Chemistry Test Chapter 3 Flashcards

Atomic Structure and History

• An atom is the smallest particle of an element that retains its identity in chemical reactions.
• Democritus first proposed the existence of atoms, believing they were indivisible and indestructible.
• The term "atamos" means indivisible; it reflects the idea that atoms cannot be divided.
• Democritus's ideas were limited as they did not explain chemical behavior and lacked experimental support.

Early Chemical Principles

• By the 1700s, two key truths were accepted: elements cannot be broken down by chemical means, and compounds formed from elements exhibit different properties.
• Controversy existed about whether elements combine in consistent ratios when forming compounds.

Laws of Chemical Combination

• The law of conservation of mass states that mass is neither created nor destroyed in ordinary chemical reactions.
• The law of definite proportions indicates that a chemical compound contains the same elements in the same mass proportions, regardless of sample size.
• The law of multiple proportions states that in different compounds made from the same elements, the ratios of the masses of one element that combine with a fixed mass of the other are small whole numbers.

Dalton’s Atomic Theory

• John Dalton further developed atomic theory using the scientific method, proposing that all elements consist of atoms and that atoms combine in whole number ratios to form compounds.
• Dalton's theory includes five key points:
  • Elements consist of atoms.
  • Atoms of the same element are identical.
  • Atoms cannot be divided, created, or destroyed.
  • Different elements can mix or combine in simple ratios.
  • Chemical reactions involve rearranging atoms.

Atomic Components

• An atom features a small nucleus at its center, containing protons and neutrons.
• Protons are positively charged particles located in the nucleus, having equal mass to neutrons, which are neutral.
• Electrons are negatively charged and orbit the nucleus in a relatively greater volume.

Key Discoveries in Atomic Theory

• J.J. Thomson discovered the electron using cathode rays, concluding they consist of negatively charged particles.
• Eugen Goldstein identified the proton by observing rays in a cathode ray tube moving opposite to electrons.
• James Chadwick discovered the neutron.
• Robert Millikan determined the charge of an electron and helped establish the mass of a proton.

Thomson and Rutherford Models

• Thomson's plum-pudding model depicted electrons embedded in a positive charge but was eventually replaced due to Rutherford's findings.
• Rutherford’s experiment revealed that mass and positive charge are concentrated in the nucleus, with most of the atom being empty space.

Atomic Identity and Structure

• Atomic number indicates the number of protons and determines element identity; it is a whole number, while atomic mass can be a decimal.
• Isotopes are variants of elements with the same number of protons but different numbers of neutrons, resulting in different atomic masses.
• Isotopes remain chemically alike since they share identical proton and electron counts.

Atomic Symbols and Notation

• A nuclear symbol includes the mass number (superscript) and atomic number (subscript).
• Hyphen notation represents isotopes by their element name followed by the mass number.

Ions and Compounds

• An ion is formed when an atom gains or loses electrons, resulting in a negatively charged anion (gains electrons) or a positively charged cation (loses electrons).
• Coefficients indicate the number of molecules in a compound, while subscripts specify the number of atoms present.

Calculating Atomic Mass and Moles

• Average atomic mass is calculated by considering the mass and natural abundance of isotopes with the formula: mass = neutrons + protons.
• Atomic mass unit (amu) is defined as 1/12th the mass of a carbon-12 atom, and it is used to compare the relative masses of atoms.
• Avogadro's number defines the number of particles in one mole of a substance, and molar mass refers to the mass of one mole of a pure substance.

Historical Context of Chemistry

• By 1790, chemists focused on quantitative analysis of chemical reactions, emphasizing measurement in chemical processes.

This quiz focuses on key concepts from Chapter 3 of Chemistry, including the definition of atoms and the historical context provided by Democritus. Test your knowledge on the fundamental principles of atomic theory and its origins with these flashcards.