Chapter 2: Atomic Structure And Interatomic Bonding

Chapter 2: Atomic Structure andInteratomic Bonding Atomic StructureElectron ConfigurationPeriodic TablePrimary Bonding– Ionic– Covalent– Metallic Secondary Bonding or van der Waals Bonding– Three types of Dipole Bonding MoleculesChapter 2 -

Why Study Atomic Structure andInteratomic Bonding?Two Allotropes of CARBONGraphite- Relatively soft- Greasy feel to it- Reasonably goodconductor of electricityDiamond- The hardest known material- Poor conductor of electricityThe disparities in properties are attributed to a type ofinteratomic bonding found in graphite that does not exist inChapter 2 diamond

Atomic Structure (Freshman Chem.) atom –electrons – 9.11 x 10-31 kgprotons-27 kg1.67x10neutrons} atomic number # of protons in nucleus of atom # of electrons of neutral species A [ ] atomic mass unit amu 1/12 mass of 12CAtomic wt wt of 6.023 x 1023 molecules or atoms1 amu/atom 1g/molCH12.0111.008 etc.Chapter 2 - 3

Atomic ModelsChapter 2 -

Timeline of Atomic TheoryDemocritus’ Aristotle and PlatoatomosDaltonJ.J. Thomson460 BC360 BC180818971909 1913 1923E. RutherfordN. BohrWave MechanicalChapter 2 5

400 BC - Democritus Ancient Greek philosopher Democritus coined the term átomos whichmeans "uncuttable" or "the smallest indivisibleparticle of matter".Structure of MatterPhysical world“VOID BEING”Chapter 2 -

Democritus Democritus postulated that atomswere completely solid, hard and smallparticles with no internal structure andhas an infinite variety of shapes andsizes. This theory was ignored for more than2000 years!Chapter 2 - 7

The Atomic Theory of MatterAristotle and Plato There can be no ultimatelyindivisible particles. Believed that fire, earth, air andwater were the four mainelements that world was madeup of.Chapter 2 - 8

1803 – John Dalton English instructor and naturalphilosopher “Each element consists ofatoms of single unique typeand can join to form chemicalcompounds.” Beginning the modern atomictheoryChapter 2 -

The Atomic Theory of MatterDalton’s Atomic Theory Postulates1. Elements are composed ofextremely small particles calledatoms.2. All atoms of a given element areidentical, having the same size,mass and chemical properties. Theatoms of different elements aredifferent. Different elements havedifferent atomic properties such asatomic mass.Chapter 2 - 10

Assist. Prof. Dr. İlkay KALAYThe Atomic Theory of MatterDalton’s Atomic Theory3. Atoms of an element are neither created nor destroyedby any chemical reactions. Chemical reactions onlyinvolve the combination, rearrangement or separation ofatoms.4. Compounds are formed from thecombination of atoms of more than oneelement. A compound always has thesame relative number and kind ofatoms.Chapter 2 - 11

1869 - Mendeleev Building upon earlierdiscoveries by scientists,Mendeleev published the firstfunctional PERIODIC TABLE. Certain chemical properties ofelements repeat periodicallywhen arranged by atomicnumber. The periodic table was firstdeveloped by Mendeleev andMeyer on the bases ofsimilarity in chemical andphysical properties exhibitedon certain elements.Chapter 2 -

Periodic TableDraft of the first periodic table, Mendeleev, 1869Chapter 2 -

1869 Chapter 2 -

Today: Periodic Table of theElementsChapter 2 -

1897 – Sir J. J. Thomson Although Dalton had postulated that atomswere indivisible, the studies have shown amore complex structure for an atom. J. J. Thomson conducted a series ofexperiments which showed that the atomswere not indivisible Discovered the electron (1906 Nobel Prizein Physics).Chapter 2 -

Atom is composed of even smaller particles buthow the particles fit together?J. J. Thomson proposed the atom consisted of a uniform positivesphere of matter in which the electrons were embedded.Plum Pudding (1904): “Theatom as being made up ofelectrons swarming in a sea ofpositive charge.J. J. Thomson's "plum-pudding"model of the atomChapter 2 - 17

1909 – E. Rutherford Tested and disproved the PlumPudding Model.Rutherford's experiment on the scatteringof α particles by metal foil Results:– Majority of a particles transmitted (pass through) ordeflected through small angles– Tiny fraction deflected through large anglesChapter 2 -

1909 – E. Rutherford Conclusion:– Disproved the Plum-Pudding Model– Large amount of the atom's charge and mass isconcentrated into a small region– Atom was mostly empty space Objections to Rutherford model– The laws of classical mechanics predict that theelectron will release electromagnetic radiation whileorbiting a nucleus. Because the electron would loseenergy, it would gradually spiral inwards, collapsinginto the nucleus.– This atom model is unsuccessful, because it predictsthat all atoms are unstable.Chapter 2 -

The Modern View of Atomic StructureThe structure of an (amu)ElectronNegative (-1)5.486 x10-4ProtonPositive ( 1)1.0073NeutronNeutral (0)1.0087ParticleCharge (C)Electron-1.6022 x 10-19Proton1.6022 x 10-19Neutron0Chapter 2 - 20

Assist. Prof. Dr. İlkay KALAYThe Modern View of Atomic StructureThe structure of an atomThe atoms are small and the atomic dimensions are expressedin terms of Angstrom (Å) unit.1 Å 10-10 mNucleus 10-4 Å1-5 ÅSchematic view of anChapter 2 - 21

Assist. Prof. Dr. İlkay KALAYThe Modern View of Atomic StructureThe structure of an atomü All atoms of an element have the same number of protons in thenucleus.ü The number of protons in the nucleus of an atom is calledatomic number (Z).ü Because the atoms have no net electrical charge, # of protons # of electrons.ü The total number of protons and neutrons in a nucleus is calledmass number (A).Mass number AAtomic number Z126CChapter 2 - 22

Assist. Prof. Dr. İlkay KALAYThe Modern View of Atomic Structureü All nuclei of the atoms of a particular element have the sameatomic number, but the the number of neutrons and so massnumbers may be different. Atoms of a given element with adifferent mass numbers are called isotopes.ü Most elements occur in nature as mixtures of isotopes.Symbol116126136146CCCC# ofelectrons# of protons# ofneutrons665666667668An atom of a specific isotope is called anuclide.Chapter 2 - 23

The Periodic TableThe periodic table, developed in 1869, shows the arrangement of elementswith similar chemical and physical properties in order of increasing atomicnumber.Periodic table presentsatomic number, atomicsymbol and atomicweight for eachelement.40Atomic number91.224Atomic symbolAtomic weightZrChapter 2 - 24

1912 – N. Bohr Many phenomena involvingelectrons in solids could not beexplained in terms ofCLASSICAL MECHANICS. We need QUANTUM MECHANICS Chapter 2 -

LineSpectraandTheBohrModelBohr’s ModelBohr proposed a model of the hydrogen atom that explains its line spectrum.Bohr’s postulates Rutherford atom is correct Classical EM theory not applicable to orbiting e Newtonian mechanics applicable to orbiting e Eelectron Ekinetic Epotential Only orbits of certain radii, corresponding to certain definite energies arepermitted for electrons in an atom. An electron in a permitted orbit has a specific energy and is in allowedenergy state. An electron in an allowed energy state will not radiateenergy and therefore will not spiral into the nucleus. During the transition of an electron from one allowed energy state toanother, energy is only emitted or absorbed by an electron. This energy isemitted or absorbed as a photon, ΔE Ef-Ei hν hc/λ where c νλChapter 2 26

LineSpectraandTheBohrModelEnergy States of the Hydrogen - ‐ llowedenergystatetoanother.- ‐ Thestatesshownarethoseforwhichn 1throughn 6,andthestateforn ,forwhichtheenergy,E,equalszero.Chapter 2 27

Bohr’sModelThe lower the value of n, the smallerthe radius of the orbit, and the lowerthe energy level.n 1 ground state (lowest energy state)(orbit closest to the nucleus)n 2, 3, or higher excited ydrogenatomfromtheBohrmodel.Chapter 2 28

BOHR ATOMorbital electrons:n principalquantum number12n 3Adapted from Fig. 2.1,Callister 6e.Nucleus: Z # protonsN # neutronsAtomic mass A Z NChapter 2 - 2

1913 - Sommerfeld German theoretical physicist Modified the Bohr Model “suppose we have plurality of orbits” – a shellcontaining multiple orbits: ORBITALS How to capture these new ideas quantitatively? We need new quantum numbers: n, l, m, sn principal quantum number, distance of anelectron from the nucleusl subshell, describes the shape of the subshellm number of energy states in a subshellChapter 2 s spin moment

Wave mechanics to arrive at same place:E E(n,l,m,s) The Bohr model – significant limitations Resolution: Wave-mechanical model(electron is considered to exhibit both wave-like andparticle-like characteristics).– De Broglie: “If a photon which has no mass, canbehave as a particle, does an electron which has masscan behave as a wave (1920)?” λ h/p h/mv– Heisenberg: Uncertainty Principle“I donʼt know where any of one of electrons is, but Ican tell you an average where any of one of them islikely to be”– SchrodingerChapter 2 -

Beyond Bohrʼs ModelDeBrogliewavelength:h! letoknowsimultaneouslytheexactposi7onandh!x.!m! "4" itsspeed(largerΔu)andviceversa.Wavemo aConincorporatesboththewave- ‐likeandpar.cle- ‐likebehavioroftheelectron.Wave function, ψ,Ψ2 provides information about an electron’s location when it s in anChapter 2 allowed energy state

QuantumMechanicsandAtomicOrbitalsElectron density: probability of the electronbeing at a pointHigher density of dots region: largervalues of Ψ2Electron- Chapter 2 -

Bohr Model vs. Wave Mechanical Atom ModelWith WM model electron nolonger treated as a particlemoving in a discrete orbital,rather position is described bya probability distributionBohrWMChapter 2 -

QuantumNumbersAccording to quantum mechanics, each electron in an atom is describedby four different quantum numbers, three of which (n, l, ml) specify thewave function that gives the probability of finding the electron at variouspoints in space.n . The principal quantum number (K, L, M, N, O and so on that correspond to 1, 2,3,4, 5 As n increases:- The orbital becomes larger- The electron spends more time farther from the nucleus- The electron has a higher energy and is therefore less tightly bound to the nucleus.l . The second quantum number l can have integral values from 0 to n-1 for each value of n defines the shape of the orbital The value of l of a particular orbital is designated by s, p, d and f, corresponding lvalues of 0, 1, 2 and 3 respectivelys: sharp p:principal d:diffuse f: fundamentalml . The magnetic quantum number Have integral values between l and –l including zero. Describes the orientation of the orbital in spaceChapter 2 -

What is the filling sequence of electrons inorbitals by n, l, m, s is not adequate?AUFBAU PRINCIPLE3 principles:1. Pauli Exclusion Principle:only one electroncan have a given set of four quantum numbers.2. Electrons-have discrete energy states-fill orbitals from lowesten. to highest en.3. Hundʼs ruleChapter 2 -

ElectronConfigura onsHund’s rule states that the lowest energy is attained by maximizing thenumber of electrons with the same electron spin.For example, for a carbon atom to achieve its lowest energy, theChaptertwo 2p237electrons will have the same spin.

Quantum Numberslmlms ½Chapter 2 -

Quantum NumbersElectrons fill quantum levels in order of increasingenergy ( only n and l make significant differencesin energy configurations).1s, 2s, 2p, 3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d, .When all electrons are at the lowest possibleenergy levels ground stateExcited states do exist such as in glow dischargesetc Valence electrons occupy the outermost filled shell.Valence electrons are responsible for all bonding !Chapter 2 -