Chapter 16: The Properties of Atoms and the Periodic Table Section 1 – The Structure of Atoms Section 2 – Masses of Atoms Section 3 – The Periodic Table

Section 1 – The Structure of Atoms The Development of the Atomic Model The word atom was first used by the Greek scientist/philosopher Democritus 2400 years ago He believed that the atom was the smallest particle of matter possible Only four atoms: fire, water, earth, and air The Greeks thought that an atom of fire would look like a very small flame, water like a very small bit of water, and so on All matter consisted of these atoms or mixtures of the atoms The problem with the Greek idea of the atom is that they could not test to see if it was correct In the early 1800’s, English scientist John Dalton provided the basic theory of the scientific atomic model with four basic statements All matter is composed of indivisible atoms. These atoms retain their identity during chemical reactions An element is a type of matter composed of only one type of atom, each kind of atom have the same properties A compound is a type of matter composed of atoms of two or more elements chemically combined in fixed proportions, Ex.: water → A chemical reaction consists of the rearrangement of the atoms present in the reacting substances to give new chemical combinations present in the substances formed by the reaction Ex.: formation of water: One problem: no one could prove the existence of atoms

Section 1 – The Structure of Atoms The Development of the Atomic Model Dalton was wrong about one thing: atoms can be divided into smaller particles In 1897 J.J., Thomson demonstrated the existence of negatively charged particles smaller than a hydrogen atom through a series of experiments using the cathode ray tube (CRT) Scientists knew that charging a CRT would result in the formation of a beam. Was beam a series of waves or a stream of particles? Thomson used a CRT and a magnet to show that the beam was deflected by the magnetic field so it must consist of a stream of particles Next, he placed two oppositely charged plates in the CRT and discovered that the particles were deflected towards the (+) charged plate. After further tests and calculations, Thomson concluded the particles were much smaller than a hydrogen atom and carried a (-) charge—the electron.

Section 1 – The Structure of Atoms The Development of the Atomic Model In 1911 Ernest Rutherford demonstrated the presence of a positively charge atomic nucleus Using alpha particles and gold foil, Rutherford conducted the following experiment: (An alpha particle is the nucleus of a helium atom which has a positive charge—at the time they just knew that it was a positively charged particle) The stream of alpha particles would leave the lead block and some would pass through a small hole in another lead block The Thomson atomic model consists of a positively charged ball with negatively charged particles evenly embedded in the ball

Section 1 – The Structure of Atoms The Development of the Atomic Model The particles would then strike a piece of gold foil Most of the particles would pass through the foil, although their path was deflected. But some particles would rebound back toward the source. Rutherford deduced the atoms of gold must have a massive positively charged nucleus (called the proton) Continued experimentation by Rutherford revealed: The nucleus contains 99.95% of the mass of an atom, but occupies a very small space in the atom If the nucleus was the size of a golf ball, the atom would be about 3 miles in diameter It would take about 2000 electrons to equal the mass of one proton The Rutherford atomic model proposed the majority of the mass of an atom was in the positively charged nucleus, that there was empty space between the nucleus and the electrons, and the electrons circled the nucleus

Section 1 – The Structure of Atoms The Development of the Atomic Model The location of electrons was also an issue In 1920, Niels Bohr theorized that electrons circled the nucleus much like the planets orbit the sun Bohr called these orbits energy levels. That is, each electron has a very specific path, and you could determine both the location and motion of the electron In 1926, Werner Heisinberg, based on quantum mechanics, demonstrated it was impossible to know both the motion and location of an electron at the same time Heisenberg proposed that the electrons form a cloud around the nucleus of an atom. In the electron cloud were regions called orbitals where the electrons were likely to be found

Section 1 – The Structure of Atoms The Development of the Atomic Model There was a problem with Rutherford’s model of the atom. The proton alone could not account for the mass of the nucleus, there was a missing particle In the early 1900’s, scientists knew that hydrogen consisted of one proton and 1 electron, and that helium contained 2 protons and 2 electrons The ratio of the mass of helium to the mass of hydrogen should have been 2 to 1, but the actual ratio was 4 to 1 In 1932 James Chadwick performed a series of experiments similar to Rutherford’s. He found that an unidentified, high energy radiation was given off Chadwick was able to prove this radiation was composed of neutrally charged particles with about the same mass as the proton Thus, the neutron was discovered The Modern Atomic Model Consists of the Atomic Nucleus which contains protons and neutron, and the Electron Cloud which contains electrons that are arranged in energy levels and whose motion and position cannot be described at the same time.

Section 2 – Masses of Atoms The identification of an element Elements are identified by the number of protons in the nucleus of an atom of the element As the number of protons in the nucleus changes, so does the element The number of protons in the nucleus of an atom is called its atomic number The mass number of an element is the sum of the number of protons and neutrons in the nucleus of an atom of the element The number of neutrons in the nucleus of an atom can be determined mathematically: Sometimes the number of neutrons does not equal the number of protons. When this is the case, we have isotopes of the same element Remember, the number of protons determine the element, so you can more (or less) neutrons than protons and have the same element Some isotopes are radioactive, and this property can be used in different ways Example: Carbon-14 (C14) – Scientists know that Carbon-14 decays at a regular rate. Using that fact, they can determine the age of something by the comparing the ratio of Carbon-12 to carbon-14

Section 2 – Masses of Atoms The mass of an atom can be measured using relative units The unit used to express the relative mass of atoms is the atomic mass unit (amu) 1 amu is defined as one-twelveth (1/12) the mass of a Carbon-12 atom Protons and neutrons have about the same mass: nearly 1 amu mP = 1.673 x 10-24-kg, mN = 1.675 x 10-24-kg

Section 3 – The Periodic Table The Periodic Table is an representation of all the known elements and is arranged by increasing atomic number When the elements are arranged in this order, the elements in a the same group (a vertical column on the Table) have similar properties Each element has its own “box” on the Periodic Table that describes the properties of the element. The properties include: the atomic number, the state of matter of the element at room temperature, the symbol used to identify the element, the name of the element, and the element’s average atomic mass. There are four basic types of elements: metals, nonmetals, metalloids, and synthetic (man-made) Metals typically are solid at room temperature, and are good conductors of heat and electricity Nonmetals typically are gases at room temperature and are poor conductors of heat and electricity Metalloids have some of the properties of both metals and nonmetals The synthetic elements are typically constructed in particle accelerators and most do not stay in existence for long periods – they breakdown into smaller atoms and/or subatomic particles