2. The experimental evidence that led to the Rutherford model was the results of bombarding a thin metal foil with an alpha particle beam. The beam was mostly undeflected, as expected; however, a small but significant number of alpha particles were deflected—some, through very large angles.
3. (a) Rutherford inferred that the nucleus was very small (compared to the size of the atom) because very few alpha
particles were deflected at all—so the vast majority had to be completely missing whatever in the atom was “solid.”
(b) Rutherford inferred that the nucleus was positively charged because the mathematics of the angles of deflection
of the alpha particles was consistent with Coulomb’s Law of repulsion of similar charges—and alpha particles
were known to be positively charged.
4. (a) The experimental evidence used in the discovery of the proton was the study of the behaviour of positive rays in
a modified cathode ray tube.
(b) A proton with an electric charge of 1+ is a small massive subatomic particle found in the atomic nucleus.
5. (a) The experimental evidence used in the discovery of the neutron was the effects of alpha bombardment of materials—
and the fact that protons and electrons could not account for all of the observed mass of atoms.
(b) A neutron is a small massive subatomic particle found in the atomic nucleus and has no electric charge.
6. A “black box” is a system that cannot be directly observed and that must be understood by indirect interpretation of evidence. Atomic structure is an example of a concept built from indirect evidence.
1. The two most important experimental observations leading to the quantum theory of light were: Max Planck’s observation that electromagnetic radiation emission could only be explained by hypothesizing that such energy release must occur in discrete amounts, or quanta; and Albert Einstein’s observation that the photoelectric effect could be explained by assuming that light energy travels in discrete packages of given energy, which he called “photons.”
2. Max Planck may be considered the father of quantum theory because he was the first to advance the concept that energy like matter, is quantized, and not continuous.
3. The photoelectric effect is the emission of electrons from the surface of a substance when electromagnetic (light) energy strikes the surface. The experiment requires a light source that can be varied in intensity (brightness) and also in frequency (wavelength/colour); and equipment to measure the (photo)electron flow rate (current) and relative electron energy (voltage).
4. Quantum is Planck’s term for a small, discrete, indivisible quantity. Photon is Einstein’s term for a discrete quantity, or quantum, of light.
For answers to questions 6 & 7 Click here
4. (a) Emission spectra consist of light emitted by a sample of a substance. Absorption spectra consist of the missing
frequencies (colours) of light after the source light passes through a sample of a substance.
(b) Bohr explained emission spectra as light emitted by an atom when its electrons drop from higher to lower energy
states. The dark lines in absorption spectra correspond to the light absorbed by electrons jumping from lower to
higher energy states.
7. Atomic numbers give the number of protons, and for neutral atoms, the number of electrons as well. The period number shows how many levels of electron energy there are, and the last digit of the group number shows how many electrons are in the valence (highest) energy level.
8. Bohr’s theory was considered a success because it explained the known emission spectral lines for hydrogen, and predicted sucessfully some lines in the infrared light spectrum. Bohr’s theory also provided a better understanding of the arrangement of elements in the periodic table.
9. One significant problem with the Bohr theory was that it did not predict correctly the spectral lines for atoms with more electrons than hydrogen. (It also could not explain some simple observations such as some emission lines were brighter than others.) For question 13 Click here
1. The main kind of evidence used comes from atomic line spectra, particularly the splitting of lines.
2. The first quantum number describes the main energy level; the second quantum number describes small energy level steps within the main energy level corresponding to different shapes of “orbits”; the third quantum number describes the orientation in space of the electron “orbits”; and the fourth quantum number describes the “spin” of electrons.
3. (a) For l = 0, 1, 2, and 3, there are 0, 3, 5, and 7 possible values of ml , respectively.
(b) Each number is the next greater odd integer (or 2l + 1 for all ls except l = 0).
(c) From the answer to (b), the number of possible values for ml for l = 4 must be 9 (the next odd integer).
4. The fourth quantum number is ms, and it is necessary to explain magnetic properties of atoms.
5.
Table 4 Summary of Quantum Numbers
Ask about this if you are not sure.
6. It takes four quantum numbers to describe fully an electron in an atom. An example listing labels and values of each quantum number might be n = 2, l = 1, ml = –1, and ms = +1/2 . This might describe an electron in a hydrogen atom in an “excited” state.
2. A periodic table can be used to help complete energy level diagrams because it is arranged according to electron energy levels, sublevels, and orbitals
4. (a) 3p ?? ?? ??
3s ??
2p ?? ?? ??
2s ??
1s ??
potassium ion, K+ chloride ion, Cl–
(b) An atom of the noble gas argon, Ar, has the same electron orbital
energy level diagram as do these two ions.
8. fluorine [He] 2s2 2p5 chlorine [Ne] 3s2 3p5 bromine [Ar] 4s2 4p5 iodine [Kr] 5s2 5p5 astatine [Xe] 6s 6p5 Each halogen configuration ends with two s and three p orbitals. Other chemical families, such as the alkali metals, also have similar valence orbital configurations.
9. fluoride ion 1s2 2s2 2p6 sodium ion 1s2 2s2 2p6 10. Isoelectronic means having the same number of electrons.
11. zinc ion [Ar] 3d10
cadmium ion [Kr] 4d10
mercury(II) ion [Xe] 4f14 5d8
2. An electron orbital describes the three-dimensional region of space occupied by an electron, that is, in which we calculate a high probability (usually > 90%) of detecting an electron of a specific energy. An orbit is a simplified (incorrect, but useful) idea describing electrons as orbiting nuclei in circular or elliptical paths.
3. Quantum mechanics provides both the general shape (volume of space), and the electron probability density, within an orbital.
4. Quantum mechanics theory says nothing about either the position or about the motion of an electron within an orbital.
5. The 1s and 2s orbitals are spherical in shape, with the 2s orbital considerably larger and having two concentric regions of high probability density. A 2p orbital is shaped like a dumbbell, with two areas of high probable density, one on each side of the nucleus.
1. (a) Rutherford interpreted the deflection of alpha particles travelling through a thin foil to mean that atoms had tiny,
massive nuclei.
(b) Bohr interpreted the bright-line spectrum of hydrogen to mean that electrons exist only at specific energy levels.
2. The Rutherford model explained nothing about the nature of electrons. The Bohr model did not make acceptable predictions for atoms larger than hydrogen.
3. Orbit and orbital are terms that both refer to electrons within atoms. An orbit is a simplistic representation of a small particle in a circular path, used in the Bohr–Rutherford model. An orbital is a probability density for a wave function that “occupies” a volume of space, used in the visualizing of the quantum mechanical model.
4. The main kind of experimental work used to develop the concepts of quantum mechanics was spectroscopy, specifically the analysis of bright-line spectra.
5. (a) Quantum is a term referring to a smallest unit or part of something.
(b) Orbital is a term describing a volume of space that is “occupied” by an electron.
(c) Electron probability density describes the calculated likelihood of locating an electron at any point within a given
volume of space.
(d) Photon is a quantum of electromagnetic energy— a smallest “piece” or “package” of light.
6, 8, 10 ,12 do not fit as text; ask in class for any assistance
7. The idea of electron spin comes from observations of line spectra influenced by a magnetic field as well as evidence from different kinds of magnetism.
9. According to quantum mechanics, an element’s properties relate to its position in the periodic table because its position is directly related to the orbital configuration of its atoms.
11. (a) All of the alkali metals are soft, metallic solids with low melting and boiling points. They have high chemical reactivity, readily forming +1 ions. (b) We explain properties, using their electron configurations. All have a single s electron in the highest energy orbital, which is easily removed by the attraction of other atoms. The nearly empty valence shell creates the metallic properties—conductivity, shininess, and so on.
13. (a) [Kr] 5s2 4d1 (b) [Kr] 5s2 4d10 5p3 (c) [Xe] Ba2+ 14. Aluminum and titanium should be paramagnetic because these two atoms have unpaired electrons. Beryllium and mercury have atoms with filled orbitals. 15. (a) arsenic atom, As (b) rubidium ion, Rb+ (c) iodide ion, I- (d) holmium atom, Ho 16. (a) 2e- (b) 8e- (c) 18e- (d) 32e-
17. A 2px orbital is identical to the 2py and 2pz orbitals, except for orientation. It lies at 90o to the other two.
18. (a) Max Planck explained that electromagnetic energy could be released only in smallest given amounts, which he
called “quanta,” with the amount determined by the frequency of the radiation.
(b) Louis de Broglie suggested that particles could have properties and characteristics of waves, and that this effect
would be significant for tiny, fast-moving particles like electrons.
(c) Albert Einstein proposed that light (electromagnetic energy) actually travels as quanta, which he called
“photons,” and he used this concept to explain the evidence of the phenomenon called the photoelectric effect.
(d) Werner Heisenberg hypothesized that electron behaviour cannot ever be exactly described, but only discussed as
a probability system, within limits imposed by his “uncertainty principle.”
(e) Erwin Schrödinger explained electron behaviour within the atom structure as a wave phenomenon, described by
a wave mechanical equation.
19. (a) Both sodium and chlorine atoms have unfilled electron energy levels. When an electron transfers from a sodium
atom to a chlorine atom, both attain the same electron configuration as a noble-gas atom. The noble gases are
quite unreactive, which is thought to be due to their completely filled electron energy levels.
(b) The occupied and empty energy levels for lithium and sodium are quite different. Therefore, electron transitions
would be different, producing different colours. (It is not possible to explain or predict the specific colours in this
course.)
(c) Both sodium and silver atoms can obtain a more stable electron arrangement of filled electron orbitals if one electron
is removed from an atom and it forms a 1+ ion. A sodium ion becomes [Ne] and a silver ion becomes [Kr]
4d10. Combined with a chloride ion (1–), the formulas are therefore similar.
(d) A tin atom has the electron configuration [Kr] 5s2 4d10 5p2. This atom could lose its 5p2 electrons to form a 2+
ion or lose both the 5s2 electrons and the 5p2 electrons to form a 4+ ion.