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Let’s have some fun with CHEMISTRYYYY!

Let’s have some fun with CHEMISTRYYYY!. Let’s have some fun with CHEMISTRYYYY!. F. BASIC STRUCTURES OF MATTER. ATOMS basic building blocks of matter STRUCTURE: protons (p) : positively charged sub-atomic particle electron (e-) : negatively charged sub-atomic particle

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Let’s have some fun with CHEMISTRYYYY!

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  1. Let’s have some fun with CHEMISTRYYYY!

  2. Let’s have some fun with CHEMISTRYYYY!

  3. F. BASIC STRUCTURES OF MATTER • ATOMS • basic building blocks of matter • STRUCTURE: • protons (p): positively charged sub-atomic particle • electron (e-): negatively charged sub-atomic particle • neutron (n): electrically neutral

  4. F. BASIC STRUCTURES OF MATTER • Mass Number A M C Z • Charge • Atomic Symbol • Atomic Number Atoms of the same element having the same number of protons (same atomic number) but different number of neutrons and different atomic mass are called isotopes. • M = (# of p+) + (# of n0) • C = (# of p+) - (# of e-) • A = # of p+

  5. ORBITALS AND QUANTUM NUMBERS • Quantum numbers: describe certain aspects of the atom • Orbital: specific distribution of space orientations between electrons • Electron shell: orbitals with the same value n/1stor principal quantum number • Subshell: same set of n and l (2nd quantum number)

  6. TYPES OF QUANTUM NUMBERS • First/Principal (n): energy level of the electron (the higher, the more energy/size) [for the maximum number of orbitals = n2) • Secondary/Azithumal (l = 0 until n-1): sublevel within the energy level = shape of the of the orbital [s = 0 = spherical, p = 1 = dumbell, d = 2 = clover , f = 3 = more complex] • Third/Magnetic (ml): the particular orbital and its space orientation [ s = 1, p = 3, d = 7, f = 10] • Fourth/Spin (ms): direction of the spin [+ ½ or – ½ ]

  7. RELATIONSHIP BETWEEN QUANTUM NUMBERS

  8. PAULI’S EXCLUSION PRINCIPLE • A maximum of two electrons can occupy an orbital (with different spins) • No two electrons can have the same quantum numbers ELECTRON CONFIGURATION

  9. PERIODIC LAW • Physical and chemical properties of elements are the periodic functions of their atomic number • Elements with similar properties = similar arrangement of outer shell electrons/same group • Valence electrons: found in the outermost shell of elements in their ground state

  10. PERIODIC TRENDS • Molecules: two or more atoms tightly bonded together • Ion: atom with a gained or loss electron (anion: negatively charged; cation: positively charged)

  11. INTRAMOLECULAR FORCES OF ATTRACTION • Covalent bond: between 2 atoms; by electron sharing (nonmetallic + nonmetallic) • Ionic bond: between + and – ions; form crystal structures; by electron transfer (nonmetallic + metallic) • Metallic bond: between free-floating valence electrons = high conductivity and luster (metallic + metallic)

  12. VAN DER WAAL’S FORCES OF ATTRACTION (IMFA’s) • Between atoms of different molecules • London dispersion forces: weakest, between nonpolar atoms and molecules (nonpolar + nonpolar) (CH3-CH3) • Dipole-dipole forces: between polar molecules (CO) • Hydrogen bond: between a hydrogen atom of a polar molecule + (any) F/O/N atoms in another polar molecule (H20, NH3) • Ion-dipole forces: strongest, between polar molecules/ions in solutions (ionic molecules + ionic salts/polar solvents)

  13. WRITING CHEMICAL FORMULAS • Chemical formulas: denote the number and kinds of atoms in a compound (reacting substances) • Represent with symbols (cation first, then anion). • Indicate oxidation states (to the upper right, put parentheses for polyatomic ions). • Write the subscript equal to the oxidation number of the other element (‘switch’) [omit the ‘1’s’]. • Reduce subscripts to their lowest terms.

  14. FORMULAS TO REMEMBER! • FORMULA WEIGHT/MASS or MOLECULAR WEIGHT/MASS • sum of all atomic weight (AW) (g/mol) 2. PERCENTAGE COMPOSITION

  15. FORMULAS TO REMEMBER! 3. EMPIRICAL AND MOLECULAR FORMULA a.) convert mass in grams to mole ** assume 100 g of any sample so that % composition = value of mass in g b.) get the smallest whole number ratio of the elements ** divide mole value of each element by the smallest mole value THIS WILL BE YOUR EMPIRICAL FORMULA (EF) c.) ** get EFM the same way you get FM

  16. FORMULAS TO REMEMBER! 4. STOICHIOMETRY ** Avogrado’s number = 6.02 x 1023 molecules ** Molar Mass = mass in grams of 1 mole of a substance CONVERSION GRAMS MOLES MOLECULES

  17. FORMULAS TO REMEMBER! 4. REACTION STOICHIOMETRY CONVERSION use molar mass of Ause molar mass of B grams of moles of moles of grams of substance A substance A substance B substance B use coefficients of A and B from theBALANCED EQUATION

  18. FORMULAS TO REMEMBER! 4. SOLUTION STOICHIOMETRY a.) Molarity (M) a.) Molality (m) 5. DILUTION

  19. FORMULAS TO REMEMBER! e.) Avogrado’s Law f.) Daton’s Law … + g.) Graham’s Law h.) Ideal Gas Law R = 0.0821 L-atm/mol-K 6. GAS LAWS a.) Boyle’s Law b.) Charles’ Law c.) Gay- Lussac’s Law d.) Combined Gas Law

  20. FORMULAS TO REMEMBER! 7. GAS DENSITIES AND MOLAR MASS At STP, standard temperature: 0°C = 273.15 K = 32°F standard pressure: 1 atm = 760 torr = 760 mmHg molar volume of gas: 1 mol = 22.4 L

  21. KINETIC MOLECULAR THEORY OF GASES This states that ideal gases: 1. the movement of particles is in continuous random motion 2. pressure of gas in due to the bombardment of the molecules to the container 3. collisionsbetween or among particles are elastic 4. the kinetic energyof the system is proportional to temperature 5. the wide separation between molecules cause the attractive and repulsive forces to be negligible

  22. PRACTICE QUESTIONS!

  23. A substance composed of two or more elements chemically united is called • an isotope • an element • a mixture • a compound • When energy, like light and heat, is liberated during a reaction such as burning of fuels, this type of reaction occurs. • endothermic • nuclear reaction • exothermic • both a and b

  24. 3. Chemical action may involve all of the following EXCEPT • combination of atoms of all elements to form a molecule • breaking down of compounds into elements • reacting a compound and an element • separation of the components of a mixture 4. When decomposed chemically, 73 g of a sample of HCl produces 71 g of Cl2 and 2 g of H2, while 34 g of H2S sample produces 32 g of S and 2 g of H2. This is an example of: • Law of Conservation of Mass • Law of Thermodynamics • Law of Multiple Proportions • Law of Definite Proportions

  25. 5. The theory states that no 2 electrons within the same atom can have the same set of quantum numbers. • Hund’s Rule • Aufbau’s Principle • Pauli’s Exclusion Theory • Dalton’s Rule 6. Elements with the most stable configuration belong to: • transition metals • inert gases • alkali metals • halogen family

  26. 7. In the periodic table, which can also indicate the highest energy level or the highest principal quantum number of an atom? • group • atomic number • period • atomic mass 8. The total number of orbitals in the fourth energy level is: • 4 • 16 • 8 • 18

  27. 9. Sodium chloride, most commonly known as table salt, has a pH of: • = 7 • < 7 • > 7 • = 0 10. What do you call the solid left after the separation of mixtures? • solvent • precipitate • mother liquor • supernatant

  28. 11. From the given statements below, which is INCORRECT? • All atoms of a given element are identical. • Atoms combine in small whole number ratio. • All atoms change their chemical identity during a chemical reaction. • Elements can combine to form compounds. 12. Isotopes, which are the reasons why elements have no whole mass numbers, have: • the same number of protons but different number of electrons • the same atomic number but different mass numbers • the same number of protons but different number of neutrons • both b and c

  29. 13. The intermolecular forces that exist between the molecules of CCl4, a non-polar compound, is/are: • dipole-dipole • London forces • Hydrogen bond • both a and b 14. If Magnesium forms a compound with chlorine, what is the formula of the molecule? • Mg2Cl • MgCl2 • Mg2Cl3 • MgCl3

  30. 15. Which of the following statements is FALSE for a neutral atom? • number of neutrons is not always equal to the number of protons • number of electrons in the atom is equal to the number of protons • nucleus has a positive charge • nucleus has a neutral charge 16. The maximum number of electrons that can be placed in the second principal energy level, n=2, of an atom is: • 6 • 8 • 4 • 2

  31. 17. Why are you advised not to heat a tightly closed vessel? • heat increases the volume of the vessel • pressure might decrease causing the vessel to implode • pressure will build up causing the vessel to explode • heat might be trapped inside the vessel 18. Gases will approach ideality (ideal state) at: • low pressure and high temperature • low pressure and low temperature • high pressure and high temperature • high pressure and high temperature

  32. 19. What should be the last 2 coefficients when the following reaction is balanced? FeS2 + O2 ___ Fe2O3 + ___ SO2 • 2, 3 • 2, 4 • 2, 6 • 2, 8 20. Gas flows from an area of _____ pressure to an are a _____ pressure. • higher-lower • equal-equal • medium-high • lower-higher

  33. 21. What type of reaction occurred in the following: 4 NiCO3 + O2 2 Ni2O3 + 4 CO2 + heat? • double displacement • single replacement • combustion • decomposition 22. The reaction of ethylene, C2H4(g) and hydrogen chloride, HCl(g) to form ethyl chloride C2H5Cl(g) is an example of what type of reaction? • substitution • synthesis • neutralization • decomposition

  34. 23. A real (or non-ideal) gas will: • have a higher kinetic energy than ideal gases • have a higher boiling than ideal gases • not have negligible attractive and repulsive forces between its molecules • conform to the kinetic molecular theory of gases 24. The sum of the partial pressures of gas components is equal to the total pressure of the system. This is stated in which law? • Boyle’s Law • Combined Gas Law • Dalton’s Law • Gay-Lussac’s Law

  35. 25. Which of the following are not STP conditions for 1 mole of ideal gas? • 760 torr • 273 °C • 0°C • 22.4 L 26. Which of the following is not stated in the Kinetic Molecular Theory of Gases? • temperature of gas is related to the kinetic energy of the molecules • pressure is due to the bombardment of molecules to walls of container • the larger the molecule of a gas, the greater its kinetic energy • collisions are elastic

  36. 27. Which has more atoms: 100 mol of Cu, 100 mol of Ag, 100 mol of Au, or 100 mol of Hg? • Hg • Au • Cu • none of the above 28. How many moles of oxygen atoms are there in 3 moles of Ca3(PO4)2? • 8 • 4 • 24 • 36

  37. EXERCISES! -2 127 p+ = A = n0 = M - A = 127 – 52 = e-= p+ – C = 52 – (-2) = 52 54 75

  38. EXERCISES! 75 52 54 -2 127 4 197 p+ = A = n0 = M - A = 197 – 79 = e-= p+ – C = 79 – 4 = 79 75 118

  39. EXERCISES! 118 75 52 79 54 75 -2 4 127 197 3 e-= p+ – C = 51 - 3 = M = n0 + A = 72 + 51 = 48 123

  40. EXERCISES! Some AW in amu or g/mol H = 1.0 Na = 23.0 N = 14.0 C = 12.0 O = 16.0 Al = 27.0 • Find the formula weight of: • H2O = 2 (AWH) + AWO = 2 (1) + 16 = 18 g/mol • NaOH = AWNa + AWO + AWH = 23 + 16 + 1 = 40 g/mol • Al(NO3)3 = AWAl + 3 (AWN) + 9 (AWO) = 27 + 3 (14) + 9 (16) = 213 g/mol

  41. EXERCISES! given: MWcompound = 30 g/mol 80% C ; 20% H (assume 100g) divide by their AW • MF = MF = MF = 2 MF: 2 (CH3) MF: C2H6 EF: CH3 ( 6.67 ( 1 80 g ( 3 ( 20 20 g by smallest mole value

  42. EXERCISES! Some AW in amu or g/mol H = 1.0 Na = 23.0 O = 16.0 • 1.5 molNaOH = ____ g NaOH = ____ molecules NaOH ** FW = 23 + 16 + 1 = 40 g/mol • 1.5 molNaOH = 60 g NaOH • 1.5 molNaOH = 9 x 1023 molecules NaOH

  43. EXERCISES! ** make sure the equation is BALANCED! Some AW in amu or g/mol H = 1.0 O = 16.0 ** 2 H2 + 1 O2 2 H2O • 2 mol O = ____ mol HO 2mol O = 4 mol H2O • 10 g H= ____ g HO 10 g H = 90 g H2O from FW from balanced equation from FW

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