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Study Guide - Smart Notes
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Fundamental Laws of Chemistry
Law of Conservation of Mass
The Law of Conservation of Mass states that mass cannot be created or destroyed in a chemical or physical process.
Antoine Lavoisier established this law.
During any reaction, the total mass of reactants equals the total mass of products.
Example: When hydrogen reacts with oxygen to form water, the combined mass of hydrogen and oxygen before the reaction equals the mass of water produced.
Law of Definite Proportions
The Law of Definite Proportions states that a given compound always contains the same proportion by mass of its constituent elements, regardless of its source or method of preparation.
Joseph Proust is credited with this law.
The mass fraction of an element in a compound is constant.
Example: Water (H2O) always contains 11.2% hydrogen and 88.8% oxygen by mass.
Law of Multiple Proportions
When two elements combine to 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.
John Dalton formulated this law.
Atoms combine in simple whole-number ratios when forming compounds.
Example: Carbon and oxygen form CO (carbon monoxide) and CO2 (carbon dioxide). The ratio of oxygen masses that combine with a fixed mass of carbon is 1:2.
Dalton's Atomic Theory
Dalton's Atomic Theory provided the first modern description of the nature of matter.
Each element is made up of tiny particles called atoms.
Atoms of a given element are identical in mass and properties.
Compounds are formed when atoms of different elements combine in fixed, simple whole-number ratios.
Chemical reactions involve the reorganization of atoms; atoms themselves are not changed in a chemical reaction.
Important Atom Characterization Experiments
J.J. Thomson's Cathode Ray Tube Experiment
Demonstrated the existence of negatively charged subatomic particles (electrons), common to all elements.
Measured the mass-to-charge ratio of the electron:
Proposed the "plum pudding model" of the atom, where electrons are embedded in a positively charged sphere.
Robert Millikan's Oil Drop Experiment
Determined the absolute charge of an electron.
Allowed calculation of the electron's mass using Thomson's ratio.
Ernest Rutherford's Gold Foil Experiment
Showed that the mass of the atom is concentrated in a small, dense region called the nucleus.
Concluded that atoms are mostly empty space.
Modern Atomic Theory
Structure of the Atom
Atoms are roughly spherical and electrically neutral.
Composed of three principal subatomic particles:
Protons (p+): Positively charged, located in the nucleus.
Neutrons (n0): Neutral, located in the nucleus.
Electrons (e-): Negatively charged, found outside the nucleus.
Atomic notation:
= atomic number = number of protons in the nucleus
= mass number = sum of protons and neutrons ()
For a neutral atom:
Isotopes
Atoms of the same element with different mass numbers (), due to different numbers of neutrons.
Isotopes have the same chemical and physical properties.
Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon.
Ions
Cations: Positively charged ions, formed by losing electrons (fewer electrons than protons).
Anions: Negatively charged ions, formed by gaining electrons (more electrons than protons).
Modern Atomic Theory (Updated)
Atoms are the smallest particle of matter with a unique identity.
Atoms of the same element have the same atomic number but can have different masses (isotopes).
Atoms of one element cannot be converted to atoms of another element by chemical reactions (only nuclear reactions can change elements).
Compounds are formed when atoms of different elements combine in whole-number ratios.
The Periodic Table and Periodic Law
Periodic Law
The Periodic Law states that when elements are arranged in order of increasing atomic mass (now atomic number), certain sets of properties recur periodically.
Classification of Elements
Metals: Good conductors of heat and electricity, malleable, ductile, and lustrous.
Non-metals: Poor conductors, varied physical properties.
Metalloids: Properties intermediate between metals and non-metals.
Groups of Elements
Alkali Metals: Group 1 (IA)
Alkaline Earth Metals: Group 2 (IIA)
Halogens: Group 17 (VIIA)
Noble Gases: Group 18 (VIIIA)
Transition Metals: Groups 3-12
Main Group Elements: Groups 1, 2, and 13-18
Lanthanides & Actinides: Inner transition metals
Ions and the Periodic Table
Main group metals tend to lose enough electrons to achieve the same number of electrons as the nearest noble gas (form cations).
Main group nonmetals tend to gain enough electrons to achieve the same number of electrons as the nearest noble gas (form anions).
Average Atomic Mass
The average atomic mass of an element is the weighted average of the masses of all naturally occurring isotopes, based on their relative abundances.
No atom of an element has exactly the average atomic mass, but for calculations, we use this value as a representative mass.
Example: The average atomic mass of chlorine is 35.45 u, reflecting the natural abundance of its isotopes.
Stoichiometry
Stoichiometry is the study of the quantitative relationships between the amounts of reactants and products in chemical reactions.
The mole is the SI unit for amount of substance, defined as the number of atoms in exactly 12 grams of carbon-12.
1 mole = entities (Avogadro's number).
1 atom = average mass (atomic mass unit, amu); 1 mole of atoms = average mass in grams (g).
Example: 1 mole of carbon-12 atoms has a mass of exactly 12 g and contains atoms.
