Chapter 2: Matter at the Atomic Level Practice Flashcards

A comprehensive set of vocabulary flashcards based on the Chapter 2 lecture transcript covering atomic theory, subatomic particles, laws of chemistry, isotopes, and the periodic table structure.

Atomic Theory

A theory exploring the connection between atomic structures and the manipulation of matter; it posits that matter consists of atoms, which are the smallest identifiable units of matter and cannot be destroyed.

Early Atomic Theory Postulates

1. Matter consists of indestructible atoms. 2. All atoms of the same element are identical, but different from atoms of other elements. 3. Atoms combine in small whole-number ratios to form compounds.

Law of Conservation of Matter

The principle stating that matter can neither be created nor destroyed.

Law of Conservation of Mass-Energy

The principle stating that while the total mass and energy of a system is conserved, mass and energy can interconvert.

Law of Multiple Proportions

When an element combines with another to form multiple compounds, the ratio of the masses of one element to a fixed amount of the second element is a ratio of small whole numbers (e.g., oxygen in $H_2O$ vs. $H_2O_2$ has a $2:1$ mass ratio).

Law of Definite Proportions

Also known as the Law of Constant Composition, it states that a compound will always comprise its constituent elements in the same proportion by mass (e.g., water is always $11.2\%$ hydrogen and $88.8\%$ oxygen by mass).

J.J. Thomson

The scientist who discovered the negatively charged particle known as the electron in 1897.

Electron

A negatively charged subatomic particle with an actual mass of $9.109 \times 10^{331}\,\text{kg}$ and a relative charge of $-1$.

Millikan's Oil Drop Experiment

An experiment conducted in 1909 that determined the exact charge and mass of an electron.

Thomson’s Plum Pudding Model

A model of the atom consisting of a positively charged spherical volume of space with electrons dispersed throughout, resulting in an inherently neutral atom.

Ernest Rutherford’s Gold Foil Experiment

Experiment that determined every atom has a positively charged dense center called the nucleus, where essentially all of the atom's mass is found.

James Chadwick

The English physicist who discovered the neutron in 1932.

Proton

A subatomic particle located in the nucleus with an actual mass of $1.673 \times 10^{-27}\,\text{kg}$, a relative mass of approximately $1$, and a relative charge of $+1$.

Neutron

A subatomic particle with no charge and an actual mass of $1.675 \times 10^{-27}\,\text{kg}$, which is similar to the relative mass of a proton.

Atomic Number

The number of protons within an atom's nucleus; for neutral atoms, this also equals the number of electrons.

Ions

Charged versions of atoms that result when atoms gain or lose electrons.

Cations

Positively charged ions formed when atoms lose electrons (e.g., $Na \rightarrow Na^+ + e^-$).

Anions

Negatively charged ions formed when atoms gain electrons (e.g., $Br + e^- \rightarrow Br^-$).

Isotopes

Atoms of the same element that have the same number of protons but differ in their number of neutrons.

Atomic Mass Number

The sum of the number of protons and neutrons in an isotope: $\text{Atomic mass} = \#p + \#n$.

Percent Natural Abundance

A measure of the average amount an isotope naturally occurs relative to all of the element’s isotopes (e.g., Carbon-12 is $98.93\%$, while Carbon-13 is $1.07\%$).

Atomic Mass

The weighted average of all the masses of an element's naturally occurring isotopes, calculated as：$\text{Atomic Mass} = \sum(\text{isotope fractional abundance})(\text{exact isotope mass})$.

Lothar Meyer and Dmitri Mendeleev

German and Russian chemists credited with creating the earliest versions of the periodic table between 1864 and 1871.

Metals

Elements characterized as being malleable, ductile, lustrous, and good conductors; they have high melting points, are typically solids at room temperature (except $Hg$), and tend to form cations.

Nonmetals

Elements that tend to be brittle, dull, and insulators; they have lower densities and melting points, exist in various physical states at $25^{\circ}\text{C}$, and tend to form anions.

Metalloids

Elements with intermediate characteristics of metals and nonmetals often used as semiconductors; Silicon ($Si$) and Germanium ($Ge$) are key examples.

Alkali Metals

Group 1A elements that are soft, dull, grey-white metals and react violently with water to form $+1$ cations.

Alkaline Earth Metals

Group 2A elements that are harder than Group 1A, react less violently with water, and tend to form $+2$ cations.

Halogens

Very reactive elements that naturally occur as diatomic compounds.

Noble Gases

Naturally occurring colorless, monatomic gases that are relatively inert or unreactive.