1. – the smallest particle of an element that has all the chemical properties of that element....

1. ? – the smallest particle of an element that has all the chemical properties of that element. Atoms are made of protons, neutrons and electrons.

Chemistry Vocabulary Review

2. ? – two or more atoms bonded together.

Ne

Ne

Ne

Ne

Ne

Ne Neon Gas

H

H H

H

H H

H H

H H H

H H

H

H H

Hydrogen Gas

3. ? – a pure substance made of only one type of atom

Click the left mouse button advance these slides as neededSlide 7 are notes that need to be copied onto the back of the worksheet.

4. ? – a molecule made of more than one type of atom. It has different kinds of atoms bonded together.

HH

HN

Ammonia

5. ? – a substance made of more than one type of molecule, many different molecules not bonded together, and can be separated physically.

For example, a mixture of salt and pepper can be separated by adding water. The salt will dissolve into the water but not the pepper. A strainer or filter can keep the pepper on top while the saltwater goes through. Heat can then be used to boil off the water leaving the salt in the pan. It takes too long to use tweezers to separate salt and pepper grains.

A Planetary Model of the Atom / The Bohr model of the Atom

Do you like my atom?

Before Mr. Bohr’s theory, people didn’t know how electrons moved aroundthe nucleus. They thought they all just orbited in one big cloud like this.

The Bohr Model is probably familiar as the "planetary model" of the atom.

This is the symbol for atomic energy

In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun (but the orbits of electrons are not always in the same plane like the planets in the Solar System).

The above image is not to scale. In a real atom the radius of the nucleus is about 100,000 times smaller than the radius of the entire atom.

Electrons are point particles essentially without a volume. The electrons take up insignificant amount of space.

In 1913 Bohr proposed his quantized shell model of the atom to explain how electrons can have stable orbits around the nucleus.

Bohr Model for the Hydrogen Atom

The discovery of the nuclear atom, first in the Rutherford laboratory and soon after confirmed by other groups, spawned acompletely new series of theories to explain the discrete lines in the atomic spectra. Initially, the most successful of these was byNiels Bohr, who was working in the Rutherford laboratory. Bohr devised a model that combined the ideas of Einstein, Planck,and Rutherford with classical mechanics of orbiting systems that successfully explained the spectra of hydrogenic spectra, i.e.,the spectra of atoms with one electron. The Bohr model was the first to propose 'quantized states' for atomic systems, a radicalidea at the time. It was soon replaced by 'real' quantum mechanics, but it is still useful to examine the Bohr model here, since it

provides a nice application of classical orbital mechanics

introduces us to some useful concepts in electricity

provides something fun to discuss for the last few lectures of the term that we will not need to be tested upon. The fundamental idea of the Bohr model (and models other of the time) was that the electron would orbit the much heaviernucleus in essentially circular orbits. This sounds very much like the motion of planets about the sun, for example, but theforce of gravity is much too weak to account for the energies involved in the atomic spectra and for energies involved in atomicphysics. Instead, the binding force is supplied by the electromagnetic force, which we study in more detail next term. This forceoperates between charged objects, and, like the gravitational force, is an inverse square law force: where Fel is the electrical force between two objects having charges q1 and q2 separated by a distance r. The constant k playsthe role of the universal gravitational constant, and we need not worry too much about it right now. In a hydrogenic atom -hydrogen, He+, Li++, etc. - those with one electron, the nuclear charge is +Ze, where Z is the atomic number (1 for H, 2 for He,3 for Li, etc.) and e is the 'fundamental charge, 1.6x10^-19 Coulombs. The charge of the electron is -e, so the force on theelectron in a hydrogenic atom is This looks very much like the gravitation formula, with a slightly different interpretation of the variables involved. In any case,to attain a circular orbit, in the usual way we need equate this central force with the centripetal acceleration. Schematically, theclassical orbit of a hydrogenic atom looks like: The force balance condition requires that There are two apparent flaws with such a model. Firstly, there is no obvious quantization. That is, the final equation does notquantize r, since we can find the electromagnetic analog a Kepplerian orbits for any r, given the appropriate value of v (andthus T). A related problem that does not appear (at least not in the same way) in the gravitational case is that the accelerationof the electron requires that it radiate energy away in the form of light. The mathematics behind this are pretty complex andyou won't see it for a couple of years, but the problem can be easily appreciated. An antenna radiates radio waves byaccelerating charge (i.e., electrons) back and forth along its length. A circular orbit, in this sense, is just an antenna. Thisemission of radiation is a dissipative process, as far as mechanical energy is concerned, so classically the electron would spiralinto the nucleus in a very short time.

Bohr solved these problems by wishing them away in a way that would also predict the spectral properties of hydrogenic atoms.His postulates were that

the electron could only move in certain non-radiating states, which he called stationary states. This terminology has been kept in the 'real' quantum theory.

the electrons could move from one stationary state to another only by making discrete jumps. The radiation emitted is not related to the electron's motion in either stable orbit, but rather is related by Planck's relationship to the change in energy in going from one orbit to another: That is, one of the Planck/Einstein photons is emitted when an electron moves from a high energy stationary state to a low energy stationary state.

the angular momentum of the electron in the stationary states was quantized to be an integral number of Planck's constant divided by 2*pi. Since the angular momentum of a circular orbit is mvr, it follows that the Bohr quantization condition requires that where n is an integer between 1 and infinity. The constant h-bar, defined as h divided by 2*pi, has come to be a more useful constant than h itself. It is also easier to remember, since its numerical value is very nearly 1x10^-34 J-sec. Thus the angular momentum of a Bohr orbit is quantized to be an integral number of units of h-bar.

Note that the Planck relation between a photon's energy and the frequency of a light wave is equivalent to Thus the use of h-bar rather than h really amounts to finding the angular frequency omega more useful that the cyclic frequency nu.

How did Bohr arrive at this particular quantization condition? Why does n start at 1 and not zero? Obviously, he guessed. He certainly tried many others, but this is the one that worked. At the time, there was no real justification.

Radii of Bohr Orbits

To apply this quantization condition to determine the radii of electronic orbits in atoms, first solve the force balance equation for v^2: and then square the quantization condition and again solve for v^2: and set these two equal: where ao is called the first Bohr radius: Numerically, ao is equal to 0.0529 nm = 0.529 Angstroms. We see that the angular momentum quantization postulate requites that the allowed orbits have well-defined radii on the length scale of a typical atom. This is good. Schematically, the atomic orbits might look something like

The Bohr model of a Hydrogen Atom

1 proton in the nucleus Z=1

e is the fundamental charge 1.6x10^-19 Coulombs

V = velocity vector of the electron spinning around the nucleus.

= Force of electrostatic attraction

Force of electrostatic attraction

= Centripetal Force keeps the electron away from the nucleus

Centripetal Force -- keeps the electron away from the nucleus

The kinetic energy at n=2 would be greater than n=1. Or else the electron would spiral closer to the nucleus. This is because v2 would be half as much when n=2 compared to n=1. Assuming you pushed it out with your hand perpendicular to its velocity along the radius. The force of attraction would be greater (it is independent of velocity) than the centripetal force (depends on velocity).

If you push an electron from n=3 to n=2 then the electron would have too much velocity to maintain the orbit at n=2 and the centripetal force would be greater and it would spin farther from the nucleus back into n=3.

This means that energy is released as an electron spins closer to the nucleus. The energy is released in the form of photons of light.

Bohr's model explained how the the colors of light were emitted from hydrogen gas when it was put in a glass tube containing a strong electric field. (cathode tube) This is known as the emission spectra of hydrogen. He promptly won the Nobel prize.

The Bohr view of emission and energy is shown schematically in the figure below. Electrons fall from higher energy orbits to lower ones; resulting in the emission of photons of energy of different light wave frequencies.

Bohr Atom Pictures for Elements 1 - 13NAME: _____________ per_____

NAME: _____________ NAME: _____________ NAME: _____________ NAME: _____________ NAME: _____________ NAME: _____________ NAME: _____________

NAME: _______________ NAME: _______________ NAME: _______________ NAME: _______________ NAME: _______________ NAME: _______________

Atomic Number: 1Atomic Mass: 1.0079








Atomic Number: 9Atomic Mass: 19





protons: _____neutrons: _____electrons: _____













How to read the Periodic Table: 5 Atomic Number

B Atomic Symbol

10.81 Atomic MassBoron name

This is the number of protons = 5This is also the number of electrons = 5(only for neutral atoms)

- Atomic Number = 11 - 5 = 6 neutronsround off 10.81 to 11 NAME: _B,_ Boron ______


protons: __5___neutrons: __6___electrons: __5___

Draw 5 protons as hollow circlesDraw 6 neutrons as solid circlesDraw only up to 2 electrons in the first orbital.

Draw only up to 8 electrons in the second orbital. 3 in this case = 5 total

Copy this page on back of the worksheet

The outermost orbital is called the valence

Atoms touch each other at the valence and make chemical bonds.

Also copy this page on back of the worksheet

Draw the parts of a Hydrogen atom on your worksheet template.

NAME: _ H ,_ Hydrogen



Draw one proton in the nucleus

Draw one electron in the first orbitalx

= number of protons

1

= 1 rounded off 1-10 calculate the neutrons by subtracting the protons0

Draw zero neutrons in the nucleus

the number of electrons is the same as protonsfor a neutral atom.

1

Do these steps on the front of the worksheet

NAME: _ He ,_ Helium



Draw 2 protons in the nucleus

Draw 2 electrons in the first orbitalx

= number of protons

2


Draw 2 neutrons in the nucleus


2

x

NAME: _ Li ,_ Lithium





= number of protons

3




3

x

Draw 1 electrons in the second orbitalx

NAME: _ BE, Beryllium





= number of protons

4




4

x


x

NAME: _ B, Boron





= number of protons

5




5

x


x

x

NAME: _ C, Carbon





= number of protons

6




6

x


x

x x

NAME: _ N, Nitrogen





= number of protons

7




7

x


x

x x

x

NAME: _ O, Oxygen




Draw 2 electrons in the first orbital

= number of protons

8




8

Draw 6 electrons in the second orbital

NAME: _ F, Fluorine





= number of protons

9




9


NAME: _ Ne, Neon





= number of protons

10




10


NAME: _ Na, Sodium



= number of protons

11

= 23 rounded off 23-11 12 calculate the neutrons by subtracting the protons12


11





Draw 1 electron in the third orbital

NAME: _ Mg, Magnesium



= number of protons

12



12





Draw 2 electrons in the third orbital

NAME: _ Al, Aluminum



= number of protons

13



13





Draw 3 electrons in the third orbital

1. – the smallest particle of an element that has all the chemical properties of that element....

Documents

bohr model

planetary model

byniels bohr

type of atom

real atom

orbits of electrons

quantized shell model

nuclear atom