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ATOMIC STRUCTURE ? Kenneth E. Schnobrich. A Brief History About 460 B.C. - a Greek Philosopher, Democritus, developed the idea of atoms (atomos) as small.

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Presentation on theme: "ATOMIC STRUCTURE ? Kenneth E. Schnobrich. A Brief History About 460 B.C. - a Greek Philosopher, Democritus, developed the idea of atoms (atomos) as small."— Presentation transcript:

1 ATOMIC STRUCTURE ? Kenneth E. Schnobrich

2 A Brief History About 460 B.C. - a Greek Philosopher, Democritus, developed the idea of atoms (atomos) as small indivisible particlesAbout 460 B.C. - a Greek Philosopher, Democritus, developed the idea of atoms (atomos) as small indivisible particles About 400 B.C. - a number of Greek philosophers said matter consisted of FIRE, EARTH, WATER, and AIR.About 400 B.C. - a number of Greek philosophers said matter consisted of FIRE, EARTH, WATER, and AIR.

3 A Brief History NEXT 2000 YEARS - Alchemy (a pseudoscience) dominated - they were concerned with turning base metals into gold. During this time Hg, S, and Sb were discovered. Alchemists also discovered how to make mineral acids.NEXT 2000 YEARS - Alchemy (a pseudoscience) dominated - they were concerned with turning base metals into gold. During this time Hg, S, and Sb were discovered. Alchemists also discovered how to make mineral acids Joseph Proust showed that a given compound always had the same proportions by mass. Law of Definite Proportions Joseph Proust showed that a given compound always had the same proportions by mass. Law of Definite Proportions.

4 A Brief History John Dalton (following the work of Robert Boyle) discovered that atoms can combine in more than one way. He proposed the Law of Multiple Proportions. He theorized that the basic unit was the atom John Dalton (following the work of Robert Boyle) discovered that atoms can combine in more than one way. He proposed the Law of Multiple Proportions. He theorized that the basic unit was the atom Michael Faraday worked on the electrolysis of molten salts and coined the word ion (Greek meaning wanderers) Michael Faraday worked on the electrolysis of molten salts and coined the word ion (Greek meaning wanderers).

5 A Brief History John Dalton published “A New System of Chemical Philosophy” which proposed his theory of atoms John Dalton published “A New System of Chemical Philosophy” which proposed his theory of atoms –All elements are composed of tiny, discrete, indivisible and indestructible particles called “atoms”. –All atoms of a given element are identical –Atoms of of different elements are different; they have different masses and properties. –Chemical combinations of these “atoms” compose all matter - different atoms combine differently to form compounds.

6 DALTON’S ATOM JOHN DALTON - he envisioned the atom as a hard spherical unit of matter (the ultimate unit) Lithium Chlorine Sulfur LiCl Li 2 S

7 DALTON’S ATOM JOHN DALTON - he envisioned the atom as a hard spherical unit of matter Oxygen Hydrogen H2OH2O H2O2H2O2

8 HISTORY (cont.) Eugen Goldstein discovered the existence of a charged stream from the cathode using a “Crookes Tube” and called them “Cathode Rays” Eugen Goldstein discovered the existence of a charged stream from the cathode using a “Crookes Tube” and called them “Cathode Rays”

9 HISTORY (cont.) Eugen Goldstein discovered the existence of positively charged particles he called “Canal Rays” Eugen Goldstein discovered the existence of positively charged particles he called “Canal Rays”

10 HISTORY (cont.) J.J. Thomson using a modified “Crookes Tube” determined that the “Cathode Rays” behaved like charged particles and measured the charge J.J. Thomson using a modified “Crookes Tube” determined that the “Cathode Rays” behaved like charged particles and measured the charge.

11 HISTORY (cont.) J.J. Thomson proposed his “Raisin Pudding Model” of the atom J.J. Thomson proposed his “Raisin Pudding Model” of the atom Negatively charged Electrons Positive Matrix *Atoms are neutral

12 HISTORY (cont.) Ernest Rutherford suggested the atom was “nuclear” based on a famous experiment - “The Scattering Experiment”. He also suggested that the proton was the fundamental unit of positive charge

13 HISTORY (cont.) Alpha Particles Metal Foil(Au)

14 ASSUMPTIONS RUTHERFORD’S WORK Most of the atoms mass is concentrated in the nucleus.Most of the atoms mass is concentrated in the nucleus. All of the positive charge is concentrated in the nucleusAll of the positive charge is concentrated in the nucleus Neutral atoms have equal numbers of protons and electrons.Neutral atoms have equal numbers of protons and electrons. The protons and neutrons are located in the nucleus of the atom.The protons and neutrons are located in the nucleus of the atom.

15 HISTORY (cont.) THOMSON MODEL RUTHERFORD MODEL

16 HISTORY (cont.) Chadwick discovered and determined the properties of the neutron Chadwick discovered and determined the properties of the neutron. Proton Neutron Nucleus Electrons

17 SUBATOMIC PARTICLES PARTICLECHARGEMASSLOCATIONSYMBOL PROTON+1 1 AMU NUCLEUS 1 H 1 or 1 p 1 NEUTRON0 1 AMU NUCLEUS 0n10n10n10n1 ELECTRON 1/1836 AMU OUTSIDE -1 e 0

18 LOOKING AT THE ATOMS STRUCTURE Atomic Number = # protons and electrons in a neutral atomAtomic Number = # protons and electrons in a neutral atom Atomic Mass Number* = sum of the protons and neutronsAtomic Mass Number* = sum of the protons and neutrons #Neutrons = Mass# - Atomic##Neutrons = Mass# - Atomic# *ATOMIC MASS MAY VARY (ISOTOPES)

19 ISOTOPES AVERAGE ATOMIC MASS Most of the elements on the periodic table have several Isotopes. The Mass that you see is the weighted average of known isotopes. Example: Carbon has two stable isotopes C 12 = 98.89% and C 13 = 1.108% 12(0.9889) + 13( ) = Average Atomic Mass

20 FORMING AN ION METALS - usually like to lose electrons to form positive ions called CATIONS.METALS - usually like to lose electrons to form positive ions called CATIONS. NONMETALS - usually like to gain electrons to form negative ions called ANIONS.NONMETALS - usually like to gain electrons to form negative ions called ANIONS.

21 FORMING IONS Na Na +1 + electron ClCl -1 + electron

22 Mn Electron Arrangement Atomic Number Atomic MassOxidation States

23 THE KERNEL AND VALENCE ELECTRONS 19 K Valence Electrons – those in the outer principal energy level Kernel – the nucleus and all of the electrons except those in the valence level

24 THE KERNEL AND VALENCE ELECTRONS 19 K Valence Electrons = 1 Kernel – has a charge of +1 *Now Lewis Dot Structures

25 LEWIS DOT STRUCTURES 19 K K 16 S S x x [ ] -2

26 PUTTING IT TOGETHER PARTICLEPROTONSELECTRONSNEUTRONS N N -3 Sn Sc +3 Na

27 CONTINUOUS SPECTRUM VISIBLE REGION OF THE SPECTRUM

28 HYDROGEN AND HELIUM LINE SPECTRUM HYDROGEN HELIUM More spectra

29 BRIGHT-LINE SPECTRA BRIGHT-LINE SPECTRA ARE LIKE FINGER PRINTS. EACH ELEMENT HAS ITS OWN CHARACTERISTIC SET OF BRIGHT LINES IN THE VISIBLE REGION OF THE SPECTRUM. Carbon Helium Hydrogen

30 THE BOHR MODEL Bohr’s Model was based on the simplest atom, Hydrogen. Bohr based his model on the following: (1)Electrons do not follow the rules of large macroscopic bodies. (2)Electrons in atoms have only specific energies. (3)Electrons are only in specific orbits outside the nucleus (ground state). (4)When an electron moves from one orbit to another it absorbs or releases energy of a specific frequency. (5)When electrons absorb energy they move to an excited state (higher energy orbit).

31 THE BOHR MODEL Hydrogen Excited State Ground State Energy ReleasedEnergy Absorbed

32 THE QUANTUM MODEL As the science of spectroscopy grew and the resolution of the bright-line spectra of an element improved and the dual nature of the electron was explored scientists formulated a new picture of the atom. This new model of the atom retains some of the original features but changes the concept of electron location. The electron, instead of occupying a specific orbit now is thought to occupy a region of 3-D space called the orbital.

33 THE QUANTUM MODEL Dual Nature of the Electron – the electron to this point, had been described as being “particle-like” in nature, but it also exhibits “wave-like behavior.” DeBroglie – was the first to suggest that, based on its extremely small size, the electron does have a measureable wavelength. Double-click on the You Tube video

34 THE QUANTUM MODEL After viewing the video we see that Erwin Schrodinger allows us to describe the electrons in an atom with a set of 4 Quantum Numbers. The quantum numbers help us to describe the relative energies and probable locations of the electrons. The Principal Quantum Number (n) – corresponds very closely with the energy levels described in the Bohr Model. The PQN can only have small whole number values (n = 1, 2, 3, 4, 5, 6 etc). The greater the value of “n” the greater the energy and distance from the nucleus for the electron.

35 THE QUANTUM MODEL The Sublevel Quantum Number ( l ) – describes the sublevels the electrons can occupy within a Principal Energy Level. The SQN – has values that are determined by the value of the PQN. It can have values from 0 … n-1 So, if n = 0, l = 0 If n = 2, l = 0, 1 (which means, in the second Principal Energy Level, there are two available sublevels the electron can occupy. There are also corresponding letter values for the sublevels 0(s); 1(p); 2(d); 3(f)

36 THE QUANTUM MODEL In the 2 nd PEL there were two sublevels, 0, 1 or s and p. Within a PEL, as the value of l increases the energy and distance from the nucleus increases. In the 3 rd PEL, there are three sublevels, 0, 1, and 2 or s, p, and d sublevels. The Orbital Quantum Number (m) (also sometimes called the Magnetic Quantum number) - describes the number of orbitals (3-D orientations in space) within a sublevel.

37 THE QUANTUM MODEL The OQN’s are determined by the values for l m can have values from 0… +/- l So, if l = 0, m = 0, which means that there is only one possible 3-D description (or orbital) in that sublevel. If l = 1, m = 0, +1, -1, which means in the “p” sublevel there are three, 3-D descriptions (or orbitals), in that sublevel. Along the X axis Along the Y axis Along the Z axis

38 THE QUANTUM MODEL If an orbital is located in an s sublevel it is referred to as an s-orbital and has a spherical distribution along the X, Y, and Z axes. If an orbital is located in an p sublevel it is referred to as an p-orbital and has a “dumbell” distribution”along the X, Y, and Z axes.

39 THE QUANTUM MODEL Of course there are other orbital shapes but they are complicated and for our purposes, our concerns will be limited to the s and p orbital shapes. The fourth quantum number is the Spin Quantum Number – based on the Stern/Gerlach experiment it is thought that an electron can have one of two possible spins, +1/2 and -1/2 (it spins on its axis). Since no two electrons can have exactly the same set of four quantum numbers, only two electrons can occupy an orbital, provided they have opposite spins.

40 THE QUANTUM MODEL Based on the work of many scientists, including deBroglie, Shrodinger, and Heisenberg, we now know that – we can only speak in terms of the probable location of the electrons the bright line spectra available for the elements gives us additional information on the energy associated with the electrons

41 THE QUANTUM MODEL OF HYDROGEN Note: the electron is pictured as a cloud or region of space where you will most probably find the electron. Nucleus

42 Quantum Atom Relationships Increasing Energy

43 Electron Filling When we fill the energy sublevels that are several rules we must follow – The Aufbau Principle – you must always fill from lowest energy to highest energy Hund’s Rule – you must completely half-fill an energy sublevel before you start pairing electrons Pauli Exclusion Principle – no two electrons can have the same set of four quantum numbers in a given orbital, they must have opposite spins to exist in the same orbital.

44 Filling the Sublevels & Orbitals When filling the sublevels and orbitals remember the rules It is also important to remember that for multi-electron atoms some of the sublevels do overlap from an energy standpoint. there is a simplified relationship to help us with this overlap – Sublevel Energy = n + l (n is the PQN and l is the SQN ). It is why the 4s sublevel fills before the 3d sublevel (see the diagram on the next slide). 4s = = 4 and 3d = = 5 the sublevel energy of 4s is lower than that of 3d, therefore, the 4s sublevel fills before the 3d sublevel.

45 General Sublevel Arrangement Energy Distance from Nucleus 1s 2s 2p 3s 3p 4s 3d

46 General Sublevel Arrangement 1s 2s 2p 3s 3p 3d 4s4p4d4f 5s 5p 5d 5f * The idea of sublevel overlap can be much more complicated for larger atoms

47 General Sublevel Arrangement Energy Distance from Nucleus 1s 2s 2p 3s 3p 4s 3d For 19 K 39 the sublevel filling would look like this

48 General Sublevel Arrangement Energy Distance from Nucleus 1s 2s 2p 3s 3p 4s 3d For 7 N 14 the sublevel filling would look like this Note: the sublevel is half-filled, the electrons have parallel spins (the same)

49 Electron Arrangement Let’s take a sample and show you how the electron arrangement can be written in three formats. 19 K K 39 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 19 K 39 [Ar]4s 1 4s 1 # of electrons Energy sublevel Principal Energy Level


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