Shivi Yadava

Hima Veeramachaneni

AP Chemistry Study Guide

Unit 1 – Basic Concepts in Chemistry

  • Nomenclature
  • Two Non-metals = Covalent
  • Prefix-element + prefix-element-ide
  • Ex. P2Cl6 – diphosphorous hexachloride
  • Prefixes

One / Mono
Two / Di
Three / Tri
Four / Tetra
Five / Penta
Six / Hexa
Seven / Hepta
Eight / Octa
Nine / Non
Ten / Dec
  • Metal + Non-metal = Ionic
  • Name both ions
  • Metal ions are the same

Non-metal ions

  • Binary-ide
  • Cl-, Br-, I-, F-
  • Ex. NaCl – Sodium Chloride
  • Polyatomic
  • No oxygen – “ide”
  • Ex. Mg3S – Magnesium Sulfide
  • Normal number of oxygen – “ate” (4 on PT – o/e 3)
  • Ex. MgSO4 – Magnesium Sulfate
  • One less Oxygen – “ite”
  • Ex. MgSO3 – Magnesium Sulfite
  • Two Less Oxygen – “hypo – ite”
  • Ex. MgSO2 – Magnesium Hyposulfite
  • One More Oxygen – “per-ate”
  • Ex. MgSO5 – Magnesium Persulfate
  • One more Sulfur & one less Oxygen – “Thiosulfate”
  • Ex. MgS2O3 – Magnesium Thiosulfate
  • Transition metal = indicate charge
  • Roman Numeral
  • Latin (two oxidation states)
  • Higher = “ic”
  • Ex. Sb(ClO)5 – Stibnic Hypochlorite
  • Lower = “ous”
  • Ex. CuCl – Cuprous Chloride
  • Memorize!

NH4+1 / Ammonium
NH3 / Ammonia
CrO4-2 / Chromate
Cr2O7-2 / Dichromate
C2O4-2 / Oxalate
C2H3O2-1 / Acetate
OH- / Hydroxide
MnO4-2 / Permanganate
CH4 / Methane
  • Acids
  • Name negative ion
  • Change ending
  • “ate”  “ic”
  • Ex. HNO3 – Nitric Acid
  • “ite”  “ous”
  • Ex. HClO – Hypochlorous Acid
  • “ide”  “hydro-ic”
  • Ex. HF – Hydrofluoric Acid
  • Reactions
  • Basic Guidelines:
  • All acids are aqueous unless organic
  • Only strong acids and bases break apart 100%
  • Acids: HClO4, HClO3, HCl, HBr, HI, HNO3, H2SO4
  • Bases: LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
  • Don’t write physical states in the ionic
  • Hidden Reactions (clues to look for):
  • If one compound is aqueous that means that it is in water, and the other compound might react with that water
  • If one compound is an acid or base, then the other compound then the other compound might react with water to form an acid or base
  • Non-metal oxides that react with water
  • Metal oxides that react with water
  • Immediately break apart because they don’t exist
  • H2SO3H2O + SO2
  • H2CO3 H2O + CO2
  • NH4OH  H2O + NH3
  • Ammonia reactions don’t form water
  • Things that don’t dissolve in water:
  • H2O
  • 4 gases (CO2, SO2, NH3, and H2S)
  • Anything going against the solubility rules
  • Solubility Rules
  • Soluble in water
  • Alkali metal compounds
  • Nitrates and nitrates
  • Chlorates and perchlorates
  • Acetates (except with Ag+1)
  • Ammonium compounds
  • Chlorides, Bromides, Iodides (except with Ag+1, Hg+2, Hg2+2, Pb+2)
  • Flourides (except with Group II metals, Pb+2, Fe+3)
  • Sulfates, Sulfites (except with Sr+2, Ba+2, Ca+2, Pb+2, Hg2+2, Ag+1)
  • Carbonates, Phosphates, and Chromates are only soluble with alkali metals, ammonium, CaCrO4, SrCrO4
  • Hyroxides are only soluble with alkali metals, ammonium, Sr+2, Ca+2, Ba+2
  • Sulfides are only soluble with Group I metals, Group II metals and ammonium
  • Oxides are only soluble with Group I metals and ammonium
  • Synthesis Reactions: A + X ------> AX
  • Metals react with non-metals to produce binary salts (two elements, no polyatomic)
  • Metal oxides (basic anhydrides) react with water to yield bases (metal hydroxides)
  • Non-metal oxides (acid anhydrides) react with water to yield acids (oxidation number of non-metal does not change – do an imaginary charge check!)
  • Metal oxides react with non-metal oxides to produce a polyatomic salt (Oxidation number of non-metal does not change – do an imaginary charge check!)
  • Decomposition Reactions: AX ------> A + X
  • Acids with oxygen decompose to give non-metal oxides and water Acids with oxygen decompose to give non-metal oxides and water (oxidation number of non-metal does not change – do an imaginary charge check!
  • Metallic hydroxides, or bases, decompose to give metal oxides and water
  • Metallic carbonates decompose to give metal oxides and carbon dioxide
  • Metallic chlorates decompose to give metal chlorides and oxygen
  • Metallic nitrates decompose to give metal nitrites and oxygen
  • Ammonium carbonate decomposes to give ammonia, water, and CO2
  • Sulfurous acid decompose to give water and sulfur dioxide
  • Carbonic acid decomposes to give water and carbon dioxide
  • Ammonium hydroxide decomposes to give ammonia and water
  • Binary compounds decompose to give two elements (with energy)
  • Hydrogen peroxide decomposes to give water and oxygen
  • Polyatomic salts not listed above can decompose to form the metal oxide and non-metal oxide that formed them (oxidation number of non-metal does not change – do an imaginary charge check!)
  • Single Replacement Reactions: A + BX ------> AX + B
  • Active metals replace less active metals in ionic compounds in aqueous solutions
  • Active metals replace H in water to form metal hydroxides (bases) and H2
  • Active metals replace H in acids to form hydrogen gas and a salt
  • Active non-metals replace less active non-metals in ionic compounds in aqueous solutions
  • Non-aqueous replacement reactions – reductions of metal oxides by hydrogen or other gases:
  • H2 + CuO → Cu + H2O (occur at high temperatures!)
  • CO + Fe2O3→ Fe + CO2
  • Double Replacement Reactions: AX + BY ------> AY + BX
  • Formation of a precipitate (solid) governed by the solubility rules
  • Formation of a gas
  • Common gases are H2S, CO2, SO2, NH3
  • Any sulfide (S-2) plus any acid forms H2S gas and a salt
  • Any carbonate (CO3-2) plus any acid forms CO2, HOH, and a salt
  • Any sulfite (SO3-2) plus any acid forms SO2, HOH, and a salt
  • Any ammonium (NH4+1) compound plus a soluble hydroxide form NH3, HOH, and a salt
  • Formation of a molecule – which is a compound that does not dissociate well in water, due to its covalent nature! It stays together as a molecule! Example – H2O!
  • Acid-base neutralization is one type – ACID PLUS BASE = WATER PLUS SALT
  • Hydrolysis – Reverse of an acid-base neutralization – a salt reacts with water – this will only happen with one in a trillion water molecules!
  • One in a trillion water molecules can break apart into H+1 and OH-1
  • The salt then breaks apart, and a double replacement reaction occurs, with the salt reacting with the H+1 and the OH-1
  • Produces an acid and a base every time!
  • Salts are products of neutralization, but salts that undergo hydrolysis are not neutral!
  • Salts of a strong acid and a weak base + H2O give an acidic solution
  • Salts of a weak acid and a strong base + H2O give a basic solution
  • Salts of a strong acid and a strong base do not undergo hydrolysis – their solutions are neutral!
  • Salts of a weak acid and a weak base + H2O may give an acidic, basic, or neutral solution – look at the strength of the acid or base produced (Ka or Kb)
  • Oxidation-Reduction Reactions
  • In a reduction/Oxidation reaction, one species is oxidized (loses electrons) and the other species is reduced (gains electrons)
  • The species being oxidized is called the “reducing agent” and the species being reduced is called the “oxidizing agent”
  • Many oxidation/reduction reactions will occur in either acidic or basic solution, taking advantage of H+ or OH- ions, along with H2O, to aid the reduction/oxidation
  • These reactions are written and balanced using the half-reaction method
  • Acidic Solution- balancing technique
  • Predict products
  • Balance with
  • Hydrogen= H+
  • Oxygen= H2O
  • Balance
  • Basic Solution
  • Balance H w/ H+
  • Add OH- to each side to neutralize H+
  • Form H2O with H+/ OH-
  • Balance charged elements
  • Ex. A solution of sodium bromide is added to an acidic solution of potassium bromated

NaBr(aq) + KBrO3 (aq)

6H+ + Na+1 + 5Br- + K+1 + BrO3- 3Br2 + 3H2O

  • There are obvious signs to look for in a common redox reaction:

Important Oxidizing Agents (These things are reduced!)Formed in Reaction

MnO4- (acid solution)Mn+2

MnO4- (basic solution)MnO2

MnO2 (acid solution)Mn+2

Cr2O7-2 (acid solution)Cr+3

CrO4-2 (basic solution)Cr+3

HNO3, concentratedNO2

HNO3, diluteNO

H2SO4, hot concentratedSO2

Metallic IonsMetallous Ions

Free HalogensHalide Ions

HClO4Cl-1

Na2O2OH-1

H2O2H2O

Perhalates, halates, halitesHalogens

Important Reducing Agents(These things are oxidized!)Formed in Reaction

Halide IonsHalogens

Free MetalsMetal Ions

Metallous IonsMetallic Ions

Sulfite IonsSO4-2

Free Halogens (dilute basic solution)Hypohalite Ions

Free Halogens (concentrated basic solution)Halate Ions

C2O4-2CO2

NO2-1NO3-1

Sn+2Sn+4

H2O2O2

  • Chromium: dichromate to Cr3+ in acid solution; chromate to Cr(OH)3 in basic solution.
  • Dichromate ion can turn into chromate in basic solution, and chromate ion can turn into dichromate ion in acidic solution (this is not reduction/oxidation – the Cr still retains a +6 charge)
  • Oxygen: hydrogen peroxide can acts as an oxidizing agent (reduced to water) and a reducing agent (oxidized to oxygen gas).
  • Nitrogen: nitrate ion is an oxidizing agent only in acid solution. The reduction product is NO.
  • Sulfur: sulfate ion is an oxidizing agent only in acid solution. The reduction product is SO2.
  • Complex Ion Reactions:
  • Ligand= double charge and sticks onto a transitional metal
  • Transition metal salt + ligand → complex ion
  • If the word excess is in the problem, then it is complex!
  • Ex. Excess sodium cyanide solution is added to a solution of silver nitrate

NaCN(aq) + AgNO3 (aq)

Na+1 + CN- + Ag+1 + NO3-1 Ag (CN)2-1

2CN- + Ag+1 Ag(CN)2-1

  • Aluminum salt + ligand → complex ion
  • Beryllium salt + ligand → complex ion
  • Both Zn+2 and Al+3 form Zn(OH)4-2 and Al(OH)4-1 when treated with excess hydroxide
  • Ag+1, Cu+2, Zn+2, and Cd+2 all form complexes with NH3
  • Infrequently seen, but has been on the AP, and used in lab:
  • Thiocyanate acts as a ligand and bonds to a transition metal
  • A drop of potassium thiocyanate is added to a solution of iron (III) chloride:
  • SCN-1 + Fe+3 → Fe(SCN)+2
  • Ammonia, as a ligand, gets turned into ammonium ion, and the transition metal is freed from being a complex ion
  • Dilute hydrochloric acid is added to a solution of diamminesilver (I) nitrate:
  • H+1 + Cl-1 + [Ag(NH3)2]+2 → AgCl + NH4+1
  • Notice the destruction, rather than the formation, of a complex
  • Common ligands are: I-1, Br-1, F-1, OH-1, H2O, C2O4-2, NH3, SCN-1, CN-1
  • It is a good idea to recognize the names of these ligands as well – iodo, bromo, fluoro, hydroxy, aqua, oxalato, ammine, thiocyanato or isothiocyanato, and cyano
  • To determine coordination number:
  • For aqua complexes of transition metals, C.N. = 6
  • For others, C.N. = cation charge x 2
  • Lewis Acid and Lewis Base Reactions
  • Lewis acid reacts with a Lewis base to form an adduct:
  • BF3 + NH3→ F3BNH3
  • Phosphorus (V) oxytrichloride is added to water
  • POCl3 + H2O → H3PO4 + Cl- + H+
  • Note that molecular phosphorus compounds form acids with water.
  • PCl5 + H2O→ H3PO4 + H2O + Cl- + H+
  • PCl3 + H2O→ H3PO3 + Cl- + H+
  • Organic bases that have unshared pairs of electrons can react with water or other H+ suppliers:
  • Methylamine gas is bubbled into water:
  • CH3NH2 + H2O → CH3NH3+ + OH¯
  • Give and take electrons in order to share
  • Things to Practice!
  • Empirical Formula
  • Stoichiometry
  • Pv=nrt
  • Dilution Formula: M1V1= M2V2
  • Look at hard miscellaneous reactions sheet

Unit 2 – Bonding and Molecular Structure

  • Intramolecular (Chemical)
  • Chemical properties (flammability)
  • Ionic
  • Crystalline solid (usually white)
  • Difference in Electronegativity > 1.7
  • Transfer of electrons
  • Cation (+)/ Anion (-)
  • Strong bond  high melting points
  • Don’t conduct electricity in the solid state  electrons can’t flow through
  • Do conduct electricity in water  like dissolves like
  • Strength is dependent on size of ions and charge of ions
  • Anytime you make a positive ion it becomes smaller because it loses electrons and sublevel
  • Ex. CaS would be stronger than NaF because CaS has a greater charge
  • Covalent
  • Sharing electrons (2 or more non-metals)
  • Gases
  • Can be liquids and solids if they are large molecules
  • Polar
  • 0.7 < Difference in Electronegativity < 1.7
  • Slight charge
  • Non- Polar
  • True non- polar has Difference in Electronegativity of 0 but…
  • Difference in Electronegativity < 0.7
  • No charge
  • Wax, petroleum  large but no charge
  • Metallic
  • So strong because there is a sea or web of electrons
  • By getting closer together they are able to attract electrons better and there is a better flow
  • Intermolecular Bonds (Physical)
  • Also known as Van Dar Waals Forces
  • Physical properties (boiling point, melting point)
  • Three types of bonds
  • Hydrogen Bond
  • Strongest
  • Special type of Dipole- Dipole Bond
  • δ+H bonded to F, N, O
  • Dipole- Dipole Bond
  • Attraction between the δ+ of one polar molecule and δ- end of another polar molecule
  • Ex.
  • London Dispersion Bond
  • Weakest (almost non-existent)
  • Frictional force can break the bonds
  • JELLO bond
  • Branched hydrocarbons have less London Dispersion
  • Temporary- constantly being changed
  • Lattice energy
  • Amount of energy given off when crystal forms
  • Energy
  • Positive= not-spontaneous, need energy
  • Negative= spontaneous
  • Born- Haber Cycle (Lattice Energy Problems)
  • Ex.
  • Lewis Dot Diagram
  • Tells nothing about shape!
  • Coordinates covalent bond- double bond in which outside atom has to give up two electrons in order to share with the central atom
  • Formal Charge Check
  • # of e- that should be there - # of e- on the atom
  • Limit formal charge as soon as possible
  • One bond = one electron
  • Resonance
  • Chemically identical
  • All formal charge checks are the same!
  • One atom moves around in the Lewis Dot Diagram
  • Ex.

  • VSEPR Theory
  • Valence shell electron repulsion theory
  • LE Theory
  • Localized Electron Theory
  • Electron geometry- shape of electrons around central atom
  • Molecular geometry- shape of atoms around central atom

Electron Geometry / Molecular Geometry / s prs / us prs / Hybridization / Dipole Moment
Linear / Linear / 2 / 0 / sp
Trigonal Planar / Bent / 2 / 1 / sp2
Trigonal Planar / 3 / 0 / sp2 / zero
Tetrahedral / Bent / 2 / 2 / sp3
Trigonal Pyramidal / 3 / 1 / sp3
Tetrahedral / 4 / 0 / sp3 / zero
Trigonal Bypramidal / Linear / 2 / 3 / sp3d / zero
T- Shaped / 3 / 2 / sp3d
See-Saw / 4 / 1 / sp3d
Trigonal Bypramidal / 5 / 0 / sp3d / zero
Octahedral / Square Planar / 4 / 2 / sp3d2 / zero
Square Pyramidal / 5 / 1 / sp3d2
Octahedral / 6 / 0 / sp3d2 / zero
  • Permanent dipole moment?- cancel out charge or not?
  • In determining charge  look at molecular geometry because these are the only things that have a permanent charge
  • Violations to the octet rule
  • Less than 8 electrons
  • Be, B
  • More than 8 electrons
  • Only elements that have a d sublevel
  • 3rd period or below (p, s)
  • Hybridization Theory
  • Only central atom
  • Blends its orbitals together to make new ones
  • # you get = # you blend (includes unshared pairs)
  • Sigma bonds- mixed 1st bond
  • Pi bonds- mixed 2nd or 3rd bonds
  • Just a reminder!- electrostatic bonds are physical bonds
  • Molecular Orbital Theory
  • Bond Order
  • ½ (# of bonding electrons- # of antibonding orbitals)

Unit 3 – Thermochemistry

  • Matter
  • “stuff”
  • Has mass and takes up space
  • Energy
  • Ability to do work or produce heat
  • State function – does not depend on path it takes
  • Potential = mgh
  • Kinetic = ½mv2
  • Law of Conservation of Energy
  • Can’t create or destroy energy
  • Can only be converted from one form to another
  • Temperature (T)
  • Measurement of speed or randomness of particles
  • It is a state function
  • Heat (Q)
  • Not a state function – because it depends on the path it takes
  • Work (W)
  • Not a state function – depends on path
  • Work = Fd
  • Work = PAh
  • Work = PΔV
  • When work is on the system or compressing
  • Work = - PΔV
  • When work is done by the system or compressing
  • Specific Heat (C)
  • The amount of Joules or calories needed to heat 1 gram of a substance by 1o C.
  • Enthalpy (ΔH)
  • Exothermic – release heat out of the system (-)
  • ΔH = Hproducts – Hreactants = always negative
  • Spontaneous
  • Endothermic
  • Absorbs heat into the system
  • ΔH = Hproducts – Hreactants = always positive
  • Not spontaneous- because they require constant energy
  • First Law of Thermodynamics
  • The energy in the universe is constant.
  • U= EK + EP
  • Internal Energy = U
  • ΔE= Q + W
  • Q= positive, when flowing into system
  • Q= negative, when flowing out of the system
  • W= positive, when the work is done by the system
  • W= negative, when the work is done to the system
  • If no work is being done
  • ΔH = Q
  • ΔH = ΔE + PΔV (must be in Pa x m3)
  • Pressure is constant in this equation
  • Molar Heat of Combustion = Heat for one mole
  • To figure out ΔH:
  • Calorimetry  Q=mCΔT
  • Hess’s Law
  • Heat of Formation Table
  • Be careful to look at physical states
  • Look at coefficients
  • If it’s an element its 0!
  • Joule
  • N x m = Pa x m3

Unit 4 – Atomic Theory and the Nucleus

  • Matter
  • Has mass and takes up space
  • Stuff you can touch
  • Protons – 1 amu
  • Neutrons – 1 amu (1 proton + 1 electron)
  • Electrons – 0 amu
  • Energy
  • Light
  • Electromagnetic Radiation
  • Ability to do work
  • Travels in waves
  • vλ = C
  • C = 3.00x108 m/s , v = frequency = waves/second = Hz, λ = wavelength = m
  • FM = Megahz  can move because it has more protons than AM  Kilohz
  • Nanometers  Meters = multiply by 1x10-9
  • Matter and Energy can indirect
  • Only way to move mass is with mass
  • Photoelectric Effect
  • Einstein proved that there were photons on light waves
  • Max Planck
  • E=hv  h = 6.63x10-34J x sec/waves
  • Quantum theory
  • Neils Bohr
  • Higher frequency = more excitement = more light given off
  • If a lot of frequency – electrons are ionized and are given off creating electricity
  • Hydrogen
  • Hydrogen makes red, blue-green, and two violets.
  • Balmer’s equation – 1/λ = 1.097x107 m-1 (1/22 – 1/n2)
  • This predicts the wavelength of light hydrogen emits when whole numbers are inserted for n and n cannot equal 0, 1, or 2.
  • This is only if it is going to the second energy level
  • Rydbherg’s equation – Energy = 2.18 x 10-18J x (1/n2)
  • This is at any level in a hydrogen atom
  • Energylight= RH (1/nf2 – 1/ni2)
  • =hv
  • Three Complications of Bohr’s model
  • Debroglie
  • Matter travels in waves
  • λ matter = h (constant)/mass x velocity
  • Heisenberg’s Uncertainty Principle
  • Can’t simultaneously know an electrons speed and location
  • e- interact electrically and magnetically
  • Magnetism – comes from electron spinning
  • Schroedinger
  • Variables (quantum numbers)
  • n = principle quantum #  any whole # integer ( 1 - ∞) energy level or region
  • l = angular momentum quantum #  (0 – n-1)  sublevel (s=0, p=1, d=2, f=3)
  • ml = magnetic quantum #  (-l – l )  orbital
  • ms = spin quantum #  ( + ½ , - ½ )  spin
  • Orbital Diagrams
  • Pauli Exclusion Principle
  • An orbital can hold 2 electrons max
  • Aufbau Principle
  • Lowest energy levels are filled first
  • Hund’s Rule
  • Fill empty orbital’s first
  • Electron Configuration
  • Ex. N -1s22s22p3
  • Ex. 157N  Top # = atomic mass number – protons + neutrons, Bottom # = Atomic number - # of protons\
  • Diamagnetic – not magnetic – no unpaired electrons
  • Paramagnetic – magnetic – unpaired electrons
  • To find average weight of an element = The sum of the % of certain weight x weight (amu)
  • Four patterns or trends in periodic table
  • Shielding effect
  • full energy levels blocking the nucleus and its charge
  • Shielding effect as you go P.T.
  • Shielding effect stays the same as you go across the PT
  • Ex. B and C have the same shielding effect – they both have the same number of full sublevels
  • Effective nuclear charge- how effective the nucleus is
  • #protons - #electrons in full energy levels
  • Ex. C has more effective nuclear charge than B
  • C: 5-2 = 3
  • B: 6-2 = 4
  • Atomic Radius
  • Across P.T. = Atomic Radius  effective nuclear charge (adding protons but not electrons in full energy levels)
  • Down P.T.= Atomic Radius  because of increasing shielding effect
  • Exception: it gets bigger because the p sublevel extends past the atom
  • Electron Affinity
  • Energy released or required for an atom to take an electron
  • Negative= energy released, like electrons, more stable
  • Positive= energy absorbed, doesn’t like electrons, less stable
  • Across P.T.= EA  effective nuclear charge
  • Down P.T. = EA  shielding effect increases
  • Ionization Energy
  • Energy required to remove an electron from an element
  • Across P.T.= IE  effective nuclear charge
  • Down P.T. = IE  shielding effect increases
  • Electronegativity
  • Tendency to pull on an electron on a bond
  • Across P.T.= EN  effective nuclear charge
  • Down P.T. = EN  shielding effect increases
  • Four forces
  • Weakest Force – strongest force = gravity weak  electric/magnetic  strong
  • Radiation
  • Elements past 82 protons are radioactive because the neutrons can no longer keep the protons stable
  • Types of Radiation
  • Alpha production (α)
  • Emits Helium nuclei- 42He2+
  • Deflected by an electric field
  • Emitted from unstable nuclei with high atomic number
  • Outside of a body, a paper can stop
  • Inside a body- dangerous
  • Ex.
  • Beta production (β)
  • Emit electrons
  • Fast, penetrating
  • Mass is basically nothing but it is fast (1/100th the speed of light)
  • A neutron in the nucleus is converted to an electron and a proton
  • Usually occurs when atoms have too many neutrons
  • Ex.
  • Gamma Ray Production (ϒ)
  • Photons of electromagnetic energy
  • Pure energy
  • Moves at speed of light
  • Most energetic
  • Wavelength= 10-12 meters
  • Need more protons to stabilize neutrons
  • Ex.
  • Positron Production
  • An electron with positive charge – 01e
  • It happens when u have a deficiency of neutrons for a given number protons in nature
  • Usually emits gamma radiation as well but since gamma radiation has no mass or charge it is optional whether or not to write it in the nuclear equation
  • Ex.
  • Electron Capture
  • Very rare
  • Nucleus captures an electron
  • It is the reverse of beta radiation
  • Makes a neutron
  • Releases gamma ray
  • Ex.
  • Rate of Decay
  • Rare of decay= -ΔN/ Δt
  • N= number of radioactive nuclides
  • Rate of decay is proportional to the number of nuclides of the radioactive material remaining
  • Can be written with a constant  rate of decay = kN
  • First order process
  • ln (Nf /Ni ) = -kt
  • ln Nf – ln Ni = -kt
  • t1/2 = .693/k
  • Decay Event- proportional to amount of substance
  • Nuclear Fission and Fusion
  • Fusion
  • Combining two like nuclei to form a more heavier and stable nucleus
  • Proton- proton chains
  • Release much higher energies but also take more energy
  • Splitting U and Pu nuclei is easier than splitting 2 H nuclei because in H you must overcome repulsion
  • Ex.
  • Fission
  • Splitting a heavy nucleus into two nuclei with smaller mass numbers
  • Possible with isotopes of U and Pu
  • Ex.
  • Exothermic process
  • Large, unstable, radioactive nuclei become more stable by forming smaller, more stable nuclei
  • Chain reaction- self sustaining fission process
  • Plasma
  • An ionized gas
  • Mixture of positive ions and negative ions
  • 4th state of matter
  • Neutral charge
  • Conduct heat and electricity
  • Deuterium
  • Heavy water (0.02% of all water)
  • Must have this for a fusion reaction
  • Tritium- from lithium
  • Lithium- a common metal
  • ΔE = ΔmC2
  • ΔE= change in binding energy
  • Binding energy- energy required to decompose the nucleus into its components or released when the nucleus is formed
  • Δm= mass defect- mass decreases and some of it changes to energy
  • The actual mass of any atom is always less than the mathematically predicted atom.
  • The mass defect was converted into pure energy when all atoms in the universe were initially created…
  • Add up masses of each proton and neutron that make up nucleus
  • Subtract actual mass of nucleus from the combined mass of nucleus from the combined mass of the components to get mass defect
  • Ex.
  • Actual mass of a S-32 atom = 32.066 amu
  • Predicted mass = (16 p+)(1.0078 amu) + (16 n)(1.0087 amu) = 32.264 amu
  • 32.264 amu – 32.066 amu = .198 amu lost to pure energy
  • .198 amu = mass defect
  • When a system loses or gains energy it also gains or loses a quantity of mass
  • Mass is always lost during a nuclear conversion
  • Creating a nucleus will create a mass defect
  • Nuclear Binding Energy
  • The energy required to break down
  • A nucleus into its components nucleons (particles)- kJ/mol
  • 3 Steps
  • Determine mass defect
  • Conversion of mass defect into energy
  • Expressing NBE as energy per mole of atoms, or as energy pre nucleon

Unit 5 – States of Matter