Chapter E Chemistry: Measurements, Density, Atomic Theory, Isotopes, and Mass Spectrometry

Vocabulary flashcards covering key concepts from the lecture notes.

Kilogram (kg)

SI unit of mass; 1 kg = 2.2 lb.

Meter (m)

SI unit of length; 1 m ≈ 39 inches.

Kelvin (K)

SI unit of temperature; relation to Celsius: K = C + 273.25.

Celsius (°C)

Temperature scale; C = (F − 32) / 1.8.

Fahrenheit (°F)

Temperature scale; used with conversion to Celsius.

Room Temperature

Notes list: 297 K or 75 °F.

Liter (L)

Unit of volume; about a quart.

Gallon

Unit of volume; under 4 liters (approximately 3.785 L for US gallon).

Kilo (k)

SI prefix for 10^3 (thousand).

Mega (M)

SI prefix for 10^6 (million).

Giga (G)

SI prefix for 10^9 (billion).

Tera (T)

SI prefix for 10^12 (trillion).

Peta (P)

SI prefix for 10^15.

Exa (E)

SI prefix for 10^18.

Zetta (Z)

SI prefix for 10^21.

Yotta (Y)

SI prefix for 10^24.

Nano (n)

SI prefix for 10^-9 (one-billionth).

Nonstandard Exponent Rule

If an exponent lacks a prefix, move the decimal to the nearest prefix set and then apply that prefix.

Conversion Factor

A ratio that converts one unit to another; defined to equal 1 in a correct setup.

Dimensional Analysis

Problem-solving method using conversion factors to track units across calculations.

Density (ρ)

Intensive physical property; mass per volume; d = m/v; independent of amount.

Density of Water vs Temperature

Water density increases as temperature falls to 4 °C; related to lake turnover.

1 mL = 1 cm^3

Equivalence between volume units.

Antoine Lavoisier

Father of Chemistry; proposed the Law of Conservation of Mass.

Phlogiston

Hypothetical substance thought to be released during burning.

Law of Conservation of Mass

Matter cannot be created or destroyed; mass of reactants equals mass of products.

Joseph Proust

Formulated the Law of Definite Proportions.

Law of Definite Proportions

A given compound has fixed ratios by mass of its elements.

John Dalton

Proposed the Law of Multiple Proportions and the Atomic Theory.

Law of Multiple Proportions

When two elements form two compounds, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios.

Atomic Theory (Dalton)

Elements are composed of atoms; atoms of a given element have same mass; atoms combine in simple whole-number ratios; atoms cannot change into other elements.

Electron

Negatively charged subatomic particle; discovered by J.J. Thomson via cathode rays; charge-to-mass ratio determined.

Millikan Oil Drop Experiment

Measured the charge of the electron by observing oil droplets in an electric field.

Proton

Positively charged subatomic particle; located in the nucleus; Rutherford’s work supported a nucleus.

Neutron

Uncharged subatomic particle; discovered by James Chadwick; accounts for missing mass.

Nuclear Theory (Rutherford)

Most mass and positive charge reside in the nucleus; atom is mostly empty space; electrons surround the nucleus; overall neutrality.

Proton Mass and Charge

Proton: 1.00727 amu, +1 charge.

Neutron Mass and Charge

Neutron: 1.00866 amu, 0 charge.

Electron Mass and Charge

Electron: 0.00055 amu, −1 charge.

Isotopes

Atoms with the same number of protons (Z) but different numbers of neutrons (A−Z).

Carbon Isotopes

Carbon-12, Carbon-13, Carbon-14 as examples.

Z and A

Z = atomic number (protons); A = atomic mass (protons + neutrons).

Natural Abundance

Percentage of naturally occurring isotopes of an element.

Neon Isotopes Abundances

Neon-20: 90.48%, Neon-21: 0.27%, Neon-22: 9.25%.

Mass Spectrometry

Technique to determine isotopic masses and their natural abundances by measuring mass-to-charge ratios.

Atomic Mass Unit (amu)

Unit used to express atomic and molecular masses.

Atomic Mass Calculation

Atomic Mass = sum of (amu × natural abundance %).