Atomic structure & nomenclature (Chapter 2) - Vocabulary Flashcards

Vocabulary-style flashcards covering the key terms, definitions, and concepts from the lecture notes on atomic structure, the periodic table, and naming conventions for ionic and molecular compounds.

Atomic structure

The arrangement of subatomic particles (protons, neutrons, electrons) in an atom; nucleus contains protons and neutrons; electrons occupy surrounding space.

Subatomic particles

Protons and neutrons in the nucleus; electrons surrounding the nucleus; protons are positively charged, electrons negatively charged, neutrons neutral.

Proton

Positively charged subatomic particle located in the nucleus; mass ≈ 1 amu; charge +1.

Neutron

Neutral subatomic particle located in the nucleus; mass ≈ that of a proton; no electric charge.

Electron

Negatively charged subatomic particle surrounding the nucleus; very small mass relative to protons and neutrons.

Nucleus

Dense center of the atom containing protons and neutrons; accounts for most of the atom’s mass; atom is mostly empty space.

Atomic number (Z)

The number of protons in the nucleus; defines the identity of the element.

Mass number (A)

The total number of protons and neutrons in the nucleus.

Isotopes

Atoms of the same element with different numbers of neutrons; same atomic number (Z) but different mass numbers (A).

Atomic symbol notation (A X Z)

Notation with mass number as a superscript and atomic number as a subscript around the element symbol.

Dalton’s atomic theory

Elements are composed of atoms; atoms of the same element are identical; atoms are not created or destroyed in reactions; compounds form when atoms combine.

Law of multiple proportions

Different compounds can be formed from the same elements in small whole-number ratios; different ratios yield different compounds.

Molecule

An assembly of two or more nonmetal atoms bonded together with distinct properties.

Molecular formula

The actual number of atoms of each element in a molecule.

Empirical formula

The simplest whole-number ratio of atoms in a compound.

Covalent bond

A chemical bond formed by sharing electrons between atoms.

Ionic bond

A bond formed by electrostatic attraction between oppositely charged ions in an ionic compound.

Ion

An atom or group that has gained or lost electrons, resulting in a net electrical charge.

Cation

Positively charged ion produced by the loss of electrons.

Anion

Negatively charged ion produced by the gain of electrons.

Polyatomic ion

A charged group of two or more atoms that acts as a single ion (e.g., NH4+, NO3-, CO3^2-, SO4^2-).

Hydroxide ion (OH-)

A common polyatomic anion consisting of oxygen and hydrogen with a −1 charge.

Oxyanion

Anions containing oxygen; naming patterns include ite/ate (and hypo/per for fewer/more oxygens).

Prefixes for binary molecular compounds

Mono-, di-, tri-, tetra-, penta-, etc.; used to indicate numbers of atoms; do not use mono for the first element.

Binary molecular naming rule

The element closer to the metals is named first; the second element ends with -ide.

Acids (general definition)

Compounds that dissolve in water to produce hydrogen ions (H+) and a corresponding anion.

Binary acids

Acids consisting of hydrogen and a nonmetal; named hydro- + root of the anion with -ide replaced by -ic, followed by 'acid'.

Oxyacids naming rule

If the oxyanion ends in -ite, the acid ends in -ous; if it ends in -ate, the acid ends in -ic; hydrogen may be added to form the acid.

Acids: examples of oxyanions

Nitrite NO2- becomes nitrous acid (HNO2); Nitrate NO3- becomes nitric acid (HNO3); similarly for other oxyanions.

Periodic table groups (common names)

Alkali metals (Group 1A), alkaline earth metals (2A), halogens, noble gases, transition metals.

Elements in groups have similar properties

Elements within the same group/family on the periodic table tend to show similar chemical behavior.

Molar composition vs. elemental composition (contextual)

Atoms form compounds by combining in ratios; ionic compounds are not discrete molecules and are represented by empirical formulas.

Empirical vs. molecular formulas (revisited)

Empirical formula shows the simplest whole-number ratio of elements; molecular formula shows the actual numbers of atoms in a molecule.

Ion formation trend (general)

Metals tend to form cations by losing electrons; nonmetals tend to form anions by gaining electrons; noble gas electron configuration is a driving factor.

Polyatomic ions (examples)

Common polyatomic ions include ammonium NH4+, nitrate NO3-, chlorate ClO3-, carbonate CO3^2-, sulfate SO4^2-.

Hydrogen oxyacids naming by oxyanion

Hydrogen addition can form hydrogen oxyanions; relate to oxyanion base names (e.g., NO3- to nitric acid HNO3).