondDouble bond Triple bond Two atoms are held together by one electron pair Single bond Double bond Is when 2 atoms share 2 pairs of electron(s) Can found in molecules of CO2 OHHOHHor OCOorOCO double bonds Triple bond Arise when 2 atoms share 3 pairs of electron(s), example N2 NNNN triple bond triple bond or Although atoms often form compounds by sharingelectrons, the electrons are not always shared equally. Polar Covalent Bonds Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end. HF F2 49 Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms H F electron rich region electron poor region FH e- riche- poor d+d- When two atoms share electrons unequally, a bond dipole results. The dipole moment, m,produced by two equal but opposite charges separated by a distance, r. Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond. Polar Covalent Bonds covalent compound are have low melting and boiling point The weakness of intermolecular force of covalent bond Can soluble in organic solvents as benzene but not in waterIts not good as conductors of electricity in all form Molecules that made up of particles are neutral not charged

Physical Properties of Covalent Compound 54 Copyright The McGraw-Hill Companies, Inc.Permission required for reproduction or display. Hybridization of Atomic Orbitals 55 Valence Bond Theory and NH3 N 1s22s22p3 3 H 1s1 If the bonds form from overlap of 3 2p orbitals on nitrogen with the 1s orbital on each hydrogen atom, what would the molecular geometry of NH3 be? If use the 3 2p orbitals predict 90o Actual H-N-H bond angle is 107.3o 56 Hybridization mixing of two or more atomic orbitals to form a new set of hybrid orbitals. 1. Mix at least 2 nonequivalent atomic orbitals (e.g. s and p).Hybrid orbitals have very different shape from original atomic orbitals. 2. Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process.3. Covalent bonds are formed by: a. Overlap of hybrid orbitals with atomic orbitals b. Overlap of hybrid orbitals with other hybrid orbitals 57 Formation of sp3 Hybrid Orbitals 58 Formation of Covalent Bonds in CH4 59 Predict correct bond angle sp3-Hybridized N Atom in NH3 60 Formation of sp Hybrid Orbitals 61 Formation of sp2 Hybrid Orbitals 62 # of Lone Pairs + # of Bonded Atoms HybridizationExamples 2 3 4 5 6 sp sp2 sp3 sp3d sp3d2 BeCl2 BF3 CH4, NH3, H2O PCl5 SF6 How do I predict the hybridization of the central atom? 1. Draw the Lewis structure of the molecule. 2. Count the number of lone pairs AND the number of atoms bonded to the central atom 63 64 sp2 Hybridization of Carbon 65 Unhybridized 2pz orbital (gray), which is perpendicular to the plane of the hybrid (green) orbitals. 66 Sigma bond (s) electron density between the 2 atoms Pi bond (p) electron density above and below plane of nucleiof the bonding atoms Bonding in Ethylene, C2H4 67 Another View of p Bonding in Ethylene, C2H4 68 sp Hybridization of Carbon 69 Bonding in Acetylene, C2H2 70 Another View of the Bonding in Ethylene, C2H4 71 Describe the bonding in CH2O. C H O H C 3 bonded atoms, 0 lone pairs C sp2 72 Sigma (s) and Pi Bonds (p) Single bond 1 sigma bond Double bond 1 sigma bond and 1 pi bond Triple bond1 sigma bond and 2 pi bonds How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? C H H CH O OH s bonds = 6+ 1 = 7 p bonds = 1 73 Molecular orbital theory bonds are formed from interaction of atomic orbitals to form molecular orbitals. O ONo unpaired e- Should be diamagnetic Experiments show O2 is paramagnetic 74 Energy levels of bonding and antibonding molecular orbitals in hydrogen (H2). A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed. 75 Constructive and Destructive Interference 76 Two Possible Interactions Between Two Equivalent p Orbitals 77 General molecular orbital energy level diagram for the second-period homonuclear diatomic molecules Li2, Be2, B2, C2, and N2.78 1. The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined. 2. The more stable the bonding MO, the less stable the corresponding antibonding MO. 3. The filling of MOs proceeds from low to high energies. 4. Each MO can accommodate up to two electrons. 5. Use Hunds rule when adding electrons to MOs of the same energy. 6. The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms. Molecular Orbital (MO) Configurations 79 bond order =1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) bond order 10 80 81 Delocalized molecular orbitals are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. Example: Benzene, C6H6 Delocalized p orbitals 82 Electron density above and below the plane of the benzene molecule. 83 Bonding in the Carbonate Ion, CO32- 84 Chemistry In Action: Buckyball Anyone? Intermolecular Forces Three States of Matter 1.5 Factors in Determining State of Matter I. Kinetic Energy II. Intermolecular Forces; Attractive Forces between different molecules neglible in gases important in solids, liquids GAS-Effect of Kinetic Energy Overwhelms Attractive Force Liquid Molecules have Enough Kinetic Energy to Slide Past One Another. Liquid Molecules have Enough Kinetic Energy to Slide Past One Another. Solid Molecules Kinetic Energy isNOT Strong Enough to Allow Molecules to Slide Past One Another Intermolecular Forces (Attractive Forces, van der Waal Forces) Part II Intermolecular Forces 11.2 Intermolecular forces; attractive forces between diff. Molecules which bring the molecules in contac with eac other Intramolecular forces hold atoms together in a molecule. Intermolecular vs Intramolecular 41 kJ to vaporize 1 mole of water (inter) 930 kJ to break all O-H bonds in 1 mole of water (intra) Generally, intermolecular forces are much weaker than intramolecular forces. Measure of intermolecular force boiling point melting point DHvap DHfus DHsub Types of Intermolecular Forces 1. Dipole/ Dipole Forces (Polar Molecules) Hydrogen Bonds 2. London/ Dispersion Forces (NonPolar Molecules) Intermolecular Forces I. Dipole-Dipole Forces Attractive forces between polar molecules Orientation of Polar Molecules in a Solid 11.2 Intermolecular Forces II.Dispersion Forces Attractive forces that arise as a result of temporary dipoles induced in atoms or nonpolar molecules 11.2 Polarizability Ease at which the electron distribution in an atom or molecule can be distorted and a temporary dipole induced More electrons (greater Molar Mass) leads to greater polarizability. Formation of Temporary Dipoles 1. Random movement of electrons 2. ion-induced dipole interaction 3. dipole-induced dipole interaction S What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces.There are also dispersion forces between HBr molecules. CH4 CH4 is nonpolar: dispersion forces. SO2 SO2 is a polar molecule: dipole-dipole forces.There are also dispersion forces between SO2 molecules. 11.2 Boiling Point- Temperature at which there is enough Kinetic Energy to Overcome Intermolecular Forces Liquid Has Intermolecular Forces Gas No Intermolecular Forces Boiling Point Increases with.. 1. Stronger Intermolecular Force 2. If same Intermolecular Force; increasing Molar Mass, higher boiling point Explain why the Higher Molar Mass Compound, CF4,has a Lower Boiling Point than H2Se CF4 Boiling Point; -150.0C Molar Mass ~ 88 g/mole H2Se Boiling Point; -42.0 C Molar Mass ~ 81 g/mole Intermolecular Force; Dispersion Force Intermolecular Force; Dipole-Dipole Force Boiling Points ofPolar Hydrogen Compounds Approximate Molar Mass; g/mole Boiling Point; C H2O18+ 100 H2S34-60 H2Se81-42 H2Te130-2 Hydrogen Bond Strong Type of Dipole-Dipole Force.This Type of Intermolecular Force Happens When H is directly bonded to O, N, or F. High Strength of H- bond 1. Large electronegativity difference between H and N, O, or F. 2. Small size of H atom allows it to get close to another molecule Hydrogen Bond O-H Covalent Bond that Makes H-bonding Possible Which of the Two Polar Molecules Has a Higher Boiling Point ? diethyl etherCHHH OCHHHCHHH CHHOHethanolImportance of H-Bonds in H2O 1. Very high boiling point for water (H2O(l)) for its Molar Mass. 2. The solid form of the material is less dense than liquid form Ice Floats on liquid water. Water expands as it freezes Ice Cubes float on water (Left) Solid benzene sinks to the bottom of liquid benzene (right) Importance of H-Bonds in H2O 1. Very high boiling point for water (H2O(l)) for its Molar Mass. 2. Ice floats on liquid water. The solid form of the material is less dense than liquid form Water expands as it freezes 3. High specific heat of Water Determining Type of Intermolecular Force Polar Molecules ? Dispersion/ London Forces H Directly Bonded to O, N, or F ? Dipole/Dipole Force Hydrogen Bonding Increasing Strength of Intermolecular Force NO YES NOYES