Where are we? Check-In ü Building Blocks of Matter ü Moles, molecules, grams, gases,

solutions, and percent composition ü Empirical and Molecular formulas ü The periodic table, Electron

Configuration, Aufbau Principle, Pauli Exclusion Principle, and Hund’s Rule

ü Periodic Trends – atomic radius (cations & anions), ionization energy, and electronegativity

ü Quantum Theory ü The Bohr Model •  Coulomb’s Law •  Photoelectron Spectroscopy •  Electron configuration and PES •  Dalton, Thomson, Millikan,

Rutherford, and Heisenberg TEST TIME! Big Idea One Test à

MONDAY, OCTOBER, 10th & TUESDAY, OCTOBER 11th

PHOTOELECTRIC EFFECT & PES October 3, 2016

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

Atom

e-

Ephoton = hv

KE = mv2

2 IEelectron = Ephoton - KE

Monochromatic Beam of X-Rays

PES Process

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.

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

PES •  Albert Einstein considered electromagnetic energy to be

bundled into little packets called photons. Energy of photon = E = hv

h = Planck constant ( 6.626 x 10-34 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.

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

The Photoelectric Effect • The photoelectric effect is interpreted with photons and the conservation of energy with the equation:

• KEelectron = ½ mv2

PES Spectrum

PES Spectrum •  http://cbc.arizona.edu/chemt/Flash/photoelectron.html

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.

PES AP Questions

Whichpeaksinthephotoelectronspectrumarerepresenta2veofthebindingenergyofporbitalelectrons?

a.  Conly c.CandEb.  Donly d.B,CandD

COULOMB’S LAW

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

Coulomb’s Law

F= force of attraction or repulsion k = constant q1 and q2 = the charges of the two bodies r = radius between the charges

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.

Coulomb’s Law

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.

SCIENTIST WHO SHAPED QUANTUM THEORY

What is Today’s Model? Dense, Positively Charged Nucleus

Mostly Empty Space

Negatively Charged Electron Cloud

Most Probable Location of the

Electrons

Composed of Protons, Neutrons, and Electrons

Timeline of Development of Current Atomic Model

1913 450 BC

Democritus proposed

the idea of atomos.

1802

Beginning of Modern

Atomic Theory

1897

Discovery of the

Electron

1911

Discovery of the

Nucleus

The Idea of Energy Levels for Electrons

was Proposed.

1930

Introduction of the wave mechanical

model

Discovery of the Proton

Discovery of the

Neutron

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.

Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.

Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Crookes Tube

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.

J.J. Thomson

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.

Ernest Rutherford’s (1871-1937)

q  electrons embedded in a positive pudding.

Where exactly are those electrons?

Thomson’s Theory: “Plum Pudding”

q  Shoot something at them to see where they are.

Rutherford’s idea:

Rutherford’s Conclusion (1911)…

•  Small, dense, positive nucleus.

•  Equal amounts of (-) electrons at large distances outside the nucleus.

Neils Bohr’s Atomic model (1913)

•  Small, dense, positive nucleus.

•  Equal amounts of (-) electrons at specific orbits around the nucleus.

And now we know of many other subatomic particles:

Chadwick

** James Chadwick discovered neutrons in 1932.

--

-- n0 have no charge and are hard to detect purpose of n0 = stability of nucleus

Quarks, muons, positrons, neutrinos, pions, etc.

photo from liquid H2 bubble chamber

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.

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).

Modern Atomic Theory

•  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!

•  All matter is composed of atoms.

•  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.

Models of the Atom

Dalton’s model (1803) Rutherford’s model

(1909) Bohr’s model (1913)

Charge-cloud model (present)

1800 1805 ..................... 1895 1900 1905 1910 1915 1920 1925 1930 1935 1940 1945

1803 John Dalton pictures atoms as tiny, indestructible particles, with no internal structure.

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.

1904 Hantaro Nagaoka, a Japanese physicist, suggests that an atom has a central nucleus. Electrons move in orbits like the rings around Saturn.

1911 New Zealander Ernest Rutherford states that an atom has a dense, positively charged nucleus. Electrons move randomly in the space around the nucleus.

1913 In Niels Bohr's model, the electrons move in spherical orbits at fixed distances from the nucleus.

1924 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.

1926 Erwin Schrödinger develops mathematical equations to describe the motion of electrons in atoms. His work leads to the electron cloud model.

1932 James Chadwick, a British physicist, confirms the existence of neutrons, which have no charge. Atomic nuclei contain neutrons and positively charged protons.

Thomson’s plum-pudding model (1897)

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.

MASS SPECTROMETRY

•  Separates ionized atoms or molecules by changing magnetic field or voltage.

•  Ionic equation for this: M0 + 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

How does it work?

A)  Ionization B)  Acceleration C)  Deflection D)  Detection

Four steps

A

B C

D

(A) Electron Trap not shown

Another view

• How many isotopes are there?

• How do you know? • Relative heights of lines give % abundance.

• What is the average atomic mass?

Example Spectrum

Label all five lines:

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

Label all six peaks

A

B

C

D

E F

Sample AP Questions

Sample MC Questions

MISCELLANEOUS

Visible Light

Visible Light & Wavelengths