CTJan27 Online JMSS - Atomicd Structure
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1
Which experimental observation led John Dalton's atomic theory concept of the indivisibility of atoms to be discarded?
The confirmation of the Law of Multiple Proportions.
J.J. Thomson’s identification of cathode rays as negatively charged particles.
Ernest Rutherford’s finding that most of the atom is empty space.
The phenomenon of chemical combination occurring in simple whole-number ratios.
2
J.J. Thomson's cathode ray experiment primarily determined the:
Total charge of the electron.
Mass of the electron.
Mass-to-charge ratio $\left( \frac{m}{q} \right)$ of the electron.
Positive charge distribution within the atom.
3
The fundamental flaw in Rutherford's planetary model, based on classical physics, was its inability to explain:
Why elements emit characteristic line spectra.
The exact location of protons within the nucleus.
Why the orbiting electrons do not continuously lose energy and spiral into the nucleus.
The negligible volume occupied by the electrons compared to the nucleus.
4
Niels Bohr successfully integrated the concept of quantized energy levels into the atomic model, directly addressing which prior experimental discrepancy?
The discrepancy between theoretical and measured electron masses.
The observation that highly energetic alpha particles sometimes rebounded from the gold foil.
The discovery that atoms of the same element can have different masses.
The discrete, rather than continuous, spectrum of light emitted by energized hydrogen atoms.
5
Which statement accurately describes the relative scale established by the Rutherford model?
The nucleus occupies approximately 99% of the atom's volume.
The atom's mass is concentrated in a nucleus that is about $10^{-5}$ times the atom's total diameter.
Electrons are uniformly distributed throughout the positive sphere.
The total positive charge is distributed in discrete packets orbiting the central mass.
6
Which feature distinguishes the Quantum Mechanical Model (QMM) from the Bohr model?
Electrons follow fixed, circular orbits around the nucleus.
Energy levels are defined by precise, quantifiable radii.
Electron location is described by probability distributions (orbitals) rather than fixed paths.
The atom contains both positive protons and neutral neutrons.
7
In the historical sequence of atomic models, what transition marked the definitive shift from viewing electrons as discrete particles to viewing them primarily in terms of wave mechanics and probability?
Dalton to Thomson
Thomson to Rutherford
Rutherford to Bohr
Bohr to Quantum Mechanical Model
8
The term 'quantized' as used in the Bohr model means:
Energy levels are continuously variable based on the electron's velocity.
Electrons can only exist in certain fixed energy levels, not in between.
The angular momentum of the electron is always zero.
The positive charge of the nucleus is spread evenly across the atom.
9
If a chemist states that a sample of pure iron atoms are 'all identical in mass,' which scientist's model, if strictly followed, would support this statement?
Niels Bohr (due to fixed orbits)
J.J. Thomson (due to uniform positive charge)
John Dalton (due to no recognition of isotopes)
Ernest Rutherford (due to mass concentration in nucleus)
10
Which particle was the last to be discovered and incorporated into the standard atomic model?
Electron
Proton
Alpha particle
Neutron
11
Which subatomic particle is primarily responsible for determining the volume and chemical reactivity of a neutral atom?
Neutrons, due to their stabilizing influence.
Protons, due to their identity definition.
Electrons, due to their orbital configuration and range.
Nucleons (Protons and Neutrons), due to their mass.
12
Which statement accurately compares the masses of the fundamental subatomic particles?
Proton $\approx$ Neutron, both significantly greater than Electron.
Proton $<$ Neutron, both significantly less than Electron.
Proton $\approx$ Electron, both significantly less than Neutron.
All three particles have approximately equal mass.
13
What is the net electrical charge (in elementary charge units) of the nucleus of an element with $Z=15$?
$+15$
$-15$
$-1$
The charge depends on the number of neutrons.
14
The strong nuclear force acts primarily to counteract which other fundamental force within the nucleus?
Gravity, by keeping the neutrons bound.
Electromagnetic force, by overcoming proton-proton repulsion.
Weak nuclear force, stabilizing neutron decay.
Electromagnetic force, by binding electrons to the nucleus.
15
A particle discovered to be highly penetrating, electrically neutral, and possessing a mass similar to the proton is the:
Electron
Neutron
Proton
Positron
16
If an atom has 15 protons, 16 neutrons, and 18 electrons, what is the charge of the resulting ion?
$+3$
$-3$
$+1$
$-1$
17
What is the primary role of neutrons in the nucleus of most elements heavier than hydrogen?
To attract the orbiting electrons.
To increase the overall magnetic moment of the atom.
To provide the binding force required to overcome proton repulsion.
To facilitate the transfer of energy during chemical reactions.
18
A particle with a relative mass of approximately $\frac{1}{1836}$th that of a proton and a relative charge of $+1$ is known as a:
Neutrino
Positron
Electron
Proton
19
If an electrically neutral atom of Boron ($\text{B}$) suddenly gains two neutrons, how does the resulting particle's mass number ($A$) and overall charge change?
$A$ increases by 2; charge becomes $+2$
$A$ increases by 2; charge remains $0$
$A$ remains the same; charge remains $0$
$A$ increases by 1; charge becomes $+1$
20
The number of electrons present in the energy shells of an atom is defined by:
The Mass Number ($A$).
The combined count of Protons and Neutrons.
The number of protons ($Z$) if the atom is electrically neutral.
The difference between the Mass Number and the Atomic Number ($A-Z$).
21
How does the Mass Number ($A$) relate to the atomic mass reported on the periodic table?
$A$ is the weighted average of isotopic masses, identical to the reported atomic mass.
$A$ is the total mass of the atom in grams.
$A$ represents the total count of nucleons in a specific isotope, which approximates the mass of that isotope in amu.
$A$ is always double the Atomic Number ($Z$).
22
An atom of Potassium (K) has $Z=19$ and $A=39$. How many neutrons does this specific atom contain?
$19$
$39$
$20$
$58$
23
A specific ion $\text{X}^{2-}$ contains 54 electrons and 78 neutrons. Determine the atomic number ($Z$) and the mass number ($A$) of this ion.
$Z=54, A=132$
$Z=52, A=130$
$Z=56, A=134$
$Z=50, A=128$
24
The identity of a chemical element is uniquely determined by which single subatomic count?
The total number of nucleons.
The number of electrons.
The number of neutrons.
The number of protons.
25
Consider an element $\text{Y}$ with 26 protons and 30 neutrons. If $\text{Y}$ forms an ion with a $+3$ charge, how many electrons does the ion possess?
$26$
$23$
$30$
$29$
26
A cation $\text{M}^{+}$ has a Mass Number ($A$) of 65 and 35 neutrons. Which of the following is true about its neutral atom?
It has 31 electrons.
It has $Z=35$ and 30 electrons.
It has $Z=30$ and 30 electrons.
Its atomic mass is exactly 65.0 u.
27
If an atom of Boron ($\text{B}, Z=5$) loses an electron, which property remains absolutely unchanged?
Its charge.
Its chemical reactivity.
Its Atomic Number ($Z$).
Its atomic radius.
28
If element $\text{Q}$ has $Z=18$ and $A=38$, and element $\text{R}$ has $Z=17$ and $A=38$. How are these two elements related?
They are isotopes of the same element.
They are isobars (nuclides with the same mass number).
They are isomers (nuclides with the same composition but different energy states).
They are isotopes of different elements.
29
A neutral atom has a nucleus containing 40 nucleons, and 22 of them are neutrons. What is the correct nuclear symbol for this atom?
$\sideset{^{40}}{}{_{22}\text{Ne}}$
$\sideset{^{40}}{}{_{18}\text{Ar}}$
$\sideset{^{40}}{}{_{22}\text{Ti}}$
$\sideset{^{18}}{}{_{22}\text{X}}$
30
The term 'nucleon' is used to collectively describe which group of subatomic particles?
Protons and electrons.
Protons and neutrons.
Electrons and neutrons.
Only the particles responsible for the overall charge.
31
Isotopes are defined as atoms of the same element that differ in their number of:
Protons.
Electrons.
Neutrons.
Atomic Number ($Z$).
32
Which property would differ the MOST significantly between two common isotopes of the same element?
Ionization energy.
Reactivity with oxygen.
Density.
Melting point.
33
Why do different isotopes of the same element exhibit virtually identical chemical behavior?
They have the same atomic mass.
They have the same number of neutrons in the nucleus.
They have the same number of valence electrons.
Their nuclei are always equally stable.
34
The standard atomic mass unit (amu or u) is historically defined relative to the mass of which specific isotope?
A hydrogen atom ($^{1}\text{H}$).
One twelfth the mass of a Carbon-12 atom.
The average mass of a sample of naturally occurring oxygen.
The mass of a single proton.
35
Chlorine has two main isotopes: $\text{Cl-35}$ (mass 34.969 u) and $\text{Cl-37}$ (mass 36.966 u). If the natural abundance of $\text{Cl-35}$ is 75.77% and $\text{Cl-37}$ is 24.23%, calculate the Relative Atomic Mass (RAM).
$35.000\text{ u}$
$36.000\text{ u}$
$35.453\text{ u}$
$35.968\text{ u}$
36
Element $\text{Z}$ has two isotopes: $\text{Z-60}$ (50\% abundance) and $\text{Z-62}$ (50\% abundance). What is the calculated Relative Atomic Mass (RAM) of element $\text{Z}$?
$60.5\text{ u}$
$61.0\text{ u}$
$61.5\text{ u}$
$62.0\text{ u}$
37
The measured Relative Atomic Mass of Lithium is $6.941\text{ u}$. Lithium exists naturally as $\text{Li-6}$ and $\text{Li-7}$. What can be reliably inferred about their relative abundances?
$\text{Li-6}$ is far more abundant than $\text{Li-7}$.
The abundances are almost exactly 50% each.
$\text{Li-7}$ is far more abundant than $\text{Li-6}$.
The RAM cannot be used to infer abundance unless the exact mass defect is known.
38
If two atoms have the same Mass Number ($A$) but different Atomic Numbers ($Z$), they are classified as:
Isomers
Isotopes
Isobars
Isotones (same number of neutrons)
39
Which statement correctly explains why the Relative Atomic Mass listed on the periodic table is almost never a whole number?
It accounts for the negative mass of electrons.
It is a weighted average reflecting the natural abundance of all isotopes of that element.
It must compensate for the binding energy (mass defect) in the nucleus.
Scientists have not yet precisely measured the mass of stable isotopes.
40
A hypothetical element $\text{E}$ consists of $\text{E-100}$ (99.00% abundance) and $\text{E-105}$ (1.00% abundance). What is the calculated RAM of element $\text{E}$?
$100.05\text{ u}$
$100.50\text{ u}$
$102.50\text{ u}$
$100.99\text{ u}$
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