AP Chemistry: Composition of Mixtures, PES, Periodicity, and Mass Spectrometry - Vocabulary Flashcards

Vocabulary-style flashcards covering key concepts from composition of mixtures, elemental analysis, mass spectrometry, PES, electron configuration, and periodic trends as presented in the notes.

Pure substance

A substance composed of only one type of particle (which can be an atom, molecule, or formula unit). Every distinct sample of a pure substance will exhibit identical chemical composition and consistent physical and chemical properties under the same conditions.

Mixture

A combination of two or more different pure substances (elements or compounds) that are physically combined. The proportions of the individual components can vary, and they retain their original chemical identities within the mixture.

Elemental composition by mass

The proportion, by mass, of each elemental constituent within a chemical compound. It indicates how much of the compound's total mass is attributed to each specific element, often expressed as a percentage.

Mass percent

A quantitative measure that expresses the mass of a specific component (e.g., an element) as a percentage of the total mass of the sample or compound. It is calculated formulaically as: ( \frac{\text{Mass of element in sample}}{\text{Total mass of sample or compound}} ) \times 100\%

Percent composition by mass

The mass percentages of all individual elements present within a chemical compound. This information is crucial for determining the empirical formula of the compound, which represents the simplest whole-number ratio of atoms.

Empirical formula

The simplest positive whole-number ratio of atoms present in a compound. It is determined from experimental data, such as percent composition, and is the most reduced form of the chemical formula. For some compounds, it may be the same as the molecular formula, but for many, it is a simplified representation.

Molecular formula

The exact number of atoms of each element that are present in a molecule of a compound. It is always a whole-number multiple of the empirical formula (e.g., for glucose, the empirical formula is CH2O, while the molecular formula is C6H{12}O6). It provides the true stoichiometry of the molecule.

Law of definite proportions

A fundamental chemical law stating that in any given pure chemical compound, the elements are always combined in the same proportions by mass, regardless of the compound's source or method of preparation. This means that a specific compound always has a constant elemental composition.

Empirical formula from percent composition

To determine the empirical formula from mass percentages: first, assume a 100-gram sample to convert mass percentages directly into grams of each element. Next, convert the grams of each element into moles using their respective molar masses. Then, divide the number of moles of each element by the smallest number of moles calculated to obtain preliminary mole ratios. Finally, if these ratios are not whole numbers, multiply all ratios by the smallest integer that converts them into whole numbers, which represent the subscripts in the empirical formula.

Mass spectrometry (MS)

An analytical technique used to measure the mass-to-charge ratio (m/z) of ions. By ionizing a sample and separating the ions based on their m/z values, it can determine the molecular weight of compounds, identify unknown compounds, and ascertain the isotopic composition of elements.

Mass spectrum

A graph produced by a mass spectrometer, typically plotting the mass-to-charge ratio (m/z) on the x-axis against the relative abundance (intensity) of ions on the y-axis. Each peak in the spectrum corresponds to a specific ion, often representing different isotopes or fragments of the original molecule, with the height of the peak indicating its relative abundance.

Isotope

Atoms of the same chemical element that possess the same number of protons (and thus the same atomic number) but differ in their number of neutrons. This difference in neutron count results in variations in their atomic mass and, consequently, their mass-to-charge ratio in mass spectrometry.

Average atomic mass

The weighted average of the atomic masses of all naturally occurring isotopes of an element. The weighting factors are the natural abundances of each isotope. This value is typically what appears on the periodic table and reflects the average mass of a mole of atoms of that element.

m/z (mass-to-charge ratio)

The ratio of an ion's mass (m) to its charge (z), expressed as \frac{mass}{charge}. It is the fundamental quantity measured in mass spectrometry and is plotted on the x-axis of a mass spectrum to identify different ions.

Isoelectronic

Describes atoms or ions that possess the exact same number of electrons in their electron shells. For example, the Na^+ ion (10 electrons) is isoelectronic with the Ne atom (10 electrons) and the Mg^{2+} ion (10 electrons).

Photoelectron spectroscopy (PES)

An experimental technique that measures the kinetic energy of electrons ejected from a material when irradiated with X-rays or ultraviolet light. This allows for the determination of the binding energies of electrons within atoms or molecules, providing insights into their electron configurations and subshell energy levels.

Binding energy

The minimum amount of energy required to remove an electron from an atom or ion in its gaseous state. Higher binding energies indicate electrons that are more strongly attracted to the nucleus and therefore more difficult to remove.

Ground-state electron configuration

The unique arrangement of electrons within the orbitals of an atom that results in the lowest possible total energy for that atom. This configuration dictates the atom's chemical behavior and is built according to the Aufbau principle, Pauli exclusion principle, and Hund's rule.

Aufbau principle

A fundamental principle in quantum mechanics that states that electrons fill atomic orbitals of the lowest available energy levels before occupying higher energy levels. This systematic filling sequence determines the ground-state electron configuration of an atom.

Core electrons

Electrons located in the inner electron shells of an atom, closer to the nucleus, that are generally not involved in chemical bonding. These electrons effectively 'shield' the outer valence electrons from the full attractive force of the nucleus.

Valence electrons

The electrons located in the outermost principal energy level of an atom. These are the electrons primarily involved in forming chemical bonds and are responsible for determining an atom's chemical properties and reactivity.

Shielding

The phenomenon where inner core electrons reduce the effective nuclear charge experienced by the outer valence electrons. The repulsion from the core electrons diminishes the attractive force that the positively charged nucleus exerts on the negatively charged valence electrons.

Effective nuclear charge (Z_{eff})

The net positive charge experienced by an electron in a polyelectronic atom. It is less than the actual nuclear charge (Z) due to the shielding effect of inner-shell electrons (S). The formula is commonly approximated as Z_{eff} = Z - S, where Z is the atomic number (number of protons) and S is the number of core electrons.

Coulomb’s Law

A fundamental law in physics describing the electrostatic force between two charged particles. The magnitude of this force (F) is directly proportional to the product of the magnitudes of the charges (Q1 and Q2) and inversely proportional to the square of the distance (r) between their centers, mathematically expressed as F = k \frac{Q1 Q2}{r^2}. Consequently, a shorter distance or larger charges lead to a stronger attractive or repulsive force.

Ionization energy

The minimum energy required to remove one mole of electrons from one mole of gaseous atoms or ions in their ground state. Ionization energy generally increases across a period (due to increasing Z_{eff} and decreasing atomic radius) and decreases down a group (due to increasing shielding and atomic radius).

Atomic radius

A measure of the size of an atom, typically defined as half the distance between the nuclei of two identical atoms that are bonded together. Atomic radius generally decreases across a period (due to increasing Z_{eff}) and increases down a group (due to increasing principal energy levels and shielding).

Ionic radius

The size of an ion. When an atom loses electrons to form a cation, its ionic radius is smaller than that of its neutral parent atom because of fewer electrons and sometimes the loss of an entire electron shell. Conversely, when an atom gains electrons to form an anion, its ionic radius is larger than its neutral parent atom due to increased electron-electron repulsion and a lesser effective nuclear charge.

Electron affinity

The energy change that occurs when an electron is added to a neutral gaseous atom to form a negative ion. A highly negative electron affinity indicates that an atom readily accepts an electron and releases a significant amount of energy, typically observed in nonmetals.

Electronegativity

A measure of the ability of an atom in a chemical bond to attract a shared pair of electrons towards itself. It is a concept crucial for predicting bond polarity and generally increases across a period and decreases down a group.

Periodic trends

Systematic and predictable patterns in the chemical and physical properties of elements when arranged by atomic number in the periodic table. These trends arise from the periodic recurrence of similar electron configurations and the predictable changes in effective nuclear charge, shielding, and atomic size.

Reactivity trend (metals)

For metals, reactivity generally increases as you move down a group (due to easier loss of valence electrons as atomic radius increases and ionization energy decreases) and decreases as you move across a period (as Z_{eff} increases, holding valence electrons more tightly).

Reactivity trend (nonmetals)

For nonmetals, reactivity generally increases as you move across a period (due to increasing electronegativity and ease of gaining electrons) and decreases as you move down a group (due to decreasing electronegativity and larger atomic size making it harder to attract electrons).

Mass-to-charge ratio (m/z) in MS

In mass spectrometry, the mass divided by the charge of an ion (m/z). For ions with a single positive charge (z=1), the m/z value directly represents the approximate isotopic mass of the ion, allowing for the identification of different isotopes present in a sample.

Spectral peak height

In a mass spectrum, the height or intensity of a peak is directly proportional to the relative abundance of the ions detected at that specific mass-to-charge ratio. A taller peak signifies a greater quantity of that particular ion.

Isotope abundance in MS

The relative proportion of each isotope of an element within a sample, as determined by the heights (intensities) of the corresponding peaks in a mass spectrum. The tallest peak is typically designated as the base peak and assigned a relative abundance of 100%, with other peaks scaled accordingly.

Empirical formula steps poem

A mnemonic to recall the steps for calculating the empirical formula from percent composition: 1. Percent to mass: Assume a 100-gram sample so percentages become grams. 2. Mass to mole: Convert grams of each element to moles using molar masses. 3. Divide by the smallest: Divide all calculated mole values by the smallest number of moles obtained to find preliminary mole ratios. 4. Multiply to get a whole-number ratio: If ratios are not whole numbers, multiply all by the smallest integer to achieve whole-number subscripts for the empirical formula.

Molecular vs empirical formula relationship

The molecular formula of a compound is always a whole-number multiple (n) of its empirical formula. This multiple can be found by dividing the true molar mass of the compound by the molar mass of its empirical formula: n = \frac{\text{Molar Mass of Molecular Formula}}{\text{Molar Mass of Empirical Formula}}. Once 'n' is found, multiply all subscripts in the empirical formula by 'n' to obtain the molecular formula.

Noble gas configuration

An electron configuration that results in a completely filled outermost electron shell, typically with 8 valence electrons (or 2 for helium). This configuration, characteristic of noble gases (e.g., He, Ne, Ar, Kr, Xe, Rn), is extremely stable and unreactive due to its low energy state.

Mass percent example (CO2)

For carbon dioxide (CO2): The molar mass of C is 12.01 \text{ g/mol}, and O is 16.00 \text{ g/mol}. The molar mass of CO2 is 12.01 + 2(16.00) = 44.01 \text{ g/mol}. The mass percent of carbon is ( \frac{12.01 \text{ g/mol}}{44.01 \text{ g/mol}} ) \times 100\% \approx 27.29\%. The mass percent of oxygen is ( \frac{2 \times 16.00 \text{ g/mol}}{44.01 \text{ g/mol}} ) \times 100\% \approx 72.71\%.

Homogeneous vs heterogeneous

A homogeneous mixture has a uniform composition and properties throughout, meaning its components are evenly distributed and indistinguishable, like salt dissolved in water. A heterogeneous mixture has a nonuniform composition, where its components are visibly distinct and usually present in separate phases, like sand mixed with water or a salad.

Groups and periods

In the periodic table, a group refers to a vertical column of elements, where elements within the same group typically share similar chemical properties due to having the same number of valence electrons. A period refers to a horizontal row of elements, indicating that elements within the same period have their valence electrons in the same principal energy level.

Valence electrons in representative elements

For main-group (representative) elements, the number of valence electrons an atom possesses is typically equal to its group number (for groups 1, 2, and 13-18, often simplifying by dropping the '1' from 13-18). For example, elements in Group 1 have 1 valence electron, and elements in Group 17 have 7 valence electrons.

Charge predictability in ionic compounds

In ionic compounds, metals tend to lose electrons and form positive ions (cations) with charges equal to their group number (e.g., Group 1 metals form +1 ions). Nonmetals tend to gain electrons and form negative ions (anions) with charges predictable by their distance from the noble gases (e.g., Group 17 nonmetals form -1 ions). This behavior is driven by the desire to achieve a stable noble-gas electron configuration.

Mass spectrum peaks and isotopes

In a mass spectrum, each distinct peak corresponds to an ion of a specific mass-to-charge ratio, primarily representing a different isotope of an element or a molecular fragment. The value on the x-axis (m/z) indicates the mass of the isotope/fragment, while the height on the y-axis (relative abundance) indicates its natural abundance or prevalence in the sample.

Formula unit vs molecule

A formula unit represents the simplest ratio of ions in an ionic compound (e.g., NaCl). It does not exist as a discrete molecule but rather as part of a crystal lattice. A molecule describes a discrete, covalently bonded aggregate of two or more atoms (e.g., H2O, O2), which can exist independently.

Atomic mass unit (u)

A standard unit for expressing atomic and molecular masses. It is defined as exactly one-twelfth (1/12) the mass of a single carbon-12 atom, which has a mass of exactly 12 \text{ u}. One atomic mass unit is approximately 1.6605 \times 10^{-27} \text{ kg} and is roughly equivalent to the mass of a proton or neutron.

Stepping from empirical to molecular formula

To convert from an empirical formula to a molecular formula: 1. First, calculate the molar mass of the empirical formula. 2. Next, obtain the actual molar mass of the compound (usually given experimentally). 3. Divide the actual molar mass by the empirical formula's molar mass to find a whole-number factor (n). 4. Finally, multiply all the subscripts in the empirical formula by this factor (n) to derive the full molecular formula.

Practice example: empirical formula from mass percentages

Consider a compound with 40.0% C, 6.7% H, and 53.3% O. Assume a 100-g sample: 40.0 g C, 6.7 g H, 53.3 g O. Convert to moles: C (40.0\text{ g}/12.01 \text{ g/mol} \approx 3.33 \text{ mol}), H (6.7\text{ g}/1.008 \text{ g/mol} \approx 6.65 \text{ mol}), O (53.3\text{ g}/16.00 \text{ g/mol} \approx 3.33 \text{ mol}). Divide by smallest mole value (3.33 mol): C (3.33/3.33 = 1), H (6.65/3.33 \approx 2), O (3.33/3.33 = 1). The empirical formula is CH_2O.