chem chapter 2-3

Law of constant composition

The principle stating that all samples of a given chemical compound contain the same elements in the exact same proportion by mass.

Law of conservation of mass

The principle stating that matter cannot be created or destroyed in an isolated system during a chemical reaction.

Law of multiple proportions

The principle stating that when two elements form multiple compounds, the ratio of the masses of the second element that combine with a fixed mass of the first is always a ratio of small whole numbers.

Atomic number

The value indicating the number of protons in an atom's nucleus, which uniquely defines the identity of an element.

Mass number

The total sum of protons and neutrons located inside an atom's nucleus.

Isotopes

Atoms of the same chemical element that contain the same number of protons but differing numbers of neutrons.

Cation

A positively charged ion produced when a neutral atom loses one or more electrons.

Anion

A negatively charged ion produced when a neutral atom gains one or more electrons.

Avogadro's number

The fundamental constant representing the count of particles in one mole of substance, equal to $6.022 \times 10^{23}\,\text{mol}^{-1}$.

Mole

The standard scientific unit for amount of substance, defined as containing exactly $6.022 \times 10^{23}$ elementary entities.

Covalent bond

A chemical linkage formed when pairs of electrons are shared between atoms.

Ionic bond

A chemical attraction resulting from the electrostatic forces between oppositely charged ions following electron transfer.

Empirical formula

A chemical formula expressing the simplest whole-number ratio of the elements in a compound.

Molecular formula

A chemical formula that specifies the actual number of each atom present in a single molecule of a compound.

Structural formula

A graphical representation that illustrates how atoms in a molecule are spatially arranged and bonded to one another.

Allotropes

Different structural modifications of the same element in the same physical state, such as $O_2$ and $O_3$.

Oxidation state

A number assigned to an atom in a substance indicating the hypothetical charge it would have if all bonds were completely ionic.

Isomers

Molecules that possess identical molecular formulas but differ in their structural arrangement or connectivity.

Functional group

A specific arrangement of atoms within an organic compound that determines its characteristic

Atomic nucleus contents

The dense central core of an atom containing protons and neutrons.

Electron motion in an atom

Electrons move in the space surrounding the central nucleus.

Divisibility of an atom

Atoms are divisible into smaller subatomic particles: protons, neutrons, and electrons.

Simplest atom

Hydrogen, consisting of a single proton and a single electron in its most common form.

Dalton's atomic theory

The scientific theory proposing that matter is composed of indivisible atoms, atoms of an element are identical, atoms cannot be created or destroyed, and compounds form by combining atoms in whole-number ratios.

Dalton's explanation for the law of conservation of mass

Chemical reactions involve the reorganization, combination, or separation of atoms, but atoms themselves are neither created nor destroyed.

Revision to Dalton's postulate on identical atoms

Atoms of the same element can have different masses due to the presence of isotopes with varying neutron counts.

Law of conservation of mass

The principle stating that matter is neither created nor destroyed during a chemical reaction, meaning total reactant mass equals total product mass.

Law of constant composition

The principle stating that all samples of a given chemical compound always contain the same elements combined in the exact same mass proportion (e.g., all water samples follow this).

Law of multiple proportions

The principle stating that when two elements combine to form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first are small whole numbers.

Primary organizing principle of the periodic table

Elements are arranged primarily in order of increasing atomic number.

Periodic table groups and periods

The periodic table contains $18$ vertical columns (groups) that share similar chemical properties and $7$ horizontal rows (periods).

Atomic number ($Z$)

The integer displayed on the periodic table that specifies the total number of protons in an atom's nucleus.

Mass number ($A$)

The total integer sum of protons and neutrons within an atom's nucleus.

Calculating neutron count

$$\text{Neutrons} = \text{Mass number} - \text{Atomic number}$$

Electron count in a neutral atom

Equal to the number of protons (atomic number).

Electron count in a cation

$$\text{Electrons} = \text{Number of protons} - \text{Positive charge magnitude}$$

Electron count in an anion

$$\text{Electrons} = \text{Number of protons} + \text{Negative charge magnitude}$$

Isotopes

Atoms of the same chemical element that have identical numbers of protons but differing numbers of neutrons.

Mass defect

The difference between the total mass of individual unbound nucleons and the actual mass of the nucleus, caused by the conversion of mass into nuclear binding energy.

Atomic mass unit (amu)

A standard unit of mass defined such that atomic masses are scaled close to their mass numbers, simplifying the comparison of atoms and isotopes.

Ionization

The process by which a neutral atom or molecule gains or loses electrons, converting into an electrically charged ion.

Motion of larger ions in deflection/stopping fields

Heavier or larger ions travel farther and require more time to come to a stop due to greater momentum and inertia.

Notation for ion charge

Written as a superscript to the upper right of the chemical symbol indicating magnitude and sign (e.g., $2+$, $3-$).

Cation

A positively charged ion formed when a neutral atom or group of atoms loses one or more electrons.

Anion

A negatively charged ion formed when a neutral atom or group of atoms gains one or more electrons.

Primary chemical purpose of the mole concept

To serve as a counting unit that links macroscopic laboratory masses to microscopic numbers of atoms and molecules.

Avogadro's number ($N_A$)

The constant $6.022 \times 10^{23}\,\text{mol}^{-1}$, denoting the number of constituent particles in one mole of substance.

Carbon-12 mole standard

The standard defining one mole as the number of atoms in exactly $12\,\text{g}$ of pure $^{12}\text{C}$.

Formula for calculating moles from mass

$$\text{Moles } (n) = \frac{\text{Mass } (m)}{\text{Molar mass } (M)}$$

Formula for calculating number of particles from moles

$$\text{Number of particles} = n \times N_A$$

Formula for the mass of a single atom

$$\text{Mass of single atom} = \frac{\text{Atomic mass}}{N_A}$$

Density formula

$$\text{Density} = \frac{\text{Mass}}{\text{Volume}}$$

Chemical compound

A pure substance composed of two or more distinct chemical elements held together by chemical bonds.

Covalent bond

A chemical linkage formed between nonmetal atoms through the sharing of electron pairs.

Ionic bond

An electrostatic attraction formed between oppositely charged ions following the transfer of valence electrons from a metal to a nonmetal.

Electron transfer tendencies in metals vs. nonmetals

Metals tend to lose electrons to form cations, whereas nonmetals tend to gain electrons to form anions.

Covalent bonding capacity of Carbon

Forms $4$ covalent bonds.

Covalent bonding capacity of Hydrogen and Halogens

Form $1$ covalent bond each.

Covalent bonding capacity of Oxygen

Forms $2$ covalent bonds.

Covalent bonding capacity of Nitrogen

Forms $3$ covalent bonds.

Structural formula

A graphic representation that illustrates the complete geometric connectivity and bond lines between every atom in a molecule.

Condensed structural formula

A formula that outlines the sequential arrangement of atom clusters without drawing explicit bond lines (e.g., $\text{CH}_3\text{COOH}$).

Molecular formula

A chemical notation showing the exact actual number and identity of each type of atom present in a molecule (e.g., $\text{C}<em>2\text{H}</em>4\text{O}_2$).

Empirical formula

A chemical formula expressing the simplest whole-number ratio of atoms of each element present in a compound.

Molecular mass vs. Formula mass

Molecular mass is the sum of atomic masses in a discrete molecule, whereas formula mass is the sum of atomic masses in the empirical formula unit of an ionic lattice.

Mass percent of an element in a compound

$$\text{Mass \%} = \frac{(\text{Number of atoms of element}) \times (\text{Molar mass of element})}{\text{Molar mass of compound}} \times 100\%$$

Allotropes

Different structural forms of the same element in the same physical state (e.g., $\text{O}<em>2$ and $\text{O}</em>3$).

Oxidation state

An assigned charge indicating the hypothetical charge an atom would possess if all shared electron pairs were completely assigned to the more electronegative atom.

Oxidation state of free elements

$0$ for an uncombined atom in its elemental state.

Sum of oxidation states rule

The sum of oxidation numbers must equal $0$ for a neutral molecule, or equal the net ionic charge for a polyatomic ion.

Standard oxidation state of Group 1 metals

$+1$ in all chemical compounds.

Standard oxidation state of Group 2 metals

$+2$ in all chemical compounds.

Standard oxidation state of Fluorine

$-1$ in all chemical compounds.

Standard oxidation state of Hydrogen

$+1$ when bonded to nonmetals.

Standard oxidation state of Oxygen

$-2$ in almost all typical compounds.

Binary compound

A chemical compound composed of exactly two distinct elements.

Ternary compound

A chemical compound composed of exactly three distinct elements.

Organic compound

A chemical compound containing carbon bonded to hydrogen, often alongside other heteroatoms such as oxygen or nitrogen.

Naming rule for transition metal cations

The element name followed by its positive ionic charge written in Roman numerals within parentheses.

Naming rule for monatomic anions

The base root of the nonmetal element name followed by the suffix "-ide".

Amine functional group

$$-\text{NH}_2$$

Alcohol (hydroxyl) functional group

$$-\text{OH}$$

Carboxyl functional group

$$-\text{COOH}$$

Methanoic acid

The simplest carboxylic acid, possessing the formula $\text{HCOOH}$.

Isomers

Molecules that have the identical molecular formula but differ in the sequence or spatial arrangement of their bonded atoms.