MR. bogo chem final

Honors Semester 1 Study Guide 2021-2022

Scientific Method

1. Basic Sequence of Events:

   • A lot of background research

   • Development of a hypothesis (testable prediction)

   • Test the prediction via experiment, observation, data analysis, etc.

   • Analyze the data

   • Draw a conclusion (was hypothesis correct?)

   • Share research with other scientists

   • Reformulate, more research, future experiments

Definitions and Concepts

2. Hypothesis:

   • A testable prediction

3. Difference between Theory and Hypothesis:

   • Hypothesis: A testable prediction about a particular causal relationship.

     • Example: "Giving a plant magnesium will increase its growth."

   • Theory: Provides an explanation for a large set of observations and experimental data.

     • Example: The theory of natural selection explains both giraffe neck length and bacterial resistance to antibiotics.

4. Definitions:

   • Independent Variable: The variable being tested/manipulated in an experiment.

   • Dependent Variable: The variable being measured in an experiment.

   • Controlled Variables: The variables that are kept constant in an experiment.

5. Control Group:

   • Used for comparison with the experimental group to determine the effect of the independent variable on the dependent variable.

     • Example: In COVID vaccine trials, the experimental group received the vaccine, while the control group received a placebo.

6. Importance of Communication in Science:

   • Scientific results are shared for checking flaws and ensuring reliability of data.

   • Allows repetition of experiments by other scientists or construction of further research based on shared data.

Metric System and Conversions

7. Base Units of the Metric System:

   • Length = meter (m)

   • Mass = gram (g)

   • Volume = liter (L)

8. Metric Prefixes:

   • Terra - T = $10^{12}$

   • Giga - G = $10^{9}$

   • Mega - M = $10^{6}$

   • Kilo - k = $10^{3}$

   • Hecta - h = $10^{2}$

   • Deca - da = $10^{1}$

   • Base units (m, L, g) = $10^{0}$ = 1

   • Deci - d = $10^{-1}$

   • Centi - c = $10^{-2}$

   • Milli - m = $10^{-3}$

   • Micro - μ = $10^{-6}$

   • Nano - n = $10^{-9}$

   • Pico - p = $10^{-12}$

9. Dimensional Analysis (Conversions):

   • 256 m = $2.56 imes 10^{4}$ cm

   • 97.25 cm = $9.725 imes 10^{2}$ mm

   • 952 g = $9.52 imes 10^{5}$ mg

   • 0.574 pm = $5.74 imes 10^{7}$ μm

   • 5.287 = $5.287 imes 10^{3}$ mL

   • 785.3 hg = $7.853 imes 10^{4}$ g

   • 84.363 km = $8.4363 imes 10^{6}$ cm

   • 872 km = $8.72 imes 10^{8}$ mm

   • 95,824 cm = $9.5824 imes 10^{5}$ mm

   • 8.26 kl = $8.26 imes 10^{6}$ mL

Significant Figures and Measurements

10. Significant Figure Rules for Lab Measurements:

    • Determine the smallest decimal place accurately measured by the instrument; estimate one more place smaller.

11. Example Measurement (Graduated Cylinder):

    • 36.5 mL

12. Another Example Measurement:

    • 20.39 mL

13. Rules for Significant Figures:

    • All non-zero numbers in a measurement are significant.

    • Leading zeros are not significant.

    • Zeros between non-zero numbers and trailing zeros after the decimal are significant.

      • Examples:

      • 0.000345 = 3 sig figs

      • 2.034 = 4 sig figs

      • 0.2300 = 4 sig figs

14. Rule for Significant Figures in Addition/Subtraction:

    • Final answer is rounded to the least accurate decimal place.

      • Example: $2.3 + 5.44 = 7.74
        ightarrow 7.7$ (tenths place)

15. Rule for Significant Figures in Multiplication/Division:

    • Final answer is rounded to the least number of significant figures.

      • Example: $2.3 imes 1.778
        ightarrow 4.09
        ightarrow 4$ (2 sig figs)

16. Perform Calculations with Correct Significant Figure Rules:

    • Calculation: $3.2 imes (4.05 - 0.245)/(3.3456 + 23.34534)$

      • $3.2(4.05 - 0.245)/(3.3456 + 23.34534) = 0.46$ because of significant figures

    • Next Calculation: $(23.56 imes 214)/(4.234 - 2.23) = 2.52 imes 10^{3}$

    • $23.234(4.56/3.2) + 4.56
      ightarrow 38$

    • $8.34 imes (2.3/4.88) + (1.111 - 0.0023)
      ightarrow 5.0$

    • $2((1.2223 imes 4.35) - (12.34 imes 6.66)) = -2 imes 10^{2}$

Density and Concentration

17. Density Calculation of Copper:

    • Density = Mass/Volume

      • $Density = rac{89.6000 ext{ g}}{10.00 ext{ cm}^3} = 8.960 ext{ g/cm}^3$

18. Percent Error Calculation for Hydrochloric Acid Sample:

    • $Percent ext{ Error} = rac{(0.1364 - 0.1355)}{0.1364} imes 100 ext{%} = 0.7 ext{%}$

Atomic Theory

19. Dalton's Atomic Theory Axioms:

    • All matter is made of tiny indivisible particles called atoms.

    • Atoms of the same element have identical mass and physical properties.

    • Atoms cannot be subdivided, created, or destroyed.

    • Atoms of different elements can combine in simple whole-number ratios.

    • Chemical reactions involve the rearrangement of atoms.

20. Models of the Atom:

    • Dalton's Model: Indivisible, neutral sphere (no parts).

    • JJ Thomson's Model: Positively charged sphere with negative electrons embedded; known as the "Plum Pudding" model.

    • Rutherford's Model: Positively charged protons and neutral neutrons in a very small, dense nucleus; electrons floating randomly around the nucleus; atom is mostly empty space.

    • Bohr Model: Positively charged protons and neutral neutrons in a nucleus; electrons orbit the nucleus at different energy levels.

    • Quantum (Electron Cloud) Model: Electrons exist in an electron cloud in orbitals with different shapes and energy levels; exact location of electrons cannot be known, only their probability of location.

21. JJ Thomson's Cathode Ray Tube Experiment:

    • Thomson experimented with cathode rays (beams of electricity in a vacuum).

    • He discovered that the beam was made of negatively charged particles (electrons), leading to the conclusion that atoms must contain smaller, negatively charged parts and also a positive aspect.

    • This shifted the atomic model from an indivisible sphere to a model with electrons embedded in a positive sphere.

22. Rutherford's Gold Foil Experiment:

    • Rutherford fired positive alpha particles at gold foil, expecting them to pass through (due to a spread-out positive charge).

    • Surprisingly, some particles deflected, leading him to postulate a concentrated positive nucleus.

    • This transformed the atomic model from Thomson's to one with a nucleus containing protons and neutrons.

23. Bohr's Experiments with the Hydrogen Emission Spectrum:

    • Bohr concluded that electrons can have specific energy values and orbit the nucleus at these levels.

Subatomic Particles

24. Subatomic Particles:

    • Proton: + charge; 1 amu; located in the nucleus.

    • Electron: - charge; 0 amu; located in the electron cloud.

    • Neutron: 0 charge; 1 amu; located in the nucleus.

25. Element Symbol Explanation:

    • Mass Number: Number of protons + number of neutrons.

    • Atomic Number: Number of protons.

26. Definitions:

    • Isotope: Atoms of the same element with different numbers of neutrons.

    • Ion: Charged atom;

    • Cation: Positive ion (lost electrons);

    • Anion: Negative ion (gained electrons).

27. Mass Number vs. Relative Atomic Mass:

    • Mass Number: Protons + neutrons for a specific atom.

    • Relative Atomic Mass: Weighted average of masses of all isotopes of an element.

28. Average Atomic Mass of Uranium Calculation:

    • For isotopes $^{234} ext{U}$ (0.01%), $^{235} ext{U}$ (0.71%), $^{238} ext{U}$ (99.28%):

      • Average atomic mass $= (234 imes 0.0001) + (235 imes 0.0071) + (238 imes 0.9928) = 237.98 ext{ amu} $

Quantum Mechanics

29. Quantum Numbers:

    • Principal Quantum Number (n): Energy levels (1-7).

    • Subshell (l): Orbitals with different shapes and sizes (s, p, d, f).

    • Magnetic Quantum Number (m_l): Orientation of a particular orbital in space.

    • Spin Quantum Number (m_s): Orientation of an electron relative to an external magnetic field; +½, -½.

30. Energy Levels:

    • There are 7 energy levels, which correspond to the rows on the periodic table.

31. Four Subshells:

    • s, p, d, f

32. Principles of Quantum Mechanics:

    • Aufbau Principle: Electrons fill energy levels in an atom in order from lowest to highest energy.

    • Hund's Rule: Within a particular subshell, electrons will spread into their own orbitals before pairing, reducing electrostatic repulsion; all take on the same spin.

    • Pauli Exclusion Principle: A pair of electrons in the same orbital must have opposite spins.

33. Order of Energy in Subshells:

    • s < p < d < f

34. Electron Configurations of Elements:

    • C: $1s^2 2s^2 2p^2$

    • P: $1s^2 2s^2 2p^6 3s^2 3p^3$

    • Na: $1s^2 2s^2 2p^6 3s^1$

    • Cr: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^5 4s^1$ (exception)

    • Cu: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^1$

35. Noble Gas Configurations:

    • Ca: [Ar] $4s^2$

    • Fe: [Ar] $3d^6 4s^2$

    • Pb: [Xe] $4f^{14} 5d^{10} 6s^2 6p^2$

    • Au: [Xe] $4f^{14} 5d^{10} 6s^1$ (exception)

    • Include configurations for ions. Anions gain electrons, cations lose electrons.

36. Absorption Spectrum vs. Emission Spectrum:

    • Absorption Spectrum: Electrons absorb EM radiation and jump to a higher energy level, resulting in black lines on a colored background.

    • Emission Spectrum: Electrons drop to a lower energy level and emit photons, creating colored lines on a black background.

Emission Spectrum of Hydrogen

37. Hydrogen Emission Spectrum with Series Identification:

    • Lyman Series: UV range; involves transitions to n=1.

    • Balmer Series: Visible range; involves transitions to n=2.

Radioactive Decay

38. General Equations for Radioactive Decay Types:

    • Alpha Decay:
      $A<em>X^Z ightarrow A</em>{X-2}^{Z-2} + ^{4}_{2} ext{He}$

    • Beta Negative Decay:
      $A<em>X^Z ightarrow A</em>{X}^{Z+1} + e^{-}$

    • Gamma Decay:
      $A<em>X^Z ightarrow A</em>X^Z + ext{energy}$

    • Positron Emission:
      $A<em>X^Z ightarrow A</em>{X-1}^{Z} + e^{+}$

    • Electron Capture:
      $A<em>X^Z + e^{-} ightarrow A</em>{X-1}^{Z}$

39. Specific Decay Equations:

    • Alpha decay of Radon-198:
      $^{198}<em>{86} ext{Rn} ightarrow ^{194}</em>{84} ext{Po} + ^{4}_{2} ext{He}$

    • Beta decay of Uranium-237:
      $^{237}<em>{92} ext{U} ightarrow ^{237}</em>{93} ext{Np} + e^{-}$

    • Alpha decay of Radium-226:
      $^{226}<em>{88} ext{Ra} ightarrow ^{222}</em>{86} ext{Rn} + ^{4}_{2} ext{He}$

    • Positron emission of Fluorine-18:
      $^{18}<em>{9} ext{F} ightarrow ^{18}</em>{8} ext{O} + e^{+}$

    • Beta decay of Scandium-46:
      $^{46}<em>{21} ext{Sc} ightarrow ^{46}</em>{22} ext{Ti} + e^{-}$

    • Electron capture of Aluminum-26:
      $^{26}<em>{13} ext{Al} + e^{-} ightarrow ^{26}</em>{12} ext{Mg}$

40. Definition of Half-Life:

    • Half-life = the time it takes for exactly one half of a radioisotope to decay:

      • Formula: $A = A_0 (1/2)^{t/h}$

41. Iodine-131 Half-Life Calculation:

    • Given a starting mass of 100.0 g, after 32 days:

      • $A = 100.0 ext{g} imes (1/2)^{32/8} = 6.250 ext{g}$

42. Chromium-51 Half-Life Calculations:

    • After 56 days:

      • $A = 510.0 ext{g} imes 2^{56/28} = 127.5 ext{g}$

    • After 1 year (365 days):

      • $A = 510.0 ext{g} imes (1/2)^{365/28} = 0.0673 ext{g}$

    • Amount present 168 days ago:

      • $510.0 ext{g} = (A_0) imes 2^{168/28} = 32,640 ext{g}$

Periodic Table Organization

43. Mendeleev's First Periodic Table Organization:

    • Organized by atomic mass.

44. Modern Periodic Table Organization:

    • Organized by atomic number (number of protons).

45. Distinction Between Periods and Groups:

    • Periods: Rows in the periodic table.

    • Groups: Columns in the periodic table.

46. Locations of Metals, Metalloids, and Nonmetals:

    • Metals: Left side and middle of the periodic table

    • Nonmetals: Right side of the periodic table

    • Metalloids: Located along the staircase.

47. General Properties of Metals:

    • Shiny

    • Dense

    • Malleable & ductile

    • Good conductors of electricity and heat

    • High melting and boiling points

48. General Properties of Nonmetals:

    • Dull

    • Low densities

    • Brittle

    • Poor conductors (insulators)

    • Low melting and boiling points

49. General Properties of Metalloids:

    • Intermediate properties between metals and nonmetals

    • Shiny appearance

    • Lower densities

    • Brittle nature

    • Intermediate conductors (semiconductors)

    • Intermediate melting and boiling points

50. Location of Various Groups on Periodic Table:

    • Alkali metals - Group 1

    • Alkaline earth metals - Group 2

    • Transition metals - D block

    • Post-transition metals - P block

    • Metalloids - On “staircase”

    • Halogens - Group 17

    • Noble gases - Group 18

    • Lanthanides - 4f block

    • Actinides - 5f block

51. Reason for Similar Properties in Same Group:

    • Elements in the same group have the same number of valence electrons, resulting in similar chemical properties.

52. Definition of Valence Electrons:

    • Valence electrons included s & p electrons from the highest energy level; outermost electrons.

53. Definition of Periodic Trend:

    • A periodic trend is one that repeats; trends on the periodic table repeat with each row (period).

Physical Trends in Periodic Table

54. Definitions and Trends:

    • Effective Nuclear Charge (Zeff):

      • $Zeff = ext{Protons} - ext{shielding electrons}$

      • Measure of electron attraction to the nucleus.

      • Increases left to right across a period; decreases down a group.

    • Atomic Radius:

      • Decreases from left to right across a period; increases down a group.

    • Ionization Energy:

      • Minimum energy needed to remove an electron.

      • Increases from left to right across a period; decreases down a group.

    • Electron Affinity:

      • Energy change when adding electrons.

      • Increases from left to right across a period; decreases down a group.

    • Electronegativity:

      • Tendency to attract a pair of electrons in a bond.

      • Increases from left to right across a period; decreases down a group.

    • Ionic Radius:

      • Neutral atoms to ions: metals lose electrons, decreasing radius; nonmetals gain electrons, increasing ionic radius.

      • For ionic radii comparison, it decreases left to right for both metal and non-metal ions.

    • Reactivity:

      • Metal reactivity increases to the left and down; nonmetal reactivity increases to the right and up.

55. Ranking Elements in Terms of Effective Nuclear Charge:

    • $F > Cl > Br > I$ (effective nuclear charge decreases with increasing radius)

56. Ranking Elements in Terms of Effective Nuclear Charge:

    • $Na < Al < S < Cl$

57. Ranking Elements in Terms of Effective Nuclear Charge:

    • $Rb < Mg < Si < F$

58. Ranking Elements in Terms of Atomic Radius:

    • $I > Br > Cl > F$

59. Ranking Elements in Terms of Atomic Radius:

    • $Na > Al > S > Cl$

60. Ranking Elements in Terms of Atomic Radius:

    • $Rb > Mg > Si > F$

61. Ranking Elements in Terms of Ionization Energy:

    • $I < Br < Cl < F$

62. Ranking Elements in Terms of Ionization Energy:

    • $Na < Al < S < Cl$

63. Ranking Elements in Terms of Ionization Energy:

    • $Rb < Mg < Si < F$

64. Ranking Elements in Terms of Electronegativity:

    • $I < Br < Cl < F$

65. Ranking Elements in Terms of Electronegativity:

    • $Na < Al < S < Cl$

66. Ranking Elements in Terms of Electronegativity:

    • $Rb < Mg < Si < F$

67. Ranking Elements in Terms of Electron Affinity:

    • $I < Br < Cl < F$

68. Ranking Elements in Terms of Electron Affinity:

    • $Na < Al < S < Cl$

69. Ranking Elements in Terms of Electron Affinity:

    • $Rb < Mg < Si < F$

70. Ranking Metals in Terms of Reactivity:

    • $Cs > K > Na > Li$

71. Ranking Nonmetals in Terms of Reactivity:

    • $I < Br < Cl < F$

Chemical Formulas and Bonding

72. Naming/Writing Formulas for Compounds:

    • NH3 = ammonia

    • AIP = aluminum phosphide

    • CaF2 = calcium fluoride

    • P4O6 = tetraphosphorus hexaoxide

    • CH4 = methane

    • H2S = hydrosulfuric acid

    • Cu3PO4 = copper(I) phosphate

    • HNO3 = nitric acid

    • HI = hydroiodic acid

    • NaCl = sodium chloride

    • MgO = magnesium oxide

    • O2 = oxygen

    • H2SO3 = sulfurous acid

    • P4 = phosphorus

    • FeCO3 = iron(II) carbonate

    • Na2O = sodium oxide

    • PbCr2O7 = lead (II) dichromate

    • BCl3 = boron trichloride

    • KClO3 = potassium chlorate

    • SF6 = sulfur hexafluoride

    • SnO2 = tin (IV) oxide

    • Mn3(PO4)7 = manganese (VII) phosphate

    • AuCl = gold (I) chloride

    • AgF = silver fluoride

    • Zn3P2 = zinc phosphide

    • HC2H3O2 = acetic acid

    • Na3C6H5O7 = sodium citrate

    • H2 = hydrogen

    • HCN = hydrocyanic acid

    • NH4OH = ammonium hydroxide

73. Reason Why Chemical Bonds Form:

    • Atoms form bonds to enter a lower potential energy state.

74. Ionic Bond Formation:

    • Ionic bonds form when oppositely charged ions (cations & anions) are attracted to each other through electrostatic attraction, which occurs when one atom transfers electrons to another.

75. Ionic Crystal Lattice:

    • Crystalline structure made up of ions bonded together in a repeating pattern.

76. General Properties of Ionic Compounds:

    • Brittle crystalline structures

    • High melting and boiling points

    • Solids at room temperature

    • Generally very soluble in water

    • Do not conduct electricity as solids / conduct electricity when molten or dissolved (electrolytes)

77. Definition of Electrolyte:

    • Substance that conducts electricity when dissolved in water; e.g., NaCl is an electrolyte.

78. Covalent Bond Formation:

    • Covalent bonds form when atoms share electrons to complete their valence shells.

79. Ranking of Covalent Bonds:

    • Length: single > double > triple

    • Strength: triple > double > single

80. Simple Molecular vs. Network Covalent Compounds:

    • Simple Molecular Compounds: Form individual, discrete molecules (e.g., water).

    • Network Covalent Compounds: Form networks of covalently bonded atoms; no separate molecules (e.g., diamond, graphite).

81. General Properties of Simple Molecular Covalent Compounds:

    • Low melting and boiling points

    • Liquids and gases at room temperature

    • Generally poor solubility in water

    • Do not conduct electricity in any state

82. General Properties of Network Covalent Compounds:

    • Very high melting points

    • Solids at room temperature

    • Not soluble in water

    • Generally do not conduct electricity, with exceptions for delocalized electrons (graphite).

83. Definition of Allotropes:

    • Different physical forms of the same element in the same state (e.g., diamond and graphite as allotropes of carbon).

84. Structure and Bonding in Allotropes of Carbon:

    • Diamond: Network covalent structure with each carbon bonded to 4 other carbons via single bonds, forming a crystalline structure.

    • Graphite: Network covalent structure with each carbon bonded to three others via one double bond and two single bonds, forming hexagonal rings; layers held together by weaker London dispersion forces.

85. Ranking of Intermolecular Forces:

    • London Dispersion Forces (LDF) < Dipole-Dipole (DD) < Hydrogen Bonds (HB)

86. Effect of Molecular Mass on LDF Strength:

    • In nonpolar molecules experiencing LDF, increased mass leads to stronger LDF due to more electrons.

    • Example: Halogen strength ranking: I > Br > Cl > F; iodine has the highest LDF, resulting in the highest melting/boiling point.

87. Effect of Molecular Length on LDF Strength:

    • Longer non-branching, nonpolar molecules experience stronger LDF.

      • Example: Alkanes ranked by increasing length and corresponding boiling points: methane < ethane < propane < butane < etc.

88. Conditions for a Covalent Molecule to be Polar:

    • A covalent molecule is polar if:

      1. It has polar bonds.

      2. There’s an asymmetrical distribution of electrons or its shape is asymmetrical.

89. Electron Domain Geometry Table:

    # of Domains	Electron Geometry	Molecular Geometry	Bond Angle
    1	Linear	Linear	180°
    2	Trigonal Planar	Trigonal Planar	120°
    		Bent	<120°
    3	Tetrahedral	Tetrahedral	109.5°
    		Trigonal Pyramidal	<109.5°
    		Bent	<109.5°

90. Lewis Structures:

    • For CO₂:

      • Electron Geometry: Linear

      • Molecular Geometry: Linear

      • Bond Angle: 180°

      • Intermolecular Forces: LDF

    • For H₂O:

      • Electron Geometry: Tetrahedral

      • Molecular Geometry: Bent

      • Bond Angle: <<109.5°

      • Intermolecular Forces: LDF, DD, HB

    • For CF₄:

      • Electron Geometry: Tetrahedral

      • Molecular Geometry: Tetrahedral

      • Bond Angle: 109.5°

      • Intermolecular Forces: LDF

    • For NH₃:

      • Electron Geometry: Tetrahedral

      • Molecular Geometry: Trigonal Pyramidal

      • Bond Angle: <109.5°

      • Intermolecular Forces: LDF, DD, HB

    • For CH₃O:

      • Electron Geometry: Trigonal Planar

      • Molecular Geometry: Trigonal Planar

      • Bond Angle: 120°

      • Intermolecular Forces: LDF, DD