Chapter 4 Atoms Section 1 Development of Atomic Theory.
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Transcript of Chapter 4 Atoms Section 1 Development of Atomic Theory.
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Chapter 4 Atoms
Section 1
Development of Atomic Theory
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BR. Who came up with the first theory of atoms?
• Objectives- • 1. Give an example of how new scientific
data can cause an existing scientific explanation to be supported, rejected or revised
• 2. Evaluate selected theories based on supporting scientific evidence.
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Objectives (cont.)
• 3. Cite evidence that scientific investigations are conducted for many reasons.
• 4. Identify scientific evidence that has caused modifications in previously accepted theories.
• GLE’s:
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Democritus
• Over 2000 years ago• Democritus• Universe made of indivisible units called
atoms.• Atomos- unable to be cut or divided• Did not have evidence to support theory
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Dalton
• 1808 John Dalton revised atomic theory• Atoms could not be divided• All atoms of a given element were exactly
alike; and atoms of different elements could join to form compounds.
• Based on experimental evidence
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Dalton (cont.)
• Law of definite proportions- a chemical compound always contains the same elements in exactly the same proportions by weight or mass. (supported Dalton’s theory)
• Foundation of modern atomic theory• Could not explain all experimental
evidence
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Thomson
• 1897 experiment suggested that atoms were not indivisible
• He was experimenting with electricity • studying cathode rays not atoms• Cathode ray tube experiment suggested
that cathode rays were made of negatively charged particles that came from inside the atoms
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Thomson (cont.)
• Revealed that atoms could be divided• Discovered electrons, negative particles• New model- electrons spread through-out
the atom ( plum pudding model)
Mass and positive charge evenly distributed
Electrons scattered through out
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Rutherford
• Found Thomson’s model needing revising• Proposed that most of the mass of the
atom was in the center• Conducted Gold-foil experiment where
most particles passed straight through• Some particles were deflected• Some particles came straight back• Not what he expected
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Rutherford (cont.)
• Discovered the nucleus• Nucleus was very small• Electrons orbit the nucleus (like sun and
planets)• Led to new model of the atom
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Section 2- Structure of Atom
• BR. Rutherford’s Gold foil experiment led to the discovery of what?
• Objectives:• 1. Identify the 3 subatomic particles by
location, charge, and relative mass • 2. Describe the results of loss or gain of
electrons on charges of atoms
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Objectives (cont.)
• 3. Identify valence electrons in first 20 elements.
GLE’S:
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Atom
• Three subatomic particles compared by mass, charge, and location in the atom.
• Copy chart on page 119
• Nucleus- small, dense, center of atom• Atoms are Neutral.
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Atom (cont.)
• Nucleus is made of
1. protons- + charged particle
2. neutron- neutral or no charge particle
Protons and neutrons are almost equal in size and mass.
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Electrons• Move in a dense cloud (fan blades
moving) outside the nucleus• Very tiny--- 1837 electrons = 1 neutron or
proton• Negatively charged• Exact location cannot be determined.
Speed and direction cannot be determined• Located by shading; shaded region is
orbital: darker shading better chance to find
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Protons
• Each element has a unique number of protons
• Elements are identified by the number of protons that they have in the nucleus of an atom
• Atoms are neutral because they have the same number of p and e. They cancel each other out.
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Ions
• Atoms that have lost or gained electrons.
• Lose electrons become positive.
• Gain electrons become negative
• If atoms lose or gain electrons they are not atoms, but IONS.
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Atoms
• Atoms are held together by an electric force
+ and – charges attract each other by an electric force
This attraction is what holds the atom together just like the attractive force between solids and liquids.
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Atoms and Elements• Atoms of different elements have unique
structures.
• Because atoms have different structures, they have different properties.
• Atoms of the same element can vary in structure also.
• Atoms of each element have the same number of protons, but different numbers of neutrons.
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Atomic number• Atomic number equals the number of
protons.• Since atoms are neutral, it also equals the
number of electrons.
• Neutral atom-- + = -
Atomic numbers go from 1 to 116
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Mass Number
• Equals the total number of subatomic particles in the NUCLEUS of the atom.
• Nucleus contains p and n.• Mass number is equal to p + n.• Neutrons vary so mass number can vary
for the same element.
• See figure 3 on page 121
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Isotopes
• Isotope has same atomic number but different number of neutrons.
• Isotopes have same atomic number but different mass numbers.
• Isotopes have the same number of protons but different number of neutrons
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Isotopes (cont.)
• Some isotopes are more common than others.
• Radioisotopes- unstable isotopes that emit radiation and decay into other isotopes.They continue to decay until they reach a
stable isotope.They decay at a fixed rate. (fraction of a
second to millions of years)
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Isotopes (cont.)
• Since isotopes have the same number of protons and electrons, they have the same chemical properties.
• Isotopes have different masses.
• Isotopes of an element vary in mass because their numbers of neutrons differ.
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Isotopes of Water
• Water has 3 isotopes. Each isotope has 1 proton and 1 electron.
• Protium- 1 proton and no neutrons (mass number of 1
• Deuterium-1 proton and 1 neutron (mass number of 2)
• Tritium- 1 proton and 2 neutrons mass number of 3)
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Calculating Neutrons
• Isotopes are written
35 mass number ( p+n)• Cl symbol
17 atomic number (p)
Mass number – atomic number = neutrons
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Atomic mass
• Atoms are expressed in unified atomic mass units because the mass is so small.
• Unified atomic mass = equal to 1/12th of the mass of a carbon-12 atom.
• Also called atomic mass unit
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Atomic Mass
• Average atomic mass for an element is a weighted average.
• More common isotopes have more effect
than less common isotopes of the element.
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Mole
• Mole- collection of a very large number of particles.
602,213,670,000,000,000,000,000 particles
Written 6.02 x 1023 particles and is called Avogadro’s number
Avogadro’s number= the number of atoms in 12 grams of carbon-12. (Popcorn kernels covering
the US 310 miles tall)
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Molar mass
• Molar mass- the mass in grams of 1 mole of a substance
1 mole C12 = 12 g
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Converting between Moles and Grams
Amount x molar mass of element = Mass(g Moles 1 mole of element
3 moles x 32.07 g S = 96.21 g S S 1 mole S
96.27 g S x 1 mole S = 3 mol S 32.21 g S
Work problem 1 a-d on page 126
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Compounds have Molar mass
• Add all molar masses in compounds and then work the same way.
H2O x (1.01 g H x 2) + 16 g O = 18.01 g/mol
1 mole H2O H2O
1 mol H20 = 18.02 g
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Section 3Modern Atomic Theory
• BR List the 3 subatomic particles and give their location, relative mass, and charge.
• Objectives:• Describe the results of the loss or gain of
electrons on the charges of atoms.• Identify valence electrons in first 20 elements.• Draw Bohr models of 1st 20 elements• GLE’S:
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Modern Model of Atom
• Electrons are found only in certain energy levels. NOT between levels
• Location of electrons can not be predicted precisely
• Bohr- electrons can be in only certain energy levels.
• Bohr energy level related to electron’s path around the nucleus
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Modern Model of Atom (cont.)
• Electrons must gain energy to move to a higher energy level.
• Electrons must lose energy to move to a lower energy level.
• 1925 Bohr’s model revised.
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Modern Atomic Theory (cont.)
• Old out – not like sun and planets• New in – • Electrons behave more like waves on a
vibrating string than like particles.
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Energy levels and Electrons
• Many energy levels for electron to occupy• The number of energy levels that are filled
in an atom depends on the number of electrons.
• Valence electrons- those electrons in outer most energy level
• Valence electrons determine the chemical properties of the atom.
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Energy levels
• Maximum electrons in energy level• 1st = 2• 2nd= 8• 3rd= 18• 4th= 32• Must fill 1st and 2nd energy level before
going to the 3rd energy level.
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Energy levels (cont.)
• There are 4 types of orbitals.
• Orbitals are s, p, d, f.
• Orbitals determine the number of electrons that each level can hold.
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Electron Jumping
• Electrons jump between energy levels when an atom gains or loses energy.
• Lowest energy level called ground state.
• Excited state-gains energy it moves to another level
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How Electrons Move
• Electrons gain energy by absorbing a photon and move to a higher energy level
• Photon- particle of light; each have different energies
• The electron may fall back to previous energy level when it releases a photon.
• Photons determine which level the electron will• jump to.
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Light-
• Photons determine which level the electron will jump to.
• Atoms absorb or emit light at certain wavelengths.
• Energy of photon is related to the wavelength of the light.
• High energy photons= short wavelengths• Low energy photons= long wavelengths
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Atomic Fingerprint
• Because of each element’s unique atomic structure, the wavelengths emitted depend on the particular element.
• Each element emits its own characteristic color. Neon= red blue=copper
sodium= yellow strontium= red
orange = calcium green= barium