Big Idea #1 Structure of the Atom. John Dalton (1766 – 1844) Proposed the first scientifically supported atomic theory.

Big Idea #1

Structure of the Atom

John Dalton (1766 – 1844)

• Proposed the first scientifically supported atomic theory.

Dalton’s Model of the Atom

• Dalton's model was that the atoms were tiny, indivisible, indestructible particles and that each one had a certain mass, size, and chemical behavior that was determined by what kind of element they were.

• What made us change Dalton’s Atomic Model?

J. J. Thomson (1856 – 1940)

• Credited for the discovery of the electron.

• Invented the mass spectrometer which led to his discovery of isotopes.

Scientific Inquiry

• Scientists use experimental results to test scientific models, such as the model of the atom.

• When experimental results are not consistent with the predictions of a scientific model, the model must be revised or replaced.

Lord Ernest Rutherford (1871 – 1937)

• Discovered the nucleus of the atom.

Rutherford’s Gold Foil Experiment

Rutherford’s Gold Foil Experiment

Rutherford’s Atomic ModelThe “Planetary Model” of the Atom

• The nucleus is very small, dense, and positively charged.

• Electrons surround the nucleus.• Most of the atom is empty space

Subatomic ParticlesPARTICLE SYMBOL CHARGE MASS

(amu)LOCATION

electron e- -1 0 orbit nucleus

proton p+ +1 1 inside nucleus

neutron n0 0 1 inside nucleus

Niels Bohr

• In 1913 Bohr published a theory about the structure of the atom based on an earlier theory of Rutherford's.

• Bohr expanded upon this theory by proposing that electrons travel only in certain successively larger orbits.

Bohr Model of the Atom

• Electrons orbit the nucleus in orbits that represent specific quantities of energy.

• The energies of the electrons in the atom are quantized.

• Only certain electron orbits (energy levels) are allowed.The Bohr Atom

Electromagnetic Waves

• Properties of waves include speed, frequency, wavelength and energy

• All electromagnetic waves including light travel at a speed of 3 x 108 m/s.

• However the frequency, wavelength and energy of the waves vary.

Wavelength ())

Measured in units of length: m, nm, Aº

Frequency ())Measured in cycles/second = hertz (Hz)

Visible Light

• For all waves For all waves • • = c = c

cc = the speed of light = 3.00 x 10 = the speed of light = 3.00 x 1088 m/s m/s

Electromagnetic Electromagnetic RadiationRadiation

Electromagnetic Electromagnetic RadiationRadiation

A photon of red light has a wavelength of 665 nm. What is the frequency of this light?

An x-ray has a frequency of 7.25 x 1020 Hz. What is the wavelength?

• For all waves: For all waves: EE == hh • • hh = 6.63 x 10 = 6.63 x 10-34-34 J • s J • s

This is known as Planck’s constant.This is known as Planck’s constant.

Energy of Electromagnetic Energy of Electromagnetic RadiationRadiation

Energy of Electromagnetic Energy of Electromagnetic RadiationRadiation

A photon of red light has a wavelength of 665 nm. What is the energy of this light?

An x-ray has a frequency of 7.25 x 1020 Hz. What is it’s energy?

Wavelength, frequency and energy

• Wavelength and frequency have an indirect relationship.

• Energy and frequency have a direct relationship.

• Electromagnetic radiation of short wavelength will have high frequency and high energy.

• Electromagnetic radiation of long wavelength will have low frequency and low energy.

Niels Bohr

• Bohr also described the way atoms emit radiation by suggesting that when an electron jumps from an outer orbit to an inner one, that it emits light.

Bohr Model• Electrons move around the

nucleus in orbits of definite energies.

• The energy of the orbit is related to its distance from the nucleus. The lowest energy is found in the orbit closest to the nucleus.

• Radiation is absorbed or emitted when an electron moves from one orbit to another.

Ground State

• The lowest energy state of an atom.

Excited State

• Any energy state of an atom that is of higher in energy than the ground state.

Energy Absorbed

Absorption (Dark – Line) Spectra

Energy EmittedElectron jumps to a lower orbit

What if we wanted to calculate the energy of electron transitions?

• E = - 2.178 x 10-18J (n2)

Where n is the principle energy level.

• So E = Efinal – Einitial

• Calculate the energy required to move an electron from n=1 to n=2. Calculate the wavelength of light that must be absorbed to make this transition.

Emission (Bright – Line) Spectra

Emission Spectra

The lines present in an emission spectrum are the lines missing in

an absorption spectrum.

The Heisenberg Uncertainty Principle

The Uncertainty principle• Heisenberg determined that it was

impossible to experimentally determine both the position and the speed of the electron at the same time.

• This became known as the Heisenberg Uncertainty Principle.

• It simply means that the electron is so small and moving so fast, that the simple act of trying to measure its speed or position would change either quantity.

Problems with the Bohr Model• It violates the

Heisenberg Uncertainty Principle because it considers electrons to have known orbits.

• It makes poor predictions regarding the spectra of atoms larger than hydrogen.

Schrodinger• The Austrian scientist,

Erwin Schrödinger, pursued the idea of the electron having wave properties and it seemed to him that the electron might be like a standing wave around the nucleus.

Schrodinger’s Model• This idea agreed very well with Bohr's idea

of quantized energy levels: only certain energies and therefore, wavelengths would be allowed in the atom.

• This explained why only certain colors (wavelengths) were seen in the spectrum of the hydrogen atom.

Schrodinger’s Model• Schrodinger set out to make a mathematical

model that assumed the electron was a standing wave around the nucleus.

• His solutions to that model agreed not only with the experimental evidence for hydrogen (as Bohr’s did too), but gave excellent results for all atoms when compared to their actual spectrum.

Schrodinger’s Model• This is our modern model of the atom and is

known as the Quantum Mechanical Model.

• Calculations involving the QM model of the atom can be approximated using computers.

• The solutions to these calculations is the basis for modern software that calculates the structure and reactivity of molecules.

The Quantum Mechanical Model

• The quantum mechanical model is based on quantum theory, which says matter also has properties associated with waves.

The Quantum Mechanical Model

• According to quantum theory, it’s impossible to know the exact position and momentum of an electron at the same time. This is known as the Uncertainty Principle.

The Quantum Mechanical Model

• The quantum mechanical model of the atom uses complex shapes of orbitals (sometimes called electron clouds), volumes of space where an electron is likely to be found.

The Quantum Mechanical Model

• Therefore the quantum mechanical model is based on probability rather than certainty.

Absorption and Emission of Energy by Molecules

• A photon is a particle representing a specific amount of electromagnetic energy.

• When a photon is absorbed by a molecule, the energy of the molecule is increased.

• When a photon is released by a molecule, the energy of the molecule is decreased.

Absorption and Emission of Energy by Molecules

• Different types of molecular motion lead to absorption or emission of photons in different spectral regions.

Principle Energy Levels – Sublevels - Orbitals

• Electrons in an atom are within atomic orbitals which are within sublevels which are within energy levels.

Principle Energy Level (VA) – 1-7

Sublevels (Manassas) – s,p,d,f

Orbitals (8909 Euclid Ave)

SUBLEVEL NUMBER OF ORBITALS

MAX. # OF ELECTRONS

s 1 2

p 3 6

d 5 10

f 7 14

© 1998 by Harcourt Brace & Company

sp

d (period-1)

f (period-2)67

Periodic Table and Electron Configuration

1

2

3

4

5

6

7

6

7

A. General Rules

• Pauli Exclusion Principle

–Each orbital can hold TWO electrons with

opposite spins.

Correct

A. General Rules

• Hund’s Rule

–Within a sublevel, place one e- per orbital before pairing them.

–All electrons in singly filled orbitals have the same direction of spin.

Incorrect

A. General Rules

• Aufbau Principle– Electrons always enter the lowest energy

location first– This is why every EC from He on starts 1s2

Magnetism

• Magnetism can be a complicated concept.

• We will only deal with magnetism as it is predicted by electron spin which is an oversimplified way of dealing with this subject.

Electron Spin & Magnetism

• Opposite spins produce opposite magnetic fields.

Electron Spin & Magnetism• Most materials with one or

more unpaired electrons are at least slightly magnetic. These substances are said to be paramagnetic.

• The overall magnetic field strength of atoms with all paired electrons is zero. These substances are said to be diamagnetic.

Photoelectric Effect

• The photoelectric effect is the emission of electrons from substances that are exposed to light.

• Ionization energy is the amount of energy it takes to remove an electron from an atom.

Photoelectric Effect

• More energy is required to remove successive electrons from atoms.

• This is due to Coulomb’s Law.

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 repulsionk = constantq1 and q2 = the charges of the two bodiesr = radius between the charges

Great Video integrating mulitple topics

• www.bozemannscience.com/ap-chem-004-coulombs.

• The link will be on the wiki page as well.

Photoelectric Effect• Light consists of photons of a certain

energy (EE == hh • • ).

• In the photoelectric effect, light ejects electrons from a material. This requires the photon to have sufficient energy to eject the electron.

• This is known as Photoelectron Spectroscopy (PES)

• PES data give support to the electron configurations that we use to place electrons within atoms.

Photoelectron Spectroscopy (PES)

• Photoelectron Spectroscopy (PES) determines the energy needed to eject electrons from a substance.

• Photoelectron Spectroscopy (PES) supports the shell, subshell model of the atom.

• PES data for multi-electron atoms show that certain electrons have ionization energies that are relatively close to one another. We therefore group these electrons into shells and subshells as a result.

Photoelectron Spectroscopy (PES)

• Photoelectron Spectroscopy (PES) data for atoms also helps us to confirm the number of electrons within particular shells and subshells.

• The intensity (shown as the height on a graph) of the photoemission signal at a given energy is a measure of the number of electrons in that shell or subshell. In other words the height of the PES signal is proportional to the number of electrons in that sublevel.

http://www.chem.arizona.edu/chemt/Flash/photoelectron.html

• Go to the website above and go through the first 10 elements:

• Show different sublevels and different numbers of electrons in each sublevel based on peak intensity (height).

• If you choose dual mode you can compare the PES spectra of two elements.