The Word Document for this page is located here: General Chemistry DAT Notes
Chapter 0 – Reference
– Once chemical rxn is balanced, the stoichiometric coefficients tell us the relative amount of reactants that combine and relative amount of product formed
Limiting reagent – is the reactant consumed first
Chapter 1 – Atomic Structure
Isotopes – two atoms of same element that differ in number of neutrons / same # of protons
Quantum Number – designate shell, subshell, orbital, spin
Electron Configuration – start at noble gas; remember to subtract 1 from “d block”; ex for Sc [Ar]4s23d1
Expanded Octets – elements in the 3rd or 4th periods may make compounds using an expanded octet of electrons
Emission Spectrum – BRIGHT lines that correspond to energies where atom EMITS photons
Absorption Spectrum- DARK lines that correspond to energies where atom ABSORBS photons
Alpha Decay – loses 2 protons
Beta_ Decay – gains 1 proton
Beta+ Decay – loses 1 proton
Chapter 2 – Periodic Trends and Bonding
Summary of Trends –Atomic Radius (down and to the left), Acidity (down and to the right)
Covalent Bond – b/w two atoms when each contributes one of its unpaired valence electrons
Ionic Bond – attraction b/w a cation and anion; solids at room T; ex salts NaCl
Intermolecular (Van der Waals) Forces – relatively weak interactions that take place between neutral molecules
**These interactions (esp H bonding) lead to: greater MP, BP, viscosities // LOWER vapor pressures
Types of Elements
Geometric Shapes
Types of Chemical Reactions
Chapter 3 – Phases
Phases Changes – result of weakening/ strengthening of intermolecular forces
** [SOLID] fusion→ ←crystallization [LIQUID] vaporization→ ←condensation [GAS]**
** [SOLID] sublimation → ← deposition [GAS] **
** [SOLID] == Heat Absorbed [GAS] **
Heat of Fusion / Vaporization – q = m ΔH
Calorimetry – q = mcΔT
– When a substance absorbs/ releases heat either temp will change or a phase change / NOT both at once
– Note: 1cal = 4.2J
– Note: HIGHER specific heat (c) holds on to heat longer = cool slowly
Triple Point – temperature and pressure where all three phases exist in equilibrium
Critical Point – end of liquid gas boundary
Cubic Crystal System (unit cells):
Chapter 4 – Gases
Standard Temperature and Pressure (STP) – 0˚C and 1 atm
Ideal Gas Law – PV = nRT
Dalton’s Law – total pressure is sum of the partial pressures of all the constituent gases
-Note molecules of two different gases at same temp have same average kinetic energy, BUT different average speeds… lighter one is going faster
*Ideal Gas Behavior – 1) molecules are so small compared to average spacing b/w them that they take up no volume 2) molecules experience NO intermolecular forces
**Diffusion (through a mixture) vs. Effusion (through a hole)**
Graham’s Law of Effusion – rates of effusion are proportional to average speed; lighter particles go faster
Chapter 5 – Solutions
* Solubility of most solids in liquids increases with increasing temperature
* Solubility of most gases in liquids decreases with increasing temperature
* Molality does not change with temperature
* When ionic substances dissolve they dissociate into IONS
Electrolytes – free ions in a solutions; can conduct electricity; all ionic compounds are STRONG electrolytes
Van’t Hoff Factor (i) – how many ions one unit of a substance will produce in solution
Salt Solubility Rules
Boiling Point Elevation –ΔT = k*i*m
Freezing Point Depression – ΔT = -k*i*m
*added solutes (i) affect boiling point / freezing point; (k) is a constant of the solvent
Common Ions
Chapter 6 – Kinetics
Reaction Rate
Catalysts – make reaction go FASTER by lowering activation energy; remains unchanged
Rate Law – rate at which a reactant disappears; only those involved in rate-determining step; only reactants included; ex rate = k[A][B]
** If looking at table and [C] doubles but rate is unaffected, then rxn rate does NOT depend on [C] // thus it is not included in the rate law
Rate and Half Life: for a 1st order rate constant (k) k = (.693)*(1/thalf life)
Reaction Orders:
Chapter 7 – Equilibrium
Equilibrium Constant – solids and liquids not included; if the reaction is gaseous we can use partial pressure of each gas for concentration; Keq = [Products] / [Reactants]
Quotient (Q) – same expression as above when the reaction is not at equilibrium
** Adding reactant makes reaction shift to right / removing reactant makes reaction shift left
** Adding catalyst does NOT affect equilibrium
** Adding inert gas causes NO disturbance
Solubility Constant – at equilibrium, rate of ions going into soln. = rate precipitating out
Ex. The solubility product for lithium phosphate Li3Po4 is Ksp = 2.7 x10-9. How many moles of this salt would be required to form a saturated, 1-liter aqueous soln?
Li3PO4(s) 3Li+ (aq) + PO43-(aq)
If you let x denote [PO43-], then [Li+] =3x Ksp = (3x)3*x=27x4
– solid doesn’t count in Ksp eqn
Common-Ion Effect – addition of NaOH has caused the amount of hydroxide ion – the common ion – in the solution to increase. This disturbs the equilibrium of magnesium hydroxide. By Le-Chatelier’s principle, the system will react by favoring the reverse reaction, producing solid Mg(OH)2, which will precipitate.
Chapter 8 – Acids and Bases
Common Names
Arrhenius – acids ionize in water to produce (H+) / bases ionize to produce (OH–)
Bronsted Lowry – acids are proton donors / bases are proton acceptors
Conjugate Base – form the conjugate base of an acid by removing an H+
Strong Acid – one that dissociates completely in water; common strong acids:
Weak Acid – only partially dissociates in water
Acid Ionization Constant – Ka; strength of acid is directly related to how much products are favored over reactants
Strong Base – Group 1 w/ OH (NaOH, LiOH, KOH); Metal amines NaNH2
Amphoteric – can act as acid or base; conjugate base of WEAK polyprotic acid is always amphoteric b/c it can gain an proton (returning to weak acid) or lose another proton
Ion Constant of Water – Kw = [H3O+][OH–] = 1 X 10-14 because water reacts with itself (amphoteric) // H2O(l) + H2O(l) => H3O+ + OH– // [H+] = 1 X 10-7 ; pH = 7
pH = -log [H+] thus [H+] = 10-pH
** pH + pOH = 14
** For strong acids, the hydrogen ion concentration will be the same as the concentration of acid
pH of a Weak Acid – to get the pH we need to use the equilibrium expression
Initial 0.2M 0 0
At Eq. (.2-x)M xM xM
Neutralization Reactions – when an acid and base react with one another producing salt and water; molar equivalents of any acid with any base will lead to a complete neutralization; only a strong acid with strong base will result in pH neutral
Buffer Solution – resists changing in pH when a small amount of acid or base is added; to design a buffer solution we choose a weak acid whose pKa is as close to desired pH as possible
Henderson Hasselbalch – pH= pKa – log [weak acid] / [conjugate base]
pOH= pKb – log [weak base] / [conjugate acid]
Acid Base Titration – used to determine identity of weak acid or weak base also can be used to determine CONC of ANY acid / base
Chapter 9 – Thermodynamics
Enthalpy – is a measure of heat energy released or absorbed; Bond Formed = E released
Gibbs Free Energy – ΔG = ΔH – T ΔS and when ΔG < 0 reaction is spontaneous
Thermodynamics and Equilibrium – ΔG = -2.3RT log Keq
Chapter 10 – Electrochemistry
Fe + 2HCl => FeCl2 + H2
** Can use redox reactions to spontaneously generate an electric current
** Anode always the site of Oxidation; Cathode always site of Reduction (An Ox, Red Cat)
** Electrons in circuit always move from anode to cathode
Galvanic Cells – use spontaneous reaction to create electric current; anode negative , cathode positive (normal); electrons move from anode to cathode (left to right); cations also move from left to right (via salt bridge)
Electrolytic Cells – use electric current to FORCE NON-spontaneous redox reactions to occur; anode is positive and cathode is negative (backwards); but electrons still move from anode to cathode b/c they are FORCED
ΔG = -nFE ; where n is number of electrons
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