Inorganic Chemistry: Atomic Models, Quantum Structure, and Periodic Trends

Comprehensive flashcards for reviewing atomic history, quantum numbers, electron configurations, subatomic particles, and periodic trends.

What is a condensed electron configuration?

A shorter way of writing an atom's electron configuration by replacing the inner core electrons with the chemical symbol of the immediately preceding noble gas enclosed in brackets.

What is the maximum electron capacity and orbital count for each subshell ($s$, $p$, $d$, and $f$)?

The $s$ subshell has $1$ orbital and holds a maximum of $2$ electrons; $p$ has $3$ orbitals and holds $6$ electrons; $d$ has $5$ orbitals and holds $10$ electrons; $f$ has $7$ orbitals and holds $14$ electrons.

What are the expected versus actual condensed electron configurations of Chromium ($\text{Cr}$) and Copper ($\text{Cu}$)?

For Chromium ($\text{Cr}$, $Z = 24$), expected is $\text{[Ar]} 4\text{s}^2 3\text{d}^4$ and actual is $\text{[Ar]} 4\text{s}^1 3\text{d}^5$. For Copper ($\text{Cu}$, $Z = 29$), expected is $\text{[Ar]} 4\text{s}^2 3\text{d}^9$ and actual is $\text{[Ar]} 4\text{s}^1 3\text{d}^{10}$.

What is the actual condensed electron configuration of Palladium ($\text{Pd}$)?

The actual condensed electron configuration of Palladium ($\text{Pd}$, $Z = 46$) is $\text{[Kr]} 4\text{d}^{10}$, whereas its expected configuration is $\text{[Kr]} 5\text{s}^2 4\text{d}^8$.

How did Democritus and John Dalton describe the atom in early atomic theory?

Democritus (~400 BC) proposed that atoms were solid, uniform, and indestructible particles (from the Greek word "atomos"). John Dalton (1803) proposed that atoms are the fundamental building blocks of matter and that all atoms of a given element are identical.

What three major chemical laws formed the basis for John Dalton's atomic theory?

The Law of Conservation of Mass (Antoine Lavoisier), the Law of Definite Proportions / Constant Composition (Joseph Proust), and the Law of Multiple Proportions (proposed by John Dalton).

What states the Law of Multiple Proportions?

If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

What subatomic particle did J. J. Thomson discover and what atomic model did he propose?

J. J. Thomson discovered the electron using a Cathode Ray Tube in 1904 and proposed the Plum Pudding (or Raisin Bread) model of the atom.

What were the key conclusions of Lord Ernest Rutherford's Gold Foil Experiment in 1911?

Rutherford showed that atoms consist mostly of empty space with a small, dense, positively charged nucleus at the center, surrounded by moving electrons.

What method did Robert Millikan use in 1911 and what electron value did he determine?

Robert Millikan used the Oil Drop Method to measure the charge on an electron and determine the mass of an electron as $9.11 \times 10^{-28}\,\text{g}$ ($9.109328 \times 10^{-28}\,\text{g}$).

What is the Heisenberg Uncertainty Principle formulated by Werner Heisenberg in 1926?

It states that it is impossible to simultaneously know the exact position and momentum of an object (such as an electron).

What are the mass values and relative charges of protons, neutrons, and electrons as given in Table 2.1?

Proton ($p$): Mass = $1.672622 \times 10^{-24}\,\text{g}$ ($1.007276\,\text{amu}$), charge = $+1$. Neutron ($n$): Mass = $1.674927 \times 10^{-24}\,\text{g}$ ($1.008665\,\text{amu}$), charge = $0$. Electron ($e^-$): Mass = $9.109328 \times 10^{-28}\,\text{g}$ ($5.485799 \times 10^{-4}\,\text{amu}$), charge = $-1$.

How are the Atomic Number ($Z$), Mass Number ($A$), and Charge of an Ion mathematically calculated?

Atomic Number ($Z$) = number of protons. Mass Number ($A$) = number of protons + number of neutrons ($A = \text{neutrons} + Z$). Charge of an Ion = number of protons $$-$ number of electrons.

What are isotopes?

Atoms of the same element that have the same number of protons (same atomic number $Z$) but different numbers of neutrons, resulting in different mass numbers ($A$).

What is the difference between atomic mass and atomic weight?

Atomic mass is the mass of a single atom or specific isotope, whereas atomic weight is the weighted average mass of an element based on the relative abundance of all its naturally occurring isotopes.

How did Dimitri Mendeleev organize his periodic table in 1869?

Dimitri Mendeleev arranged chemical elements by increasing atomic mass, predicted undiscovered elements, and left open spaces for them in his periodic table.

What is the Periodic Law?

It states that the physical and chemical properties of elements are not random, but depend upon the structure of the atom and vary systematically with atomic number.

What is the Principal Quantum Number ($n$) and how many maximum electrons can a principal energy level contain?

The Principal Quantum Number ($n = 1, 2, 3, \dots, \infty$) specifies the main energy level and orbital size. The total number of orbitals is $n^2$, and each principal level can contain up to $2n^2$ electrons.

What does the Angular Momentum Quantum Number ($l$) describe and what letter codes correspond to its values?

It specifies the orbital shape and subshell ($l = 0, \dots, n-1$). The letter codes are: $l = 0 \rightarrow s$, $l = 1 \rightarrow p$, $l = 2 \rightarrow d$, $l = 3 \rightarrow f$, $l = 4 \rightarrow g$, and $l = 5 \rightarrow h$.

What is the Magnetic Quantum Number ($m_l$) and how many orbitals are present in each subshell?

$m_l$ specifies the spatial orientation of an orbital, taking integer values from $-l$ to $+l$. There are $2l + 1$ orbitals in each subshell.

What is the Spin Quantum Number ($m_s$)?

It specifies the orientation of the spin axis of an electron, which can spin in only one of two directions: $+\frac{1}{2}$ or $-\frac{1}{2}$.

What are the definitions of the Aufbau Principle, Pauli Exclusion Principle, and Hund's Rule?

Aufbau Principle: Electrons occupy the lowest energy level orbital available. Pauli Exclusion Principle: No two electrons in an atom can have identical values for all four quantum numbers. Hund's Rule: Electrons occupy equal-energy orbitals singly with parallel spins before pairing.

What is the magnetic difference between diamagnetic and paramagnetic substances?

Diamagnetic substances have all electron spins paired and are not attracted to magnets. Paramagnetic substances contain unpaired electrons and are weakly attracted to magnets.

What is Nuclear Charge versus Effective Nuclear Charge ($Z_{\text{eff}}$)?

Nuclear charge ($Z$) is the total positive charge of the nucleus (number of protons). Effective nuclear charge ($Z_{\text{eff}}$) is the net positive charge felt by an electron after subtracting core electron shielding ($Z_{\text{eff}} \approx Z - S$).

What are the periodic trends for atomic radius across a period and down a group?

Atomic radius decreases across a period (left to right) because nuclear charge and $Z_{\text{eff}}$ increase, pulling electrons closer. Atomic radius increases down a group (top to bottom) because additional energy levels are added and shielding increases.

How do ionic radii of cations and anions compare to their neutral parent atoms?

A cation is smaller than its neutral atom because losing electrons reduces electron repulsion and shrinks the cloud. An anion is larger than its neutral atom because gaining electrons increases electron-electron repulsion, expanding the cloud.

What is the trend in radius across the isoelectronic series $\text{O}^{2-}, \text{F}^-, \text{Ne}, \text{Na}^+, \text{Mg}^{2+}$?

Radius decreases as proton count increases: $\text{O}^{2-} > \text{F}^- > \text{Ne} > \text{Na}^+ > \text{Mg}^{2+}$. Higher nuclear charge pulls the same $10$ electrons closer.

What is First Ionization Energy and what are its general periodic trends?

First Ionization Energy is the minimum energy required to remove the first electron from a neutral gaseous atom ($\text{X}(g) \rightarrow \text{X}^+(g) + e^-$). It generally increases across a period and decreases down a group.

Why does Beryllium ($\text{Be}$) have a higher first ionization energy than Boron ($\text{B}$)?

Beryllium ($1\text{s}^2 2\text{s}^2$) has a filled $2\text{s}$ subshell, whereas Boron ($1\text{s}^2 2\text{s}^2 2\text{p}^1$) has its outermost electron in a higher-energy $2\text{p}$ orbital, which is easier to remove.

Why does Nitrogen ($\text{N}$) have a higher first ionization energy than Oxygen ($\text{O}$)?

Nitrogen ($2\text{p}^3$) has three singly occupied $2\text{p}$ orbitals. Oxygen ($2\text{p}^4$) has one paired set of electrons in a $2\text{p}$ orbital; electron-electron repulsion in that pair makes removing an electron easier.

Why does Chlorine ($\text{Cl}$) have a more favorable electron affinity than Fluorine ($\text{F}$)?

Fluorine's very small $2\text{p}$ orbital causes high electron-electron repulsion when an incoming electron is added, whereas Chlorine's larger $3\text{p}$ orbital accommodates the added electron with less repulsion.

What is electronegativity and which element is the most electronegative?

Electronegativity is the ability of an atom in a chemical bond to attract shared bonding electrons toward itself. Fluorine ($\text{F}$) is the most electronegative element.

How do metallic and nonmetallic character trend across the periodic table?

Metallic character decreases across a period and increases down a group (towards the lower-left). Nonmetallic character increases across a period and decreases down a group (towards the upper-right).

How do metal reactivity and nonmetal reactivity trend down a group?

Metal reactivity increases down a group because larger metals lose valence electrons more easily. Nonmetal reactivity decreases down a group (e.g., $\text{F} > \text{Cl} > \text{Br} > \text{I}$) because smaller nonmetals attract electrons more strongly.

What are the master periodic trends across a period and down a group?

Across a period: Atomic radius decreases; Ionic radius generally decreases; Ionization energy increases; Electron affinity becomes more favorable; Electronegativity increases; Metallic character decreases; Nonmetallic character increases. Down a group: Atomic radius increases; Ionic radius increases; Ionization energy decreases; Electron affinity becomes less favorable; Electronegativity decreases; Metallic character increases; Nonmetallic character decreases.