Matter and Atomic Structure — Free Chemistry Review Games.

Atoms are made of protons (positive), neutrons (neutral), and electrons (negative).

Protons and neutrons are found in the dense, central nucleus of an atom.

The atomic number equals the number of protons in the nucleus and defines the element.

Electrons carry a negative charge and orbit the nucleus in energy levels.

An atom is the smallest unit of an element that maintains the chemical identity of that element.

Isotopes are variants of an element that have the same number of protons but different numbers of neutrons.

Oxygen has 8 electrons: 2 in the 1s orbital, 2 in the 2s orbital, and 4 in the 2p orbitals.

The mass number is the sum of protons and neutrons in the nucleus of an atom.

Rutherford discovered that most of an atom is empty space with a small, dense, positive nucleus at the center.

An ion is an atom with a net electric charge due to the loss or gain of one or more electrons.

The uncertainty principle states that it is impossible to precisely determine both the position and momentum of an electron at the same time.

The maximum number of electrons in energy level n is 2n^2, so for n=3: 2(9) = 18 electrons.

The Bohr model depicts electrons in fixed circular orbits, while the quantum model describes electron positions as probability distributions (orbitals).

The Aufbau principle states that electrons occupy the lowest available energy orbitals before filling higher ones.

The number of protons (atomic number) defines an element's identity; changing the proton count transforms it into a different element.

A proton carries a positive charge of \(+1\) elementary charge, which is balanced by the negative charge of electrons in a neutral atom. The distractor 'Negative' describes the electron, not the proton, since electrons and protons have opposite charges. Remembering that protons are positive and located in the nucleus is essential for understanding atomic structure and ion formation.

A neutron has no net electrical charge, which is why it is called 'neutral' and does not affect the overall charge of an atom. The distractor 'Positive' is incorrect because that charge belongs to protons, which reside alongside neutrons in the nucleus. Since neutrons add mass without affecting charge, they are key to explaining isotopic variation without changing an element's chemical identity.

The electron has a mass of about \(9.11 \times 10^{-31}\) kg, roughly 1/1836th the mass of a proton, making it by far the lightest subatomic particle. The distractor 'Proton' is incorrect because protons are nearly 1800 times heavier than electrons despite carrying an equal and opposite charge. This mass difference explains why nearly all of an atom's mass is concentrated in the nucleus rather than in the electron cloud.

One amu is defined as approximately the mass of a single proton or neutron, since both particles have nearly identical mass of about \(1.67 \times 10^{-24}\) g. The distractor 'The mass of one electron' is wrong because electrons are about 1836 times lighter and contribute negligibly to atomic mass. Using amu as a unit lets chemists compare relative atomic masses without dealing with extremely small numbers in grams.

The plum pudding model described the atom as a positively charged sphere with negatively charged electrons embedded throughout it, like fruit in a pudding, with no concentrated nucleus. The distractor 'All mass is concentrated in a tiny, dense nucleus' actually describes the later nuclear model that replaced this idea after further experimentation. This early model illustrates how scientific understanding of atomic structure evolved through successive experimental discoveries.

The symbol \(Z\) denotes the atomic number, which equals the number of protons in an atom's nucleus and defines the element's identity. The distractor 'Mass number' is incorrect because mass number, often denoted \(A\), includes both protons and neutrons rather than protons alone. Recognizing that \(Z\) is fixed for each element is fundamental to reading the periodic table correctly.

Valence electrons occupy the outermost occupied energy level of an atom and are primarily responsible for chemical bonding and reactivity. The distractor 'An electron located in the nucleus' is wrong because electrons never reside in the nucleus, which contains only protons and neutrons. Understanding valence electrons is critical for predicting how atoms interact to form molecules and compounds.

The first energy level contains only one subshell, the \(1s\) orbital, which can hold a maximum of 2 electrons according to the Pauli exclusion principle. The distractor '8' actually applies to the second energy level, which includes both \(2s\) and \(2p\) subshells. Knowing the electron capacity of each energy level helps students correctly build electron configurations.

The electron cloud model represents electrons as existing within a probability region, or cloud, around the nucleus rather than following exact, predictable paths. The distractor 'Electrons traveling in fixed circular orbits' describes the earlier Bohr model, which was later replaced due to its inability to explain more complex atomic behavior. This modern model reflects the quantum mechanical understanding that electron position can only be described in terms of probability.

All isotopes of an element share the same number of protons, since this number defines the element's atomic number and chemical identity. The distractor 'The number of neutrons' is incorrect because isotopes are defined by differing neutron counts while protons remain constant. This distinction is key to understanding why isotopes have different masses but identical chemical properties.

The number of protons in an atom's nucleus, known as the atomic number, uniquely determines which element the atom represents. The distractor 'Neutron' is incorrect because varying neutron numbers only create isotopes of the same element rather than different elements. This principle explains why changing proton number through nuclear reactions actually transforms one element into another.

An s orbital has a spherical shape, meaning the probability of finding an electron is equally distributed in all directions at a given distance from the nucleus. The distractor 'Dumbbell-shaped' actually describes the shape of a p orbital, which has two lobes on opposite sides of the nucleus. Recognizing orbital shapes helps students visualize how electron density is distributed within an atom.

An orbital is defined as the region of space around the nucleus where there is a high probability of locating an electron, based on quantum mechanical calculations. The distractor 'Shell number' refers to the principal energy level, which is a broader concept encompassing multiple orbitals rather than a specific probability region. Orbitals are the building blocks used to construct electron configurations for every element.

Nitrogen has 7 electrons, which fill the \(1s\) and \(2s\) subshells completely before placing 3 electrons in the \(2p\) subshell, giving \(1s^2 2s^2 2p^3\). The distractor \(1s^2 2s^2 2p^5\) represents fluorine, which has 9 electrons rather than 7. Building configurations by filling lower-energy subshells first according to the Aufbau principle ensures the correct ground-state arrangement.

Chlorine's valence electrons are those in the outermost energy level, \(n=3\), which includes \(3s^2\) and \(3p^5\), totaling 7 electrons. The distractor '17' is the total number of electrons in the entire atom, not just those in the outer shell responsible for bonding. Counting only the outermost shell electrons is essential for predicting an element's chemical reactivity and bonding behavior.

Hund's rule states that electrons occupy degenerate orbitals, such as the three \(p\) orbitals, one at a time with parallel spins before any orbital receives a second electron. The distractor 'Electrons always pair up in the lowest energy orbital first' contradicts this rule since pairing before singly filling all degenerate orbitals increases electron-electron repulsion. This rule explains the stability of half-filled subshells and correctly predicts orbital diagrams for multi-electron atoms.

The Pauli exclusion principle states that within a single atom, no two electrons can share an identical set of all four quantum numbers, which limits each orbital to a maximum of two electrons with opposite spins. The distractor 'All electrons in an atom must have parallel spins' is false because paired electrons in the same orbital must have opposite, not parallel, spins. This principle underlies the maximum electron capacities of orbitals, subshells, and energy levels throughout the periodic table.

Sodium has 11 electrons, and using the shorthand notation, the first 10 electrons match neon's configuration, leaving one additional electron in the \(3s\) subshell, giving \([Ne]3s^1\). The distractor \([Ar]4s^1\) actually represents potassium, which has 19 electrons rather than 11. Noble gas shorthand simplifies writing configurations by referencing the nearest preceding noble gas core.

The angular momentum quantum number, \(l\), determines the shape of an orbital, distinguishing between s, p, d, and f subshells with their characteristic geometries. The distractor 'Principal quantum number (\(n\))' instead indicates the overall size and energy level of the orbital, not its shape. Understanding each quantum number's role helps students interpret the structure of electron configurations and orbital diagrams.

The spin quantum number, \(m_s\), can take values of \(+\frac{1}{2}\) or \(-\frac{1}{2}\) and describes the intrinsic spin orientation of an electron within an orbital. The distractor 'Magnetic quantum number (\(m_l\))' instead specifies the orientation of the orbital in space, not the electron's spin direction. Spin quantum number differences allow two electrons to occupy the same orbital without violating the Pauli exclusion principle.

A p subshell consists of three degenerate orbitals, each holding a maximum of 2 electrons, resulting in a total capacity of 6 electrons for the entire subshell. The distractor '10' actually corresponds to the capacity of a d subshell, which contains five orbitals rather than three. Knowing subshell capacities is essential for correctly writing electron configurations across the periodic table.

An excited-state configuration occurs when one or more electrons occupy higher-energy orbitals than the lowest available ones dictated by the Aufbau principle, often after absorbing energy. The distractor 'The total number of electrons has changed' is incorrect because excitation does not add or remove electrons, only relocates them to different energy levels. This concept explains phenomena like atomic emission spectra, where excited electrons release energy as light when returning to the ground state.

A neutral atom has an equal number of protons and electrons, which balances the positive and negative charges so the atom carries no overall net charge. The distractor 'Protons always outnumber electrons' describes a cation, which forms only after an atom loses electrons and becomes positively charged. This proton-electron balance is the defining feature that separates neutral atoms from ions.

An extra electron relative to protons gives the ion one more negative charge than positive charge, resulting in a net charge of \(1^-\), forming an anion. The distractor '1+ (positive one)' would instead describe a cation, which forms when an atom loses an electron rather than gains one. Determining ionic charge by comparing proton and electron counts is a core skill for predicting ion formation and compound formulas.

The number of neutrons equals the mass number minus the atomic number, so \(23 - 11 = 12\) neutrons for this atom. The distractor '11' is actually the atomic number, representing the proton count, not the neutron count. This calculation, \(A - Z = \text{neutrons}\), is a fundamental formula for determining nuclear composition from mass and atomic numbers.

Calcium has 20 electrons that fill in order of increasing energy, completing through \(3p^6\) and then placing the final two electrons in the \(4s\) subshell before the \(3d\) subshell begins filling, giving \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2\). The distractor \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^2\) is incorrect because the \(4s\) subshell fills before \(3d\) according to the Aufbau principle's energy ordering. This filling order reflects the relative energies of subshells rather than strictly following increasing principal quantum number.

Electrons occupy separate orbitals of equal energy singly first because this arrangement minimizes electrostatic repulsion between negatively charged electrons, resulting in a lower-energy, more stable configuration, as described by Hund's rule. The distractor 'Paired electrons cannot exist in the same subshell under any circumstance' is false since orbitals do eventually pair once all degenerate orbitals contain one electron each. This repulsion-minimizing behavior explains observed magnetic properties and stability trends across the periodic table.

Degenerate orbitals are orbitals within the same subshell, such as the three p orbitals, that share identical energy levels under normal conditions. The distractor 'An orbital that has lost an electron' confuses orbital occupancy with the concept of orbital energy equivalence, which is unrelated to electron count. Understanding degeneracy is necessary for correctly applying Hund's rule when constructing orbital diagrams.

Due to overlapping energy levels, the \(3d\) subshell fills after \(4s\) despite having a lower principal quantum number, because \(4s\) is slightly lower in energy than \(3d\) for the relevant elements. The distractor '\(4p\)' is incorrect because the \(4p\) subshell fills only after \(3d\) is completely occupied, not immediately after \(4s\). This exception to the simple \(n\)-based filling order explains the placement of transition metals on the periodic table.

The principal quantum number \(n\) indicates the main energy level of an electron and correlates with its average distance from the nucleus, with higher values of \(n\) representing greater energy and distance. The distractor 'The specific shape of the orbital the electron occupies' is instead determined by the angular momentum quantum number \(l\), not \(n\). Recognizing that \(n\) governs overall energy level size helps students organize electron configurations by shell.

A d subshell contains five degenerate orbitals, each capable of holding two electrons, giving the subshell a total capacity of 10 electrons. The distractor '3' actually corresponds to the number of orbitals within a p subshell, not a d subshell. Knowing the number of orbitals per subshell type is necessary for correctly determining maximum electron capacities.

Since mass number represents the total count of protons and neutrons combined, subtracting the atomic number, which equals the proton count, isolates the neutron count as \(A - Z\). The distractor 'Neutrons equal the atomic number alone' confuses neutron count with proton count, which are only equal for a specific subset of lighter, stable isotopes. This relationship is essential for identifying isotopes and calculating nuclear composition from standard atomic notation.

Promoting one electron from \(4s\) to \(3d\) creates a half-filled \(3d^5\) subshell, which has extra stability due to the symmetrical distribution of electrons and reduced electron-electron repulsion, lowering overall energy. The distractor 'The \(4s\) subshell is unable to hold more than one electron in transition metals' is false because \(4s\) can hold two electrons in most other elements, showing this behavior is an exception unique to specific stability effects. This anomaly illustrates that subshell stability, not strict energy-level ordering alone, can influence real electron configurations.

Moving one electron from \(4s\) into \(3d\) completes a fully filled, highly symmetrical \(3d^{10}\) subshell, which is energetically favorable compared to a nearly full \(3d^9\) configuration, resulting in enhanced stability. The distractor 'The \(3d\) subshell can only ever hold nine electrons under normal conditions' is incorrect since \(3d\) can and does hold up to ten electrons once fully occupied. This exception, alongside chromium's, demonstrates that fully or half-filled subshells provide extra electronic stability beyond simple orbital filling rules.

The quantum mechanical model describes electron locations using probability distributions, or orbitals, since Heisenberg's uncertainty principle shows that an electron's exact position and momentum cannot both be precisely known, unlike the Bohr model's fixed circular orbits. The distractor 'It eliminates the concept of energy levels entirely' is wrong because energy levels, described by the principal quantum number, remain a central feature of the quantum mechanical model. This shift from deterministic paths to probabilistic regions represents a major conceptual advancement in understanding atomic structure.

The angular momentum quantum number \(l\) can only take integer values from \(0\) to \(n-1\), so for \(n=2\), the maximum allowed value of \(l\) is \(1\), making \(l=2\) invalid. The distractor 'The magnetic quantum number cannot be zero' is incorrect because \(m_l\) can legitimately range from \(-l\) to \(+l\), including zero. Applying the quantum number rules correctly is essential for verifying valid electron states within an atom.

For a \(d\) orbital, the angular momentum quantum number is \(l=2\), so the magnetic quantum number \(m_l\) ranges from \(-l\) to \(+l\), giving the five values \(-2, -1, 0, 1, 2\), corresponding to the five d orbitals. The distractor '\(-1, 0, 1\)' corresponds instead to a p orbital where \(l=1\), not a d orbital. Correctly deriving \(m_l\) values from \(l\) helps explain why d subshells contain five orbitals capable of holding ten electrons total.

Electron energy within an atom is quantized, meaning electrons can only occupy specific discrete energy levels rather than any arbitrary value, a principle confirmed by atomic emission and absorption spectra showing distinct spectral lines. The distractor 'Energy levels are evenly spaced and electrons fill them continuously' is incorrect because quantized energy levels are not evenly spaced and electrons jump discretely between them rather than filling continuously. This quantization explains why atoms absorb and emit light only at specific, characteristic wavelengths.

For \(n=2\), allowed values of \(l\) are 0 and 1, corresponding to one \(2s\) orbital holding 2 electrons and three \(2p\) orbitals holding 6 electrons, which together give a maximum of 8 electrons using the formula \(2n^2\). The distractor 'It contains two \(2s\) orbitals and two \(2p\) orbitals, giving 8 electrons' is incorrect because there is only one \(s\) orbital per energy level, not two. This subshell-based reasoning, rather than memorization alone, allows students to derive electron capacities for any energy level.

When excited electrons fall from higher energy orbitals back to lower ones, they emit photons whose energy exactly matches the quantized energy difference between the two levels, producing distinct spectral lines rather than a continuous spectrum. The distractor 'Electrons continuously lose energy while orbiting, producing a smooth spectrum of colors' contradicts the quantized nature of atomic energy levels, which produces discrete lines instead. This connection between electron transitions and emitted photon energy is the basis for identifying elements through spectroscopy.

Alkali metals all have a single valence electron in their outermost \(s\) orbital, which is relatively easy to remove because doing so leaves the atom with a stable, full noble gas configuration underneath, explaining the consistent \(1+\) charge across the group. The distractor 'They all share the identical total number of electrons' is false because each alkali metal has a different total electron count corresponding to its unique atomic number. This pattern demonstrates how valence electron configuration, not total electron count, governs an element's characteristic chemical behavior.

After the \(3d\) subshell becomes populated, the relative energies shift so that \(3d\) electrons drop below \(4s\) in energy, making the \(4s\) electrons higher in energy and easier to remove first when the atom ionizes. The distractor 'Transition metals never actually lose electrons from the \(4s\) orbital' contradicts experimental evidence showing that ions like \(Fe^{2+}\) form by losing both \(4s\) electrons before any \(3d\) electrons. This reversal of filling versus removal order is a key nuance that distinguishes transition metal ion formation from simple Aufbau-based predictions.

The Bohr model successfully explained hydrogen's simple spectral lines but failed to account for the fine splitting of these lines observed in multi-electron atoms or under applied magnetic fields, a phenomenon later explained by the quantum mechanical model's inclusion of electron spin and orbital shape via additional quantum numbers. The distractor 'The overall existence of a positively charged nucleus' was already established by earlier nuclear model experiments and is unrelated to spectral line splitting. This limitation highlights why the more complex quantum mechanical model, incorporating multiple quantum numbers, was necessary to fully describe atomic behavior beyond hydrogen.

Electron probability density represents the statistical likelihood of locating an electron within a particular region of space, reflecting the inherent uncertainty in electron position rather than a single, deterministic path around the nucleus. The distractor 'Probability density guarantees the electron's exact location at every instant' contradicts the very nature of probability density, which describes likelihood rather than certainty. This probabilistic framework, grounded in the Heisenberg uncertainty principle, replaced the outdated notion of electrons following fixed, classical orbits.

Sodium and potassium both have a single valence electron in their outermost \(s\) orbital, \(3s^1\) and \(4s^1\) respectively, which governs their similar tendency to lose one electron and react similarly despite having different total numbers of electrons. The distractor 'Their electron configurations are completely identical overall' is false because their full configurations differ significantly in total electrons and energy levels occupied. This principle explains why elements within the same group of the periodic table share similar chemical properties based on matching valence electron arrangements.