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Chemistry Regents Review
2
Table of Contents
Review #1-2: Introduction/Matter and Energy……………………pg 3
Review #3: Atomic Concepts…………………………………………….pg 6
Review #4: Periodic Table…………………………….……………………
Review #5: Bonding and Shapes……………………………………...pg 19
Review #6: Polarity and Intermolecular Forces…………………...pg
Review #7: Chemical Reactions…..…………………………………………
Review #8: The Mole and Stoichiometry……………………………...pg
Review #9: Equilibrium and Solutions...........................................
Review #10: Energy and Kinetics…………………..…………………………
Review #11: Acids and Bases………………………………………..…pg 48
Review #12: Redox………………………………………………………...pg 52
Review #13: Nuclear Chemistry……………………………………….pg 57
Review #14: Organic Chemistry……………………………………….
Review #15: Reference Tables……………………………………..….pg 63
Review #16: Big Ideas……………………………..……………………..
pg 3
pg 12
pg 17
pg 22
pg 26
pg 30
pg 32
pg 35
pg 39
pg 45
pg 49
pg 54
pg 59
pg 63
pg 66
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Review #1-2 Introduction/matter
403.00 _____ 8.100 x 105 _____ 0.0010 _____ 3,500,000 _____
When calculating:
Metric System:
Multiplying and dividing: round to the same number of significant digits
as the factor with the least number of significant digits
Adding and Subtracting: round to the same decimal place as the factor
with the least number of decimal places.
4
Some important formulas:
Matter and Energy
Matter: has _______ and ______ Energy: capacity to do _______ and
produce _________
The Law of Conservation of Mass:
Solids Liquids Gases
Key Idea: Temperature is a measurement of _______________ _____________ ___________. All three phases differ by the amount of movement among particles.
5
Combined Gas Law
• (note: temperature in Kelvin)
Avogradro's Law
• Same volumes have same number of particles
Graham's Law
• Lighter gases diffuse/effuse faster
Dalton's Law
• Total pressure is the sum of all the partial pressures
Kinetic Molecular Theory describes an __________ gas: Gas particles are in ____________, ___________, ____________ motion Gas particles are separated by ____________ distances relative to size. The volume of
particles is considered _____________. Gas particles have no ________________ forces between themselves. This occurs best
at _______ temperatures and _______ pressures. Collisions between the particles and the _______ of their containers create
___________ and are perfectly ___________ (transfer of energy).
In reality:
Key Idea: Kinetic Molecular Theory is a model that is used to explain the behavior of matter. It describes relationships among pressure, volume, temperature, velocity, frequency, and force of _____________.
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Changes in Phase:
Physical Change Chemical Change
Melting Vaporization Sublimation Freezing Condensation Deposition
Key Idea: A heating curve (which shows an _______________ process) demonstrates that during a phase change, the _____________ ___________ __________, or temperature, does not change.
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Element: ___________________________________________________________________
Compound: _________________________________________________________________
Mixture: ____________________________________________________________________
Homogeneous: ______________________________________________________
Heterogeneous: _____________________________________________________
Approx 118 total Found on the Periodic Table Named after people, places, Latin origins, etc Cannot be broken down physically or chemically
Can be chemically separated Fixed ratio of components Properties of compound are different than individual components Written with formulas
Not pure substances Physically combined, can be physically separated Variable ratio Individual components retain properties
Key Idea: Elements and compounds are considered ________ _______________.
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Separating Mixtures:
1.
5.
2.
6.
3.
7.
Filtration: Used to separate a solid from a liquid. WILL NOT SEPARATE A DISSOLVED
SUBSTANCE!
Distillation: Separates two liquids
based on different boiling points.
Crystallization: Separated a dissolved
substance (aq) from its solvent
Chromatography: Separates based on
polarity and solubility
Practice Regents Questions
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Particle of CO2:
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Review #3 Atomic Concepts
Define:
Atom: ____the basic unit of a chemical element_________
Atomic Number: _____ number of protons in an atom_____________
Mass Number: _____ total mass of the atom (protons + neutrons)_____________
Atomic Mass: _______ weighted average of naturally occurring isotopes_________
Isotope: ____ same element (protons), different mass (number of neutrons)_______
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Valence Electron: ____ outermost electron(s)_________________________
Orbital: _______ area of probable electron location___________________________
Principal Energy Level: ____ ‘shell’ of electrons - 7 total in atom_______________
Excited State: ______ electron has jumped to a higher energy level within atom__
Ground State: __ electrons are in lowest energy configuration (closest to nucleus)
Bright Line Spectrum: __ produced when electrons return to ground state_____
Electron Configurations: ___ denote how electrons are distributed in P.E.L.s____
Models of the Atom
Dalton’s Atomic Theory:
1. All elements are composed of ______________________ 2. All atoms of a given element are ___________________ 3. Atoms of different elements are ________________________________
J.J. Thomson:
Cathode ray tube experiments Discovered the electron Model of the atom: ___________________ Rutherford:
Gold Foil Experiment: Shot alpha particles at thin gold foil, most went straight through, some deflected
and other particles shot straight back.
1. ________________________________________________________________
2. ________________________________________________________________
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Bohr:
Planetary model of the atom Electrons are in defined ________around the nucleus. Wave-Mechanical Model:
In the wave – mechanical model (________________ ________________), the electrons are in
______________ , which are defined as regions of most ______________ electron location
(____________ __________________).
The wave- mechanical model describes the dual nature of the electron as it has properties of both a __________ and a ____________.
Subatomic Particles & Mass
3 subatomic particles make-up the atom:
1. _____________ located in ___________ charge of ____________ 2. _____________ located in ___________ charge of ____________ 3. _____________ located in ___________ charge of ____________
1 amu = __________________________ = mass of 1 proton or 1 neutron
Each atom has its own _____________ _____________ which equals the number of ___________
Key idea: The mass of each proton and each neutron is approximately equal to one atomic mass unit. An
electron is much less massive than a proton or neutron.
Electrons have a ______ charge. In a neutral atom the number of electrons = the number of
___________.
Key idea: The Atomic Number is equal to the number of ______________ in the nucleus of an atom.
The mass number is calculated by adding the number of _______________ + _______________.
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Key idea: The average atomic mass of an element is the ___________ ___________ of the masses of its
_____________ occurring ______________.
Electron Configuration & Bright Line Spectra Key Idea: The outermost electrons in an atom are called the ___________ electrons. The number of
valence electrons affects the chemical properties of an element.
Key idea: * Each electron in an atom has its own distinct amount of energy.
* When an electron in an atom gains a specific amount of energy, the electron is at a higher
energy state known as the ___________ _____________.
* When an electron returns from a higher energy to a lower energy state (_______ ________)
a specific amount of energy is released in the form of _____________.
The emitted energy can be used to identify an
element (_________________________)
1.
7.
2.
8.
3.
9.
Practice Regents Questions
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Review #4
Periodic Table
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Directions:
1. On the above table label the following: a. Group, Period, Noble Gases, Halogens, Transition Metals, Alkaline Earth Metals, Alkali
Metals. b. Place an X over the diatomic elements c. Circle the elements that are liquids d. List all of the gases: _____________________________________________________ e. Label the ions groups 1, 2, 13-18 form above the group.
Define: Atomic Radius: __distance from the nucleus to the outermost electron__________________________ Metallic Character: ___How easily an element loses an electron (reactivity)_______________________ Ionization Energy: __energy required to remove an electron___________________________________ Electronegativity: ____attraction/affinity for electrons in a bond________________________________ Facts and Trends of The Periodic Table 1. The elements on the periodic table are arranged in order of: _________________________________ 2. As you go across a period the atomic radius: ______________________________________________ 3. As you go down a group the atomic radius: ________________________________________________ 4. A positive ion has a _________ radius than the original atom 5. A negative ion has a ___________ radius than the original atom. 6. As you go across a period the ionization energy:___________________________________________ 7. As you go down a group the ionization energy: ___________________________________________
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8. As you go across a period the electronegativity: _________________________________________ 9. As you go down a group the electronegativity: ___________________________________________ 10. As you go across a period the metallic character of the elements: ____________________________ 11. As you go down a group the metallic character of the elements: ____________________________ 12. Name the two groups that are so chemically active that they occur naturally only in compounds. _________________________ 13. Name the group that contains elements in three phases of matter at room temp. ________________ 14. The majority of elements on the periodic table are in what phase at room temperature? __________ 15. The seven semimetals, found on the step line, are ______________________________ Properties of Metals and Nonmetals:
Metals Nonmetals
Location on table
Appearance
Physical properties
Phases
Only liquid example
Conductivity
Ionization
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Electronegativity
Ions formed
Diatomic Elements: _________________________________
Valence electrons are the ______________ electrons.
1.
8.
2.
9.
Key Idea: The periodic table is arranged by ____________ ___________. The properties of elements are ___________, or repeat.
Key Idea: Because each member of a __________ has the same number of valence electrons, each member of the group has ____________ _____________.
Key Idea: An ______ is an atom with a charge. Metals form ____________ ions by _________ electrons, while nonmetals form _______________ ions by ____________ electrons.
Practice Regents Questions
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Review #5 Bonding and Shapes
Lewis Dot Diagrams: depict valence electrons
Li Ca O Br-1 Mg+2
H2O NH3 MgCl2 CO2 CH4
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Define: Ionic Bond: ________________________________________________________________
_____________________________________________________Example_______________
Covalent Bond: ____________________________________________________________
_____________________________________________________Example_______________
Octet Rule: ________________________________________________________________
Naming Ionic Compounds: name the ___________ and ____________. (Table E has ________________ ions). Add an _________ ending if the anion is not from table E.
Add roman numerals, which indicate the ___________, onto the following elements:
Naming Covalent Compounds: use the _____________ listed below. You can drop the _________ prefix on the ________ atom only. Writing Formulas: ensure that the ____________ balance.
Key Idea: Atoms will gain, lose, or share ___________ to become _________. Stability is achieved with ____ valence electrons. (Remember – first shell has _____ valence electrons)
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Criss cross method:
Types of Solids:
Crystalline Solid: Amorphous Solid:
Allotrope:
1.
6.
Practice Regents Questions
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Review #6
Polarity and Intermolecular Forces
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Polar Bond: ________________________________________________________________
_____________________________________________________Example_______________
Nonpolar Bond: ____________________________________________________________
_____________________________________________________Example_______________
Shapes and Polarity: ____________________________________________________
Intermolecular Forces:
Bent:
Pyramidal:
Linear:
Trigonal Planar:
Tetrahedral:
Key Idea: The larger the _______________ in electronegativity, the more __________ the bond is.
Stronger Weaker
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Vapor Pressure: ____________________________________________________________
Key Idea: Intermolecular forces affect physical properties. The stronger a substance’s intermolecular substances, the __________ its boiling point, weaker its __________ pressure, and more _________ tension it experiences.
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Practice Regents Questions
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Review #7 Chemical Reactions
Synthesis
Decomposition
Single Replacement
Double Replacement
Single Replacement Occurs:
Double Replacement Occurs:
Balancing: remember, add ______________ only. Do the following reactions occur?
___Fe + ___NaCl ___Na + ___FeCl3
___Cu(OH)3 + ___NaCl ___CuCl3 + ___NaOH
Key Idea: The law of conservation of mass states that the number of atoms on each side of the equation must be the same.
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2.
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3.
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Practice Regents Questions
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Review #8 The Mole and Stoichiometry
Mathematical Calculations:
Gram Formula Mass: % Composition
Empirical/Molecular Formulas
The Mole:
Two most likely conversions:
Key Idea: Converting between any quantity and another is best accomplished with dimensional analysis. Ensure that your units cancel out.
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6.
2.
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3.
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Practice Regents Questions
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Review #9 Equilibrium and Solutions
Chemical Equilibrium:
Equilibrium can also be in ________________. The processes of _______________ and _______________ are opposite reactions; ______________ equilibriums also exist.
Key Idea: When a system is at equilibrium, reaction _______ are equal. Concentrations of reactants and products are ____________, but not necessarily equal.
Le Chatelier’s Principle: _________________________________________________________ Stresses on a system include a change in: 1. 2. 3. Note: The addition of a catalyst ________________ the rate of the forward and reverse reactions. However, there is NO SHIFT IN EQUILIBRIUM.
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Soluble:__ability to dissolve_______Insoluble:___unable to dissolve __________
Solute: __ is dissolved____ Solvent: ____ does the dissolving_________________
Solution: _____ solute + solvent = homogeneous mixture______________________
Unsaturated: ____ solvent can hold more solute______________________________
Saturated: _____ solvent cannot hold more solute____________________________
Supersaturated: ___ solvent is holding more solute then it can hold__________
The solubility of a substance is affected by the ___________ of the solute/solvent, ______________, and _____________ (for gases only). The rate, or how quickly a solute dissolves, is affected by ______________, ____________, and ____________ ________.
“Like Dissolves Like”
When ionic substances dissolve, the _______ become surrounded by water molecules. The oxygen of surrounding water molecules are attracted to _____ ions, while the hydrogens of surrounding water molecules are attracted to ______ ions.
Colligative Property:
Na+ Cl-
Key Idea: Boiling point is __________, freezing point is ____________ with the addition of solutes. _______ and ___________ pressure also changes with the addition of solutes.
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Practice Regents Questions
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Review #10 Energy and Kinetics
Chemical Kinetics:
Reaction rates are affected by:
Energy:
1. What is the number of joules of energy released when 25g of water are cooled from 20 C to
10C?
2. What is the total number of kilojoules of heat needed to change 150g of ice to water at 0 C?
Key Idea: Collision theory states that for a reaction to occur, a ______________ must occur between two particles. The particles must have the correct _____________ and __________.
Mechanics and rates of reactions Particles must collide to react
Capacity to do work/produce heat Always transferred/change form = conservation of energy warm cold
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Catalyst
Exothermic Endothermic
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Heat of reaction: ΔH = ____________ Entropy: ΔS = ____________
Universal Tendencies:
Table I -ΔH = Exothermic
1. ________________ created 2. _______________ increased 3. more ______________ created 4. ___________ from solids and liquids
2. If 2.0 moles of nitrogen are consumed in this reaction, how much heat is absorbed?
1. Explain, in terms of ∆S and ∆H, why this reaction is spontaneous.
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2.
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3.
7.
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Practice Regents Questions
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Review #11 Acids and Bases
Define: Arrhenius Acid: ___donates H+
_____________________________________________
Arrhenius Base: ___ donates OH-___________________________________________
Bronsted-Lowry Acid: ____ donates H+_____________________________________
Bronsted-Lowry Base: ___ accepts H+____________________________________
Neutralization: ____ acid + base H2O + salt______________________________
Indicator: _____ changes color at different pH levels___________________________
Ionize/Dissociate: ____ molecules break into ions____________________________
pH: ___ measures amount of H+ ions________________________________________
Titration: ____ lab technique to measure molarity of acid/base____________
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Electrolytes: Titration:
Bronsted-Lowry:
Key Idea: _______, __________, and _________ ________ are all electrolytes due to their __________ ___________ ______________.
MAVA = MBVB Remember!
MA = molarity of H+ ions MB = molarity of OH- ions [H2SO4] = .3 M [H+] = [Ca(OH)2] = 1.4 M [OH-] = [H3PO4] = .15 M [H+] = [Mg(OH)2] = [OH-] = 1.5 M
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Practice Regents Questions
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Review #12 Redox
“GER”
When an atom is oxidized (LOSE ELECTRONS) it acts as a ___________ ________because that electron is GAINED by another atom. When atom is reduced (GAINS ELECTRONS) it acts as an ____________ ________because the electron is LOST by another atom.
Oxidation:
Reduction:
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Vo
ltai
c C
ell
Elec
tro
lyti
c C
ell
.
1. Cu + 2AgNO3 Cu(NO3)2 + 2Ag Oxidized: ______ Oxidizing Agent: ________ Reduced: ______ Reducing Agent: ________ Oxidation ½ Reaction: ________________________ Reduction ½ Reaction: ________________________ 2. 2Mg + Pb(NO3)4 2Mg(NO3)2 + Pb Oxidized: _______ Oxidizing Agent: _________ Reduced: _______ Reducing Agent: _________ Oxidation ½ Reaction: __________________________________
Reduction ½ Reaction: __________________________________
Key Idea: A voltaic cell is a ________________ redox reaction. Oxidation occurs at the _______ (“____ ______”) and reduction at the ____________ (“ ______ _________”)
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1. Label the anode 2. Label the cathode 3. Draw in the flow of electrons 4. What is the purpose of the salt bridge? ________________________________ 5. Which electrode is gaining in mass? ___________________________________ 6. Which electrode is losing mass? ______________________________________ 7. Write the half reaction that takes place in half- cell 2: _____________________ 8. Write the half reaction that takes place in half- cell 1: _____________________ 9. Write the overall balanced redox reaction: ______________________________ 10. What type of energy conversion takes place in the cell shown? ____________
1. Label the anode 2. Label the cathode 3. Draw in the flow of electrons 4. Which electrode is gaining in mass? ___________ 5. Which electrode is losing mass?_______________ 6. What is the purpose of the battery? ___________________________________
Key Idea: An electrolytic cell is a __________________ reaction that requires an _________ energy source to occur. _______________ and ______________ are examples.
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Practice Regents Questions
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Review #13 Nuclear Chemistry
Nuclear Reactions:
239 Pu ___________________ 94
90 Sr ____________________ 38
53 Fe ____________________ 26
Key Idea: A nucleus that is ___________ will spontaneously emit radioactive particles. An atom’s nucleus is unstable if:
1. The _____________________________ is unbalanced.
2. The nucleus has more than ____________ and is simply too large to hold itself together.
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Half – Life
The amount of time it takes for half the sample to decay.
If I have a 12.5 gram sample of N-16, how much time has elapsed if the original sample was 100.0 grams?
How many grams of a 60 gram sample of Ra-226 will be left after 3198 years?
Tips for solving half-life problems:
Never divide grams by grams. You
can only multiply or divide grams
by 2.
# of H.L. = total time elapsed Length of H.L. Use Table N for specific times.
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Artifical Transmutation
Uses & Dangers of Radioisotopes
Radioactive isotopes have many beneficial uses. Several different radioactive isotopes are used
in medicine and industrial chemistry.
Tracing biological processes Detecting diseases Treating cancers/diseases Nuclear power
Fission
Fusion
Iodine – 131:
Carbon-14
Cobalt-60
U-238 to Pb-206
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Practice Regents Questions
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Review #14 Organic Chemistry
Alkanes: Single bonds, Saturated
Draw and name: C3H8
Draw: Butane Name the structure below:
Alkenes: Double bonds, Unsaturated
Draw and name: C3H6
Draw: 1-butene Name the structure below:
Alkynes: Triple bonds, Unsaturated
Draw and name: C3H4
Draw: 2-pentyne Name the structure below:
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2-bromobutane
2,3-butandiol
Pentanoic acid
Ethyl pentanoate
Name the following:
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Organic Reactions:
Combustion
Hydrocarbon + O2
Fermentation
Sugar
Substitution
(remember: involves ______________ hydrocarbons)
Addition
(remember: involves ______________ hydrocarbons)
Esterfication
Saponification
Ester + Base Alcohol + Soap
Polymerization
Addition: Condensation:
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1. Put a circle around the alkanes, a square around the alkenes, and a triangle around the
alkynes.
Butane 3-propyne 2-methyl-1-hexene
2-octene 2,3-dimethyl octane
C4H6 C3H4 C7H16 C2H4 CH4
2. Which of the hydrocarbon series are considered saturated? ___________________
3. What is wrong with the following diagram?
4. Draw the isomers of hexane – try to do it without your notes - there are 5 total
5. Put a circle around the compounds that could be used in an addition reaction, and a
square which could be used in a substitution reaction.
C4H6 C3H4 C7H16 C2H4 CH4
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Review #15 Reference Tables
Directions: Find each of the following pieces of information in the Reference Tables. 1. If I start with 100 grams of a sample and after 58.2 years I have 25 grams,
what atom did I start with?
2. What is the name of the anion in KMnO4? What is its charge?
3. If a solution turns yellow with the addition of methyl orange and bromthymol blue, what range could the pH be?
4. How much NaNO3 can be dissolved in 200 grams of water at 45°C?
5. If I overhear a discussion about the pascal, what quantity is being discussed?
6. What classification of hydrocarbon would C2H4 be?
7. What liquid boils at 80°C under normal pressure (101.3 kPa)?
8. This element has an electronegativity of 2.6 and a boiling point of 718 K.
9. What is the standard temperature of water?
10. Is NaOH soluble under normal conditions?
11. Of the following compounds, the decomposition of which one is exothermic: KNO3, NH4Cl, NaOH, or NaCl.
12. What is the formula for acetic acid?
13. Can this reaction occur: Ag + CuNO3 ??
14. What emitted particle has a mass of 4 amu?
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15. CH3 – CH2 – O – CH3 would be what class of organic compound?
16. What element has an electron configuration of 2-8-1?
17. 1.5 x 10-9 meters can also be expressed as 1.5 ____meters.
18. What is the weighted average of all the naturally occurring isotopes of oxygen?
19. What is the heat of fusion of water? Using table T, solve the following problems:
1. Give the parts per million of solute for a solution containing 25g of sodium chloride in 200g of water.
2. If the accepted value for the mass of an object is 10.3g and a student
found that the mass was 10.1g, what is the student’s percent error? 3. If a peanut is burned in a calorimeter containing 50g of water, and the
water temperature changes from 450C to 570C, how many joules of energy were released by the peanut?
4. Convert: 40 grams of NaCl into moles of sodium chloride.
5. What mass of iron has the dimensions 3.20 cm x 6.50 cm x 0.30 cm?
6. Convert: 384 K into °C.
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7. A sample of gas is in a container at 30°C and 100 mL under a pressure of
40 kPa. If the container is put under STP conditions, what is the new volume of the container?
8. If I dissolve 5 moles of NaCl into 2.50 L of water, what is the molarity of the resulting solution?
9. 28 mL of 0.500 M HCl is required to neutralize 50. mL of NaOH. What is the molarity of the sodium hydroxide?
a. How would your answer change if the base being neutralized was Ca(OH)2?
10. What is the percentage composition of oxygen in MgSO4?
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Review #16 Big Ideas
1. Atomic # = # of Protons = # of electrons in an atom
2. Isotopes: Same element (same protons) different number of neutrons
3. How light is produced: An electron gains energy and jumps to a higher principle
energy level (excited state). Then the electron falls back to the ground state emitting
energy in the form of light.
4. Know Periodic Trends in atomic radius, ionization energy, and electronegativity.
5. Ideal Gas: High Temperature and Low Pressure
6. Heat is a form of energy. Heat travels from hot to cold.
7. When a chemical bond is formed energy is released.
8. When a chemical bond is broken energy is absorbed.
9. Temperature measures Average Kinetic Energy of molecules and atoms.
10. Higher boiling points = STRONGER intermolecular forces
11. If a substance sublimes at room temperature it has WEAK intermolecular forces.
12. Gases have the most entropy
13. A saturated solution is at equilibrium
14. Electrolytes: Ions in an aqueous solution that will conduct ELECTRICITY.
15. Electrolytes include: Acids, Bases, and Aqueous Salt solutions
16. Atoms are made up of mostly empty space.
17. Run away from an added stress in a reaction at equilibrium.
18. The addition of a catalyst to a reaction lowers the energy of activation.
19. Coefficients are used to balance an equation. They also represent the ratio of moles.
20. Standard Temperature Pressure (Table B).
21. Ionic Bond: Between a metal and a nonmetal
22. Covalent Bond: Between two nonmetals
23. Like polarity will dissolve in like polarity. (polar in polar, etc)
24. As temperature increases the solubility of a gas decreases.
25. Pressure affects the solubility of gases NOT solids.
26. Acids donate protons (H+ )
27. Bases accept protons (H+ )
28. Arrhenius Acids: Yield H+ ions
29. Arrhenius Bases: Yield OH-1 ions
30. Fission: (Look for NEUTRONS) ( )
31. Fusion: Combining of Hydrogen to make helium.(
1H + 3H 4He + Energy)
32. All reactions show a conservation of MASS, CHARGE and ENERGY
33. LEO: Lose electrons oxidation (get more +)
34. GER: Gain electrons reduction (get more -)
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35. An Ox: Anode Oxidation
36. Red Cat: Reduction Cathode
37. Alpha particles have a mass of 4 amu and a positive charge
38. Beta particles have no mass and a negative charge
39. Particle Diagrams:
Elements Compounds Heterogeneous Mixture
40. Collision Theory: If you add more reactants there will be more effective collisions
between the reactants causing the reaction to shift right and make more products.
41. H2 O2 F2 Br2 I2 N2 Cl2 : Diatomic elements
42. Valence Electrons: Are used in bonding and give elements their unique chemical
properties
43. If you see the word TITRATION or NEUTRALIZATION use the formula: MA VA =
MB VB
44. Mole – Mole problems: set up a proportion using coefficients and the # of moles
given
45. Mole – Gram or Gram – Mole conversions: Moles = Given Mass
Gram formula mass
46. Hydrocarbon: Alkane, Alkene, Alkyne
47. Class of organic compounds and functional groups: Table R
48. Organic Reaction: Substitution
Two reactants and two products
49. Organic Reaction: Addition
Two reactants and one product
50. Organic Reaction: Esterification
Organic Acid + Alcohol
51. Heating Curve for Water
q = mc ∆T
q = mHf
q = mHv
