Lesson 3.1 Introduction to Atoms Vocabulary: Atomic Theory

Lesson 3.1 Introduction to Atoms Vocabulary: Atom - Electron – Nucleus – Proton – Energy level – Neutron – Atomic number – Isotope – Mass number - Atomic Theory Studying atoms is difficult because atoms are so small. The smallest visible speck of dust may contain about 10 million billion atoms. Scientists have created models to describe atoms, but the models have changed many times as technology improved. Around 430 BC, a Greek philosopher Democritus suggested that matter was formed of small pieces that could not be cut into smaller parts. He used the word “atomos” meaning “uncuttable,” for these smallest possible pieces. The atom is the smallest particle that still can be considered an element The idea of atoms began to develop again in the 1600s. Atomic theory grew as a series of models that developed from experimental evidence. As more evidence was collected, the theory and models were revised. Dalton’s Atomic Theory in 1803 He thought that atoms were like smooth, hard billiard balls that could not be broken into smaller pieces. Dalton’s Theory: 1. All elements consist of atoms that cannot be divided 2. All atoms of the same element are exactly alike and have the same mass. Atoms of different elements have different masses 3. An atom of one element cannot be changed into an atom of a different element by a chemical reaction 4. Compounds are formed when atoms of more than one element combine in a specific ratio Thomson’s Model in 1897 One important change to Dalton’s model was that atoms have even smaller parts. In 1897, JJ Thomson discovered that atoms contain negatively charged particles called electrons. But scientists knew that atoms had NO electrical charge, so that meant that atoms must also contain some sort of positive charge to balance out the negative electrons. Rutherford’s Model in 1911 Ernest Rutherford was one of Thomson’s students. His research team aimed a beam of positively charged particles at a thin sheet of gold foil. They predicted that the charged particles would pass straight though the foil, or bend slightly – and most did. To their surprise, a few particles were deflected back at them and sideways. Based on these results, Rutherford suggested that the atom is mostly empty space, but has a positive charge at its center. Like charges repel, so the atom’s positive charge must be packed at the center, called a nucleus. Any particle that was deflected back had been repelled by the positive nucleus. Later research suggested that the nucleus had one or more positively charged particles, and he called them protons. Rutherford’s new model looked more like a cherry, with the pit representing the nucleus, and the fruit space taken up by electrons. Bohr’s Model in 1913 Niels Bohr was one of Rutherford’s students. Bohr revised the atomic model and suggested that electrons are found only in certain orbits, like planets orbiting the sun. Bohr Model above Cloud Model above Cloud Model in the 1920s Scientists determined that electrons do not orbit the nucleus like planets, but move rapidly within a cloud-like region around the nucleus. An electron’s movement is related to its energy level, or the specific amount of energy it has. Electrons at different energy levels are likely to be found in different places. This was proposed by Erwin Schrodinger – he took Bohr’s idea and improved upon it. Modern Model of the Atom In 1932, English scientist James Chadwick showed that another particle exists in the nucleus of the atoms. This particle, called a neutron, was hard to find because it has no electric charge. At the center of the atom is a tiny, dense nucleus containing protons and neutrons. Surrounding the nucleus is a cloudlike region of moving electrons. Most of an atom’s volume is the space in which the electrons move. This space is huge compared to the space taken up by the nucleus. If the nucleus was the size of a green pea in the middle of a football stadium, the electrons would reach as far away as the top row of seats! Protons are shown with a plus sign + Electons are shown with a negative sign – The number of protons equals the number of electrons, so the charges balance, making the atom neutral. Neutrons don’t affect the charge of an atom because they have a charge of zero. Comparing Particle Masses Electrons are so tiny that it takes almost 1,840 electrons to equal the mass of one proton. A proton and a neutron are about equal in mass. Together, the protons and neutrons make up almost all the mass of an atom. Atoms are too small to be described by grams and kilograms, so scientists use “atomic mass units” or “amu.” A proton OR a neutron is equal to one amu. Atomic Number – The number of protons in the nucleus of an atom is the atomic number. Isotopes – All atoms of an element have the same number of protons, but sometimes the number of neutrons can vary. Atoms with the same number of protons, but different numbers of neutrons are called isotopes. An isotope is identified by its mass number, which is the sum of the protons and neutrons in the atom. About 99% of naturally occurring carbon is carbon 12. Despite their different mass numbers, all three carbon isotopes react the same way chemically. .

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