大化 Midterm Review

Chapter 1

Classification of Matter

• Heterogeneous
• Homogeneous
  • Mixture(Solution)
  • Pure Substance
    • Element
    • Compound

Properties

• Intensive
• Extensive

Measurements

• Accuracy
• Precision

Significant Number

• +/- : Lowest significant decimal places
• \times / \div : Lowest significant numbers

Chapter 2

Atom

• Proton
• Electron
• Neutron

Isotope and Atomic Weight

Formula

• Empirical / molecular

Ion

• Cation
  • Just use the element name
  • Label the oxidation state (e.g., copper (II)) if necessary
• Anion
  • -ide / -ate / -ite / per- / hypo- / hydro- … , for different type of anions

Molecular Compounds

• Prefix:
  • Mono-
  • Di-
  • Tri-
  • Tetra-
  • Penta-
  • Hexa-
  • Hepta-
  • Octa-
  • Nona-
  • Deca-

Chapter 3

Formula weight and molecular weight

%Element=(number of atoms) \times (atomic weight) / (formula weight) \times 100%

Limiting Reactants / Excess Reactants

Theoretical Yield / Actual Yields

• %Yield

Chapter 4

Solutions

• Solutes
• Solvent
• Aqueous Solution
  • Dissociation
    • Electrolyte
    • Nonelectrolyte

Solubility of Ionic Compounds

• Common Soluble Compounds:
  • \ce{NO_{3}- / CH_{3}COO-}: No exception
  • \ce{Cl- / Br- / I- :Ag+, Hg_{2}^2+, Pb^2+} Precipitation
  • \ce{SO_{4}^2- : Sr^2+, Ba^2+, Hg_{2}+, Pb^2+} Precipitation
• Common Insoluble Compounds
  • \ce{S^2- / OH- : NH_{4}+, \text{alkali metal cations}, Ca^2+, Sr^2+, Ba^2+} Soluble
  • \ce{CO_{3}^2- / PO_{4}^3- : NH_{4}+, \text{alkali metal cations}} Soluble

Reactions Types

• Precipitation Reaction
• Exchange Reaction
• Displacement Reaction
• Neutralization Reaction
• Oxidation / Reduction Reaction
• Combination / decomposition
• Combustion

Acid / Base

• Arrhenius (increase \ce{[H+] / [OH-]}) / Brønsted (Donate / Accept Protons)
• Strong / Weak
  • Strong Acids: \ce{HCl, HBr, HI, HClO_{3}, HClO_{4}, HNO_{3}, H_{2}SO_{4}}
  • Strong Bases: \ce{LiOH, NaOH, KOH, RbOH, CsOh, Ca(OH)2, Sr(OH)2, Ba(OH)2}

Oxidation Numbers

Molarity

Dilution

Titration

Chapter 6

Wavelength of different radiation(increasing order)

• X-ray
• Ultraviolet
• Infrared
• Microwave
• Television
• Radio

Isoelectronic

Photoelectric Effect E=h\nu = E_{k}+I

Continuous / Line Spectra

• Element - Line spectrum

Wave Nature of Matter \lambda = \frac{h}{mv}

Heisenberg’s uncertainty principle (\Delta x)(\Delta p)\geq \frac{h}{4\pi}

\psi^2 gives electron density

Quantum Numbers

• Principle quantum number - shells/orbitals - n
• Angular quantum number - subshells - l(0\leq l\leq n-1)
• Magnetic quantum number - m_{l} (-l\leq m_{l} \leq l)
• Spin quantum number - m_{s}(\frac{1}{2} / -\frac{1}{2})

Electron Configuration / Orbital Diagram

• Pauli Exclusion Principle
• Hund’s Rule
• Aufbau principle
• Condensed Electron Configuration
  Exception \ce{[Ar] 3d^5 4s^1, [Ar] 3d^10 4s^1}

Chapter 7

Effective Nuclear Charge Z_{\text{eff}}=Z-S

• Calculation is not required , as Chapter 8 & 9 are harder

Size of atom

• Anion > Parent Atom > Cation
• Isoelectronic Series <- Judge from Nuclear Charge

Ionization Energy

• I_{n} Energy required to remove the nth election
• Judge group of a element (e.g. Al has a sharp increase from I_{3} to I_{4})
• Always endothermic

Electron Affinity

• Energy change with addition of electron
• Can be endothermic or exothermic

Metals, nonmetals and metalloids

• Metals: lustrous, malleable; oxides tend to be basic
• Nonmentals: dull; gain electron
• Metals and nonmentals tend to form ionic compound

Groups

• Alkali metals, Alkaline earth metals, Boron group, Carbon group, Nitrogen group, Chalcogens, Halogens, noble gases

Chapter 8

Chemical bonds

• Ionic bonds
• Covalent bonds
• Metallic bonds

Lewis Symbols

• Octet rule

Lattice Energy

• \ce{NaCl(s)-> Na+(g) + Cl-(g)}, E_{el}=\frac{\kappa Q_{1}Q_{2}}{d}
• The energy required to break the ionic bonds
• \Delta H_{lattice}=\Delta H\degree_{f}[\ce{Na(g)}] + \Delta H\degree_{f}[\ce{Cl(g)}] + I_{1}(\ce{Na})+E(\ce{Cl}) - \Delta H\degree_{f}[\ce{NaCl(s)}]
  • \Delta H\degree_{f}[\ce{X}]: The energy change when forming X
  • The addition part: \ce{Na(s) + \frac{1}{2}Cl_{2}(g) -> Na+(g) + Cl-(g)}
  • The subtraction part: \ce{Na(s) + \frac{1}{2}Cl_{2}(g) -> NaCl(s)}

Bond Polarity

• Dipoles \mu=Qr
  • Units: \text{Debyes (D)} (carefully convert to \text{C} \cdot \text{m})

Lewis Structure

• Assign the electrons with the Octet Rule (only for period 1 and 2)
  • Exceptions
    • odd number of electrons (\ce{NO})
    • less than an octet of valence electron(\ce{BF3})
    • more than eight valence electrons (an expanded octet) (\ce{PF5})
• Formal Charge = Valence electrons - 1/2(bonding electrons) - all nonbonding electrons
• Minimized Formal Charge gives the most stable structure

Resonance Structures

• Delocalized electrons

Bond Enthalpy

• \Delta H_{\text{rxn}}=\sum \Delta H(\text{bonds broken}) +\sum \Delta H(\text{bonds formed})

Chapter 9

VSEPR Theory

Electron Domains	Electron-Domain Geometry	Predicted Bond Angle
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	120°, 90°
6	Octahedral	90°

Number of Electron Domains	Nonbonding Domains	Molecular Geometry	Example
2	0	Linear	\ce{CO_{2}}
3	0	Trigonal planar	\ce{BF_{3}}
3	1	Bent	\ce{NO_{2}-}
4	0	Tetrahedral	\ce{CH_{4}}
4	1	Trigonal pyramidal	\ce{NH_{3}}
4	2	Bent	\ce{H_{2}O}
5	0	Trigonal bipyramidal	\ce{PCl_{5}}
5	1	Seesaw	\ce{SF_{4}}
5	2	T-shaped	\ce{ClF_{3}}
5	3	Linear	\ce{XeF_{2}}
6	0	Octahedral	\ce{SF_{6}}
6	1	Square pyramidal	\ce{BrF_{5}}
6	2	Square planar	\ce{XeF_{4}}

Valence Bond Theory and Hybrid Orbitals

• \sigma bond: head-to-head
• \pi bond: side-to-side (weaker), not hybridized orbitals
• Hypervalent: more than an octet of electrons

Molecular Orbital Theory

• Bonding orbitals and Antibonding orbitals
• Bond Order = 1/2 (Bonding electrons - Antibonding electrons)
• HOMO : highest occupied molecular orbital
• LUMO : lowest unoccupied molecular orbital
• Paramagnetism: have unpaired electrons
• Diamagnetism: no unpaired electrons