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  1. Regents Review Unit 1 – Matter and Measure Unit 2 – Atomic Theory and Structure Unit 3 – Nuclear Chemistry Unit 4 – Periodic Table

  2. Unit 1Matter and Measure Chapters 1-3

  3. Chemistry • What is Chemistry? • Study of matter and the changes it undergoes • Branches • Organic • Physical • Analytical • Biochemical • Inorganic

  4. Matter • Anything that has mass and volume • Classified into 2 categories • Pure Substance • Sometimes just referred to as substances • Mixtures

  5. Matter

  6. Pure Substances • Element • simplest form of matter that has a unique set of properties. • Can’t be broken down by chemical means • Compounds • substance of two or more elements chemically combined in a fixed proportion • Can be broken down by chemical means

  7. Mixtures • Physical blend of two or more substances • Two Types: • Homogeneous • Composition is uniform throughout • Heterogeneous • Composition is not uniform throughout

  8. Separating Mixtures • Differences in physical properties can be used to separate mixtures • Filtration – Separates solids from liquids in heterogeneous mixtures • Distillation – Separates homogeneous liquid mixtures based on different boiling points • Evaporation – evaporate away liquid to leave solid • Chromatography – separation of substances based on polarity and solubility

  9. Density • Amount of matter in a given amount of space • Amount of mass in a given volume

  10. Identifying Substances • Physical Property • Property of a substance that can be observed or measured without changing the substance’s composition • Ex: Color, shape, size, mass • Physical Change • some properties change, but the composition remains the same • Ex: melting, freezing, tearing

  11. Identifying Substances (cont) • Chemical Change • change that produces matter with a different composition than the original matter • Ex. burning, rusting, decomposing, exploding, corroding • Chemical property • property that can only be observed by changing the composition of the substance. • Ex: Reacting with, forming a new substance

  12. Significant Figures (cont) • If the decimal point is written, start on the LEFT side, go until you get to the first non-zero digit, count that one and every one from there to the end 1 2 3 0.00310 (3 sig. figs.)

  13. Significant Figures (cont) • If the decimal point is not written, start on the RIGHT side, go until you get to the first non-zero digit, count that one and every one from there to the end 3 2 1 31,400 (3 sig. figs.)

  14. SigFigs for Math • Addition and Subtraction • Answer has to have the same number of decimal places as least decimal places in what you are adding or subtracting • Multiplication and Division • Answer has to have same number of Sigfigs as least number of Sigfigs in what you are multiplying or dividing

  15. Unit 2Atomic Theory and Structure Chapters 4-5

  16. Atom • Atoms are made of subatomic particles • Protons • Neutrons • Electrons

  17. Electron • Discovered first • Negative charge (-1) • Approx mass ~ 0u • Found outside of nucleus • Valence Electron • Electrons in the outermost energy level

  18. Proton • Discovered second • Positive charge (+1) • Approx mass ~ 1u • Found inside nucleus

  19. Neutron • Discovered last • No charge (0) • Approx mass ~ 1 atomic mass unit (u) • Just slightly larger than a proton • Found inside nucleus

  20. Atomic Structure • Atoms have no net charge • # of electrons = # of protons • Nucleus • Center of atom, contains protons and neutrons • Positive charge

  21. Atomic Structure • Atomic Number • Number of protons • All atoms of the same element have the same number of protons • Mass Number • Number of protons and neutrons in an atom • # of Neutrons = Mass Number – Atomic Number

  22. Chemical Symbols Mass Number • Cl-35 • Chlorine-35 Atomic Number

  23. Atomic Structure • Isotope • atoms of the same element with different number of neutrons • Ion • Atom or group of atoms that have gained or lost one or more electrons • Have a charge

  24. Average Atomic Mass • Atomic Mass • Weighted average based on the relative abundance and mass number for all naturally occurring isotopes • Relative Abundance • Percent of each naturally occurring isotope found in nature

  25. Atomic Mass • C-12 98.9% • C-13 1.1% • Carbon = 0.989*12 + 0.011*13 = 12.011u

  26. Atomic Theories • Dalton’s Atomic Model • Also called Hard Sphere Model • First model • Plum Pudding Model • Uniform positive sphere with negatively charged electrons embedded within. • Came as a result of discovery of electron

  27. Rutherford Gold Foil Experiment • Shot alpha particles at gold foil • Most went through, some were deflected back • Conclusions • Atom is Mostly Empty Space • Dense positive core (nucleus)

  28. Atomic Theories • Rutherford Model • Dense positive core (nucleus) • Electrons moving randomly around nucleus • Bohr Model • Dense positive core (nucleus) • Electrons in specified circular paths, called energy levels

  29. Bohr Model • Each energy level can only hold up to a certain number of electrons • Level 1  2 electrons • Level 2  8 electrons • Level 3  18 electrons • Level 4  32 electrons

  30. Atomic Theories • Wave Mechanical Model • Dense positive core (nucleus) • Electrons in orbitals • Regions of space where there is a high probability of finding an electron • Modern (current) Model • AKA Quantum Mechanical Model, Electron Cloud Model

  31. Energy Level Transitions • Electrons can move between energy levels • Gaining energy will move an electron outward to a higher energy level • When an electron falls inward to a lower energy level, it releases a certain amount of energy as light

  32. Electron Configuration • The way in which electrons are arranged in the atom • Ground State • When the electrons are in the lowest available energy level (shown on reference tables) • Excited State • When one or more electrons are not in the lowest available energy level

  33. Valence Electrons • Electrons in the outermost energy level

  34. Unit 3Nuclear Chemistry Chapter 25 Reference Tables N, O

  35. Radioisotopes • Nuclei of unstable isotopes are called radioisotopes. • An unstable nucleus releases energy by emitting radiation during the process of radioactive decay • Mass and/or energy

  36. Radiation • Late 1800’s – discovery of radiation • Three Types • Alpha • Beta • Gamma

  37. Radiation • Three Types

  38. Symbols Alpha Beta Gamma

  39. Nuclear Stability • For smaller atoms a ratio of 1:1 neutrons to protons helps to maintain stability • C-12, N-14, O-16 • For larger atoms, more neutrons than protons are required to maintain stability • Pb-207, Au-198, Ta-181

  40. Transmutations • Any reaction where one element is transformed into a different element • Nuclear Reactions • Natural • Has one reactant • Alpha and Beta Decay • Artificial • Has more than one reactant • Particle Accelerators

  41. Radioactive Decay • Radioisotopes will undergo decay reactions to become more stable • Alpha Decay • Beta Decay

  42. Half Life • Amount of time for half of a sample to decay into a new element • Parent Atoms • Undecayed atoms • Daughter Atoms • Decayed atoms

  43. Half Life Equations t = total amount of time elapsed T = half-life

  44. Half Life Equations t = amount of time elapsed T = half-life

  45. Half Life Equations Mass Left = Original Mass

  46. Example • How many half lives does it take for a sample of C-14 to be 11430 yrs old?

  47. Example • What fraction of P-32 is left after 42.84days?

  48. Example • How long will a sample of Rn-222 take to decay down to 1/4 of the original sample? 7.646d

  49. Practice • How much Carbon-14 was originally in a sample that contains 4g of C-14 and is 17145 years old? 32g

  50. More Practice • How much 226Ra will be left in a sample that is 4797 years old, if it initially contained 408g? 51g