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AP Chem Atomic Structure

Terms in this set (60)

Developed rubber ball model of atom
Dalton
Developed plum pudding model of atom
Thompson
Discovered proton although it was identified and named later
Thompson
Determined mass of proton and electron in his oil drop experiment
Millikan
Accidentally discovered radioactivity
Becquerel
Discovered nucleus with his gold foil experiment
Rutherford
Discovered neutron
Chadwick
Determined actual charge of electron
Millikan
When using cathode ray tube, what 5 things determine amount of deflection of the beam
1. Mass of particle
2. Velocity of particle
3. Electric charge of particle
4. Strength of magnet
5. Amount of charge in plates
Principle Quantum Number
n
-main energy levels and size of orbitals
Azimuthal Quantum Number
l
-sublevels and shapes of orbitals
Magnetic quantum number
m
-orbitals and spatial orientation
Spin Quantum Number
s
-Represents electrons and determines their spin
Describe the significance of the bright line spectrum of elements and how it indicates that electrons have quantized energies
The significance of the bright line spectrum of an element is that each element has a unique line spectrum- it is often used to identify elements. The energy for a given wavelength of light can be found by the equation:

E=hc/wavelength

Therefore, each wavelength of light has a specific amount of energy associated with it. Furthermore, only certain lines show up in the line spectrum. If the lines are caused by the photons given off by the electron losing energy within the atom, then the electrons can only have specific amounts of energy in order to produce these lines.
List the evidences for supporting the:
a) particle properties of light
b) wave properties of matter
a) Planck showed that the wavelength of light given off by glowing gases have discrete amounts of energy (E=hc/wavelength), instead of having just any amount. Einstein stated that light is made up of photons having these discrete amounts of energy. Furthermore, Einstein shows that energy can be associated with mass (e=mc^2). By combining these two equations, Einstein indicated that photons could have a mass (m=h/wavelength*c). Einstein also showed that light behaves as particles with the photoelectric effect in which light can know electrons off a metal surface. Compton performed experiments in which collisions took place between x-rays and electrons and that photons do in face have the mass indicated in Einstein's equations.

b) De Broglie shows that matter can have wave properties by showing the relationship between the mass of an object and wavelength (wavelength=h/mv), where m is the mass of an object in kg and v is its velocity in m/s. also, experiments were carried out that showed that a beam of electrons exhibit the same diffraction pattern a beam of x-rays (diffraction is a wave property) and therefore showed particles having wave properties.
Describe Bohr's model of the atom and his support
It included:
-electron in hydrogen atom moving around the nucleus in circular orbits (energy levels)- the electrons are quantized
-the energy of the energy levels in Bohr's atom increases as the distance between the nucleus and the energy level increases
-stated an equation to find the energy of an electron when it jumps from one energy level to another

His theory is supported by the bright line spectrum for hydrogen. The change in energy determined by the above equation matched up perfectly with the energies of the different wavelengths of light that appear on this spectrum.
Why didn't Bohr's model work?
-Since electrons are under constant acceleration, they should lost energy and end up in the nucleus, but this doesn't happen
-It only worked for atoms with 1 electron
What did Schrodinger assume about an electron when coming up with the quantum theory?
That they travel in standing waves
What is an orbital and from what was it derived?
An orbital is a region about the nucleus with a specific amount of energy in which an electron may be found; derived from the wave function
What is the quantum mechanics theory based upon?
A 90% probability that electrons will be where expected
What is meant for something to be quantized in energy?
That it exists in specific energy levels
What are degenerate orbitals? Give an example
Degenerate orbitals are alike in all ways except for spatial orientation. This means they have the same shape, size, energy, etc. Example: 3px, 3py, 3pz (any orbitals in the same sublevel will be degenerate)
How are the 3d and 4d orbitals similar and different
Similar
-same shape
-both hold up to 2 electrons
Different
-4d is larger
-4d has more energy
Why can't we account exactly for the repulsion among electrons in polyelectronic atoms?
The amount of repulsion between electrons depends on the distance between the electrons. From the Heisenberg's uncertainty principle, we can't know precisely the position of each electron. Thus, we can't know precisely the distance between the electrons to determine the amount of repulsion involved.
What does it mean for a 4s electron to be more penetrating than a 3d electron?
There is a higher probability of finding the 4s electron closer to the nucleus than the 3d electron
What is significant about valence electrons?
When atoms interact with each other, it will be the valence electrons that are involved in these interactions. Furthermore, how tightly the nucleus holds these valence electrons determines atomic size, ionization energy and other properties of the atom. In other words, the valence electrons determine the chemical properties of the atom.
Why do elements in the same group on the periodic table have similar properties?
They have the same number of valence electrons
Why is the second ionization energy for Li larger than that of Be?
After losing the first electron, Li has a noble gas configuration (like He) and is so stable that it would take a
tremendous amount of energy to remove the second electron. For Be, after losing the first electron, it can still
lose another electron to become like He. Therefore, it will take more energy to remove the second electron from
Be than its first electron, but not as much as it does for Li.
Why does the element N not have a stable electron affinity whereas C dos?
The orbital diagram for N,

2s 2p
, shows that N is somewhat stable already and that any other electrons
added to it will cause extra electron repulsion forces that cannot be overcome by the nuclear charge, thus making
the atom unstable. For C, the orbital diagram,
2s 2p

, shows that C can accept another electron without
excess electron repulsion since it has an empty orbital in which to accept the electron.
Why is first electron affinity exothermic and endothermic for each electron thereafter?
Generally, adding an electron doesn't cause excess repulsion forces to occur—the nuclear charge can handle the
extra electron in this case—and can form a stable anion. Therefore, since the atom is becoming more stable,
energy would be released to do this. Every electron added thereafter will cause extra repulsion forces to occur.
Thus, the atom must gain energy to hold onto these extra electrons
Explain why the first ionization energy tends to increase as one proceeds across left to right across a period; why the first ionization energy of Al is lower than Mg and that of S is lower than P?
Generally, the smaller the atom the more ionization energy is needed to remove an electron. Since the atomic size
of the atoms get smaller as one proceeds across the period of the periodic table, the first ionization energies will
increase from left to right across the period. As for the slight dip in IE between Mg and Al and the dip that occurs
between P and S, one needs to look at the orbital diagrams for these elements.
Mg →
3s

Al →
3s 3p

The Mg has a filled 3s sublevel—therefore a stable sublevel. The Al has only a single electron in its 3p sublevel. It
will require less energy to remove the single 3p electron in Al than to remove one of the electrons in Mg's 3s
sublevel.
P →

3s 3p
S →

3s 3p
The configuration for P is quite stable with each 3p orbital containing a single electron. With the addition of
another electron in the 3p sublevel of S, additional repulsion forces occur between the two electrons in the same
orbital. Thus, it would require less energy to remove the additional electron in 3p sublevel of S than to remove an
electron from P's 3p sublevel.
Explain why successive ionization energies of an atom always increases.
After an atom loses an electron, the remaining electrons move closer to the nucleus since the nuclear charge is
now distributed over fewer electrons. This means it would require more energy to remove the 2nd electron from an
atom since the nucleus has a tighter hold on the electrons than when it is neutral. With each successive electron
removed, the fewer electrons the nuclear charge is distributed over and therefore more energy is needed to
overcome the addition attraction between the nucleus and the electrons.
Why is there a larger gap between the third and fourth ionization energies for Al that what exists between the first three ionization energies for Al?
After Al loses 3 electrons, it is like a noble gas (Ne). Therefore, it would require a tremendous amount of energy to
remove another electron from a very stable ion. As for the gap between the 1st, 2nd and 3rd ionizations, each time
Al loses an electron, its nucleus binds the remaining electrons tighter to the nucleus. Thus it will require more
energy to remove the next electron from the cation, but none nearly as much as it would to remove the 4th
electron from this atom.
What is the most reactive nonmetal? metal?
nonmetal → F metal → Fr
Why does the atomic size of the atoms tend to decrease from left to right across the period and increase from top to
bottom down the group?
The atomic size decreases as you go across the period (left to right) due to the increased nuclear charge having a
tighter hold on electrons that are all in the same energy level. The atomic size increases as you go down the group
due to larger energy levels being occupied and increased shielding effect of the core electrons on the valence
electrons.
Explain why the most reactive metals are on the left side of the periodic table and the most reactive nonmetals are on
the right side
The metals on the left side of the periodic table have low ionization energies. This allows these atoms to more
easily lose electrons to form ions and be involved in chemical reactions. The nonmetals to the right side of the
periodic table have a higher electronegativity. This allows them to more easily gain electrons to form ions and be
involved in chemical reactions.
Why do nonmetals have higher ionization energies, the more exothermic electron affinity, and higher
electronegativity than metals?
The nonmetals are typically smaller than metals. This means more IE is needed to remove electrons. Also, it
means the nonmetallic atoms are more apt to gain electrons (more exothermic electron affinity and higher
electronegativity) since electrons outside the atom are closer to the nucleus in these atoms than for metallic
atoms.
Explain why metals are reducing agents and nonmetals are oxidizing agents
Metals are reducing agents because they have lower ionization energies and can easily lose electrons (oxidize). By
doing this, they cause another substance to gain electrons and are therefore reducing agents. Nonmetals are
oxidizing agents since they have high electronegativities and can easily gain electrons (reduce). By doing this, they
cause another substance to lose electrons and are therefore oxidizing agents.
Give the increasing order of reducing power for solid alkali metals reacting with nonmetals and explain why this
occurs.
Fr > Cs > Rb > K > Na > Li
This order is the reverse order of increasing first ionization energies for these elements. The lower the ionization
energy, the more reducing power the element has.
Explain the relationship between the alkali metals (from Li - Cs) and their density, melting point, boiling point, first
ionization energies, and size (atomic and ionic).
. As you go down the group, the:
density generally increases (except K is less than Na)
melting point decreases
boiling point generally decreases (Cs is higher than Rb)
first ionization energy decreases
atomic size increases
ionic size increases
When using the cathode ray tube, what 5 things determine the amount of deflection of the beam?
-the mass of the particle—the heavier the particle the less it can be bent
-the velocity of the particle—the faster the particle the less it can be bent
-the electric charge of the particle—the higher the charge the more it can be bent
-the strength of the magnet—the stronger the magnet the more the beam can be bent
-the amount of charge in the plates—the stronger the char
Explain how the threshold frequency of light affects the amount of electrons being emitted from a metal surface in the
photoelectric effect
-below threshold frequency: no electrons emitted regardless of light intensity
-greater than threshold frequency: number of electrons emitted increases with light intensity
-greater than threshold frequency: kinetic energy of emitted electrons increases linearly with the frequency of
light
amplitude
the height of a wave's crest
degenerate orbitals
Orbitals that are identical except for their spatial orientation
diamagnetic
has no unpaired electrons and will not be attracted to a magnetic field
electron affinity
energy involved with attracting electrons into an atom
electronegativity
attraction an atom has for an electron
frequency
the number of complete wavelengths that pass a point in a given time
ground state
The lowest energy state of an atom
ionization energy
The amount of energy required to remove an electron from an atom
isoelectronic
Having the same number of electrons (same electron configuration)
node
A point of zero amplitude on a standing wave
orbital
regions of space around the nucleus of an atom where an electron is likely to be found
paramagnetic
unpaired electrons, attracted to magnetic field
photoelectric effect
refers to the emission of electrons from a metal when light shines on the metal
quantized in energy
only certain energy levels are allowed
standing wave
a pattern of vibration that simulates a wave that is standing still
valence electron
Electrons on the outermost energy level of an atom
valence shell
outermost electron shell
wavelength
Horizontal distance between the crests or between the troughs of two adjacent waves
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