General Chemistry DAT Notes

General Chemistry Dental Admission Test Notes

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

  • 2Al + 6HCl → 2AlCl3 + 3H2
  • 2 moles of Al react w/ 6mol of HCl to form 2mol AlCl3 and 3mol of H2

Limiting reagent – is the reactant consumed first

  • Given grams of each, convert to moles. Use this information to see how much product each reactant could potentially produce with that starting amount of grams; using mol ratios of rxn.
  • Divide moles of the reactant by its coefficient in the equation to see how many “rounds of the rxn” it could do. The one that can do the least is limiting!

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

  • 1stshell (n): any whole number starting with one (# of period)
    • Diff. in E b/w 3&4 < 2&3
  • 2ndsubshell (l): 0,1,2… n-1
    • 0 = s, 1 = p, 2 = d, 3 = f
    • 3d has more E than 4s
  • 3rdorbital (ml): -L, 0, L
    • 2 e- per orbital
  • 4thspin number (ms): ½ , – ½ (+ = higher E)

Electron Configuration – start at noble gas; remember to subtract 1 from “d block”; ex for Sc [Ar]4s23d1

  • If a higher level subshell is almost half-filled or full, an electron will be promoted
  • Cr: [Ar]4s13d5 & Cu: [Ar]4s13d10 / the families of Cr and Cu are EXCEPTIONS

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

  • Coordinate Covalent: one atoms donates BOTH 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

  • Dipole Forces: dipole-dipole and ion-dipole
  • London Dispersion: instantaneous dipoles in non polar molecules; very weak
  • Hydrogen Bonding: b/w H’s of N, O, F; ex water

**These interactions (esp H bonding) lead to: greater MP, BP, viscosities // LOWER vapor pressures

Types of Elements

  • Metals – can be deformed w/o breaking (malleability) & drawn into wires (ductility)
  • Nonmetals – generally brittle in solid state, located @ upper right of periodic table
  • Metalloids – B, Si, Ge, As, Te, At (diagonal from BAt)

Geometric Shapes

  • ex. BeCl2  linear  180o (2 subs.)
  • Ex. BH3  trigonal planar  120o (3 subs.)
  • Ex. CH4  tetrahedral  109.5o (4 subs.)
  • Ex. PCl5  trigonal bipyramidal  90o, 120o, 180o (5 subs.)
  • Ex. SF6  octahedral  90o & 180o (6 subs.)

Types of Chemical Reactions

  • Combination reaction – 2 reactants  1 product
  • Decomposition reaction – 1 reactant  2 products
  • Single displacement reaction – 1 atom/ion of 1 compound replaced by atom of another element (ex. redox rxns)
  • Double displacement reaction – AKA metathesis rxns; elements from 2 diff. compounds displace each other to form 2 new compounds
    • ex. neutralization rxns (acid-base)

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):

  1. simple cubic – atom @ each point of the cube, each is shared by 8 other cells  (1/8)*8 corners = 1 atom per cell
  2. body-centered cubic – 1 atom @ center of the cube + an atom @ each point of the cube (shared by 8 other cells) [ (1/8)*8 corners]+1 body = 2 atoms per cell
  3. facecentered cubic – 1 atom embedded @ the center of each cube face (shared with 1 other cell) + an atom @ each point of the cube (shared by 8 other cells)  (1/8) * 8 corners + (1/2) * 6 faces = 4 atoms per cell

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

  • P = Pa + Pb + Pc
  • To get partial pressure, multiply the mole fraction of the gas by its pressure

-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

  • 1/r2 = (MM2/MM1)1/2

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

  • C6H12O6 is non-ionic; does not dissociate; i=1 (same for most biomolecules)
  • NaCl dissociates into Na and Cl; i=2
  • CaCl2 dissociates into Ca and 2Cl; i=3

Salt Solubility Rules

  1. All Group I (Li, Na, K, Rb, Cs) and ammonium (NH4 ) salts SOLUBLE
  2. All nitrate (NO3 ), perchlorate (ClO4), and acetate (C2H3O2) salts SOLUBLE
  3. Silver (Ag), lead (Pb), and mercury (Hg) salts INSOLUBLE unless they have rule 2

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

  • Fe2+ = Ferrous Cu+ = Cuprous
  • Fe3+ = Ferric Cu2+ = Cupric
  • S2- = Sulfide N3- = Nitride P3- = Phosphide
  • NO2 = Nitrite SO32- = Sulfite
  • NO3 = Nitrate SO42- = Sulfate
  • ClO = Hypochlorite ClO2 = Chlorite
  • ClO3 = Chlorate ClO4 = Perchlorate
  • HCO3 = hydrogen carbonate AKA bicarbonate
  • HSO4 = hydrogen sulfate AKA bisulfate

Chapter 6 – Kinetics

Reaction Rate

  1. LOWER activation energy, FASTER reaction rate
  2. GREATER concentration of reactants, FASTER reaction rate
  3. HIGHER the temperature, FASTER 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)

  • 2nd order reaction rates are related to concentration & time

Reaction Orders:

  • Zero-order rxns = constant rate = k
  • First-order rxns
    • Rate = k[reactant]
    • [At] = [Ao]e-kt (A=reactant)
    • t1/2 = .693/k
  • Second-order rxns
    • Rate = k[A][B] or rate = k[A]2
    • Units = M-1sec-1
  • Higher-order rxns = order is greater than 2
  • Mixed-order rxns = fractional order (ex. rate=k[A]3/2)

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]

  • Keq < 1 favors reactants ; Keq = 1 equilibrium; Keq > 1 favors products
  • Equilibrium means rxn no longer changes w/ time (ex saturated)

Quotient (Q) – same expression as above when the reaction is not at equilibrium

  • Q < Keq reaction proceeds forward; Q > Keq reaction proceeds in reverse

** 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

  • HClO (hypochlorous acid), HClO2 (chlorous acid), HClO3 (chloric acid), HClO4 (perchloric)
  • HNO2 (nitrous acid), HNO3 (nitric acid)

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+

  • conjugate bases of strong acids have no basic properties in water (no reverse rxn)
  • conjugate base of weak acid is weak base

Strong Acid – one that dissociates completely in water; common strong acids:

  • Hydroiodic acid (HI), Hydrobromic acid (HBr), Hydrochloric acid (HCl)
  • Perchloric acid (HClO4), Chloric acid (HClO3)
  • Sulfuric acid (H2SO4)
  • Nitric acid (HNO3)

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

  • Ka = [H3O+][A] / [HA]
  • Ka >1 then products favored (strong acid) // Ka < 1 reactants favored (weak acid)

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

  • Ex) .01M solution of HCl means [H+] = .01M

pH of a Weak Acid – to get the pH we need to use the equilibrium expression

  • Ex) say you add 0.2 mol of HCN to water to create a 1L solution. Initially, you have [HCN] 0.2M; Let x = moles of HCN dissociated at equilibrium; so at equilibrium [HCN] = (0.2-x)M; each mol of HCN dissociates to 1mol H+ AND 1mol CN // (1mol HCN => 1mol H+ and 1mol CN)
  • HCN => H+ + CN

Initial 0.2M 0 0

At Eq. (.2-x)M xM xM

  • Ka = [H][CN]/[HCN] = x2 / (0.2 – x)
  • Since the Ka for HCN is 4.5 X 10-10 (very small) not much H+ will dissociate and the value (0.2 – x) will be approximately 0.2
  • x2 / 0.2 = 4.9 X 10-10 ; x = 1 X 10-5 = [H+] ; pH = 5

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

  • HCl + NaOH => NaCl + H2O
  • MaVa = MbVb ( how much acid / base needs to be added to neutralize)

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

  • Ex) add 0.1 mole of acetic acid (CH3COOH) and 0.1 mol of acetate (CH3COO) to water to obtain 1L solution
  • CH3COOH + H2O => H3O+ + CH3COO
  • [H3O+] = Ka [CH3COOH] / [CH3COO] = 1.75 X 10-5 // pH = 4.76
  • If we add 0.005 mol of HCL it will dissociate to => .005 mol H+ & .005 mol Cl; Cl will have no effect on equilibrium; adding H+ will (just like adding any product, it will shift reaction to left) // now there are 0.005 moles MORE of CH3COOH and 0.005 moles LESS of CH3COO
  • [H3O+] = Ka [1.005] / [0.095] // pH= 4.71 ** hardly any change!
  • If we had added strong base; it would have been like adding reactant => reaction would have shifted right; the moles would have been added to acetate and subtracted from acetic acid

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

  • HF titrated by NaOH (NaOH is being added to the HF)
  • NaOH + HF => Na+ + F- + H2O
  • Initially the pH is that of the original HF, as OH is added the acid will slowly be neutralized and the pH will rise. At first it goes slowly because the solution is behaving like a buffer (appropriately named the buffering domain)
  • Solution suddenly loses buffering capability and the pH increases dramatically (HF neutralized => every new OH molecule remains in solution)
  • OH concentration increases rapidly until it is not that much different from the NaOH concentration

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

  • Oxidation half-reaction: Fe => Fe2+ + 2e
  • Reduction half-reaction: 2H+ + 2e => 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

  • If the cell voltage (E) is positive, reaction is spontaneous
  • If the cell voltage (E) is negative, reaction is non-spontaneous
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