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On Earth, matter usually can be found as a solid, liquid, or gas. BEC & Plasma are usually found in space. What’s Matter? Volume & Mass Classifications of Matter Substances always the same composition (makeup) Elements (1 kind of atom) 92 Natural About 120 Total see Periodic Table Examples: Gold (Au), Lead (Pb), Mercury (Hg) * 8 Elements make up 98% of crust – Al, Fe, Ca, Na, K, Mg, Si 27%,O 47% Molecules (2 or more kinds of atoms chemically bonded that act as a unit) Examples: Hydrogen (H), Bromine (Br), Sugar Compounds (2 or more elements chemically joined in a specific combination) Examples: Water (H2o),Carbon Dioxide CO2 Mixtures Matter that varies in composition. 2 or more substances that are blended but not bonded. Heterogeneous Mixtures - not evenly mixed Examples: Trail Mix, Granite, Smoke Homogeneous Mixtures - evenly mixed/no bond! Examples: Brass, Natural Gas, Windex Solutions – not bonded mixture Composition Changes in composition Properties of parts Solutions Substances evenly mixed together Compounds Atoms bonded together in the same combination Solution still the Changes same but relative composition substances differ. makes a new compound with new properties. Substances keep Properties of the their own compound are properties when different than mixed. atoms that make it up. The Greeks named these particles atoms, a term that means ―cannot be divided.‖ • Early philosophers didn’t try to prove their theories by doing experiments as scientists now do. • Their theories were the result of reasoning, debating, and discussion—not of evidence of proof. 400 B.C. - Democritus thought matter could not be divided indefinitely. • This led to the idea of atoms in a void. fire Democritus earth air water • 350 B.C - Aristotle modified an earlier theory that matter was made of four “elements”: earth, fire, water, air. • Aristotle was wrong. However, his Aristotle theory persisted for 2000 years. • During the eighteenth century, scientists, especially the French, began debating the existence of atoms once more. Newton proposed held together by force. • They found that certain substances couldn’t be broken down into simpler substances. • Scientists came to realize that all matter is made up of elements. • An element is matter made of 1 kind of atom. Dalton’s Concept 1800 -Dalton an English schoolteacher proposed a modern atomic model based on experimentation not on pure reason. Dalton pictured an atom as a hard sphere that was the same throughout. • • • • All matter is made of atoms. Atoms of an element are identical. Each element has different atoms. Atoms of different elements combine in constant ratios to form compounds. • Atoms are rearranged in reactions. • His ideas account for the law of conservation of mass (atoms are neither created nor destroyed) and the law of constant composition (elements combine in fixed ratios). Scientific Evidence • In 1870, the English scientist William Crookes did experiments with a glass tube that had almost all the air removed from it. • The glass tube had two pieces of metal called electrodes sealed inside. • The electrodes were connected to a battery by wires. Crookes’s tube is known as a cathode-ray tube, or CRT. • An electrode is a piece of metal that can conduct electricity. • One electrode, called the anode, has a positive charge. The other, called the cathode, has a negative charge. • Crookes hypothesized that the green glow in the tube was caused by cathode rays. Models of the Atom Cathode Rays • Many scientists were not convinced that the cathode rays were streams of particles. • In 1897, J.J. Thomson, an English physicist, tried to clear up the confusion. • He placed a magnet beside the tube from Crookes’s experiments. • Thomson concluded that the beam must be made up of (-) negatively charged particles of matter that came from the cathode. Models of the Atom Thomson’s Model • These negatively charged particles are now called electrons. Electrons were 1st proposed by proposed in 1874 by G. Johnstone Stoney Thomson also inferred that electrons are a part of every kind of atom. • If atoms contain one or more negatively charged particles, then all matter, which is made of atoms, should be negatively charged as well. Thomson’s Atomic Model • The negatively charged electrons were spread evenly among the positive charge. • The negatively charged electrons and the unknown positive charge would then neutralize each other in the atom. • The atom is neutral. Thomson’s Atomic Model • Later discovered not all atoms are neutral. #Electrons within an element can vary. • More positive charge than negative electrons = overall positive charge. • More negative electrons = overall negative charge. • Rutherford wanted to see what would happen when they fired fast-moving, positively charged bits of matter, called alpha particles +, at a thin film (400nm) of a metal such as gold surrounded by a fluorescent screen. Models of the Atom Rutherford’s Results Fail! • His prediction -speeding alpha particles would pass right through the foil and hit the screen on the other side. • Rutherford reasoned - the thin, gold film did not contain enough matter to stop the speeding alpha particle or change its path. Rutherford was shocked when his students Hans Geiger & Ernest Marsden rushed in to tell him that some alpha particles were veering off at large angles. The Model Fails Positively charged alpha particles moving with such high speed that it would take a large positive charge to cause them to bounce back. The Proton • The actual results did not fit this model, so Rutherford proposed a new one. • He hypothesized – most of the mass of the atom and its positive charge is in the nucleus. Models of the Atom The Proton • In 1920 scientists identified the positive (+) charges in the nucleus as protons. • Rutherford’s new model of the atom fits the experimental data. • Most alpha particles could move through the foil with little or no interference. Electricity is called “cathode rays” when passed through an evacuated tube. These rays have a small mass and are negative. Thompson noted that these negative subatomic particles were a fundamental part of all atoms. Dalton’s “Billiard ball” model (1800-1900) Atoms are solid and indivisible. Thompson “Plum pudding” model (1900) Negative electrons in a positive framework. Rutherford - Atoms are mostly empty space. Negative electrons orbit a positive nucleus. The Neutron • According to Rutherford’s model, the only other particle in the atom was the proton. • He proposed that another particle must be in the nucleus to account for the extra mass. • The mass of most atoms is at least twice as great as the mass of its protons. • The particle, which was later call the neutron, would have the same mass as a proton and be electrically neutral. Models of the Atom The Neutron • 20 Years Later in 1932 -The model of the atom was revised again to include the newly discovered neutrons in the nucleus by James Chadwick. • The nuclear atom has a tiny nucleus tightly packed with positively charged protons and neutral neutrons. Bohr Models • Then, electrons would travel in orbits around the nucleus. • A physicist named Niels Bohr even calculated energy levels for the hydrogen atom. • However, scientists soon learned that electrons are in constant, unpredictable motion and can’t be described try & explain the orbits. Electrons orbit the nucleus in “shells” Electrons can be bumped up to a higher shell if hit by an electron or a photon of light. The Electron Cloud Model • The electrons are more likely to be close to the nucleus. They are attracted to the positive charges of the proton. • Electrons travel in a region surrounding the nucleus, called the electron cloud. Identifying Numbers • The atomic number of an element is the number of protons in the nucleus. • Atoms of an element are identified by the number of protons because this number never changes without changing the identify of the element. Mass Number • The total masses of the protons and neutrons in an atom make up most of the mass of an atom. The mass number isotope is neutrons plus protons. • You can find the # of neutrons in an isotope by subtracting the atomic # from the mass #. • Electron 1/1,800 of a proton. (We don’t calculate!) Number of Neutrons • These 3 kinds of carbon atoms are called isotopes. Isotopes are atoms of the same element that have different numbers of neutrons, but the same # of protons. • Radioactive Decay • Many atomic nuclei are stable when they have about the same number of protons and neutrons. Some nuclei are unstable because they have too many or too few neutrons. This is especially true for heavier elements such as uranium and plutonium. • The release of nuclear particles and energy is called radioactive decay. • In these nuclei, repulsion builds up. The nucleus must release a particle to become stable. Radioactive Decay • When the particles that are ejected from a nucleus include protons, the atomic number of the nucleus changes. When this happens, one element changes into another (Transmutation). • A smoke detector makes use of radioactive decay. It contains americium-241, undergoes transmutation by ejecting energy and an alpha particle. Radioactive Decay • The fast-moving alpha particles enable the air to conduct an electric current. As long as the electric current is flowing, the smoke detector is silent. • The alarm is triggered when the flow of electric current is interrupted by smoke entering the detector. Changed Identity Alpha Particles • When americium expels an alpha particle, it’s no longer americium, it becomes the element that has 93 protons, neptunium. Loss of Beta Particles • During this different kind of transmutation, a neutron becomes unstable and splits into an electron and a proton. The proton, however, remains in the nucleus. • The electron, or beta particle, is released with a large amount of energy, so the atomic # of the element that results is greater by one.. • A beta particle is a high-energy electron that comes from the nucleus, not from the electron cloud. • Rate of Decay • Radioactive decay is random. • The rate of decay of a nucleus is measured by its half-life. • The half-life of a radioactive isotope is the amount of time it takes for half of a sample of the element to decay. The radioactive decay of unstable atoms is steady, unaffected by conditions (weather, pressure, magnetic or electric fields, and even chemical reactions. Radioactive Decay • Carbon-14 is used to determine the age of dead animals, plants, and humans. • In a living organism, the amount of carbon14 remains in constant balance with the levels of the isotope in the atmosphere or ocean. • This balance occurs because living organisms take in and release carbon. • Geologists examine the decay of uranium & potassium argon to age rocks. Uses of Radioactive Isotopes • Tracer elements are used to diagnose disease and to study environmental conditions. • The radioactive isotope is introduced into a living system such as a person, animal, or plant. • It then is followed by a device that detects radiation while it decays. • The isotopes chosen for medical purposes have short half-lives, which allows them to be used without the risk of exposing living organisms to prolonged radiation. The isotope iodine-131 has been used to diagnose problems with the thyroid. Also used to detect cancer, digestion problems, & circulation. Environmental Uses • Tracers such as phosphorus-32 are injected into the root system of a plant. • A detector then is used to see how the plant uses phosphorus to grow and reproduce. • Radioisotopes also can be placed in pesticides or the water cycle and followed to see the impact to the ecosystem. 1789 Antoine Lavoisier grouped elements into categories of metals, nonmetals, gases, & Earth. Before 1830 55 elements were known to early man & used to make weapons and jewelry. Examples: Gold, Silver, Iron, Copper, Tin 1869 Dimitri Mendeleev – Periodic Table – The elements were ordered by atomic mass & arranged according to similarities in properties. Not all elements were known so he left blank spaces for later discovery. Early 20th Century Henry Mosely - Rearranged the table by atomic number rather than mass. Organized by increasing atomic number Periods (rows) are labeled 1-7 (# of electrons) 18 Columns of elements contains a group (family) of elements that have similar properties. (Properties Repeat -Periodic Law) The number of an “A” group element matches the number of valence electrons for an element in that group. ( Valence electrons are the electrons that are in the highest occupied energy level of an atom. Valence electrons are the ones involved in chemical bonding.) Divided into Zones. - Representative Elements 1&2 & 13-18 *Metals *Metalloids *Nonmetals - Transition Elements 3-12 *Lanthanide Series & *Actinide Series 1 or 2 letter abbreviation - may be for the Latin name, discoverer’s name, or place Examples: O = Oxygen Sc = Scandium 3 letter abbreviation for New Names – Must be approved by the IUPAC International Union of Pure and Applied Chemistry Keys on the Periodic Table usually show the information represented in the periodic table. Atomic Number Element name Symbol Often Atomic Mass number (amu – atomic mass unit 1/12 of carbon -12) Sometimes state of matter at room temp. (Solid, Liquid, or Gas) Sometimes Natural (1-92) or Synthetic (93+) Using the Periodic Table Step 1 - Gather Information Go to the Periodic Table of Elements. Use the Table of Elements to find your element's atomic number and atomic weight. The atomic number usually the number located in the upper left corner and the atomic weight is the number located on the bottom, as in this example for krypton: Step 2 - The Number of Protons is... The atomic number is the number of protons in an atom of an element. In our example, krypton's is 36. This tells us that an atom of krypton has 36 protons in its nucleus. Every atom of krypton contains 36 protons. If an atom doesn't have 36 protons, it can't be an atom of krypton. Adding or removing protons from the nucleus of an atom creates a different element. For example, removing one proton from an atom of krypton creates an atom of bromine. Step 3 - The Number of Electrons is... Atoms have no overall electrical charge (balanced). Atoms must have equal numbers of protons and electrons. In our example, an atom of krypton must contain 36 electrons since it contains 36 protons. An atom can gain or lose electrons, becoming what is known as an ion. An ion is nothing more than an electrically charged atom. Adding or removing electrons from an atom doesn’t change it, just its net charge. Step 4 - The Number of Neutrons is... The atomic mass the total number of particles in an atom's nucleus. The atomic mass is a weighted average of all of the naturally occurring isotopes of an element relative to the mass of carbon-12. All you really need to find is something called the mass number. To find the mass number, round the atomic mass to the nearest whole number. Krypton's mass number is 84 since its atomic mass, 83.80, rounds up to 84. The mass number is a count of the number of particles in an atom's nucleus. Remember that the nucleus is made up of protons and neutrons. http://videos.howstuffworks.com/hsw/14741-simplyscience-the-history-of-the-periodic-table-video.htm http://videos.howstuffworks.com/hsw/5966chemistry-of-life-atoms-and-elements-video.htm In Summary...For any element: Number of Protons = Atomic Number Number of Electrons = Number of Protons = Atomic Number Mass Number = (Number of Protons) + (Number of Neutrons) Number of Neutrons = Mass Number Atomic Number The # of electrons increases from L to R across a period on the table. The last element in each period has 8 in its outer energy level so it is complete. Halogens have 7 in their outer shells. Noble gases have a stable electron configuration. That’s why they don’t combine! Each has 8 valence electrons, with the exception of helium, which has 2. The nucleus containing protons & neutrons is surrounded by the electron cloud. The # & arrangement of electrons determines the chemical & physical properties. More electrons can fit into orbitals farther away from the nucleus. (see overhead) When the highest occupied energy level of an atom is filled with electrons, the atom is stable and not likely to react. Ions are atoms that are no longer neutral. + & - attract each other to form an ionic bond An ionic bond is formed when a cation & anion bond. A nonmetal & metal bond! Metals form positive (+) ions (Cations) by losing electrons Nonmetals form negative (-) (Anions) by gaining electrons Ionization Energy – amount of energy used to remove an electron. Ionization Energy – increases from L to R across the period. Ionization Energy – decreases from the top of a group to the bottom. Covalent Bond – 2 atoms sharing a pair of valence electrons. Ex. P 166 Covalent Bond – 2 nonmetal elements Single bond – 1 pair of electrons Double bond – 2 pairs of electrons Triple bond – 3 pairs of electrons Covalent bond – Sharing of electrons Molecule is a neutral (no charge) group of atoms joined by a covalent bond. Diatomic means 2 atoms (7 elements – are only covalent – Br, H, N,O,F,I,Cl Remember - Brighties Have NO Fights In Class L side of periodic table has a greater attraction for electrons. Top of the group has greater attraction. Ionic Compounds – can be shown by a chemical formula. Shows the ratio (3:2) of the compounds. Ionic Crystals – shape determined by the arrangement of ions in its lattice. Ionic Crystals – classified by shape Have a high MP & good conductors. Properties explained by strong attractions among ions within a crystal lattice. Ex. P Polar covalent bonding – electrons unevenly shared Polar bonds have opposite sides like a magnet + / Polar covalent bonding means the atom with the greater charge has a partial negative charge. Greek symbols are used to show this. Nonpolar covalent bonding means the atoms share the electrons equally. (Diatomic molecule do this). Like atoms form nonpolar bonds. There are forces of attraction between molecules. (Ex. Water surface tension) Polar molecules have stronger attractions than nonpolar molecules. Symbols used to represent atoms & compounds. Elements use 1,2, & 3 letters. Compounds are symbols & numbers. In formulas the small # after the letter is a subscript This shows the number of atoms in the molecule. A superscript may show the charge if an ion is formed. A chemical formula tells which elements are present & how much of the atom.