Atoms, Molecules, and Ions. Search for:. Learning Objectives Describe the early developments leading to the modern concept of the atom. Key Takeaways Key Points The ancient Greek philosophers Democritus and Leucippus recorded the concept of the atomos , an indivisible building block of matter, as early as the 5th century BCE. The idea of an indivisible particle was further elaborated upon and explored by a number of scientists and philosophers, including Galileo, Newton, Boyle, Lavoisier, and Dalton.
John Dalton, an English chemist and meteorologist, is credited with the first modern atomic theory based on his experiments with atmospheric gases.
Key Terms atom : The smallest possible amount of matter that still retains its identity as a chemical element, now known to consist of a nucleus surrounded by electrons. The Law of Conservation of Mass The law of conservation of mass states that mass in an isolated system is neither created nor destroyed.
Learning Objectives Define the law of conservation of mass. Key Takeaways Key Points The law of conservation of mass states that mass in an isolated system is neither created nor destroyed by chemical reactions or physical transformations. According to the law of conservation of mass, the mass of the products in a chemical reaction must equal the mass of the reactants.
The law of conservation of mass is useful for a number of calculations and can be used to solve for unknown masses, such the amount of gas consumed or produced during a reaction. Key Terms law of conservation of mass : A law that states that mass cannot be created or destroyed; it is merely rearranged.
Also, a molecule before it undergoes a chemical change. The Law of Definite Composition The law of definite composition states that chemical compounds are composed of a fixed ratio of elements as determined by mass. Learning Objectives Define the law of definite composition. Key Takeaways Key Points The law of definite composition was proposed by Joseph Proust based on his observations on the composition of chemical compounds. Proust proposed that a compound is always composed of the same proportions of elements by mass.
Key Terms element : Any one of the simplest chemical substances that cannot be decomposed in a chemical reaction or by any chemical means, and are made up of atoms all having the same number of protons. Joseph Proust : Portrait of Joseph Proust. The Law of Multiple Proportions The law of multiple proportions states that elements combine in small whole number ratios to form compounds.
Since the deflections occurred a small fraction of the time, this charge only occupied a small amount of the space in the gold foil. Analyzing a series of such experiments in detail, Rutherford drew two conclusions:.
Most pass through the relatively large region occupied by electrons, which are too light to deflect the rapidly moving particles. With one addition, which you will learn next, this nuclear model of the atom, proposed over a century ago, is still used today. Another important finding was the discovery of isotopes.
During the early s, scientists identified several substances that appeared to be new elements, isolating them from radioactive ores.
However, a more detailed analysis showed that mesothorium was chemically identical to radium another decay product , despite having a different atomic mass. This result, along with similar findings for other elements, led the English chemist Frederick Soddy to realize that an element could have types of atoms with different masses that were chemically indistinguishable.
These different types are called isotopes —atoms of the same element that differ in mass. Soddy was awarded the Nobel Prize in Chemistry in for this discovery. One puzzle remained: The nucleus was known to contain almost all of the mass of an atom, with the number of protons only providing half, or less, of that mass. Different proposals were made to explain what constituted the remaining mass, including the existence of neutral particles in the nucleus.
As you might expect, detecting uncharged particles is very challenging, and it was not until that James Chadwick found evidence of neutrons , uncharged, subatomic particles with a mass approximately the same as that of protons. The existence of the neutron also explained isotopes: They differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons.
This will be explained in more detail later in this unit. The ancient Greeks proposed that matter consists of extremely small particles called atoms. Dalton postulated that each element has a characteristic type of atom that differs in properties from atoms of all other elements, and that atoms of different elements can combine in fixed, small, whole-number ratios to form compounds.
Samples of a particular compound all have the same elemental proportions by mass. When two elements form different compounds, a given mass of one element will combine with masses of the other element in a small, whole-number ratio. During any chemical change, atoms are neither created nor destroyed. Although no one has actually seen the inside of an atom, experiments have demonstrated much about atomic structure.
Millikan discovered that there is a fundamental electric charge—the charge of an electron. Chadwick discovered that the nucleus also contains neutral particles called neutrons. Soddy demonstrated that atoms of the same element can differ in mass; these are called isotopes. Austin State University with contributing authors. Have feedback to give about this text? These atoms couldn't be destroyed or created, only rearranged and combined in different ways.
This became the conservation of mass, which is part of our current understanding of a chemical reaction. Dalton also made major contributions to our knowledge of chemical compounds and formulae, measuring the relative masses of elements which he found reacted together to make new chemical substances. JJ Thomson is credited with discovering the electron , as a small electrically charged part of an atom.
Positively-charged alpha particles were fired at high-speed at a thin gold foil sheet and the way they deflected was recorded. Scientists create a hypothesis or prediction to test the observation. Scientists design an experiment which applies the hypothesis, where a measurable result will tell you if it is true or not. The hypothesis is about deflection of particles, so the scientists measure deflection of particles.
If the hypothesis is true, there would be almost no deflection. If it is not true, there will be significant deflection Instead of virtually no deflection in all the alpha particles, while most particles passed through unaffected, some had huge deflection angles. Some particles even scattered back towards the source. To add to the above, the observation made from the plum pudding model did not make sense anymore.
How could the densely charged high velocity alpha particles get knocked back and away by the sparse charge of the gold atoms? The emprical evidence did not back up the hypothesis made before the test.
When this happens, the earlier observation is incorrect. In its place, we have the new observation, the results of this experiment. From this failed hypothesis, Rutherford developed his own atomic theory. What did some of the alpha particles collide with that caused such a huge deflection in their path? This is also where most of the atomic mass is found. Positive alpha particles colliding with a positive nucleus would cause strong charge repulsion and radically deflect the particles from their path.
In , J. Thomson discovered the electron. He believed atoms could be divided. Because the electron carried a negative charge, he proposed a plum pudding model of the atom, in which electrons were embedded in a mass of positive charge to yield an electrically neutral atom.
Ernest Rutherford, one of Thomson's students, disproved the plum pudding model in Rutherford found that the positive charge of an atom and most of its mass were at the center, or nucleus, of an atom. He described a planetary model in which electrons orbited a small, positive-charged nucleus.
Rutherford was on the right track, but his model couldn't explain the emission and absorption spectra of atoms, nor why the electrons didn't crash into the nucleus.
In , Niels Bohr proposed the Bohr model, which states that electrons only orbit the nucleus at specific distances from the nucleus. According to his model, electrons couldn't spiral into the nucleus but could make quantum leaps between energy levels. Bohr's model explained the spectral lines of hydrogen but didn't extend to the behavior of atoms with multiple electrons. Several discoveries expanded the understanding of atoms.
In , Frederick Soddy described isotopes, which were forms of an atom of one element that contained different numbers of neutrons. Neutrons were discovered in This, in turn, led to Werner Heisenberg's uncertainty principle , which states that it's not possible to simultaneously know both the position and momentum of an electron.
Quantum mechanics led to an atomic theory in which atoms consist of smaller particles. The electron can potentially be found anywhere in the atom but is found with the greatest probability in an atomic orbital or energy level. Rather than the circular orbits of Rutherford's model, modern atomic theory describes orbitals that may be spherical, dumbbell-shaped, etc. For atoms with a high number of electrons, relativistic effects come into play, since the particles are moving at a fraction of the speed of light.
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