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Questions and Answers
What is the maximum value of l in an atom?
Which letter designates the orbital with l = 1?
How many orbitals are present in the d subshell?
What is the formula to calculate the total number of orbitals in a shell?
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Which magnetic quantum number values are possible for a p orbital?
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In the 3d subshell, what is the value of l?
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How many individual orbitals does an s subshell contain?
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Which subshell corresponds to an l value of 3?
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What is the energy of one photon emitted by a laser with a frequency of $4.69 \times 10^{14} , s^{-1}$?
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If a laser emits a pulse of energy containing $5.0 \times 10^{17}$ photons, what is the total energy of that pulse?
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How many photons are emitted during a pulse if the laser emits $1.3 \times 10^{-2} , J$ of energy?
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What characterizes a continuous spectrum?
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Which of the following statements about line spectra is correct?
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Which color in the hydrogen line spectrum has the shortest wavelength?
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What occurs when light from a source is separated into its different wavelength components?
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What happens when a high voltage is applied to a gas at low pressure?
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How many subshells are present in the shell with principal quantum number n = 3?
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What is the maximum number of orbitals in a subshell designated as 4d?
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What is the total number of orbitals in the shell with principal quantum number n = 2?
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Which of the following statements about s orbitals is true?
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For n = 4, which designation of subshell has the largest number of orbitals?
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Which of the following is the correct relationship for the total number of orbitals and the elements in the periodic table?
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How many possible values of ml exist for the 2p subshell?
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What is the principal quantum number for the subshell designated as 3p?
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What happens to the orbit radius of an electron in a hydrogen atom as the principal quantum number n increases?
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What is the energy of an electron when n reaches infinity?
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Which equation represents the change in energy of an electron as it transitions between two energy states?
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What limitation of the Bohr model is highlighted regarding its applicability to atoms other than hydrogen?
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What assumption did Bohr make to explain why electrons do not fall into the nucleus?
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In the equation derived from Bohr's theory to calculate energy transitions, what does the term Ephoton represent?
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Which constant is used in the formula to express the energies of the allowed orbits in the hydrogen atom?
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What does the term ΔE = hν = -hc(1/n²_f - 1/n²_i) describe?
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What is the primary role of outermost electrons in an atom?
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How does electron affinity differ from ionization energy?
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What characterizes an ionic bond?
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In covalent bonding, which of the following is true?
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What do Lewis symbols represent?
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Which of these statements regarding metallic bonds is correct?
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What type of bond typically forms between metals and nonmetals?
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Which statement about inner electrons is correct?
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Study Notes
Understanding the Bohr Model
- Bohr calculated the energies corresponding to each allowed orbit for the electron in the hydrogen atom.
- Each orbit corresponds to a different value of n, and the radius of the orbit gets larger as n increases.
- The radius increases as n^2, reaching a point where the electron is completely separated from the nucleus.
- When n = ∞, the energy is zero:
- E = (-2.18 × 10-18 J) (∞^2) = 0
- When an electron transitions from a state of energy Ei to a final state of energy Ef, the change in energy is:
- ΔE = Ef - Ei = Ephoton = hν
- The equation derived from Bohr's theory corresponds to the Rydberg equation obtained using experimental data:
- (1/λ) = -RH (1/n^2f - 1/n^2i)
Limitations of the Bohr Model
- The Bohr model explains the line spectrum of the hydrogen atom, but it cannot explain the spectra of other atoms.
- Bohr avoided the problem of why the negatively charged electron would not fall into the positively charged nucleus by simply assuming it would not happen.
- Bohr assumed that an electron is located at a definite distance from the nucleus and is revolving around it with a definite velocity, i.e. associated with a fixed value of momentum.
Understanding Quantum Numbers
- The principal quantum number, n, describes the electron shell, which specifies the energy level of the electron.
- The angular momentum or azimuthal quantum number, l, describes the shape of the electron's orbital.
- The value of l for a particular sub-level is generally designated by the letters s, p, d, and f, corresponding to l values of 0, 1, 2, and 3, respectively.
- l typically does not go beyond 3 in atoms.
- The magnetic quantum number, m1, describes the orientation of the orbital in space.
- For a sub-level with quantum number l, there are 2l+1 individual orbitals.
- They are distinguished by the magnetic quantum number, ml, which can have the 2l +1 integer values from -l down to +l.
- The spin quantum number, ms, describes the intrinsic angular momentum of an electron.
- It is an intrinsic property of the electron, not a consequence of its motion.
- The direction of the intrinsic angular momentum is spin up or down, giving a value of +1/2 or -1/2, respectively.
Understanding Electron Shells and Subshells
- The collection of orbitals with the same value of n is called an electron shell.
- All the orbitals that have n = 3, for example, are said to be in the third shell.
- The set of orbitals that have the same n and l values is called a subshell.
- Each subshell is designated by a number (the value of n) and a letter s, p, d or f (corresponding to the value of l).
- For example, the orbitals that have n = 3 and l = 2 are called 3d orbitals and are in the 3d subshell.
Electron Configuration Rules
- Aufbau Principle: Electrons are added one at a time to the lowest-energy orbitals available.
- Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.
-
Hund's Rule: Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied.
- All electrons in singly occupied orbitals have the same spin (e.g., all spins up).
Atomic Orbitals
- The s orbitals are spherically symmetric.
- The p orbitals are dumbbell-shaped, with a node at the nucleus.
- The d orbitals have more complex shapes, with two nodes.
- The f orbitals have even more complex shapes.
Predicting the Quantum Number Combinations
- The restrictions on the possible values of the quantum numbers give rise to the following important observations:
- The shell with principal quantum number n will consist of exactly n subshells.
- Each subshell consists of a specific number of orbitals. Each orbital corresponds to a different allowed value of ml.
- The total number of orbitals in a shell is n^2, where n is the principal quantum number of the shell.
Chemical Bond Types
- Ionic bonds: Electrostatic forces between ions of opposite charge.
- Covalent bonds: Sharing of electrons between two atoms.
- Metallic bonds: Bonding electrons are relatively free to move throughout the three-dimensional structure of the metal.
Lewis Symbols and the Octet Rule
- Valence electrons are the outermost electrons in an atom.
- Lewis symbols represent the valence electrons of an atom, and the chemical symbol plus a dot for each valence electron.
- The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons (like a noble gas). This rule is a guideline, not a hard and fast rule.
Electron Affinity
- Electron affinity measures how easily an atom gains an electron.
- Most atoms release energy when an electron is added.
- Ionization energy measures how easily an atom loses an electron.
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Description
This quiz delves into the Bohr model, focusing on the energy levels of electrons in hydrogen atoms and the limitations of Bohr's theory. Explore how the radius and energy of orbits are calculated and understand the significance of the transitions between energy states. Test your knowledge on the Rydberg equation and the model's applicability to different atoms.
