Atoms and Elements

Law of Conservation of Mass

The Law of Conservation of Mass states that the total mass remains constant during a chemical reaction. This principle is foundational in chemistry, ensuring that the mass of reactants equals the mass of products.

• Key Point: Mass of substances before the reaction = mass of substances after the reaction.

• Example: When sodium reacts with chlorine to form sodium chloride, the total mass of sodium and chlorine equals the mass of sodium chloride produced.

Law of Definite Proportions (Law of Constant Composition)

The Law of Definite Proportions states that all samples of a given compound, regardless of their source or preparation, have the same proportions of their constituent elements. This law is crucial for understanding chemical formulas and compound identity.

• Key Point: The ratio of elements in a compound is always fixed.

• Example: Water (H2O) always decomposes into 16.0 g of oxygen and 2.0 g of hydrogen for every 18.0 g of water, giving a mass ratio of oxygen to hydrogen as 8:1.

Law of Multiple Proportions

The Law of Multiple Proportions applies when two elements form more than one compound. The masses of one element that combine with a fixed mass of the other element can be expressed as ratios of small whole numbers.

• Key Point: The ratios of masses are simple whole numbers.

• Example: Carbon dioxide (CO2) and carbon monoxide (CO) have mass ratios of oxygen to carbon as 2.67:1 and 1.33:1, respectively. The ratio between these two is 2:1.

Dalton’s Atomic Theory

Dalton’s Atomic Theory provided the first modern description of atoms and their behavior in chemical reactions.

• Key Point 1: Each element is composed of tiny particles called atoms.

• Key Point 2: All atoms of a given element have the same mass and properties.

• Key Point 3: Atoms combine in simple, whole-number ratios to form compounds.

• Key Point 4: In chemical reactions, atoms change the way they are bound together, but are not created or destroyed.

Discovery of Electrons

J.J. Thomson discovered the electron through experiments with cathode rays, determining the charge-to-mass ratio of the electron.

• Key Point: Charge-to-mass ratio of electron:

Millikan’s Oil Drop Experiment

Millikan determined the mass and charge of a single electron using the oil drop experiment.

• Electron mass:

• Electron charge:

Rutherford’s Gold Foil Experiment and the Nuclear Atom

Rutherford’s experiment led to the discovery of the atomic nucleus, disproving the Plum Pudding model and establishing the nuclear model of the atom.

• Key Point: Atoms consist of a dense nucleus (protons and neutrons) surrounded by an electron cloud.

Subatomic Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons. Atomic mass units (amu) are used to express their masses.

• Proton (p+): Positively charged, found in nucleus.

• Neutron (n0): Neutral, found in nucleus.

• Electron (e-): Negatively charged, found in electron cloud.

• 1 amu: Defined as the mass of a carbon atom with 6 protons and 6 neutrons.

Atomic Number (Z) and Chemical Symbols

The atomic number (Z) is the number of protons in the nucleus and identifies the element. In a neutral atom, the number of electrons equals the number of protons.

• Key Point: Atomic number determines element identity and electron count in neutral atoms.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons. The mass number (A) is the sum of protons and neutrons and must be a whole number. Many elements have naturally occurring isotopes with varying natural abundances.

• Example: Neon has three isotopes: Neon-20 (90.48%), Neon-21 (0.27%), Neon-22 (9.25%).

Ions

Atoms can lose or gain electrons during chemical changes, forming ions. Cations are positively charged (loss of electrons), and anions are negatively charged (gain of electrons).

• Example: Na (Z = 11) forms Na+ (cation); Cl (Z = 17) forms Cl- (anion).

Periodic Law and the Periodic Table

The Periodic Law states that when elements are arranged by increasing atomic number, elements with similar properties occur at regular intervals. The periodic table organizes elements into groups (vertical columns) and periods (horizontal rows).

• Group 1A: Alkali metals

• Group 2A: Alkaline earth metals

• Group 7A: Halogens

• Group 8A: Noble gases

Atomic Mass and Isotopic Abundance

Atomic mass is a weighted average of the masses of all isotopes of an element, based on their natural abundance.

• Example: Chlorine-35 (34.97 amu, 75.77%), Chlorine-37 (36.97 amu, 24.26%)

The Mole and Avogadro’s Number

The mole (mol) is the chemist’s counting unit, representing particles. The value of the mole is defined as the number of atoms in exactly 12 grams of carbon-12.

• Key Point: particles

• Conversion: C atoms

Molar Mass

An element’s molar mass in grams per mole is numerically equal to its atomic mass in atomic mass units (amu). This allows conversion between mass and number of moles.

• Key Point: Molar mass (g/mol) = Atomic mass (amu)

Quantitative Relationships: Moles, Mass, and Atoms

Quantitative relationships allow conversion between moles, grams, and number of atoms using molar mass and Avogadro’s number.

• Formula:

• Formula:

Example Calculations

Examples include converting between moles and atoms, grams and moles, and atoms and grams for various elements.

• Example: Calculate the number of lead (Pb) atoms in 6.03 mol of lead.

• Example: Calculate the mass of one copper (Cu) atom in nanograms.

• Example: How many gold (Au) atoms are in a gold bar of given dimensions and density?

Summary Table: Subatomic Particles

Additional info: These notes cover foundational concepts from General Chemistry Chapter 2 (Atoms & Elements), including laws of mass conservation, atomic theory, isotopes, ions, and quantitative relationships. All images included are directly relevant to the explanation of the adjacent paragraph.