Measurement, Physical and Chemical Change, and Atomic Theory: General Chemistry Study Notes

Chapter 1: Measurement, Physical and Chemical Change

Chemical Changes and Physical and Chemical Properties

Chemistry seeks to understand the properties and behavior of matter by studying atoms and molecules. Matter can undergo both physical and chemical changes, each with distinct characteristics.

• Physical property: A characteristic that can be observed or measured without changing the substance's chemical composition (e.g., boiling point, melting point, density).

• Chemical property: Describes how a substance reacts with other substances, resulting in a change in chemical composition (e.g., flammability, reactivity with oxygen).

• Physical change: A change that alters the state or appearance of matter without changing its composition (e.g., boiling water: H2O(l) → H2O(g)).

• Chemical change: A change that alters the composition of matter, resulting in the formation of new substances (e.g., rusting of iron: Fe + O2 → Fe2O3).

• Example: Boiling water is a physical change; rusting iron is a chemical change.

Energy: A Fundamental Part of Physical and Chemical Change

Physical and chemical changes are often accompanied by energy changes. The total energy of an object is the sum of its kinetic and potential energy.

• Potential energy: Energy due to position or composition.

  • Gravitational

  • Elastic

  • Chemical (stored in chemical bonds)

• Kinetic energy: Energy due to motion.

  • Movement of atoms, molecules, or electrons

  • Heat

  • Mechanical

• Law of Conservation of Energy: Energy is neither created nor destroyed, only transferred or transformed.

• Systems with high potential energy tend to change in a way that lowers their potential energy, often releasing energy in the process (e.g., combustion of hydrocarbons).

Measurement in Chemistry

The Units of Measurement

Scientific measurements use the International System of Units (SI units) to ensure consistency and accuracy.

• Base SI Units:

  • Length: meter (m)

  • Mass: kilogram (kg)

  • Time: second (s)

  • Temperature: kelvin (K)

  • Amount of substance: mole (mol)

  • Electric current: ampere (A)

  • Luminous intensity: candela (cd)

• SI Prefixes: Used to express very large or very small quantities. For example:

  • kilo- (k): $10^3$

  • centi- (c): $10^{-2}$

  • milli- (m): $10^{-3}$

  • micro- (μ): $10^{-6}$

  • nano- (n): $10^{-9}$

• Derived Units: Formed by combining base units (e.g., volume: $m^3$, density: $kg/m^3$).

Significant Figures and Scientific Notation

The precision of a measurement is reflected in the number of significant figures reported.

• Significant figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Rules for significant figures:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Scientific notation: Used to express very large or small numbers and clarify significant figures (e.g., $2.00 \times 10^3$ has three significant figures).

• Calculations:

  • Multiplication/division: Result has the same number of significant figures as the measurement with the fewest significant figures.

  • Addition/subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

Accuracy and Precision

• Accuracy: How close a measured value is to the true value.

• Precision: How close a series of measurements are to each other.

Solving Chemical Problems: Dimensional Analysis

Dimensional analysis is a method for converting between units and solving problems in chemistry.

• Use conversion factors to ensure units cancel appropriately.

• Set up calculations so that unwanted units cancel, leaving the desired unit.

• Example: Converting 1 meter to inches:

  • $1 \text{ m} \times \frac{100 \text{ cm}}{1 \text{ m}} \times \frac{1 \text{ in}}{2.54 \text{ cm}} = 39.37 \text{ in}$

Chapter 2: Atoms and Elements

Early Ideas About the Building Blocks of Matter

The concept of atoms as the fundamental building blocks of matter has evolved over centuries, influenced by scientific discoveries and the development of the scientific method.

• Key contributors: Leucippus, Dalton, Copernicus, Bacon, Newton, and others.

• Development of atomic theory and the laws of chemistry.

Modern Atomic Theory and the Laws That Led to It

• Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction.

• Law of Definite Proportions: All samples of a given compound have the same proportions of their constituent elements.

  • Example: Water (H2O) always has a mass ratio of 8:1 for oxygen to hydrogen.

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

• Dalton's Atomic Theory:

  • Elements are composed of tiny, indestructible particles called atoms.

  • All atoms of a given element have the same mass and properties.

  • Atoms combine in simple, whole-number ratios to form compounds.

  • Atoms of one element cannot change into atoms of another element in a chemical reaction.

Atomic Structure

Discoveries in the late 19th and early 20th centuries revealed the internal structure of the atom.

• Electron: Discovered by J.J. Thomson using the cathode ray tube; has a negative charge and a very small mass.

• Millikan's Oil Drop Experiment: Determined the charge of the electron ($-1.6 \times 10^{-19}$ C).

• Plum-Pudding Model: Early model of the atom with electrons embedded in a positively charged sphere.

• Rutherford's Gold Foil Experiment: Showed that atoms have a small, dense, positively charged nucleus.

• Subatomic particles:

  • Proton: Positive charge, mass ≈ 1 amu, located in the nucleus.

  • Neutron: No charge, mass ≈ 1 amu, located in the nucleus.

  • Electron: Negative charge, very small mass, located outside the nucleus.

Particle	Mass (kg)	Mass (u)	Charge (C)
Proton	1.67262 × 10-27	1.00727	+1.60 × 10-19
Neutron	1.67493 × 10-27	1.00866	0
Electron	9.10938 × 10-31	0.00054858	-1.60 × 10-19

Isotopes and Ions

• Isotopes: Atoms of the same element (same number of protons) but different numbers of neutrons.

• Atomic number (Z): Number of protons in the nucleus.

• Mass number (A): Total number of protons and neutrons.

• Ions: Atoms that have gained or lost electrons.

  • Cation: Positively charged ion (loss of electrons).

  • Anion: Negatively charged ion (gain of electrons).

Atomic Mass and the Mole

• Atomic mass: The weighted average mass of an element's isotopes.

• Calculation: $\text{Atomic mass} = \sum_i (\text{fractional abundance})_i \times (\text{isotope mass})_i$

• The mole (mol): The amount of substance containing as many entities as there are atoms in 12 g of carbon-12 ($6.022 \times 10^{23}$, Avogadro's number).

• Molar mass: The mass of one mole of a substance, numerically equal to its atomic or molecular mass in grams.

• Example: $1 \text{ mol C} = 12.01 \text{ g}$; $1 \text{ mol He} = 4.003 \text{ g}$

The Periodic Table

Development and Structure of the Periodic Table

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Elements are arranged in rows (periods) and columns (groups or families).

• Elements in the same group have similar chemical properties.

• Mendeleev's periodic table left gaps for undiscovered elements, predicting their properties.

Classification of Elements

Category	Properties	Examples
Alkali metals (Group 1)	Very reactive, react explosively with water	Li, Na, K
Alkaline earth metals (Group 2)	Fairly reactive, react with water but less vigorously	Mg, Ca
Halogens (Group 17)	Very reactive nonmetals	F, Cl, Br
Noble gases (Group 18)	Very unreactive gases	He, Ne, Ar

• Metals tend to lose electrons to form cations.

• Nonmetals tend to gain electrons to form anions.

Origin of Elements

Elements were formed in the early universe and in stars through nuclear fusion and supernova explosions, creating all naturally occurring elements up to uranium.

Additional info: Some context and examples were inferred and expanded for clarity and completeness, especially in the explanation of atomic structure, isotopes, and the periodic table.