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Atoms & Molecules

Introduction

This study guide covers the foundational laws of chemistry, the development of atomic theory, and key experiments that led to our modern understanding of atomic structure. These concepts are essential for understanding the behavior of matter at the atomic and molecular levels.

Fundamental Laws of Chemistry

Law of Conservation of Mass

The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction. This principle was established by Antoine Lavoisier and is fundamental to all chemical processes.

• Definition: The total mass of reactants equals the total mass of products in a chemical reaction.

• Example: The reaction of glucose with oxygen to form carbon dioxide and water:

Glucose + Oxygen → Carbon Dioxide + Water 180 g + 192 g → 264 g + 108 g 372 g before = 372 g after

Law of Definite Proportions

The Law of Definite Proportions (or Law of Constant Composition), proposed by Joseph Proust, states that a given compound always contains the same proportion of elements by mass, regardless of its source.

• Definition: The composition of a compound is fixed and does not vary.

• Example: Calcium carbonate (CaCO3) from marble, coral, or chalk always contains the same mass ratio of calcium, carbon, and oxygen.

Element	Mass in 20.0 g sample	Mass Fraction	% by Mass
Ca	8.0 g	0.40	40%
C	2.4 g	0.12	12%
O	9.6 g	0.48	48%
Total	20.0 g	1.00	100%

Application Example: Mass Fraction Calculation

Given a sample of pitchblende (an oxide ore of uranium) with a total mass of 84.2 g containing 71.4 g of uranium, the mass fraction of uranium is:

$\text{mass fraction} = \frac{\text{mass of U in #1}}{\text{total mass of #1}} = \frac{71.4\ g}{84.2\ g} = 0.8479$

To find the mass of uranium in a second sample (total mass = 102 kg):

$\text{mass of U in #2} = \text{total mass #2} \times \text{mass fraction} = 102\ kg \times 0.8479 = 86.5\ kg$

Law of Multiple Proportions

The Law of Multiple Proportions, formulated by John Dalton, states that 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.

• Definition: The ratios of the masses of the second element that combine with 1 gram of the first element can be reduced to small whole numbers.

• Example: Nitrogen and oxygen form several compounds. The mass of nitrogen that combines with 1 g of oxygen in three compounds is:

Compound	Mass of N with 1 g O
A	1.750 g
B	0.8750 g
C	0.4375 g

Calculating ratios:

$\frac{A}{B} = \frac{1.750}{0.8750} = 2$ $\frac{B}{C} = \frac{0.8750}{0.4375} = 2$ $\frac{A}{C} = \frac{1.750}{0.4375} = 4$

These are small whole numbers, illustrating the law.

Dalton's Atomic Theory

John Dalton's atomic theory provided a scientific explanation for the fundamental laws of chemistry. The main postulates are:

• Each element is composed of tiny, indivisible particles called atoms.

• All atoms of a given element are identical in mass and properties.

• Atoms of different elements differ in mass and properties.

• Chemical compounds are formed when atoms of different elements combine in simple whole-number ratios.

• Chemical reactions involve the reorganization of atoms; atoms themselves are not changed in a chemical reaction.

Characterizing the Atom: Key Experiments

Several classic experiments established the structure of the atom and the existence of subatomic particles.

Thomson's Cathode Ray Tube Experiment

• Demonstrated the existence of negatively charged subatomic particles, later called electrons.

• Measured the charge-to-mass ratio of the electron.

• Showed that atoms must also contain positive charges to balance the negative charge of electrons.

Experimental setup: High voltage is applied to a partially evacuated glass tube with metal electrodes, producing a stream of negative particles (electrons) from the cathode to the anode.

Millikan's Oil Drop Experiment

• Determined the absolute charge of a single electron.

• Allowed calculation of the electron's mass when combined with Thomson's charge-to-mass ratio.

Rutherford's Gold Foil Experiment

• Discovered the nucleus of the atom: a small, dense, positively charged center.

• Showed that most of the atom is empty space, with electrons orbiting the nucleus at a distance.

Structure of the Atom

Atoms are electrically neutral, roughly spherical, and composed of three principal subatomic particles:

• Electrons (e-): Negatively charged, found outside the nucleus.

• Protons (p+): Positively charged, found in the nucleus. The number of protons defines the atomic number (Z).

• Neutrons (n0): No charge, found in the nucleus, similar in mass to protons.

The nucleus is extremely small and dense, containing nearly all the atom's mass.

Atomic Notation

• Element symbol: E

• Mass number (A): Total number of protons and neutrons ($A = p^+ + n^0$)

• Atomic number (Z): Number of protons ($Z = p^+$)

• Number of neutrons: $n^0 = A - Z$

Isotopes

• Atoms of the same element (same Z) with different numbers of neutrons (different A).

• Isotopes have nearly identical chemical properties but different masses.

• Most elements exist as mixtures of isotopes.

Example: Carbon-12 ($^{12}$C), Carbon-13 ($^{13}$C), and Carbon-14 ($^{14}$C) are isotopes of carbon.

Ions

• Atoms that gain or lose electrons become charged particles called ions.

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

Modern Atomic Theory (Dalton Updated)

• Atoms are divisible into subatomic particles (protons, neutrons, electrons).

• Atoms are the smallest unit of matter with a unique identity.

• Atoms of one element cannot be converted to atoms of another by chemical reactions (but can by nuclear reactions).

• Atoms of an element have the same number of protons but may have different numbers of neutrons (isotopes).

• Compounds are formed by chemical combination of two or more elements.

The Periodic Table

The periodic table arranges elements in order of increasing atomic number, revealing periodic trends in their properties.

• Elements are grouped into metals and nonmetals.

• Groups (columns) have special names (e.g., Alkali metals, Alkaline earth metals).

• Elements in the same group have similar chemical properties.

Predicting Ion Charges

• Group 1A metals: Tend to lose one electron to form +1 cations (e.g., Na+).

• Group 2A metals: Tend to lose two electrons to form +2 cations (e.g., Ca2+).

• Nonmetals: Tend to gain electrons to achieve the electron configuration of the nearest noble gas, forming anions (e.g., Cl-).

Additional info: The periodic table's structure allows prediction of element properties, reactivity, and common ion charges based on group and period trends.