Where are we? Check-In

Transcription

1 Where are we? Check-In ü Building Blocks of Matter ü Moles, molecules, grams, gases, ü The Bohr Model solutions, and percent composition Coulomb s Law ü Empirical and Molecular formulas Photoelectron Spectroscopy ü The periodic table, Electron Electron configuration and PES Configuration, Aufbau Principle, Dalton, Thomson, Millikan, Pauli Exclusion Principle, and Rutherford, and Heisenberg Hund s Rule ü Periodic Trends atomic radius (cations & anions), ionization energy, and electronegativity ü Quantum Theory TEST TIME! Big Idea One Test à MONDAY, OCTOBER, 10 th & TUESDAY, OCTOBER 11 th

2 PHOTOELECTRIC EFFECT & PES October 3, 2016

3 Electron Configuration Is there any direct evidence that this diagram is accurately showing potential energy of electrons on the atom?

4 PES Process E photon = hv Atom Monochromatic Beam of X-Rays IE electron = E photon - KE KE = mv 2 e - 2

5 Photoelectron Spectroscopy (PES) Photoelectron spectroscopy (PES) is a technique used for determining the ionization potentials of molecules. Two types: 1. Ultraviolet photoelectron spectroscopy (UPS): focuses on ionization of valence electrons 2. X-ray photoelectron spectroscopy (XPS): ionizes core electrons and prys them away.

6 PES Process When a sample surface is irradiated with photons of energy hν, electrons are emitted from the sample surface.

7 PES Albert Einstein considered electromagnetic energy to be bundled into little packets called photons. Energy of photon = E = hv h = Planck constant ( x J s ) & v = frequency (Hz) of the radiation Photons of light hit surface electrons and transfer their energy E = hv = B.E. + K.E. The energized electrons overcome their attraction (to the nucleus) and escape from the surface (binding energy) Photoelectron spectroscopy detects the kinetic energy of the electron escaped from the surface.

8 Photoelectric Effect Refers to the phenomenon in which electrons are emitted from the surface of a metal when light strikes it: 1. Electrons are not emitted below a threshold frequency 2. Electron are not emitted below the threshold frequency regardless of the intensity of the light 3. For light above the threshold frequency, the # of electrons increases with light intensity 4. The kinetic energy of emitted electrons increases linearly with the frequency of light when above the threshold frequency

9 The Photoelectric Effect The photoelectric effect is interpreted with photons and the conservation of energy with the equation: KE electron = ½ mv 2

10 PES Spectrum

11 PES Spectrum photoelectron.html

12 PES Conclusions PES determines the energy needed to eject electrons from the material. Measurement of these energies provides a method to deduce the shell structure of an atom. The intensity of the photoelectron signal at a given energy is a measure of the number of electrons in that energy level.

13 PES AP Questions Which peaks in the photoelectron spectrum are representa2ve of the binding energy of p orbital electrons? a. C only c. C and E b. D only d. B, C and D

14 COULOMB S LAW

15 Photoelectric Effect & Coulomb s Law More energy is required to remove successive electrons from atoms. This is due to Coulomb s Law. Coulomb s Law quantifies the general rule of electrostatics that opposites charges attract and like charges repel. The electrostatic force between two charged bodies is proportional to the product of the amount of charge on the bodies divided by the square of the distance between them

16 Coulomb s Law F= force of attraction or repulsion k = constant q 1 and q 2 = the charges of the two bodies r = radius between the charges

17 Coulomb s Law If the bodies are oppositely charged, one positive and one negative, they are attracted toward one another; if the bodies are similarly charged, both positive or both negative, the force between them is repulsive.

18 Coulomb s Law

19 Coulomb's law helps describe the forces that bind electrons to an atomic nucleus. Based on Coulomb s Law, the force between two charged particles is proportional to the magnitude of each of the two charges and inversely proportional to the square of the distance (radius) between them.

20 SCIENTIST WHO SHAPED QUANTUM THEORY

21 What is Today s Model? Dense, Positively Charged Nucleus Composed of Protons, Neutrons, and Electrons Negatively Charged Electron Cloud Mostly Empty Space Most Probable Location of the Electrons

22 Timeline of Development of Current Atomic Model 450 BC Discovery of the Proton Discovery of the Neutron 1930 Democritus proposed the idea of atomos. Beginning of Modern Atomic Theory Discovery of the Electron Discovery of the Nucleus The Idea of Energy Levels for Electrons was Proposed. Introduction of the wave mechanical model

23 Summary for Dalton s Atomic Theory (Father of the Modern Atomic Theory) All atoms of a single element have the same mass Atoms of different elements are different. Atoms can t be divided, created or destroyed. Atoms of different elements combine in simple whole-number ratios to form compounds.

24 Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Crookes Tube Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

25 J.J. Thomson J.J. Thomson He proved that atoms of any element can be made to emit tiny negative particles. From this he concluded that ALL atoms must contain these negative particles. He knew that atoms did not have a net negative charge and so there must be something positive that balances the negative charge.

26 William Thomson s (Sir Kelvin) Atomic Model (1910) Thomson believed that the electrons were like plums embedded in a positively charged pudding, thus it was called the plum pudding model.

27 Ernest Rutherford s ( ) Where exactly are those electrons? Thomson s Theory: Plum Pudding q electrons embedded in a positive pudding. Rutherford s idea: q Shoot something at them to see where they are.

28 Rutherford s Conclusion (1911) Small, dense, positive nucleus. Equal amounts of (-) electrons at large distances outside the nucleus.

29 Neils Bohr s Atomic model (1913) Small, dense, positive nucleus. Equal amounts of (-) electrons at specific orbits around the nucleus.

30 ** James Chadwick discovered neutrons in n 0 have no charge and are hard to detect -- purpose of n 0 = stability of nucleus Chadwick And now we know of many other subatomic particles: photo from liquid H 2 bubble chamber Quarks, muons, positrons, neutrinos, pions, etc.

31 Quantum Mechanical Model -electron cloud model- -charge cloud model- Schroedinger, Pauli, Heisenberg, Dirac (up to 1940): According to the QMM, we never know for certain where the e are in an atom, but the equations of the QMM tell us the probability that we will find an electron at a certain distance from the nucleus.

32 Quantum Mechanical Model Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).

33 Modern Atomic Theory All matter is composed of atoms. Atoms of the same element are chemically alike with a characteristic average mass which is unique to that element. Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions! Atoms of any one element differ in properties from atoms of another element The exact path of electrons are unknown and e- s are found in the electron cloud.

34 Models of the Atom Dalton s model (1803) 1803 John Dalton pictures atoms as tiny, indestructible particles, with no internal structure. Thomson s plumpudding model (1897) 1897 J.J. Thomson, a British scientist, discovers the electron, leading to his "plum-pudding" model. He pictures electrons embedded in a sphere of positive electric charge. Rutherford s model (1909) 1911 New Zealander Ernest Rutherford states that an atom has a dense, positively charged nucleus. Electrons move randomly in the space around the nucleus. Bohr s model (1913) 1913 In Niels Bohr's model, the electrons move in spherical orbits at fixed distances from the nucleus. Charge-cloud model (present) 1926 Erwin Schrödinger develops mathematical equations to describe the motion of electrons in atoms. His work leads to the electron cloud model Hantaro Nagaoka, a Japanese physicist, suggests that an atom has a central nucleus. Electrons move in orbits like the rings around Saturn Frenchman Louis de Broglie proposes that moving particles like electrons have some properties of waves. Within a few years evidence is collected to support his idea James Chadwick, a British physicist, confirms the existence of neutrons, which have no charge. Atomic nuclei contain neutrons and positively charged protons.

35 Millikan Oil Drop Experiment The oil drop experiment was performed by Millikan and Fletcher in 1909 to measure the electric charge of the electron. The experiment entailed observing tiny charged droplets of oil between two metal electrodes.

36 MASS SPECTROMETRY

37 How does it work? Separates ionized atoms or molecules by changing magnetic field or voltage. Ionic equation for this: M 0 + 1e- M + + 2e- Mass Spectrometry (a.k.a. MS or mass spec) a method of separating and analysing ions by their mass-to-charge ratio

38 Four steps A) Ionization B C D B) Acceleration A C) Deflection D) Detection (A) Electron Trap not shown

39 Another view

40 Example Spectrum How many isotopes are there? How do you know? Relative heights of lines give % abundance. What is the average atomic mass?

41 Label all five lines:

42 Fragmentation: Define this term based on the peaks below.

43 Label all six peaks C A B D E F

44 Sample AP Questions

45 Sample MC Questions

46 MISCELLANEOUS

47

48 Visible Light

49 Visible Light & Wavelengths