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Chemistry, Chapter 4, Theories and Structures.
Terms in this set (40)
Dalton's Atomic Model
An atom represented as a sphere.
Dalton's Atomic Theory
All matter is made of atoms.
Dalton's Atomic Experiment/Proof
1. An element is made up of atoms. All atoms of a given element are identical.
2. Atoms cannot be destroyed or created (atoms can be rearranged so technically they can be created).
3. All atoms of one elements have the same mass. Atoms of two different elements have different masses.
4. Atoms of different elements may combine in the ratio of small, whole number to form compound.
Dalton's Atomic Shortcomings
1. The indivisibility of an atom was proved wrong: an atom can be further subdivided into protons, neutrons and electrons. However an atom is the smallest particle that takes part in chemical reactions.
2. According to Dalton, the atoms of same element are similar in all respects. However, atoms of some elements vary in their masses and densities. These atoms of different masses are called isotopes. For example, chlorine has two isotopes with mass numbers 35 and 37.
3. Dalton also claimed that atoms of different elements are different in all respects. This has been proven wrong in certain cases: argon and calcium atoms each have an atomic mass of 40 amu. These atoms are known as isobars.
4. According to Dalton, atoms of different elements combine in simple whole number ratios to form compounds. This is not observed in complex organic compounds like sugar (C12H22O11).
5. The theory fails to explain the existence of allotropes; it does not account for differences in properties of charcoal, graphite, and diamond.
Thompson's Atomic Model
The pink portion represents a spherical cloud of positive charge and the yellow "dots" represent electrons with a negative charge. You can think of a chocolate chip cookie to remember this model.
Thompson's Atomic Theory
An atom consists of a sphere of positive charge with negatively charged electron embedded in it. The positive and the negative charges in an atom are equal in magnitude, due to which an atom is electrically neutral. It has no overall negative or positive charge. He also discovered that the atom can be divided into four sections.
Thompson's Atomic Experiment/Proof
Thomson discovered electron by the cathode ray tube. It has been previously seen that if a electric current is passed through a vacuum tube, a stream of glowing material was formed. Thomson found that the mysterious glowing stream would bend toward a positively charged electric plate. Thomson atomic theory proved that the stream is made up of small particles which is piece of the atom and is negatively charged. Thomson named these particles as electrons.
Thompson's Atomic Shortcomings
Thomson's atomic model explained the overall neutrality of an atom. Its assumption that the total mass of an atom is uniformly distributed all over the atom was inconsistent with some of the experimental results. It was later found that the plum pudding atomic model was insufficient to explain the structure of an atom (Rutherford's gold foil experiment).
Rutherford's Atomic Model
Larger red "dot" in the middle represents a positively charged nucleus, green "dots" scattered about represent negatively charged electrons, and the dashed line represents the "circumference" of an atom and suggests that the walls aren't perfectly circular.
Rutherford's Atomic Theory
Atoms have a uniform structure with protons found in the nucleus and electrons orbiting around the nucleus.
Rutherford's Atomic Experiment/Proof
He shot high velocity alpha particles (helium nuclei) at an atom at a thin film of gold atoms. He expected the alpha particles to go right through the gold foil, but to his amazement a few alpha particles rebounded almost directly backwards. Rutherford reasoned that the only way the alpha particles could be deflected backwards was if most of the mass in an atom was concentrated in a nucleus. He thus developed the planetary model of the atom which put all the protons in the nucleus and the electrons orbited around the nucleus like planets around the sun.
Rutherford's Atomic Model Shortcomings
1. He did not say anything about the arrangement of electrons in an atom which made his theory incomplete.
2. Accelerated charged particles emit electromagnetic radiations and hence an electron revolving around the nucleus should emit electromagnetic radiation. This radiation would carry energy from the motion of the electron which would come at the cost of shrinking of orbits. Ultimately the electrons would collapse in the nucleus, meaning Rutherford's model was not completely correct.
Bohr's Atomic Model
Labels below explain the model.
Bohr's Atomic Theory
Electrons orbit in shells around the nucleus of an atom, the shells having discrete energy levels.
Bohr's Atomic Experiment/Proof
Niels Bohr proposed a theory for the hydrogen atom based on quantum theory that energy is transferred only in certain well defined quantities. Electrons should move around the nucleus but only in prescribed orbits. When jumping from one orbit to another with lower energy, a light quantum is emitted.
Bohr's Atomic Shortcomings
1. The Bohr atomic model theory considers electrons to have both a known radius and orbit i.e. known position and momentum at the same time, which is impossible according to the Heisenberg Uncertainty Principle.
2. The Bohr atomic model theory made correct predictions for smaller sized atoms like hydrogen, but poor spectral predictions are obtained when larger atoms are considered.
3. It failed to explain the Zeeman effect when the spectral line is split into several components in the presence of a magnetic field.
4. It failed to explain the Stark effect when the spectral line gets split up into fine lines in the presence of electric field.
Quantum Atomic Model
Labels below explain the model.
Quantum Atomic Theory
Describes electrons in atoms based on Schrodinger's equation.
Quantum Atomic Experiment/Proof
The proof was developed from Schrodinger's equation.
Quantum Atomic Shortcomings
Didn't explain well what the size of an atom was.
Proton Mass = 1 Amu
Proton Charge = +1
Proton Location = Nucleus
Neutron Mass = 1 Amu
Neutron Charge = 0
Neutron Location = Nucleus
Electron Mass = 1/1840 Amu
Electron Charge = -1
Electron Location = Not Nucleus
Very dense region consisting of a positive charge
at the center of an atom.
The system of electrons surrounding the nucleus of an atom.
A mathematical function that describes the wave-like behavior of either one electron or a pair of electrons in an atom.
The number of protons found in the nucleus of an atom of that element.
The average mass of all of the isotopes of an element (a group of atoms).
Atomic Mass Unit
One twelfth of the mass of an unbound neutral atom of
carbon-12 in its nuclear and electronic ground state (approximately the mass of one nucleon and is equivalent to 1 g/mol).
The total number of protons and neutrons (together known as nucleons) in a nucleus (a single atom).
Smallest constituent unit of ordinary matter that has
the properties of a chemical element.
Variants of a particular chemical element which
differ in number of neutrons, although all isotopes of a
given element have the same number of protons.
An atom or molecule in which the total number of
electrons is not equal to the total number of protons, giving
the atom or molecule a net positive or negative electrical
A positively charged ion, i.e., one that would be attracted to the cathode in electrolysis. A cation has more protons than electrons.
A negatively charged ion, i.e., one that would be attracted to the anode in electrolysis. An anion has more electrons than protons.
Two atoms that have the same numbers of electrons or the same electronic structure.
How to determine ionic charge.
(# protons) - (# electrons) = (ion charge)
How to determine the number of protons.
(atomic #) = (proton #)
How to determine the number of neutrons.
(mass #) - (atomic #) = (# neutrons)
How to determine the number of electrons.
(# electron) = (# proton) = (atomic #)
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