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Transcript
Chemistry Review: Antoine Lavoisier (1743-94) defined the term “element” and identified 23 known substances as elements. He based his investigations on careful measurements of mass in chemical reactions. Allessandro Volta (1745-1827) invented the voltaic pile, now known as a battery. It provided scientists with a source of electricity for their investigations. Humphry Davy (1778-1829) used electrolysis to decompose compounds and identified many metallic elements. Michael Faraday (1791-1867) worked with Davy and coined the term “electrolysis” Joseph Proust (1754-1826) decomposed various compounds into elements and measured the mass of the elements produced. The patterns he discovered led to the law of definite proportions. John Dalton (1766-1844) proposed the first atomic theory. His theory extended the particle theory to provide a model for the properties of elements in terms of their atoms, and the formation of compounds. Sir Francis Bacon (1561-1626) used the scientific method to carry out investigations. He was one of the first scientists to develop new knowledge as a result of experimentation. Robert Boyle (1627-1691) believed that the Greek philosophers’ fourelement theory (earth, air, fire and water) could be improved upon, and he helped to lay the foundation for the concept of elements and compounds. Ernest Rutherford Carried out his experiment in 1919 with alpha particles scattering, identified the “proton” as an elementary particle. Density = mass ÷ volume (Flow chart): pure substances can be elements and compounds, which are made up of atoms (which cannot be created or destroyed chemically), which make up molecules. (which can be created or destroyed chemically) Molecule (latin: molecula) means “little mass” Parts of the Atom: proton found in the nucleus (+ charge), neutron found in the nucleus (neutral) and the electron in motion around the nucleus (- charge) Bohr-Rutherford model: both Bohr and Rutherford made important contributions to the structure of the atom. Their model is used today as an introduction to studying chemistry. Definitions: Isotopes: atoms of the same element that differs in mass. Ex there are 3 isotopes of hydrogen, with mass numbers of 1, 2 and 3. These atoms have different masses because they have different numbers of neutrons. However, since they have the same number of protons, they are the same element. Most elements have more than 1 naturally occurring isotope. However, the relative abundance of the isotopes follows no patterns. For example, 99.98% of all hydrogen atoms have a mass number of 1 (H-1), less than 0.02% are H-2, and even less are H-3! Properties of isotopes: Since all isotopes of an element have the same number of electrons, their chemical properties will be exactly the same. However, they have different masses due to the different number of neutrons; therefore, they will have slightly different properties that relate to mass. (Ex. Density , melting and boiling point). Also. Since the neutron/proton ratio is different for isotopes of an element, some may be radioactive while others are stable. Atomic masses: The atomic masses listed on the periodic table take into account the number and abundance of all naturally occurring isotopes of each element and are a weighted average of the atomic masses. Carbon-12 is the only isotope for which the atomic mass is a rounded number. Atomic Number: Is the number of electrons or protons in an atom. The number of neutrons is determined by subtracting the atomic number from the mass number of the element. The elements in the periodic table are arranged in increasing atomic number. Atomic Mass: Is the mass of an atom of an element, often expressed in atomic mass units. (u) The elements are arranged on the periodic table in order of increasing atomic number. Periodic Table: Elements are arranged in order of increasing atomic number. Groups/Families: are the vertical rows, each member has similar chemical and physical properties. There is an extra shell/energy level of electrons as you go down the rows/period. Each element has the same number of electrons in the outer shell. (Except He, which is full with 2 electrons and the heavier elements) Periods: are the horizontal rows, as you go to the right each successive element has one more electron in the outer shell/energy level. Once the outer shell is “full” the next electron must go into a new outer shell. Bonding: Ionic bonding: electrons are transferred from one atom to another. This requires the formation of ions which creates a positive or negative charge. (All atoms are neutral to begin with because they contain the same number of electrons (-) and protons (+). Once this change occurs the atom is going to be “attracted to the opposite”! Ex. When fluorine atoms gain 1 electron from sodium it becomes negative (F-) and is ready to bond with a now positive sodium(Na+). Covalent bonding: 2 or more atoms “share” electrons.