Chapter 4

Democritus

Proposed that matter is made up of tiny, indivisible particles called "atomos" (from Greek meaning "indivisible").

Aristotle

Rejected the idea of atoms and promoted the belief that matter was composed of the four classical elements: earth, water, air, and fire.

Dalton's Atomic Theory

Key Ideas: Matter is composed of indivisible particles called atoms; All atoms of a given element are identical in size, mass, and chemical properties; Atoms of different elements have different properties; Atoms combine in simple whole-number ratios to form compounds; In chemical reactions, atoms are rearranged, separated, or combined but are not created or destroyed.

Law of Conservation of Mass

Dalton's theory explained that chemical reactions involve the rearrangement of atoms, consistent with the Law of Conservation of Mass.

Limitations of Dalton's Theory

Atoms are Divisible: Later discoveries showed that atoms are made of subatomic particles (electrons, protons, neutrons); Isotopes: Atoms of the same element can have different masses, which contradicts Dalton's notion of identical atoms.

Size of Atoms

Atoms are extremely small, with a diameter around 1.28 × 10^-10 meters for a copper atom.

Visualization of Atoms

If an atom were as large as an orange, the nucleus would be the size of a pea in the center.

Cathode-Ray Tubes

Sir William Crookes discovered cathode rays, which were streams of charged particles.

Electrons

J.J. Thomson identified these particles as electrons, proving that atoms contained smaller subatomic particles.

Mass-to-Charge Ratio

J.J. Thomson calculated the charge-to-mass ratio of electrons, proving they were much lighter than atoms.

Millikan's Oil-Drop Experiment

Robert Millikan determined the charge of an electron (1.602 × 10^-19 coulombs) and calculated the mass of an electron (9.1 × 10^-28 g).

Thomson's Plum Pudding Model

Proposed that the atom was a sphere of positive charge with negatively charged electrons embedded within it like raisins in a pudding.

Rutherford's Gold Foil Experiment

A beam of alpha particles was directed at gold foil, and the scattering was observed.

Rutherford's Conclusion

Rutherford concluded that the atom has a small, dense nucleus that contains most of its mass and positive charge, with electrons orbiting outside.

Rutherford Nuclear Model

This model describes the atom as consisting mostly of empty space, with the nucleus at the center and electrons surrounding it.

Proton

A positively charged particle within the nucleus.

Neutron

A neutral particle in the nucleus that has a mass similar to the proton.

Electron

A negatively charged subatomic particle located outside the nucleus.

Nucleus

The small, dense, positively charged center of the atom.

Atom

The smallest unit of an element, composed of protons, neutrons, and electrons.

Cathode Ray

A stream of negatively charged particles observed in cathode-ray tubes.

Plum Pudding Model

J.J. Thomson's model of the atom, which proposed a positive charge with electrons embedded like raisins in pudding.

Gold Foil Experiment

Rutherford's experiment that led to the discovery of the atomic nucleus.

Atomic Number

The number of protons in an atom, which determines the element's identity.

Isotope

Atoms of the same element that have the same number of protons but a different number of neutrons, which results in a different mass number.

Mass Number

The total number of protons and neutrons in an atom.

Quantum Mechanics

The branch of physics that describes the behavior of particles at the atomic and subatomic levels.

Electrons in Orbitals

Electrons exist in 'orbitals' or regions of probability, rather than in defined paths.

Neutrality of the Atom

The number of protons in the nucleus equals the number of electrons orbiting the nucleus, making the atom electrically neutral.

Size of the Nucleus

A typical atom's diameter is 10,000 times the diameter of the nucleus.

Protons Calculation

Protons = Atomic number.

Electrons Calculation

Electrons = Number of protons (since atoms are neutral).

Neutrons Calculation

Neutrons = Mass number - Atomic number.

Atomic Mass

The weighted averages of the masses of all naturally occurring isotopes of an element, considering their relative abundance.

Isotope

Atoms of the same element with the same number of protons but different numbers of neutrons.

Mass Number

The sum of protons and neutrons in an atom's nucleus.

Atomic Mass Unit (amu)

A unit of mass used to express atomic and molecular weights, where 1 amu is defined as 1/12th the mass of a carbon-12 atom.

Atomic Mass

The weighted average of the atomic masses of all isotopes of an element.

Atomic Number

The atomic number of an element tells you how many protons (and electrons) it has.

Example of Atomic Number

Hydrogen has an atomic number of 1, so it has 1 proton and 1 electron.

Example of Gold Atomic Number

Gold (Au) has an atomic number of 79, so it has 79 protons and 79 electrons.

Neutrons Calculation

If an element has an atomic number of 82 (Pb, Lead), it has 82 protons and 82 electrons (since it's neutral). If the mass number is 207, the number of neutrons is calculated by subtracting the atomic number from the mass number: Neutrons = 207 - 82 = 125 neutrons.

Isotopes of Potassium

Potassium (K) has isotopes with 20, 21, or 22 neutrons, but all have 19 protons.

Isotopes of Copper

Copper-63 has 29 protons and 34 neutrons, while Copper-65 has 29 protons and 36 neutrons.

Isotope Notation

The mass number is the sum of protons and neutrons. Isotope notation includes the element's symbol and the mass number.

Example of Isotope Notation

Copper-63 = Cu-63 or Cu with mass number 63.

Atomic Mass Calculation

To calculate atomic mass, multiply the mass of each isotope by its relative abundance and add the results.

Example Calculation for Copper

Copper has two isotopes: Cu-63 (abundance = 69.2%, mass = 62.930 amu) and Cu-65 (abundance = 30.8%, mass = 64.928 amu). Atomic mass of copper = (62.930 × 0.692) + (64.928 × 0.308) = 63.546 amu.

Challenge Example for Boron

Boron has two isotopes: Boron-10: mass = 10.013 amu, abundance = 19.8% and Boron-11: mass = 11.009 amu, abundance = 80.2%. Atomic mass of Boron = (10.013 × 0.198) + (11.009 × 0.802) = 10.81 amu.

Radioactive Decay

Radioactive decay occurs when an unstable atomic nucleus loses energy by emitting radiation, transforming into a more stable form.

Alpha Radiation

Consists of alpha particles (helium nuclei, 2 protons, 2 neutrons) and is deflected toward a negatively charged plate.

Example of Alpha Radiation

Radium-226 (Ra) decays to form Radon-222 (Rn) by emitting an alpha particle.

Beta Radiation

Consists of fast-moving electrons (beta particles, -1 charge) and is deflected toward a positively charged plate.

Example of Beta Radiation

Carbon-14 decays into Nitrogen-14 by emitting a beta particle.

Gamma Radiation

High-energy radiation with no mass and no charge, which is not deflected by electric or magnetic fields.

Atomic Number

The number of protons in an atom's nucleus, determining the identity of the atom.

Isotopes

Variations of elements that differ in the number of neutrons.

Atomic Mass

The weighted average of the masses of an element's isotopes, which is not always a whole number.

Nuclear Reactions

Processes that change the nucleus of an atom, such as radioactive decay.

Alpha Radiation

Radiation consisting of alpha particles, which have 2 protons and 2 neutrons with a 2+ charge.

Beta Radiation

Radiation consisting of beta particles, which are electrons with a 1- charge.

Gamma Radiation

High-energy radiation with no mass or charge, often accompanying alpha or beta decay.

Mass Number

The total number of protons and neutrons in an atom's nucleus.

Isotope Notation

A way to represent isotopes, typically including the element symbol and mass number.

Atomic Mass Calculation

The process of determining the atomic mass using the formula: (Mass of Isotope 1 × Abundance of Isotope 1) + (Mass of Isotope 2 × Abundance of Isotope 2).

Example of Atomic Number

Hydrogen (H) has an atomic number of 1.

Example of Atomic Number

Helium (He) has an atomic number of 2.

Example Problem - Gold

Gold (Au) has an atomic number of 79, so it has 79 protons and 79 electrons.

Isotopes Example

Potassium-39, Potassium-40, and Potassium-41 have the same number of protons (19) but different numbers of neutrons (20, 21, and 22, respectively).

Neon-22

An isotope with an atomic number of 10, mass number of 22, and 12 neutrons.

Chlorine Atomic Mass

Chlorine has an atomic mass of 35.453 amu, the weighted average of chlorine-35 and chlorine-37 isotopes.

Copper Atomic Mass Calculation

Cu-63: 69.2% abundance, 62.930 amu mass; Cu-65: 30.8% abundance, 64.928 amu mass.

Radioactivity

The process by which unstable atomic nuclei lose energy by emitting radiation.

Nuclear Reaction

A reaction that changes an atom's nucleus, leading to the transformation of one element into another.

Alpha Decay of Radium-226

Radium-226 decays into radon-222 and emits an alpha particle.

Challenge Problem

An atom has a mass number of 55. Its number of neutrons is the sum of its atomic number and five.