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Atomic Structure Review AC Physical Science 2008 Atom, Molecule, or Compound ?? • Atoms - smallest particle of an element • Element – cannot be broken down by any means possible • Molecule – 2 or more atoms that are bonded together. For example O2 and H2O. • Compounds – substance made of 2 or more elements. For example: H2O • ALL compounds are also molecules, BUT all molecules ARE NOT compounds… WHY?? Elements… • Cannot be separated into simpler substances by any chemical or physical means. • Pure Substances are substances where there is only one type of particle – for example pure oxygen. • Elements have their own unique properties, but can share characteristics. Classifying Elements • Classified by their properties… • Categories are defined as metals, nonmetals, and metalloids. • Samples of metals = Pb, Cu, Sn • Samples of nonmetals = S, Ne, I • Sample of metalloids = Si, B, Sb Molecules • 2 or more atoms that are combined together • Example = O2 Compounds • A Compound is a pure substance composed of 2 or more elements that are chemically combined. • Examples: CO2 Mixtures • Combination of 2 or more substances that are not chemically combined • Solutions – mixture that appears to be a single substance – composed of 2 or more substances that are evenly distributed among each other. • Solute – substance that is dissolved • Solvent – substance in which the solute is dissolved in Mixtures continued… • Concentration – measure of the amount of solute dissolved in a solvent • Solubility – solutes ability to dissolve in the solvent at a certain temperature • Suspension – mixture in which particles of a material are dispersed throughout a liquid or gas (example Italian Salad Dressing) • Colloid – mixture consisting of tiny particles that are dispersed throughout, but are not heavy enough to settle out (Example = jello or milk) Atoms.. • Atom = the smallest particle an element can be divided into and still be the same substance. • Atom is from the Greek word atomos, meaning “not able to be divided”. Dalton’s Atomic Theory • John Dalton, a British chemist published his theory in 1803. • His theory states the following ideas: – All substances are made of atoms. Atoms are small particles that cannot be created, divided, or destroyed. – Atoms of the same element are exactly alike, and atoms of different elements are different. – Atoms join with other atoms to make now substances. Not quite correct… • 1897, J.J. Thomson showed that there was a mistake in Dalton’s theory. • He discovered that there are smaller particles inside the atom. • Which means that the atom can be divided into even smaller parts. • Thomson discovered negatively charged particles that are called electrons. Rutherford’s Atomic “Shooting Gallery” • In 1909, Ernest Rutherford tested Thomson’s Theory. • He aimed a beam of small positively charged particles at a thin sheet of gold foil. • In 1911, Rutherford revised the atomic theory. He made a new model. • He proposed that the center of an atom is tiny, extremely dense, and positively charged part called the nucleus. • Because like charges repel, a particle that headed straight for a nucleus would be pushed almost straight back and in the direction in which it came. Bohr’s Electron Level • In 1913, Niels Bohr, worked with Rutherford, And studied that way that atoms react to light. • The results led him to propose that electrons move around the nucleus in certain paths or energy levels. • Bohr’s model was a valuable tool in predicting some atomic behavior, but the atomic theory still had room for improvement. The Modern Atomic Theory • Erwin Schrödinger and Werner Heisenberg work and further explained that nature of electrons in an atom. • The determined that the exact path of an electron cannot be predicted. • According to the current theory, there are regions inside the atom where electrons are likely to be found. • These regions are called electron clouds. Electron Clouds… • Use the formula 2(N2) to determine how many electrons fit in each cloud… Practice, Practice… Parts of an Atom • Nucleus: the small dense positively charged center of the atom. – This contains most of the atom’s mass. • Protons: are positively charged particles in the nucleus of an atom • Neutrons: are particles in the nucleus that have no charge. • Electrons: negatively charged particles found in the electron clouds outside the nucleus. Subatomic Particles… • • • • Protons are positive + Neutrons are neutral Electrons are negative – Protons and neutrons are in the center of the nucleus • Electrons are situated in electron clouds spinning around the nucleus Things to remember… • Atomic Number = the number of protons and electrons • The difference between the atomic number and the atomic mass = the number of neutrons The Nucleus • Protons: are positively charged particles in the nucleus of an atom • The SI unit used to express the masses of particles in atoms is the: atomic mass unit (amu) – Each proton has a mass of about 1amu • Neutrons: are particles in the nucleus that have no charge. Outside the Nucleus • Electrons: negatively charged particles found in the electron clouds outside the nucleus. • Compared with protons and neutrons, electrons are very small in mass. • The charges of protons and electrons are opposite, but equal, so they cancel each other out. • Because an atom has no overall charge, it is neutral. • What happens if the numbers of the electrons and protons are not equal?... Well…Ions… • The atom becomes a charged particle called an ion. • An ION is an atom or group of atoms that carry an electric charge. • So…if an atom has more protons it is??? • If there are more electrons?? • This is how weather patterns are created… Isotopes • An atom that has the same number or protons (or the same atomic number) as other atoms of the same element do but that has a different number of neutrons (and thus a different atomic mass). • You can tell Isotopes apart by the mass number which is the sum of the numbers of protons and neutrons in the nucleus of an atom. Example of an Isotope = Carbon-14 used to date rocks… How do I find Neutrons again? • The difference between the atomic number and the atomic mass = the number of neutrons