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Chemistry Midterm Review Presentation!
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Chemistry Midterm Review Presentation!

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  1. Chemistry Midterm Review Presentation! • Aligned to the New York State Standards and Core Curriculum for “The Physical Setting-Chemistry” • Can be used in any high-school chemistry class!

  2. Outline for Review 1) The Atom (Nuclear, Electron Config) 2) Matter (Phases, Types, Changes) 3) Bonding (Periodic Table, Ionic, Covalent) 4) Compounds (Formulas, Reactions, IMAF’s) 5)Math of Chemistry (Formula Mass, Gas Laws, Neutralization, etc.)

  3. The Atom 1) Nucleons 2) Isotopes 6) Electron Configuation 7) Development of the Atomic Model

  4. Nucleons • Protons: +1 each, determines identity of element, mass of 1 amu, determined using atomic number, nuclear charge • Neutrons: no charge, determines identity of isotope of an element, 1 amu, determined using mass number - atomic number (amu = atomic mass unit) • 3216S and 3316S are both isotopes of S • S-32 has 16 protons and 16 neutrons • S-33 has 16 protons and 17 neutrons • All atoms of S have a nuclear charge of +16 due to the 16 protons. (c) 2006, Mark Rosengarten

  5. Isotopes • Atoms of the same element MUST contain the same number of protons. • Atoms of the same element can vary in their numbers of neutrons, therefore many different atomic masses can exist for any one element. These are called isotopes. • The atomic mass on the Periodic Table is the weight-average atomic mass, taking into account the different isotope masses and their relative abundance. • Rounding off the atomic mass on the Periodic Table will tell you what the most common isotope of that element is.

  6. Weight-Average Atomic Mass • WAM = ((% A of A/100) X Mass of A) + ((% A of B/100) X Mass of B) + … • What is the WAM of an element if its isotope masses and abundances are: • X-200: Mass = 200.0 amu, % abundance = 20.0 % • X-204: Mass = 204.0 amu, % abundance = 80.0% • amu = atomic mass unit (1.66 × 10-27 kilograms/amu)

  7. Most Common Isotope • The weight-average atomic mass of Zinc is 65.39 amu. What is the most common isotope of Zinc? Zn-65! • What are the most common isotopes of: • Co Ag • S Pb • FACT: one atomic mass unit (1.66 × 10-27 kilograms) is defined as 1/12 of the mass of an atom of C-12. • This method doesn’t always work, but it usually does. Use it for the Regents exam. (c) 2006, Mark Rosengarten

  8. Electron Configuration • Basic Configuration • Valence Electrons • Electron-Dot (Lewis Dot) Diagrams • Excited vs. Ground State • What is Light? (c) 2006, Mark Rosengarten

  9. Basic Configuration • The number of electrons is determined from the atomic number. • Look up the basic configuration below the atomic number on the periodic table. (PEL: principal energy level = shell) • He: 2 (2 e- in the 1st PEL) • Na: 2-8-1 (2 e- in the 1st PEL, 8 in the 2nd and 1 in the 3rd) • Br: 2-8-18-7 (2 e- in the 1st PEL, 8 in the 2nd, 18 in the 3rd and 7 in the 4th) (c) 2006, Mark Rosengarten

  10. Valence Electrons • The valence electrons are responsible for all chemical bonding. • The valence electrons are the electrons in the outermost PEL (shell). • He: 2 (2 valence electrons) • Na: 2-8-1 (1 valence electron) • Br: 2-8-18-7 (7 valence electrons) • The maximum number of valence electrons an atom can have is EIGHT, called a STABLE OCTET. (c) 2006, Mark Rosengarten

  11. Electron-Dot Diagrams • The number of dots equals the number of valence electrons. • The number of unpaired valence electrons in a nonmetal tells you how many covalent bonds that atom can form with other nonmetals or how many electrons it wants to gain from metals to form an ion. • The number of valence electrons in a metal tells you how many electrons the metal will lose to nonmetals to form an ion. Caution: May not work with transition metals. • EXAMPLE DOT DIAGRAMS (c) 2006, Mark Rosengarten

  12. Example Dot Diagrams Carbon can also have this dot diagram, which it has when it forms organic compounds. (c) 2006, Mark Rosengarten

  13. Excited vs. Ground State • Configurations on the Periodic Table are ground state configurations. • If electrons are given energy, they rise to higher energy levels (excited state). • If the total number of electrons matches in the configuration, but the configuration doesn’t match, the atom is in the excited state. • Na (ground, on table): 2-8-1 • Example of excited states: 2-7-2, 2-8-0-1, 2-6-3 (c) 2006, Mark Rosengarten

  14. What Is Light? • Light is formed when electrons drop from the excited state to the ground state. • The lines on a bright-line spectrum come from specific energy level drops and are unique to each element. • EXAMPLE SPECTRUM (c) 2006, Mark Rosengarten

  15. EXAMPLE SPECTRUM This is the bright-line spectrum of hydrogen. The top numbers represent the PEL (shell) change that produces the light with that color and the bottom number is the wavelength of the light (in nanometers, or 10-9 m). No other element has the same bright-line spectrum as hydrogen, so these spectra can be used to identify elements or mixtures of elements. (c) 2006, Mark Rosengarten

  16. Development of the Atomic Model • Thompson Model • Rutherford Gold Foil Experiment and Model • Bohr Model • Quantum-Mechanical Model (c) 2006, Mark Rosengarten

  17. Thompson Model • The atom is a positively charged diffuse mass with negatively charged electrons stuck in it. (c) 2006, Mark Rosengarten

  18. Rutherford Model • The atom is made of a small, dense, positively charged nucleus with electrons at a distance, the vast majority of the volume of the atom is empty space. Alpha particles shot at a thin sheet of gold foil: most go through (empty space). Some deflect or bounce off (small + charged nucleus). (c) 2006, Mark Rosengarten

  19. Bohr Model • Electrons orbit around the nucleus in energy levels (shells). Atomic bright-line spectra was the clue. (c) 2006, Mark Rosengarten

  20. Quantum-Mechanical Model • Electron energy levels are wave functions. • Electrons are found in orbitals, regions of space where an electron is most likely to be found. • You can’t know both where the electron is and where it is going at the same time. • Electrons buzz around the nucleus like gnats buzzing around your head. (c) 2006, Mark Rosengarten

  21. Properties of Phases • Solids: Crystal lattice (regular geometric pattern), vibration motion only • Liquids: particles flow past each other but are still attracted to each other. • Gases: particles are small and far apart, they travel in a straight line until they hit something,they bounce off without losing any energy, they are so far apart from each other that they have effectively no attractive forces and their speed is directly proportional to the Kelvin temperature (Kinetic-Molecular Theory, Ideal Gas Theory) (c) 2006, Mark Rosengarten

  22. Solids The positive and negative ions alternate in the ionic crystal lattice of NaCl. (c) 2006, Mark Rosengarten

  23. Liquids When heated, the ions move faster and eventually separate from each other to form a liquid. The ions are loosely held together by the oppositely charged ions, but the ions are moving too fast for the crystal lattice to stay together. (c) 2006, Mark Rosengarten

  24. Gases Since all gas molecules spread out the same way, equal volumes of gas under equal conditions of temperature and pressure will contain equal numbers of molecules of gas. 22.4 L of any gas at STP (1.00 atm and 273K) will contain one mole (6.02 X 1023) gas molecules. Since there is space between gas molecules, gases are affected by changes in pressure. (c) 2006, Mark Rosengarten

  25. Types of Matter • Substances (Homogeneous) • Elements (cannot be decomposed by chemical change): Al, Ne, O, Br, H • Compounds (can be decomposed by chemical change): NaCl, Cu(ClO3)2, KBr, H2O, C2H6 • Mixtures • Homogeneous: Solutions (solvent + solute) • Heterogeneous: soil, Italian dressing, etc. (c) 2006, Mark Rosengarten

  26. Elements • A sample of lead atoms (Pb). All atoms in the sample consist of lead, so the substance is homogeneous. • A sample of chlorine atoms (Cl). All atoms in the sample consist of chlorine, so the substance is homogeneous. (c) 2006, Mark Rosengarten

  27. Compounds • Lead has two charges listed, +2 and +4. This is a sample of lead (II) chloride (PbCl2). Two or more elements bonded in a whole-number ratio is a COMPOUND. • This compound is formed from the +4 version of lead. This is lead (IV) chloride (PbCl4). Notice how both samples of lead compounds have consistent composition throughout? Compounds are homogeneous! (c) 2006, Mark Rosengarten

  28. Mixtures • A mixture of lead atoms and chlorine atoms. They exist in no particular ratio and are not chemically combined with each other. They can be separated by physical means. • A mixture of PbCl2 and PbCl4 formula units. Again, they are in no particular ratio to each other and can be separated without chemical change. (c) 2006, Mark Rosengarten

  29. Bonding 1) The Periodic Table 2) Ions 3) Ionic Bonding 4) Covalent Bonding 5) Metallic Bonding (c) 2006, Mark Rosengarten

  30. The Periodic Table • Metals • Nonmetals • Metalloids • Chemistry of Groups • Electronegativity • Ionization Energy (c) 2006, Mark Rosengarten

  31. Metals • Have luster, are malleable and ductile, good conductors of heat and electricity • Lose electrons to nonmetal atoms to form positively charged ions in ionic bonds • Large atomic radii compared to nonmetal atoms • Low electronegativity and ionization energy • Left side of the periodic table (except H) (c) 2006, Mark Rosengarten

  32. Nonmetals • Are dull and brittle, poor conductors • Gain electrons from metal atoms to form negatively charged ions in ionic bonds • Share unpaired valence electrons with other nonmetal atoms to form covalent bonds and molecules • Small atomic radii compared to metal atoms • High electronegativity and ionization energy • Right side of the periodic table (except Group 18) (c) 2006, Mark Rosengarten

  33. Metalloids • Found lying on the jagged line between metals and nonmetals flatly touching the line (except Al and Po). • Share properties of metals and nonmetals (Si is shiny like a metal, brittle like a nonmetal and is a semiconductor). (c) 2006, Mark Rosengarten

  34. Chemistry of Groups • Group 1: Alkali Metals • Group 2: Alkaline Earth Metals • Groups 3-11: Transition Elements • Group 17: Halogens • Group 18: Noble Gases • Diatomic Molecules (c) 2006, Mark Rosengarten

  35. Group 1: Alkali Metals • Most active metals, only found in compounds in nature • React violently with water to form hydrogen gas and a strong base: 2 Na (s) + H2O (l)  2 NaOH (aq) + H2 (g) • 1 valence electron • Form +1 ion by losing that valence electron • Form oxides like Na2O, Li2O, K2O (c) 2006, Mark Rosengarten

  36. Group 2: Alkaline Earth Metals • Very active metals, only found in compounds in nature • React strongly with water to form hydrogen gas and a base: • Ca (s) + 2 H2O (l)  Ca(OH)2 (aq) + H2 (g) • 2 valence electrons • Form +2 ion by losing those valence electrons • Form oxides like CaO, MgO, BaO (c) 2006, Mark Rosengarten

  37. Groups 3-11: Transition Metals • Many can form different possible charges of ions • If there is more than one ion listed, give the charge as a Roman numeral after the name • Cu+1 = copper (I) Cu+2 = copper (II) • Compounds containing these metals can be colored. (c) 2006, Mark Rosengarten

  38. Group 17: Halogens • Most reactive nonmetals • React violently with metal atoms to form halide compounds: 2 Na + Cl2 2 NaCl • Only found in compounds in nature • Have 7 valence electrons • Gain 1 valence electron from a metal to form -1 ions • Share 1 valence electron with another nonmetal atom to form one covalent bond. (c) 2006, Mark Rosengarten

  39. Group 18: Noble Gases • Are completely nonreactive since they have eight valence electrons, making a stable octet. • Kr and Xe can be forced, in the laboratory, to give up some valence electrons to react with fluorine. • Since noble gases do not naturally bond to any other elements, one atom of noble gas is considered to be a molecule of noble gas. This is called a monatomic molecule. Ne represents an atom of Ne and a molecule of Ne. (c) 2006, Mark Rosengarten

  40. Diatomic Molecules • Br, I, N, Cl, H, O and F are so reactive that they exist in a more chemically stable state when they covalently bond with another atom of their own element to make two-atom, or diatomic molecules. • Br2, I2, N2, Cl2, H2, O2 and F2 • The decomposition of water: 2H2O  2 H2 + O2 (c) 2006, Mark Rosengarten

  41. Electronegativity • An atom’s attraction to electrons in a chemical bond. • F has the highest, at 4.0 • Fr has the lowest, at 0.7 • If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion. • If the two atoms have an END of less than 1.7, they will share their unpaired valence electrons…covalent bond! (c) 2006, Mark Rosengarten

  42. Ionization Energy • The energy required to remove the most loosely held valence electron from an atom in the gas phase. • High electronegativity means high ionization energy because if an atom is more attracted to electrons, it will take more energy to remove those electrons. • Metals have low ionization energy. They lose electrons easily to form (+) charged ions. • Nonmetals have high ionization energy but high electronegativity. They gain electrons easily to form (-) charged ions when reacted with metals, or share unpaired valence electrons with other nonmetal atoms. (c) 2006, Mark Rosengarten

  43. Ions • Ions are charged particles formed by the gain or loss of electrons. • Metals lose electrons (oxidation) to form (+) charged cations. • Nonmetals gain electrons (reduction) to form (-) charged anions. • Atoms will gain or lose electrons in such a way that they end up with 8 valence electrons (stable octet). • The exceptions to this are H, Li, Be and B, which are not large enough to support 8 valence electrons. They must be satisfied with 2 (Li, Be, B) or 0 (H). (c) 2006, Mark Rosengarten

  44. Metal Ions (Cations) Note that when the atom loses its valence electron, the next lower PEL becomes the valence PEL. Notice how the dot diagrams for metal ions lack dots! Place brackets around the element symbol and put the charge on the upper right outside! • Na: 2-8-1 • Na+1: 2-8 • Ca: 2-8-8-2 • Ca+2: 2-8-8 • Al: 2-8-3 • Al+3: 2-8 (c) 2006, Mark Rosengarten

  45. Nonmetal Ions (Anions) Note how the ions all have 8 valence electrons. Also note the gained electrons as red dots. Nonmetal ion dot diagrams show 8 dots, with brackets around the dot diagram and the charge of the ion written to the upper right side outside the brackets. • F: 2-7 • F-1: 2-8 • O: 2-6 • O-2: 2-8 • N: 2-5 • N-3: 2-8 (c) 2006, Mark Rosengarten

  46. Ionic Bonding • If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion. • The oppositely charged ions attract to form the bond. It is a surface bond that can be broken by melting or dissolving in water. • Ionic bonding forms ionic crystal lattices, not molecules. (c) 2006, Mark Rosengarten

  47. Example of Ionic Bonding (c) 2006, Mark Rosengarten

  48. Covalent Bonding • If two nonmetal atoms have an END of 1.7 or less, they will share their unpaired valence electrons to form a covalent bond. • A particle made of covalently bonded nonmetal atoms is called a molecule. • If the END is between 0 and 0.4, the sharing of electrons is equal, so there are no charged ends. This is NONPOLAR covalent bonding. • If the END is between 0.5 and 1.7, the sharing of electrons is unequal. The atom with the higher EN will be d- and the one with the lower EN will be d+ charged. This is a POLAR covalent bonding. (d means “partial”) (c) 2006, Mark Rosengarten

  49. Examples of Covalent Bonding (c) 2006, Mark Rosengarten

  50. Metallic Bonding • Metal atoms of the same element bond with each other by sharing valence electrons that they lose to each other. • This is a lot like an atomic game of “hot potato”, where metal kernals (the atom inside the valence electrons) sit in a crystal lattice, passing valence electrons back and forth between each other). • Since electrons can be forced to travel in a certain direction within the metal, metals are very good at conducting electricity in all phases. (c) 2006, Mark Rosengarten