Libretto - Chem Review
Name____________________
Chemistry Regents Review Concept Sheets
Table of Contents
Topic 1 – Math Skills and Relationships – pg. 2
Topic 2 – Physical Behavior of Matter – pg. 3
Topic 3 – The Atom – pg. 5
Topic 4 – Nuclear Chemistry – pg. 7
Topic 5 – The Periodic Table – pg. 9
Topic 6 – Bonding – pg. 11
Topic 7 – Formulas and Equations – pg. 13
Topic 8 – The Mathematics of Formulas and Equations – pg. 15
Topic 9 – Kinetics and Equilibrium – pg. 16
Topic 10 – Properties of Solutions –pg. 18
Topic 11 – Acids, Bases, and Salts – pg. 20
Topic 12 – Oxidation-Reduction – pg. 22
Topic 13 – Organic Chemistry – pg. 24
Mr. Libretto
www.elibretto.wikispaces.com
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Libretto - Chem Review
Topic #1 – Math Skills and Relationships
Math Concepts/Skills to Master:
1. Graphing
2. Understanding Relationships (direct/indirect etc.)
3. % error =
measured value−accepted value
accepted value
x 100. (See Reference Table T).
4. Significant Figures Rules:
a. Any number 1 through 9 is significant
b. Any zeroes in between 1 through 9 numbers are significant
c. No decimal  Any zeroes at end are not significant
d. Decimal  Find the first number 1 through 9 all the way to the left. All numbers from
there to the right including 0 are significant. (Remember Pacific decimal point present
and Atlantic decimal point absent, start counting from that side with the first nonzero
number.)
Ex: 435 (3 sig figs); 4035 (4 sig figs); 403500 (4 sig figs); 4020. (4 sig figs); 4020.0 (5 sig figs);
4020.06 (6 sig figs); 0.0402 (3 sig figs); 0.4020 (4 sig figs); 0.402010 (6 sig figs).
Decimal
Present
Decimal
Absent
5. K= °C + 273 (Reference Table A) Ex: 10°C = 283 K; -273°C = 0K
6. Unit Conversions (Reference Table C):
mass
7. Density = volume
8. Melting Point and Boiling Points on Reference Table S
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Topic #2 – Matter and Energy
1. Properties of Solids, liquids, gases: Solid – definite shape and volume;
Liquid – definite volume, no definite shape; Gas – no definite shape or
volume
2. Element – one kind of a substance – all atoms have the same atomic
Solid
Gas
Liquid
number (homogeneous), which cannot be decomposed. Ex: Copper
(Cu)
3. Compound – two or more elements chemically combined in a definite proportion by weight
(homogenous), which can be decomposed. Ex: CuCl2, H2O
4.
Mixture – two or more substances physically mixed but not chemically
combined (may be heterogeneous or homogenous). Ex: NaCl(aq). Mixtures of
solids are heterogeneous but solutions are homogenous.
5. Physical Change: no change in the identity of substance. Ex – melting, boiling,
evaporating, freezing, condensation, sublimation. H2O(g)  H2O(l)
6. Chemical Change: a substance changes into a different substance with different properties (a
chemical reaction). Ex – Decomposition of water (2H2O  2H2 + O2).
7. Joule Problems: Know formula: Q=mcΔT. m= mass in grams; ΔT = change in temp.; c = specific
heat
8. Joule: 4.18 J changes the temp. of 1 gram of water by 1°C. KJ = 1000 J
9. Temperature: a measure of the average kinetic energy of the molecules of a substance.
Temperature scales: Kelvin (absolute) and Celcius. °K = °C + 273 but 1°K = 1°C. 10°C has
higher kinetic energy than 5°C.
10. Fixed Point on Thermometer: 0°C (273 K) = melting point/freezing point of H2O and 100°C (373
K) = boiling point/condensation point of H2O. Absolute zero = 0 K (-273°C).
11. Gas Laws Problems: P1V1/T1 = P2V2/T2. a) Drop what is constant. b) Temperature must be in K!
12. Boyle’s Law: (constant Temp.) P and V vary inversely. P1V1 = P2V2
13. Charles’s Law: (constant P) V and T vary directly V1/T1 = V2/T2.
14. Activation Energy: minimum amount of energy needed to start a reaction (all reactions need
activation energy).
15. S.T.P. : standard temp. and pressure = 101.3 kPa (1 atm) and 0°C (273 K). (Table A.)
16. Density: mass/volume.
17. Phase Change Diagrams: during a phase change (ex – melting, vaporization, freezing) the temp.
stays constant. No change in K.E., only P.E.
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Heating Curve
Cooling Curve
Endothermic: sl, lg (forward phase change)
Exothermic: gl, ls (reverse phase change)
18. Kinetic Molecular Theory: Ideal Gas – no attractive forces between molecules, molecules have no
volume. Real Gas – (ex. H2, O2) has attractive forces, have volume.
19. Ideal and Real gases most alike (deviate the least) when attractive forces are the weakest (need
molecules apart to have the weakest forces). P.L.I.G.H.T – low pressure, ideal gas, high
temperature.
20. The Smallest molecule is the most ideal – He, and other monatomic Group 18 molecules. Next
most like ideal are diatomic molecules (BrINClHOF)
21. Heat of Fusion: # of joules required to melt one gram of solid at its melting point. H2O = 334 J/g
(Table B). q = mHf
22. Heat of Vaporization: # of joules required to vaporize one gram of liquid at its boiling point. H2O =
2,260 J/g (Table B). q = mHv.
23. Boiling Point: the temperature at which the vapor pressure of a substance equals atmospheric
pressure. (Normal B.P. = B.P corresponding to the air pressure = 1atm, 101.3 kPa).
24. Vapor Pressure: dependent on: a) temp. of the liquid – the higher the temp. the higher the v.p.; b)
strength of the intermolecular forces of attraction – the stronger the attractive forces, the lower
the v.p (Ref Table H).
25. Substances that evaporate readily have high v.p. and low b.p. Sublimation – weak attractive
forces, high vapor pressure. Higher the b.p., the stronger the IMAF are.
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Topic #3 – The Atom
1. Fundamental particles of the atom:
Proton = +1 charge, mass = 1 a.m.u
Electron = -1 charge, mass = ~0 (1/1836 of a.m.u)
Neutron = 0 charge, mass = 1 a.m.u.
1 a.m.u. = 1 atomic mass unit = 1/12 mass if Carbon-12.
2. Nucleons: protons and neutrons (the particles in the nucleus).
3. Nucleus: contains most of the mass of the atom, and has a positive charge. Charge of the
nucleus = the number of protons (called the nuclear charge). Ex: if an atom has 3 protons and 4
neutrons the nuclear charge = +3; an atom with 7 protons and 5 neutrons the nuclear charge =
+7.
Proton
Neutron
Charge
+1
0
Mass #
1
1
Electron
-1
0
Location
Nucleus
Nucleus
Outside
Nucleus
4. Atoms are neutral: # protons = # electrons. Ex = If an atom has 14 protons, how many electrons
will it have? Answer = 14.
5. Atomic Number = # protons in the nucleus. All atoms of the same element have the same
number of protons. This identifies the element. Ex: All atoms of sodium must have 11 protons.
6. Mass Number = the # of protons and the # of neutrons in the nucleus (a whole # not found on
the Periodic Table).
𝑝+𝑛
7. Nuclear Notation: shows Mass Number = p+n; Atomic Number = p as: 𝑝 X
Ex. You are given 15
C, how many neutrons are there? Answer = top # - bottom # = 15 – 8 = 7
8
8. Isotopes are atoms of the same element (same # of protons) but a different number of neutrons.
Isotopes of the same element have the same number of protons and electrons. Ex: = 12
C=6
6
protons, 6 neutrons, and 6 electrons but
14
6
C = 6 protons, 8 neutrons, and 6 electrons.
9. Average atomic mass of an element (on periodic table) is determined by taking the weighted
average mass of the naturally occurring isotopes of the element. Because of this, the atomic
mass listed on the periodic table is a decimal.
Average Atomic Mass = (% isotope1)(mass isotope1) + (% isotope2)(mass isotope2)…etc.
Ex: 20X = 80%, 22X = 20% therefore avg. atomic mass = (.80)(20)+(.20)(22) = 20.4 amu
10. Ernest Rutherford – gold foil experiment shows a) atoms are mostly empty space;
b) atoms have a positively charged, dense center.
11. Bohr’s Model of the Atom: an electron in the ground state can absorb energy and jump to higher
energy levels or an excited state; the energy is then released as wavelengths of light as the
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electrons fall back to the ground state. The wavelengths of light emitted are unique to each
element, so an element’s identity can be found by studying the wavelength (color) of light it gives
off. This is the emission spectra of light.
12. Valence Electrons: #e in the outermost principle energy level. Ex. 9 F has 9 total electrons, and 7
valence electrons.
13. Electron Dot Diagram: uses dots to show the number of valence electrons.
Ex. Fluorine 9 F = 2-7
Carbon 6 C = 2-4
14. Know the following chart:
Principal Energy
Levels (P.E.L.)
n=1
n=2
n=3
n= 4
Number of e- in
the P.E.L.
2
8
18
32
The maximum number of electrons in each energy level = 2n2
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Topic #4 – Nuclear Chemistry
1. Radioisotope = an isotope that is radioactive and unstable.
2. Stability of nuclei – nuclei are composed of protons and neutrons. The
ratio of protons to neutrons determines the stability of the atom.
a. If the ration of protons to neutrons is not stable, the isotope is
radioactive (radioisotope).
b. Any element with an atomic number great than 83 (83 or more
protons) is radioactive.
3. Transmutation: different elements from reactant side to product side.
226
222
4
Ex: 88 Ra  88 Rn + 2He
4. Chemical Reaction: Has the same elements on both sides.
Ex: 2H2 + O2  H2O
5. Artificial Transmutation: Bombard a nucleus with high-energy particles that change it from one
element to another. Transmutation reactions have two things on reactant side, and the
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4
17
1
elements change from reactants to products. Ex: 7 N + 2He  8 O + 1H
6. Natural Transmutation: (one reactant on left side, elements change from left side to right
side):
a. Alpha, Beta, positron, gamma decay (see pg. 217 of review book or Reference Table
O).
b. Alpha Decay:
Atomic # decreases by 2, Mass # decreases by 4.
c. Beta Decay:
232
0
232
Th  −1e + 91 Pa; Mass # remains the same; atomic number increases by 2.
90
d. Positron Decay:
37
0
37
K +1e + 18Ar; Mass # remains the same; atomic number decreases by 1.
91
e. Gamma Decay:
3
3
0
He  2He + 0 ϒ; no change in mass # or atomic # - high-energy photon released.
3
Reference Table N has decay modes.
7. Nuclear Fission: splitting of an atom to produce energy (exothermic).
235
1
87
146
1
U
+
n

Br
+
La
+
3
n
92
0
35
57
0
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8. Nuclear Fusion: joining light molecule together to produce energy.
2
2
4
H + H  He; difficult to initiate because positive nuclei repel each other.
1
1
2
9. Half-Life: the amount of time it takes for a radioactive sample to half its
radioactive mass (Reference Table N)
a. Half-life not affected by anything, including temperature or pressure.
b. Reference Table T Equations – Radioactive Decay:
𝑡
1𝑇
Fraction Remaining = 2 ; t = total time; T = half-life.
𝑡
Number of half-life periods = 𝑇; t = total time; T = half-life. (Use when given two different
times).
1 (if fraction)
or
Mass in kg
Half-Life Time
Fraction
or
Final Mass (kg)
10. Uses of radioisotopes: - used to trace the path of a chemical reaction (tracers)
a. C-14: used for dating organic substances (anything that was once living)
b. I-131: Thyroid diagnosis
c. Co-60: treating cancer
d. U-230: used in geological dating
e. Tc-99: used to trace cancer cells
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Topic #5 – The Periodic Table
1. The arrangement of the Periodic Table is based on atomic number.
2. The Periodic Law states that the chemical properties of elements are periodic functions
of their atomic numbers.
3. Elements are classified as (1) metals (2) non-metals (3) metalloids. More than 2/3 of all
the elements are metals. Trends:
4. Metalloids: have properties of both metals and non-metals.
5. Metals: lose electrons to form positive ions, and become smaller. Ex: Na+ is smaller then
Na0. The ion will generally have the electron structure of an inert gas. (Metals are solid
at room temperature except Hg = liquid).
6. Non-metals: - gain electrons to form negative ions, and become larger. Ex: Cl- is larger
than Cl0 The ion will generally have the electron structure of an inert gas. The only liquid
non-metal is Bromine (Br2).
7. Group 1 metals: called alkali metals (form strongest bases); so active only exist in
compounds.
8. Group 2 metals: called alkaline earth metals; less active than alkali metals.
9. Group 18 nonmetals: called Inert (noble or rare) Gases.
10. Group 16, Group 15 and Group 14: contain both nonmetals and metalloids.
11. Group 17 non metals: called halogens. This group exhibits all three phases of matter at
room temperature (F2 and Cl2 are gases, Br2 is a liquid, I2 is a solid).
12. Groups 3-12 are the transition metals.
13. Elements found in the same period have the same number of energy levels.
14. Elements found in the same group have the same number of valence electrons and
therefore similar chemical properties.
15. The most active metals are in the lower left corner of The Periodic Table (excluding Li)
and the most active non-metals are in the upper right corner of the Periodic Table
(excluding inert gases).
16. The most active elements form the most stable compounds: Ex. Group 1 Rb and Group
17, F = RbF. Very Stable.
17. Monatomic molecules – contain 1 atom per molecule = Group 18: He, Ne, Ar, Kr, Xe, Rn
(generally nonreactive)
Allotropes – different forms of the same element ex – O3, S8
18. Diatomic Molecules – 2 atoms per molecule = Br2, I2, N2, Cl2, H2, O2, F2 (BrINClHOF).
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19. Transition Elements – (1) Group B elements and Group VIII can have electrons from two
outermost shells involved in a reaction. (2) Form colored ions in compounds or solutions
(3) have multiple oxidation states (4) Remember Cu2+ is blue, thus CuSO4 is blue in
solution.
20. Xe and Kr, although expected to be inert (like the other Group O elements), can form
compounds with F and O under special conditions. Ex: XeF4, XeF6 exist.
21. Van der Waals forces increase down a group (due to increasing molecular size), and thus
b.p. and m.p. increases.
Examples:
He
Ne
F2
As Van der Waals Increase
Ar
Cl2
F.P and B.P. Increase
Kr
Br2
Xe
I2
Rn
The smallest molecule has the lowest melting point
22. Formulas: Ex: X2O3
X is in Group 3, has a +3 charge (oxidation number)
XD
X is in Group 2, has a +2 charge (oxidation number)
23. States of elements: Metals are all solids at room temperature except Hg is a liquid.
Nonmetals that are gases H2, N2, O2, F2, Cl2 and Group 18.
Liquid = Br2; Solids = C, P, S, I2 (nonmetals)
24. Atomic radius decreases across period due to increase in # of protons (greater nuclear
charge) in nucleus (more pull with same # of energy levels).
Atomic radius increases down each group due to the addition of an extra shell for each
successive element (increases size substantially).
25. Carefully examine charts – Periodic Table, and Table S for atomic radius, ionization
energy and electronegativity trends by element
.
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Topic #6 –Bonding
1. When a bond is made energy is released (exothermic). When a bond is broken, energy is
absorbed (endothermic). Remember “Karate Chop.”
2. Atoms bond together to become more stable (usually to get a stop octet ending).
3. Metals tend to lose electrons to form positively charged ions (the metal ion is smaller
than the atom), and have low electronegativity values. Ex: Na+ is smaller than Na0.
4. Non-metals tend to gain electrons to form negatively charged ions (the ion is larger than
the atom), and have high electronegativity values (F is the highest) Ex: Cl- is larger than
Cl0.
5. A chemical bond: results from the simultaneous attraction of electrons to two nuclei.
(Usually between two non-metals.)
6. Ionic Bonds: formed between metal and non-metal atoms; created by a transfer of
electrons. Ex: NaCl; MgBr2. (MINT – Metal Ionic Nonmetal Transfer of e- from metal to
nonmetal).
7. Covalent Bond: formed by the sharing of two electrons between two nuclei. (Usually
between two non-metals.) Ex. CO2, NH3.
8. Electronegativity: the ability of an atom to attract the electron in a bond. See Reference
Table S for electronegativities. Electronegativity based on a scale of 0.0 - 4.0 By
subtracting the electronegativity values of the two elements, the electronegativity
difference (END) is calculated, and can be used to determine the type of bond:
 Ionic Bond –electrons are transferred (metal and a nonmetal)
 Polar Covalent Bond – electrons are shared unequally – one element has partial
negative charge, the other a partial positive charge
 Non-Polar Covalent Bond – electrons are shared equally – diatomic molecules
 Large END = more ionic, less covalent
 Small END = more covalent, less ionic
Ex: NaCl – Electronegativity difference = 3.2-0.9 = 2.3. This is IONIC
 MgCl2 = 3.2-1.3 = 1.9. This is IONIC. H2O = 3.4- 2.2 = 1.2 this is Polar Covalent.
Diatomic Molecules (BrINClHOF) have non-polar covalent bonding (END=O). Ex.
N2 has a triple bond
9. In electron dot digrams of covalent compounds all atoms need 8 electrons around them,
except for H, which has 2 electrons around it.
10. A molecule is defined as a particle, which has covalent bonds, and is the smallest unit
that shows the properties of the substance. Ex: H2; H2O; CO2.
11. Polyatomic Compounds have both covalent and ionic bonds:
Ex: Na2SO4 has an ionic bond between N-SO4; and a covalent bond between the SO4.
12. K+1 and Cl-1 have the same number of electrons, and the same electron
configuration.
K+1 = 18 electrons Cl- = 18 electrons
13. Ionization Energy: the amount of energy required to remove the most loosely bound
electron. Ionization Energies are list on Reference Table S. Ex: Ionization energy of Li
= 520 kJ/mol, and F=1681kj/mol; this means it takes less energy to take an electron
from Li than from F.
14. Ionic Solids: Ex - NaCl, K2O, etc.
a. Have high melting points, high boiling points
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b. Are hard
c. Do not conduct electricity as solids, but do conduct when
dissolved in water, melted, or evaporated.
15. Metallic Solids: Ex – Ag, Zn, etc.
a. Mobile electrons (“sea of electrons”)
b. Conductors in solid phase
c. Malleable, ductile (bendable, made into wires)
d. Liquid metal at room temp = Hg
16. Molecular Solids: Ex – H2O, CH4, etc.
a. Held together by weak attractive forces (Van der Waals)
b. Have low melting point, low boiling point
c. Are soft
d. Are poor conductors
17. Network Solids: Ex – SiO2, SiC, diamond, graphite
a. Held together by strong covalent bonds
b. Often made of group 14 elements
c. High melting point, high boiling point
d. Are electrical insulators
18. Van der Waals Forces: exist between non-polar molecules – all diatomic, monatomic, and
molecules that are not dipoles. Ex: CO2, CH4, etc.
a. Van der Waals Forces depend on the size of molecules and distance between
molecules.
b. Larger molecules = stronger Van der Waals Forces.
c. Closer the molecules = stronger Van der Waals Forces.
d. Ex: C17H36 is a liquid at room temp, by CH4 is a gas. The larger molecule has
stronger Van der Waals Forces, which requires a higher temp. to turn it
into a gas.
19. Hydrogen Bonding: between molecules, which contain Hydrogen bonded to an
atom of small radius and high electronegativity (HFON). Strongest hydrogen bonds
between molecules of H2O, HF, NH3. Hydrogen bonds responsible for high b.p. of
H2O, HF, and NH3.
20. Polar Molecule: a dipole – polar bonds in an asymmetrical molecule shape; ends have a partial
charge. Polar Molecules:
21. Non-Polar Molecule: molecule has no charged ends – molecule is symmetrical.
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Topic #7 –Formulas and Equations
1. Reference Table S – has the names and symbols of the elements on the Periodic Table.
2. Diatomic Molecules: (BrINClHOF) – elements that pair up when not combined with other
elements – Br2, I2, Cl2, H2, O2, F2.
3. Polyatomic Ions: See Reference Table E – ions made of more than one element and act as a
unit in compounds.
4. Binary Compounds – have only two elements in the formula; Polyatomic Compounds: have
more than two elements in the formula
5. Formula Writing:
a. Criss-cross oxidation numbers to write subscripts
b. Any 1’s for subscript don’t get written. Ex: KCl; MgCl2
c. Keep polyatomic ions in parentheses unless there is a
“1” outside it.
Ex1: Mg(NO3)2 needs parentheses, telling you there are
2 Nitrogens and 6 Oxygens.
Ex2: NaNO3 does not need parentheses
d. If “ate”, “ite”, or ammonium are in the name use Table
E. Otherwise it will be a binary compound. Exceptions
of “-ide” include peroxide, hydroxide, and cyanide. Be
careful of chlorine polyatomics; there are four very
similar ones.
e. Examples of compounds with polyatomic ions from
Table E: Ammonium sulfite (NH4)2SO3; Magnesium
Nitrate Mg(NO3)2; Aluminum hypochlorite Al(ClO)3.
Examples of binary compounds: Hydrogen Peroxide
H2O2; Barium Fluoride BaF.
6. Formulas and The Stock System:
a. When Roman numerals are given in the name of the compound, it is the oxidation
number of the first element.
b. Roman Numerals: I (+1), II (+2), III (+3), IV (+4), V (+5), VI (+6).
Ex: Copper (I) sulfide = Cu2S; Iron (III) chloride = FeCl3; Tin (IV) sulfate Sn(SO4)2;
Lead (II) nitrate Pb(NO3)2
c. Naming Compounds: Make sure the first element has only one oxidation number when
naming without the stock system (roman numerals)
d. Naming Binary Compounds: end in –ide. Ex: MgCl2 is Magnesium Chloride, no roman
numeral since Mg can only have a +2 oxidation number.
e. Naming Polyatomic Compounds: keep the suffix of the polyatomic ion if it is second.
Ex: NH4NO3 ammonium nitrate.
f. Peroxides: the subscripts are not reduced. Ex: Hydrogen Peroxide H2O2; Sodium
Peroxide Na2O2
7. Naming Compounds with the Stock System:
a. Solve for the oxidation number of the first element. FeI2 – oxidation number of Fe
could be +2 or +3, so the number must be written in the name = Iron (II) Iodide. Ex:
CuSO4 = Copper (II) sulfate; Sn3(PO4)2 = Tin (II) Phosphate; SrBr2 = Strontium Bromide
(No stock necessary since Sr has only one oxidation number; no stock number ever
needed for second part of compound.)
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8. Types of Reactions:
a. Synthesis (A + B  AB): one product only.
Ex: 2H2(g) + O2(g)  2H2O (g)
b. Decomposition (AB  A + B): one reactant only
Ex: H2O(l)  H2(g) + O2(g)
or
4Fe(s)+ 3O2(g)2Fe2O3(s)
CaCO3(s)CaO(s) + CO2(g)
c. Single Replacement (A+BC  AC + B): 1 element and 1 compound on both sides.
Ex: Cu(s)+2AgNO3 Cu(NO3)2(aq) + 2Ag(s) or
Mg(s) + 2HCl(aq)  H2(g) + MgCl2 (s)
d. Double Replacement Reaction (AB + CD  AD + CB): 2 compounds on both sides.
Ex: AgNO3 + NaCl  AgCl + NaNO3
e. Neutralization Reaction: Acid + Base  Salt + Water
Ex: H2SO4 + NaOH  Na2SO4 + H2O
(Write H2O as HOH) and balance:
H2SO4 + 2NaOH  Na2SO4 + 2HOH
9. Balancing Equations: number of atoms for each element must be the same on both sides.
a. Fe + Al2O3  Al + FeO. Balance single elements first:
9Fe + 4Al2O3 8Al + 3Fe3O4
b. Keep polyatomic ions together when balancing:
Al2(SO4)3 + ZnCl2  AlCl3 + ZnSO4
Balanced:
Al2(SO4)3 + 3ZnCl2  2AlCl3 + 3ZnSO4
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Topic #8 –The Mathematics of Formulas and Equations
1. Formula Mass: use the atomic mass of each element and round off to the nearest whole number.
Ex: Mg(NO3)2 = 148 g/mol
Ex: CuSO4 5H2O = 250 g/mol
Element # Atoms Atomic
Formula
Mass
Mass
Element
#
X
Formula
Mg
1
x 24
= 24
Atoms Atomic
Mass
Mass
N
2
x 14
= 28
Cu
1
x 64
= 64
O
6
x 16
= 96
S
1
x 32
= 32
=148
O
4
x 16
= 64
𝑝𝑎𝑟𝑡
2. Percent Composition = (𝑤ℎ𝑜𝑙𝑒) *100%
H2O
5
x 18
= 90
 Find % composition of N in Mg(NO3)2 =
= 250
28
x 100 =19%
148
 Find % composition of H2O in CuSO4 5H2O
90
=250x 100 = 36%.

𝐺𝑖𝑣𝑒𝑛 𝑀𝑎𝑠𝑠 (𝑔)
Mole Calculations (See reference Table T): number of moles = 𝑔𝑟𝑎𝑚−𝑓𝑜𝑟𝑚𝑢𝑙𝑎 𝑚𝑎𝑠𝑠
𝑥
3.
4.
5.
6.
 Find the mass of 2.5 moles of Mg(NO3)2. 2.5 moles = 148 𝑔/𝑚𝑜𝑙. x = 370g.
Molecular Formula: shows the actual number of atoms in the molecule. Ex. C6H12O6.
Empirical Formula: show the simplest whole number ration of atoms. Ex. The empirical formula
of C6H12O6 is CH2O (cannot be divided anymore and have a whole number for each).
Finding Molecular Formulas from Empirical Formulas:
 Gram-molecular weight = whole #. Then multiply subscripts of empirical formula by that # to
get molecular form.
Ex. A compound has a molecular mass of 180 a.m.u. and an empirical formula CH2O. What is
the molecular formula: ANS:
C: 1 x 12 = 12
 Formula Mass = 12 + 2 + 16 = 30.
180
H: 2 x 1 = 2
 30 = 6x’s
O: 1 x 16 = 16
 C1x6H2x6O1x6 = C6H12O6
Mole Relations in Balanced Equations
Ex: Find the number of moles of oxygen produced when 1.5 moles of KClO3 decomposes given
1.5 𝑚𝑜𝑙𝑒𝑠
𝑥 𝑚𝑜𝑙𝑒𝑠
the following reaction: 2KClO3  2KCl + 3 O2.
= 3 𝑚𝑜𝑙𝑒𝑠
2 𝑚𝑜𝑙𝑒𝑠
1.5KClO3
X O2
x = 2.25 moles of O2
𝑚𝑎𝑠𝑠
7. Density = 𝑣𝑜𝑙𝑢𝑚𝑒. Density of elements is listed on Reference Table S.
8. Avogadro’s Hypothesis: Equal volumes of different gases at the same temperature
and pressure contain the same number of molecules.
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Libretto - Chem Review
Topic #9 – Kinetics and Equilibrium
1. Bond-Breaking is endothermic (energy must be absorbed to break a bond). When a bond is
made, energy is released – exothermic (karate chop).
2. The heat of reaction (ΔH) is the difference between the potential energy of products and the
potential energy of reactants.
3.
Endothermic
Exothermic
Energy is absorbed
Energy is released (liberated)
Reactants have less energy than products
Products have less energy than reactants
ΔH = +kJ
ΔH = -kJ
Form less stable products in comparison to
Form more stable products than reactants
reactants
4. Exothermic reactions are self-sustaining, because the reaction releases enough energy to keep it
going. Endothermic reactions are not self-sustaining therefore they continually need added
energy to keep it going.
5. ΔH (heat of reaction) is measured in kJ (kilojoules). The ΔH values are expressed for the
compounds formed, and ΔH values for common reactions are found on Table I. To find the ΔH for
the reverse reaction of an equation listed, reverse the sign of the ΔH value.
6. Exothermic Reactions:
A+BC+D
ΔH = -kJ
OR
If you
A+BC+D+kcal
reverse the
reaction,
7. Endothermic Reactions:
reverse the
A+BC+D
ΔH = +kJ
PE diagram
OR
A+B+kJC+D
8. A –ΔH tells you the reaction is exothermic in that direction, if you switch the reaction, the
opposite direction will be endothermic and have a +ΔH.
Ex: For N2 +3H2  2NH3, the ΔH=-91.8kJ and is exothermic. For the reverse, 2NH3 N2 +3H2
the ΔH = +91.8kJ and is endothermic.
Ex: How many kJ are required to decompose 1 mole of NH3? Ans: +91.8kJ/2 moles NH3 = 45.9kJ.
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Libretto - Chem Review
9. Factors Affecting Reaction Rate (speed): (Effective Molecular Collisions)
a. Catalyst: speeds up a reaction by lowering the reaction energy required to start the
reaction.
b. Increasing Concentration of one or more reactants increases the reaction rate due to an
increasing number of collisions.
c. Increasing Temperature: increases the rate of all reactions by increasing the number of
collisions and the effectiveness of the collisions.
d. For gases only: Increasing the pressure will increase the reaction rate by increasing the
number of collisions (molecules are closer together).
e. Increase surface area (solids): break solids into smaller pieces to increase the reaction
rate.
10. Entropy(ΔS): the amount of disorder, randomness, or lack of organization of a system.
a. Solids have the least entropy, gases the most.
b. Compounds have less entropy than free elements. Ex. H2O(g) has less entropy than H2(g)
and O2(g).
c. A compound dissolved in water increases in entropy. Ex. C6H12O6(s) + H2O C6H12O6(aq).
11. Exothermic Reactions (-ΔH) are favored in nature; Reactions with an increase in Entropy (+ΔS) are
favored in nature.
12. Equilibrium shifts: Shifting the equilibrium point to the left makes more reactants, shifting to the
right makes more products. Think of a seesaw; the direction the reaction shifts is the side you
push on to make the seesaw level again. Example:
C3H8(g)+5O2(g) 3CO2(g)+4H2O(g)+2219 kJ
a. If O2 is added, then the left side of the seesaw goes down; push on the right side to make
it level = reaction shift right, which causes the CO2 and H2O to increase, and the C3H8 to
decrease.
b. Pressure: an increase in pressure will cause the reaction to shift to the side with fewer gas
molecules (add coefficients of gas on each side). Ex: increase in pressure will shift the
reaction to the left. A decrease in pressure would do the opposite. Increase in pressure:
c. Temperature: An increase in temperature shift equilibrium away from the kJ, a decrease
towards the kJ. Ex: The reaction in exothermic, so increasing the temperature would shift
the reaction to the left.
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Libretto - Chem Review
Topic #10 – Properties of Solutions
1. Solution – a homogenous mixture. Ex – NaCl(aq).
2. Properties of Solutions:
 Are clear and do not disperse light
 Can have a color (transition elements)
 Will not settle on standing
 Will pass through a filter
3. Rate of Solution (How Quickly it Dissolves):
 Decrease the size by crushing to increase the surface area
 Stir
 Increase the temperature
4. “Likes Dissolve Likes”
 Polar solvents dissolve polar solutes. Ex – NaCl in water.
 Nonpolar solvents dissolve nonpolar solutes (oil paint in turpentine)
5. Temperature Change
 When dissolving, the reaction is endothermic if the water gets colder
 When dissolving, the reaction is exothermic if the water gets warmer.
6. Solubility – maximum quantity of a solute that can be dissolved in a certain amount of solvent or
solution at a specific temperature.
7. Factors Affecting Solubility:
 Nature of substance (Type of substance)
 Temperature - solids increase with higher temp; gases decreases
with higher temp. (Ex - CO2 comes out of soda in fridge)
 Pressure – only affects gases; solubility increases when pressure
increases. (Ex – put cap on soda).
8. Unsaturated Solution – holds less solute than maximum. This means
more can still be dissolved.
9. Saturated Solution (at equilibrium) – dissolves maximum amount;
cannot dissolve any more solute without changing another factor (like
temp, etc.)
10. Supersaturated Solution – temporary state that is dissolving more solute
than it should; very likely will precipitate out.
11. Concentrated – holds large amount of solute relative to the amount of
solvent
12. Dilute – holds small amount of solute relative to the amount of solvent.
13. Solubility Curves: See Reference Table G
14. Units of Concentration: (See Ref. Table T)
moles of Solute
a. Molarity = Liters of Solvent;
b. # moles = Molarity x Volume in L
c. #grams = Molarity x Volume in L x GMW
Ex: What is the molarity of a solution containing 82.0 g of Ca(NO3)2 in 2.0 L of solution:
moles
Ans: GMW of Ca(NO3)2 = 164 g/mol  0.5 moles. M=𝐿𝑖𝑡𝑒𝑟𝑠 =
0.5 𝑚𝑜𝑙𝑒𝑠
2.0 𝐿
= .250 M
15. Colligative Properties: Solutes added to water will raise the boiling point and lowers the freezing
point of water. Same strength electrolytes (acid, base, or salt) will change the b.p and f.p. more
than a nonelectrolyte (such as organic substances or alcohol).
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Libretto - Chem Review
a. A combination of ions and concentration determines how much the f.p. decreases and the
b.p increases.
Ex: Freezing Point Highest to Lowest: 1M C6H12O6, 2M C6H12O6, 1M Mg(NO3)2, 2M Mg(NO3)2;
Boiling Point Lowest to Highest: 1M C6H12O6, 2M C6H12O6, 1M Mg(NO3)2, 2M Mg(NO3)2.
Reason: Mg(NO3)2 breaks un into three ions: Mg2+ and 2 NO3-. **Count subscripts of ions and
multiply by molarity. **C6H12O6 does not break into ions (covalently bonded – other examples: CH4,
O2, C6H8O7, etc.)
grams of solute
16. Parts Per Million (ppm) = on reference Table T. ppm = grams of solution x 1,000,000.
(Grams of solution = grams of solute + grams of solvent).
Ex: Approximately 0.0043g of oxygen can be dissolved in 100 g of water. Express the concentration
in ppm.
0.0043𝑔
Ans: ppm = 100.0043𝑔x 1,000,000 = 43ppm
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Libretto - Chem Review
Topic #11 – Acids, Bases, Salts
1. Electrolytes: A compound when melted, vaporized, or dissolved in H2O will conduct electricity
(form ions). Ex: Acids, Bases, Salts.
2. Non-electrolytes: will not conduct electricity under the above conditions (does not form ions). Ex
– Organic Compounds (Except CH3COOH)
3. Strong Electrolyte: strong acids, soluble bases and salts (See Table F for solubility – nitrates vs.
carbonates).
4. Acids: begin with H. Ex: HCl, HC2H3O2, H2SO4, CH3COOH (acetic acid is an exception), See Ref.
Table K for a list. Form H+ (same thing as H3O+) ions.
Bases: end on –OH. Ex: NaOH, NH4OH, Ca(OH)2, etc. All bases dissociate except for NH3(aq)
which ionizes. Bases increase the [OH-].
Salts: ionic compounds that do not begin with H or end in –OH. Ex: NaCl, MgS, MgCO3
Organic: begin with C. Ex: C6H12O6 etc. Exceptions include CH3COOH (acid). Alcohols are a type
of organic compounds that start with C and end in OH (like CH3OH).
5. Properties of Acids:
a. Turn blue litmus red (colorless in phenolphthalein)
b. Have a pH less than 7
c. React with metals above H2 on Table J to form salt and
H2 gas
d. React with bases to form a salt and water
(neutralization).
e. Taste sour
f. Conduct electricity in relation to the degree of their
ionization (more soluble = more ions = better conductor).
g. Acidic solutions contain more H+ (H3O+) ions than OHions.
6. Properties of Bases:
a. Turn red litmus blue, pink in phenolphthalein
b. Have a pH greater than 7
c. React with acids to form a salt and water (neutralization)
d. Taste bitter, and feel slippery
e. Conduct electricity in relation to their solubility (more soluble = more ions = better
conductor).
f. Basic solution contain more OH- than H+.
7. pH scale: Compares the [H3O+] to [OH-]. H3O+ is called the hydronium ion and OH- is called the
hydroxide ion.
[H3O+] x [OH-] = 10-14 (a constant.)
Use this expression to find either [H If the [H3O+] increases, the [OH-] decreases, a solution
becomes acidic and pH < 7. If the [H3O+] decreases, the [OH-] increases, a solution becomes
basic and pH > 7. In water the [H3O+] = [OH-], and the pH = 7 which is neutral.
Adding an acid decreases the pH, adding a base increases the pH.
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Libretto - Chem Review
8. Neutralization:
a. Acid + Base  salt + water
b. H+ + OH-  H2O
or
+
c. H3O + OH  H2O
Ex: 2NaOH + H2SO4 Na2SO4 + 2H2O or HCl + NaOH  NaCl + H2O
9. In neutralization reactions for every mole of H+ that reacts, one mole of OH- reacts. Ex: If 2 moles
of H+ ions are neutralized, how many moles of OH- ions are needed? Ans. = 2 moles
10. Formula for titration of acids and bases: when different concentrations of an acid and base are
mixed, we can use the titration formula to find an unknown (a version of the formula is also found
on Table T, just add #H+ and #OH- to it):
(#H+) MA VA = MB VB (#OH-)
MA = molarity of the acid; VA = volume of the acid; #H = number of H+ in acid formula
MB = molarity of the base; VB = volume of the base; #OH = number of OH- in base formula
Example: What volume of 2.0M NaOH is needed to neutralize 30. mL of 4.0M H2SO4?
Answer: (2)(4.0M)(30.mL) = (2.0M)(VB)(1)
VB = 120mL
11. Definitions of Acids and Bases:
a. Arrhenius Theory: acids yield H+ ions, bases yield OH- ions. Stronger acids produce
more H+ ions than weaker acids. Strong bases produce more OH- ions than weak
bases.
b. Bronsted-Lowry Theory: acid = proton (H+ donor); base = proton acceptor.
Ex. NH4+ is a BL acid: NH4+ + H2O  NH3 + H+ (NH4+ donates proton to H2O)
Ex. NH3 is a BL base: NH3 + H2O  NH4+ + OH- (NH3 accepts proton from H2O)
Amphoteric – a substance that can either accept or donate H+ depending on the
conditions. Ex: H2O can become either H3O+ or OH11. The larger the [H+], the stronger the acid, the lower the pH, and the better the conductivity.
12. Strongest bases have Group 1 with OH Ex – NaOH, KOH, etc.
13. Acids react with any metal higher than H2 on Reference Table J spontaneously. Ex: Cu will not
react with an acid, but Mg will.
Metal + Acid  Salt + Hydrogen gas
Mg + HCl  MgCl2 + H2
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Libretto - Chem Review
Topic #12 – Oxidation-Reduction
1. Know rules for determining oxidation #’s (oxidation numbers are found on periodic table.
2. The sum of the oxidation numbers in a compound is zero. Elements by themselves have an
oxidation state of zero. Ex: Na0, Mg0, Cl20.
Ex: Find the oxidation number of S in Na2S2O3. Ans: Na2+1S2+4O3-2.
3. Oxidation= a loss of electrons. Reduction= a gain of electrons. Remember:
LEO the lion goes GER = “Losing Electrons is Oxidation” and “Gaining Electrons is Reduction” and
OIL RIG = “Oxidation is Losing electrons” and “Reduction is Gaining electrons”.
4. Oxidation Number:
a. Oxidation = increasing oxidation number
b. Reduction = decrease in oxidation number.
Ex: Mg0 + Zn0Cl2+2 Mg+2Cl2-1 + Zn0.
Mg0 (0+2) is oxidized and the reducing agent
Zn+2 (+20) is reduced and the oxidizing agent.
5. In a redox reaction, there must be a change in oxidation numbers. In a double replacement there
is no redox, there is always a redox reaction in a single replacement reaction (change in oxidation
numbers).
6. As the number of particles oxidized during a reaction increases, the number of particles reduced
also increases.
7. Know the reaction:
Cu + 2 AgNO3  Cu(NO3)2 + 2 Ag
Net Reaction: Cu0 + 2 Ag+1  Cu+2 (blue color) + 2 Ag0
NO3- is the spectator ion in the above reaction (oxidation state did not change).
8. Writing half reactions:
a. Equations must be balanced according to both mass and charge.
9. The number of each element must be the same on both sides of the arrow, and the sum of the
charges must be same on both sides of the arrow. Ex:
10. Some substances (Ex: Sn+2) can act as both an oxidizing and reducing agent:
a. Sn+2Sn+4 + 2 eSn+2 is oxidized = reducing agents
b. Sn+2 + 2e- Sn0
Sn+2 is reduced = oxidizing agent
11. To find the number of moles of electrons transferred, find the number of moles of etransferred from oxidized to reduced part of equation.
Ex: Find the total number of moles of e- transferred from 2Al(s) to 3 Cu+2 in:
2 Al(s) = 3 Cu+2(aq)  2 Al+3 + 3 Cu(s)
Use Half Reaction:
Al0Al+3 + 3 etherefore
2 Al0  2Al+3 + 6e- which means 6 moles of etransferred
12. Reference Table J shows Metals from most to least likely oxidized and lose electrons;
and nonmetals from most to least likely reduced and gain electrons.
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Libretto - Chem Review
13. Electrochemical Cell (Uses chemical redox reaction to create
electricity):
a. Produces electricity by means of a spontaneous reaction
b. Be able to write half-reactions for the cell. The more active
metal (Closer to top of Table J) is oxidized, the less active
metal (closer to bottom of Table J) is reduced.
c. Electrons flow from the more active metal to the less active
metal.
d. Oxidation occurs at the negative electrode (anode)
e. Be able to write overall reaction for the cell and the moles
of electrons transferred
f. The salt bridge functions to allow the migration of ions. If
the salt bridge is removed the cell voltage = 0 volts.
“An Ox, Red Cat” = Anode is oxidized, cathode is reduced.”
14. Electrolytic Cell (Electricity added produces redox reaction)
a. Uses electricity to force a non-spontaneous reaction to
occur
b. Oxidation occurs at the positive electrode (anode)
c. Be able to write half reactions at each electrode
d. In electrolysis a compound is broken down into its free
elements.
Ex: Hoffman Apparatus- 2H2O 2H2 + O2.
“Jump start a car – needs a battery, Nephew, have you ever
seen a Red Cat? NEGATIVE.
Have you ever seen an ox? POSITIVEly.”
For Both types of cell: “An Ox, Red Cat” – anode is oxidized, cathode is reduced.
15. Predicting if a redox reaction will occur:
a. Use Table J: If the free element is higher than the first element in the compound, the
reaction will take place. If not it will not take place.
Ex: 2 Al +3 CuCl2  3 Cu + 2AlCl3; the reaction will take place since Al is higher up than Cu.
Sayings to Remember for Redox Reactions:
 “LEO the lion goes GER” – loss of electrons is oxidation
 “OIL RIG” - oxidation is losing electrons, reduction is gaining electrons
 “An Ox and Red Cat” – anode is oxidized, cathode is reduced.
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Libretto - Chem Review
Topic #13 – Organic Chemistry








propene




Anything organic has to have the element CARBON.
All carbon atoms need four bonds. Hydrogen and Halides (group 17) may combine with
Carbon and can form only 1 bond.
Properties of Organic Molecules:
a. Non-electrolytes (except acids “COOH”).
b. Low boiling and melting points (Ex – Sugar C6H12O6)
c. Compounds usually insoluble in water
d. Compounds react slowly
e. Are molecular in structure
Homologous Series – successive members differ by CH2 group
Hydrocarbon – contains both Carbon and Hydrogen only. Ex – hexane; propene; 2
heptyne.
Saturated = all single bonds; Unsaturated = at least one double or triple bond.
Table P gives organic prefixes, Table Q summarizes alkanes, alkenes, alkynes (“n” = #
of carbons).
Alkanes: CnH2n+2 - contain only single bonds (saturated). Ex: CH4, C2H6, C3H8, C4H10, etc.
Names end in “-ane” = methane, ethane, propane, butane…
Alkenes: CnH2n – contain one double bond (unsaturated); starts with ethene (no meth-).
Ex:C2H4, C3H6, C4H8, C5H10, etc. Names end in “ene” = ethene, propene, butene,
pentene…
Alkyne: CnH2n-2 – contain one triple bond (unsaturated). Starts with ethyne (no meth-).
Ex: C2H2, C3H4, C4H6, etc. Name ends in “-yne” = ethyne, propyne, butyne…
As the molecular mass of organic compounds in each series increases, the boiling point
increases due to Van der Waals forces.
Alkyl Radicals: formed from each alkane by removing H. Ex: methyl group CH3
Ethyl group - C2H5 .

Naming Organic Compounds:
a. find longest consecutive carbon chain
b. find closest side chain to first carbon
c. use prefix for more than one of same side chain with di-, tri-, tetra- etc.
d. End with suffix of parent. Ex –ane for alkane, -ene, for alkene, -yne for
alkyne.
Example: 3,4 dimethyl heptane
and 2,2 diethyl-4-methylpentane

Isomers: compounds that have the same molecular formulas but different structural
formulas (different arrangement of atoms.) For hydrocarbon molecules you need at least
four carbon atoms to have an isomer. The more carbons, the more possible isomers
(Isomers have different properties). Ex –
Pentane
Methyl Butane
Boiling Point = 36°C
Boiling Point = 30°C
2,2 Dimethyl Propane
Boiling Point
= 30°C
24
propane
ethyne
Libretto - Chem Review
15. Functional Groups: (Use Table R). Know how to draw out each example in Table R:
 Alcohols: R-OH. Monohydroxyl (1-OH), dihydroxyl (2-OH), trihydroxyl (3-OH)
a. Primary Alcohol: the “C” attached to the “OH” has 1 Carbon directly attached to it
(on the last “C”)
b. Secondary Alcohol: the “C” attached to the “OH” has 2
“C’s” directly attached to it.
c. Tertiary Alcohol: the “C” attached to the “OH” has 3 “C’s”
directly attached to it
Some Functional Group Sayings to help you remember them:


Organic Acids: they’re “COOH”.
Ketone: Think of a dog howling “O” straight up. (Means “O” in middle
has a double bond.

Ester:
“
” is she ugly, “O” is she ugly (to side means
single bond, up down is double bond.” She is a COOC
 Ether:
or
there is always a COC
 Aldehyde: “Al the cheerleader saying the H and O are “hyding” at the
end. Al dates a lot of women at the end –HO”.
 Amine: Nitrogen but No Oxygen.
 Amide: Nitrogen Does have Oxygen
16. Organic Reactions
 Addition: involves unsaturated hydrocarbons
Ex: C2H4 + Cl2  C2H4Cl2. Alkenes are unsaturated, and just like in math,
one answer when you add.
 Substitutions: involves saturated hydrocarbons. (S in Saturated, S in substitution).
Ex: C2H6 + Cl2  C2H5Cl + HCl. Alkanes are saturated, and there are two products. Like
sports, one thing is switched with another.




Esterification: acid + alcohol  ester + H2O.
Saponification: can’t spell saponification w/o SOAP. Fat + Base  soap + glycerol.
Fermentation: C6H12O6  2C2H5OH + CO2.
Polymerization: Small chains put together to form long chains. This forms natural
compounds such as nylon, rayon, starches, cellulose; and artificial compounds such as
Teflon, polystyrene, etc.
a. Addition Polymerization: joining monomers (small chains) of unsaturated
compounds by opening n(C2H4)  (C2H4)n. Chains linked together.
b. Condensation Polymerization: long chains formed and water as a byproduct.
“Dehydration synthesis”.
25