Name: ______________________Chemistry 1 Notes,

Topic 2 – Atomic TheoryTopics: 1.1.1 – Analyze the structure of atoms, isotopes and ions.1.1.2 – Analyze an atom in terms of the location of electrons.1.1.3 – Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model.

I. Historical Development of Atomic Theory

EARLY CONTRIBUTIONS: DEMOCRITUS & DALTON

♦ By 400 BC, a Greek philosopher named Democritus had postulated that there must be some basic unit of matter that could not be divided any further. He called this basic unit an atomon (from the Greek, meaning indivisible units). From this term we derive the English word atom.

There was disagreement over this idea for the next 2200 years, until the invention of the chemical balance – the tool needed to study composition of pure substances quantitatively.

Watershed event – 1803 – John Dalton postulated his atomic theory:

1) All matter is composed of extremely small particles called atoms.2) Atoms of a given element are identical in size, mass, and other properties;

atoms of different elements differ in size, mass, and other properties.3) Atoms cannot be subdivided, created, or destroyed.4) Atoms of different elements can combine in simple, whole-numbered ratios

to form chemical compounds.5) In chemical reactions, atoms are combined, separated, or rearranged.

We now know there are exceptions to Dalton’s theory: Atoms can be subdivided (fission, radioactive decay), created (fusion, etc.) and destroyed (all of

these are complicated process, sometimes requiring much energy). Several different types of atoms for a given element can exist, all with different masses. These

are called isotopes).

Still, Dalton’s theory = cornerstone in thinking about chemistry.

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Democritus

John Dalton

Name: ______________________Chemistry 1 Notes,

EXPERIMENTS – THE ELECTRON

♦ CRT (Cathode Ray Tube) Experiments – Sir Joseph John Thomson

In 1897, J.J. Thomson built a cathode ray tube, which generated a steady stream of electrons. These electrons traveled down the glass tube, causing a fluorescent glow which could be easily seen and measured. Thomson concluded several things from his experiments: Thomson placed two oppositely charged plates on either side of the CRT

(by hooking the plates up to a strong battery). When the electron beam passed between the two plates, it was deflected toward the positive plate. Thus, Thomson reasoned that electrons must be negative.

Thomson also placed north and south magnetic poles on opposite sides of the CRT. When the electron beam (travelling perpendicular to the poles) passed between the magnetic poles, it was deflected in a particular direction. (Positively charged particles are deflected in the opposite direction.)

When the oppositely charged plates and magnetic poles were used simultaneously, Thomson could cause the electron beam to continue travelling in a straight line.

Thomson could not quantify the exact mass or charge (only that it was negative) of an electron, but he could measure how much it was deflected in a magnetic field. Thus, Thomson was able to calculate the mass to charge ratio (m/e) for an electron.

Big picture: The 2 most important things Thomson discovered in the CRT Experiments: The electron is negatively charged. The mass to charge ratio (m/e) for an electron.

♦ An animation of Thomson’s experiments: http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/animations_center.html#

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J.J. Thomson

Name: ______________________Chemistry 1 Notes,

♦ Oil-Drop Experiment – Robert Millikan

This led to the determination of actual mass and charge of the electron.

Between 1908 and 1917, Robert Millikan measured the charge on an electron with the apparatus shown above. In these experiments, the atomizer from a perfume bottle was used to spray oil droplets into a sample chamber. Some of these droplets fell through a pinhole between two plates of an electric field, where they could be observed through a microscope.

A source of X-rays was then used to ionize the air in the chamber by removing electrons from the molecules in the air. Droplets that did not capture one of these electrons fell to the bottom of the chamber due to the force of gravity. Droplets that captured one or more electrons were attracted to the positive plate at the top of the viewing chamber and either fell more slowly or rose toward the top.

By carefully studying individual droplets, Millikan was able to show that the charge on a drop was always an integral multiple of a small, but finite value. When his data are converted to SI units, the charge on a drop is always some multiple of 1.59 x 10–19 C. Combining this value for the charge on a single electron with the mass to charge ratio (m/e) for the electron confirms Thomson's hypothesis. The mass of an electron is at least 1000 times smaller than the lightest atom.

Good animations of Millikan’s experiment: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/004_MILLIKANOIL.MOV

http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/animations_center.html#

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Robert Millikan

Name: ______________________Chemistry 1 Notes,

EXPERIMENTS – THE NUCLEUS (PROTONS & NEUTRONS)♦ Gold Foil Experiment – Ernest Rutherford

Rutherford’s experiments led to discovery of the atomic nucleus, and that it is small, dense and positively charged. Fast-moving positively-charged particles (-particles) were shot through

a thin layer of gold foil, only a few atoms thick. Most passed straight through the gold foil in a straight line. Some were

slightly deflected from their straight-line course. But, roughly 1 in 8000 ricocheted back toward the source.

Rutherford recognized that these -particles must be repelled by a very dense object in the middle of the atom, which takes up a very small amount of space. Thus, atoms must be mostly empty space. The nucleus must also have a positive charge (because like charges repel).

Rutherford concluded that the atomic nucleus was the positively charged, dense central portion of the atom that contains nearly all of its mass, but takes up only an insignificant fraction of its volume

Another very good animation of Rutherford’s experiment:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/006_RUTHERFORD.MOV

A good animation of Rutherford’s gold foil experiment: http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/animations_center.html#

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Ernest Rutherford

Name: ______________________Chemistry 1 Notes,

♦ Beryllium–Wax Experiments – James Chadwick

Following Rutherford’s experiments, scientists knew that there was “extra” mass somewhere in the atom (they knew electrons have essentially no mass, and the mass of an atom was greater than the apparent number of protons in the atom). But scientists couldn’t precisely characterize the source of this mass. In the early 1930s, James Chadwick smashed alpha particles into beryllium, causing it to release a different type of radiation (neutrons), which hit another target: paraffin wax. When the beryllium radiation hit hydrogen atoms (protons) in the wax, the atoms (protons) were dislodged and sent into a detecting chamber. In physics, it is known that only a particle having almost the same mass as a hydrogen atom could affect hydrogen in that manner. Chadwick’s experiment showed that a collision with beryllium atoms would release these neutral particles, which Chadwick named neutrons. This provided the answer for hidden mass in atoms. Chadwick published his findings in the journal Nature in 1932.

Chadwick proposed the following reaction, where 01 n represents the "new" particle, the neutron:

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James Chadwick

This is the detector used by James Chadwick.

Name: ______________________Chemistry 1 Notes,

II. Atomic Arrangement Within the nucleus are protons and neutrons:

Protons (p+) have a positive charge equal to the negative charge of an electron. Protons are about 2000x heavier than electrons.

Neutrons are electrically neutral particles with effectively the same mass as protons.

Electrons (e–) are negatively-charged subatomic particles that orbit the nucleus of an atom in “clouds”. We need to discuss this electron cloud in terms of probability model of where the electron(s) are most likely to be found at any given point in time…

atomic number (Z) – the number of protons in the nucleus of each atom of that element # of protons = # of electrons (since all atoms are electrically neutral)

mass number (A) – the total number of protons + neutrons in the nucleus of an isotope

# of neutrons = mass # – atomic #

Isotopes are atoms of the same element that have different masses. Isotopes have different masses due to different #s of neutrons.

There are three isotopes of hydrogen:protium (99.985%), deuterium (0.015%), and tritium (very rare, radioactive)

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mass chargeproton 1 +1neutron 1 0electron 1/2000 –1

Name: ______________________Chemistry 1 Notes,

2 ways to name isotopes:

1) nuclear symbol: atomic#mass# symbol

= ZA E

ex: 92238 U

2) hyphen notation: element–mass # ex: uranium–238 or U–238

carbon–12 or C–12

So, the three isotopes of hydrogen could be named:

protium: 11 H or hydrogen–1

deuterium: 12 H or hydrogen–2

tritium: 13 H or hydrogen–3

atomic mass unit (u) – 1/12 the mass of a carbon-12 atom

relative atomic mass – mass of an atom expressed in atomic mass units (it’s relative to the mass of carbon-12)

Non-integer atomic masses (average atomic masses) result from the existence of several different isotopes of that element.

average atomic mass – the weighted average of the masses of the naturally occurring isotopes of an element

For example, there are 2 main isotopes of chlorine: 1735 Cl & 17

37 Cl

75% of all chlorine is 1735 Cl , with a mass of 35u and 25% is 17

37 Cl , with a mass of 37u

So the average atomic mass for chlorine is (.75 x 35u) + (.25 x 37u) = 35.5u

Take carbon as a second example:

carbon-12 = 12.0u 98.90%, in naturecarbon-13 = 13.003355u 1.10%, in nature

average atomic mass for carbon = (0.9890 x 12.0u) + (0.0110 x 13.003355u) = 12.011 u

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Name: ______________________Chemistry 1 Notes,

III. Electrons and Light

WAVE-PARTICLE NATURE OF LIGHT AND ELECTRONS

At the time of Rutherford (~1900), electrons were pictured as particles, and light was pictured as waves. But this doesn’t explain all properties of electrons or light. In the early 1900’s the wave-particle theory was proposed to fully explain the properties of

electrons.

Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. Shown below is the full electromagnetic spectrum, arranged according to decreasing wavelength. It is a continuous spectrum, showing all wavelengths of the full electromagnetic spectrum. Note wavelength () and frequency () are inversely proportional. There is also an inset showing the continuous spectrum of visible light (ROY G. BIV, like a

rainbow).

All electromagnetic radiation travels at 3.0 x 108 m/s (the speed of light, c). c = (speed of light = wavelength x frequency) Again, note that wavelength and frequency are inversely proportional. Since their product (c)

is a constant, if one goes up, the other must go down.

A line spectrum shows only certain, specific wavelengths of light (like a fingerprint). Every element has its own unique line spectrum. That is how we can observe starlight through a spectrophotometer, look at the wavelengths of

light, and infer what elements are contained in that star.

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Name: ______________________Chemistry 1 Notes,

The emission spectrum of an element is the unique line spectrum produced by an element when it is burned in a flame. Shown here is the emission spectrum of hydrogen, and a gas discharge tube filled with hydrogen, with an electric current running through it, producing the characteristic lavender color of hydrogen. (We’ll look at how these spectra form next…)

♦ Here is a good animation of several flame tests for different metals: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/039_FlameTestsMet.MOV

QUANTUM THEORY – MAX PLANCK

♦ Here is a quick synopsis of Quantum Theory, first suggested by Max Planck in 1900:

Quantum Theory describes the particles that make up matter and how they interact with each other and with energy. The name “quantum theory” comes from the fact that it describes matter and energy in the universe in terms of single indivisible units called quanta.

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Max Planck

Name: ______________________Chemistry 1 Notes,

Quantum theory is different from classical physics. Classical physics is an approximation of the set of rules and equations in quantum theory, and can help explain the motion of a car accelerating or of a ball flying through the air.

Quantum theory, on the other hand, can accurately describe the behavior of the universe on a much smaller scale, that of atoms and smaller particles. The rules of classical physics do not explain the behavior of matter and energy on this small scale. Quantum theory is more general than classical physics, and in principle, it could be used to predict the behavior of any physical, chemical, or biological system. However, explaining the behavior of the everyday world with quantum theory is too complicated to be practical.

Quantum theory not only specifies new rules for describing the universe but also introduces new ways of thinking about matter and energy. The tiny particles that quantum theory describes do not have defined locations, speeds, and paths like objects described by classical physics. Instead, quantum theory describes positions and other properties of particles in terms of the chances that the property will have a certain value. For example, it allows scientists to calculate how likely it is that a particle will be in a certain position at a certain time (see the Heisenberg Principle below).

Quantum theory describes all of the fundamental forces—except gravity—that physicists have found in nature. The forces that quantum theory describes are the electrical, the magnetic, the weak, and the strong.

One of the striking differences between quantum theory and classical physics is that quantum theory describes energy and matter both as waves and as particles. For example, classical physics considers light to be only a wave, and it treats matter strictly as particles. Quantum theory acknowledges that both light and matter can behave like waves and like particles.

Max Planck said that when an object (like an electron) loses energy, it does not do so continuously (which it would if radiation were in the form of waves).

Energy is radiated off in small specific amounts called quanta. A quantum is a finite quantity of energy that can be gained or lost by an atom. A photon is a quantum of light – a particle of radiation. Radiation is absorbed and emitted only in whole numbers of photons.

According to Planck, there is a relationship between the frequency of a particular radiation and the energy with which it is associated.

E = h , where “h” is known as Planck’s constant (h = 6.626 x 10–34 Js) As frequency increases, the energy of the radiation increases (frequency and energy are

directly proportional).

Using the equation: c = (speed of light = wavelength x frequency), you can rearrange the first equation to relate energy with wavelength:

E =

hcλ

As wavelength increases, the energy of the radiation increases (wavelength and energy are inversely proportional).

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Name: ______________________Chemistry 1 Notes,

Consider the emission spectrum of hydrogen: it consists of only 4 lines (red, blue-green, blue, and violet).

Remembering the Bohr model of the H atom, now we can explain why these are the only colors produced in an emission spectrum.

The state of an atom in which its electrons exist at their normal energy levels is called the ground state, or E0. When electrons absorb energy (when they are heated, or two atoms collide, etc.), they get “excited” or “promoted” up to higher energy levels (excited states or E’). An electron must gain an amount of energy equal to the energy difference between the two energy levels, in order to move from its ground state to the higher energy level.

When these excited electrons drop down to lower energy levels, they give off the excess energy they no longer need as photons of electromagnetic energy. Specifically, they must release an amount of energy equal to the energy difference between these two energy levels as a photon of electromagnetic radiation.

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Name: ______________________Chemistry 1 Notes,

Look at the Bohr model for the hydrogen atom shown below, and make sure you understand the specific details of the process. And note that we only use the Bohr model to describe the hydrogen atom. Other atoms, with more electrons, become too complex to describe.

This basic idea explains why each element (with its own unique arrangement of electrons and energy levels) would have its own unique emission spectrum.

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Remember that these energy levels are discrete. Electrons must exist within a given energy level at any point in time – they cannot exist between energy levels, etc.

Name: ______________________Chemistry 1 Notes,

EVOLUTION OF THE ATOMIC MODEL

♦ Good website that further summarizes development of atomic theory, including animations:http://www.broadeducation.com/htmlDemos/AbsorbChem/HistoryAtom/page.htm

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♦ In 1911, Rutherford suggested that electrons orbit the atomic nucleus like planets round the Sun (to explain his gold foil experiment results).

In 1914, Niels Bohr modified Rutherford's model by introducing the idea of energy levels:1. Electrons move around nucleus in

orbits of definite potential energy. 2. The lowest energy orbit is the one

closest to the nucleus; the orbit with the most energy is farthest away from the nucleus.

3. An electron can neither gain nor lose energy in its orbit, but could move up or down into other orbits.

Schrodinger’s quantum model of the atom (wave mechanical model) – electrons move around the nucleus in orbitals (clouds) of increasing potential energy, where they are most likely to be found (wave-like behavior).

Heisenberg Uncertainty Principle – You cannot simultaneously measure the position and velocity of an electron

Werner Heisenberg

Year Scientist Model

1803-07 Dalton Billiard BallModel

1898 Thomson Plum PuddingModel

1910 Rutherford

1913 Bohr PlanetaryModel

1925 SchrodingerWave

MechanicalModel

Name: ______________________Chemistry 1 Notes,

♦ We mentioned earlier that both light and electrons exhibit particle-like and wave-like behavior. The particle part is easy. An electron has mass and takes up space – not much of either, but it’s there. The wave part is a bit more convoluted. Suffice it to say that a scientist named Louis de Broglie first

postulated the idea. He reasoned that if light could have wave-particle duality, then so should matter. He applied this idea to electrons specifically, and his prediction received experimental confirmation from two scientists experimenting with electron diffraction in crystals. Thus, de Broglie’s theory has become cemented into the foundations of modern physics.

As illustrated in this diagram, you can even apply the concept of wave-like behavior to an electron as it orbits around the nucleus of an atom. This concept is important in more advanced theories of molecular bonding.

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Name: ______________________Chemistry 1 Notes,

IV. Quantum numbers Quantum numbers – specify the properties of atomic orbitals and their electrons

There are four quantum numbers: 1) Principal 2) Orbital 3) Magnetic 4) Spin

♦ Principal quantum number (n): indicates the energy levels (shells) surrounding the nucleus Values of n are whole #s only: 1-7 As n increases, the distance of e- from the nucleus increases & so does the energy of the electrons

in that energy level.

Orbital quantum number (l): indicates orbital shape (subshells or sublevels) orbital – 3-dimensional region around the nucleus where a particular e- can be located

s = sphere p = peanut d = double peanut f = flower

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an s-orbital

Name: ______________________Chemistry 1 Notes,

s, p, d, and f actually stand for sharp, principal, diffuse and fundamental, but we use sphere, peanut, double peanut and flower to help remember and describe their shape.

Electrons with a given orbital quantum # (s, p, d, and f) can always be found orbiting the nucleus somewhere in that shape/region. These orbitals are really the shapes of electron clouds, then. They describe paths electrons of

a given energy sweep out as they move around the nucleus. With this in mind, the actual motion of the electron cannot be defined. Remember the

Heisenberg uncertainty principle? We can know either the electron’s position or its velocity, but not both simultaneously.

How an electron exists in its orbital is essentially unknown. In fact, it only occupies this “shape” about 90% of the time. The other 10%, the electron is elsewhere, somewhere in the universe (typically limited by relativity) but not in the orbital. So these shapes can actually be described as probability models, and electrons of very specific energy are found within these orbitals about 90% of the time.

Different orbitals have slightly different energies, so different electrons cannot move between orbitals (unless the energy levels are degenerate, like the px, py, and pz orbitals, for a given energy level) without absorbing or giving off energy.

Magnetic quantum number (ml): indicates the orientation of an orbital around the nucleus s – only one orientation (it’s round) p – three axes (x, y, z) three possible orientations – px, py, pz

p-orbital takes the letter of the axis it’s centered on (see previous page) d – five possible orientations (see previous pages) f – seven possible orientations (see previous pages)

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Name: ______________________Chemistry 1 Notes,

Spin quantum number (ms): indicates the two possible states of an electron in an orbital The two values of the spin quantum number are +½ and –½. Each orbital can hold no more than two electrons, which must have opposite spins. These electrons can be symbolized with an up arrow () and a down arrow ().

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Name: ______________________Chemistry 1 Notes,

V. Electron Configurations Electron configuration – the arrangement of electrons in atoms

Aufbau principle – an electron occupies the lowest-energy orbital that can receive it Electrons, like everything else in nature, want to be stable – arranged with the lowest possible

energy.? Which orbitals have the lowest energy? – the ones closest or farthest away from the nucleus?

(closest – 1s, then 2s, then 2p, etc.)

Hund’s rule – orbitals of equal energy are each occupied by one electron before any one orbital receives a second electron, and all electrons in singly occupied orbitals must have the same spin

Ex: each of the three p orbitals gets one electron before any one gets a second electron. Each of these unpaired p electrons will have the same spin, either +½ or -½.

? How many unpaired electrons can there be (max) in a d sublevel before any of the d orbitals gets a second electron?

Pauli exclusion principle – no two electrons in the same atom can have the same set of four quantum numbers

They are all unique – we name each of them differently, as they cannot have the same set of quantum numbers.

Electrons in the same orbital have different spin quantum #s.

♦ This website has a great animation explaining these rules / principles and some of the notations on the next page of notes:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/043_ElectronConfig.MOV

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Name: ______________________Chemistry 1 Notes,

There are several notations that can be used to indicate electron arrangement in an atom:

1) orbital notation An unoccupied orbital is represented by a line ___ and up or down arrows (since a pair of

electrons in an orbital will have opposite spins). Underneath the lines will be the principle quantum # (shell) and orbital (subshell).

2) electron configuration notation Write the principle quantum # and subshell and add a superscript indicating how many

electrons are present in that orbital. Shorthand electron configuration notation uses noble gases enclosed in brackets to

indicate what which orbitals have already been filled.

The electron arrangement Mg could also be represented as: 2.8.2 or 2,8,2to indicate the number of electrons occupying each successive energy level. Know this arrangement, because it will definitely be on the IB exam.

3) electron-dot notation Shows only the electrons occupying the highest (or outermost) main energy level.

bromine =

Only s and p orbitals are in the highest energy level !! Therefore, there can only be a max of 8 electrons in the highest energy level. An atom with a full 8 electrons has a full octet. All atoms would like to have a full octet, because that is the most stable arrangement.

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