Download Topic 7. 1 Atomic Structure

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Topic 7. 1 Atomic Structure
7.1.1 Describe a model of the atom that features a small nucleus surrounded by electrons.
 The modern atom has gone through a few stages of development
1) Dalton’s Atomic Theory – idea of an atom
2) JJ Thompson – 1890 – negative charge (electrons)
3) Earnest Rutherford – 1911 - positive nucleus (protons)
4) Niels Bohr – 1913 – orbital shells
5) Chadwick – 1932 – neutrons
 This is a VERY simplified idea of the atom
 Nucleus
 Protons – positive charge – 1.6 x 10-19C
 Neutrons – no charge
 Diameter order of 10-15m
 Electron “cloud”
 Electrons – negative charge – 1.6 x 10-19C
 Diameter order of 10-10m
 The nucleus is about 100,000 times smaller than the electron orbits.
 Imagine a pea in the center of a football field with the track being the orbits.
 Protons and Neutrons have very similar mass.
 Protons and Neutrons are about 1800 times bigger than electrons.
****IB DATA*** See the data book for actual values.
7.1.2 Outline the evidence that supports a nuclear model of the atom
 Dalton’s Atomic Theory
1) All matter is composed of extremely small particles called atoms.
2) All atoms of a given element are identical.
3) Atoms cannot be created, divided into smaller particles, or destroyed. (This part proven wrong)
4) Different atoms combine in simple whole number ratios to form compounds.
5) In a chemical reaction, atoms are separated, combined or rearranged.
6) Subatomic Particles and the Atom
 J. J. Thomson – 1890-1900
 Used cathode ray tube to prove existence of electron.
 Proposed “Plum Pudding Model”
 Cathode ray tube
 Stream of charged particles (electrons).
 Plum Pudding  Ernest Rutherford
 Gold Foil experiment
 Used to prove the existence of a positively charged core (Nucleus)
 Fired alpha particles(2protons and 2 neutrons) into very thin gold foil.
 The results were “like firing a large artillery shell at a sheet of paper and having the shell come back and
hit you!”
 What should have happened
 What did happen
 After performing hundreds of tests and calculations, Rutherford was able to show that the diameter of
the nucleus is about 105 times smaller than the diameter of the atom
 Subatomic Particles and the Atom
 James Chadwick – 1932
 Worked with Rutherford.
 Noted there was energy in the nucleus, but wasn’t the protons.
 Concluded that neutral particles must aslo exist in nucleus.
 Bombarded a beryllium target with alpha particles
 Alpha particles are helium nucleus
 Discovered that , carbon was produced with another particle.
 **** Write reaction on board****
 Concluded this particle had almost identical mass to proton but no charge.
 Called it a neutron
 Three main particles:
 Proton
 Positive
 In nucleus
 Neutrons
 Neutral
 In nucleus
 Electrons
 Negative
 Orbiting the nucleus (not inside)
7.1.3 Outline one limitation of the simple model of the nuclear atom.
7.1.4 Outline evidence for the existence of atomic energy levels.
 If low-pressure gases are heated or current is passed through them they glow.
 Different colors correspond to their wavelengths.
 Visible spectrum 400nm(violet) to 750nm(red)
 When single element gases such as hydrogen and helium are excited only specific wave lengths were
emitted.
 These are called emission line spectra
 If white light is pass through the gas the emerging light will show dark bands called absorption lines.
 They correspond to the emission lines.
LIMITATION
 Rutherford’s model didn’t explain why atoms emitted or absorbed only light at certain wavelengths.
 1885 JJ Balmer showed that hydrogen’s four emission lines fit a mathematical formula.
 This “Balmer series” also show the pattern continued into non-visible ultra-violet and infra-red.
 Bohr called these “energy levels”
 Reasoned that the electrons do not lose energy continuously but instead, lose energy in discrete amounts
called “quanta”.
 He agreed with Rutherford that electrons orbit the nucleus but only certain orbits were allowed.
 Bohr explained the emission and absorption line spectra with the idea that electrons absorbed only certain
quantity of energy that allowed it to move to a higher orbit or energy level.
 Each element has its own “finger print”.
7.1.5 Explain the terms nuclide, isotope and nucleon
7.1.6 Define nucleon number A, proton number Z, and neutron number N.
IB Definitions
 Nucleon – any of the constituents of a nucleus. Protons and neutrons.
 Atomic Number – The number of protons in the nucleus.
 Nucleon Number – The number of nucleons in the nucleus. AKA the mass number. (protons +
neutrons)
 Isotope – Nuclei which contain the same number of protons but different numbers of neutrons.
 Nuclide – the nucleus of an atom. The nuclides of isotopes are different, even though they are the
same element.
 Atomic Number (proton number), Z
 How many protons there are.
 This is what defines the element.
 Ex. Hydrogen Z =1, Oxygen Z = 8 Carbon Z = 6
 Nucleon Number (mass number), A
 How many nucleons there are.
 Protons + neutrons
 Number of neutrons, N
 Mass number = atomic number + number of neutrons
 A=Z+N
.
 Standard notation is: A over Z in front of element(X)
 *****Draw on board*****
 Isotopes
 More evidence for neutrons is the existence of isotopes.
 When nuclei of the same element have different numbers of neutrons.
 Carbon has 6 isotopes: Carbon-11, Carbon-12, Carbon-13, Carbon-14, Carbon-15, Carbon-16.
 All have 6 protons but each has different number of neutrons.
 The different isotopes don’t exist in nature in equal amounts.
 Carbon:
 C – 12 is most abundant (98.9%)
 C – 13 is next (1.1%)
 This is where atomic mass comes from. It’s the weighted average mass of all the different isotopes.
 Nuclei of different atoms are known as nuclides.
 Ex. C – 12, C – 14
 Both are carbon but different isotopes
 Their nuclei have different numbers of neutrons.
 These are different nuclides.
7.1.7 Describe the interactions in a nucleus
 How do like charge (protons), stay stuck together?
 We already know that like charges repel each other.
 We have also seen that they are stronger than gravitational forces.
 Strong Force – The force that binds the nucleus together.
 It is an attractive force that acts between all nucleons.
 Short – range interactions only (up to 10-15m)