Download Bucher Regents Review

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Mr. Bucher’s Regents Chemistry Crib Sheet
Topic 1: Atomic Concepts
1.
The modern model of the atom has evolved over a long period of time through
the work of many scientists.
Know the gold foil experiment; especially the two conclusions Rutherford
made: (+) charged nucleus, atom is mostly empty space.
The law of conservation of mass (Lavoisier) often appears in equation and ½
reaction problems.
2.
Each atom has a nucleus, with an overall positive charge, surrounded by
negatively charged electrons.
Nuclear particles (nucleons) and electron notation can be found on Table O.
3.
Subatomic particles in the nucleus include protons and neutrons.
4.
The proton is positively charged and the neutron has no charge. The electron
is negatively charged.
5.
Protons and electrons have equal but opposite charges.
6.
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 a
neutron.
A common question: which particle is most/least massive: alpha, beta, proton,
neutron. (Table O)
7.
In the wave-mechanical (electron cloud) model, the electrons are in orbitals,
which are regions of most probable electron location (ground state)
Know the difference between orbits and orbitals.
8.
Each electron in an atom has its own distinct amount of energy.
9.
After an electron in an atom gains a specific amount of energy, the electron is
at a higher energy level (excited state)
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Be able to distinguish between ground state electron configurations (periodic
table) and excited state: Mg 2-8-2 ground state, 2-7-1-1, 2-8-1-1, etc. excited
state.
10.
When an electron returns from a higher energy state to a lower energy state,
energy is emitted. This emitted energy corresponds to a specific wavelength in
the electromagnetic spectrum.
11.
Wavelengths can be used to identify a substance. Each kind of atom or
molecule can gain or lose energy in discrete amounts, and thus can absorb or
emit energy only at wavelengths corresponding to these amounts.
Be able to distinguish elements by their emission spectrum (bar code
questions)
12.
In general, the outermost electrons in an atom are called valence electrons.
The number of valence electrons determines the chemical properties of an
element. The chemical reactivity of an atom is dependent on its size.
The group numbers of the representative elements (s and p block) will tell you
the number of valence electrons. For groups 13-17, the second number is the
number of valence electrons.
The most active metals are the largest metals (bottom left). The most active
non-metals are the smallest (top right)
13.
Atoms of an element that contain the same number of protons but a different
number of neutrons are called isotopes of that element.
Be able to determine the number of neutrons in an isotope (nuclide) of an
element.
Note that the mass number of a nuclide (example, 12C) is not the same as the
atomic mass found on the periodic table (carbon =12.011 amu). Do not look
on the periodic table for mass numbers!!
14.
The average atomic mass of an element is the weighted average of the masses
of its naturally occurring isotopes.
Be able to set up an atomic mass problem on the free response section. You
take the mass of the isotope (in amu, not the mass number), multiply it by the
abundance, sum the products, and divide by 100.
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Topic 2: The Periodic Table
1. The placement or location of elements on the periodic table gives an indication of
physical and chemical properties of that element. The elements on the periodic
table are arranged in order of increasing atomic number.
A few elements have smaller atomic masses than the elements that precede them
(K, Ni, I). Be alert for a question that compares atomic masses of consecutive
elements.
2. The number of protons in an atom(atomic number) identifies the element. The sum
of the protons and neutrons in an atom (mass number) identifies an isotope.
Examples of common notations that represent isotopes include: 614C, 14C, carbon14, C-14.
The phrase “greatest nuclear charge” is often used instead of atomic number.
Remember, mass number ≠ atomic mass!
3. Elements can be classified as metals , nonmetals, metalloids, and noble gases.
Know all the ways to distinguish between metals and non-metals.
Metalloids rarely appear on the regents. You need to be able to identify them (on
the step ladder) and know that they have properties of metals and non-metals.
4. Elements can be differentiated by their physical properties. Physical properties of
substances, such as density, conductivity, malleability, solubility, and hardness,
differ between elements.
Metals: very dense, conductors, malleable, usually hard
Non-metals: not very dense (often gases), poor conductors, soft or brittle solids
5. Elements can be differentiated by chemical properties. Chemical properties
describe how an element behaves during a chemical reaction.
Metals lose electrons (+ oxidation #s) Non-metals gain electrons (- oxidation #s)
6. Some elements exist as allotropes. Allotropes are two or more forms of the same
element that differ in their molecular or crystalline structure, and hence in their
properties.
The two examples they use are diamond/graphite for carbon, and ozone and O2.
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7.
For groups 1,2 and 13-18, elements within the same group have the same number
of valence electrons (helium is an exception) and therefore similar reactivity.
They often ask questions about elements with similar properties. The answer
usually has two elements in the same group.
Be able to draw Lewis diagrams for any element (# dots = # valence electrons)
8. The succession of elements within the same group demonstrates characteristic
trends,e.g., differences in atomic radius, ionic radius, electronegativity, first
ionization energy. These trends can be explained in terms of atomic
structure(size).
Atomic radius increases down the group (larger size)
Ionic radius also increases down the group (larger size)
Electronegativity decreases down the group (elements become more metallic)
Know the definition for electronegativity!!
Ionization energy decreases down the group (valence electrons are shielded from
the nucleus by lower energy levels)
9. The succession of elements across the period demonstrates characteristic trends,
e.g., differences in atomic radius, ionic radius, electronegativity, first ionization
energy. These trends can be explained in terms of atomic structure (nuclear
charge).
Atomic radius decreases across the period (more effective nuclear charge).
Electrons from lower periods can “block” the nucleus from the valence electrons.
The number of blockers does not change across the period, but the nuclear charge
does!
Ionic radius depends on # valence electrons:
Non-metal ions are negatively charged, and are larger than their atomic
radii. Metal ions are positively charged, and are smaller than their atomic
radii. This is a common question on the exam!!!
Electronegativity increases across the period. Non-metals attract electrons more
than metals.
1st ionization energy generally increases across the period. More effective nuclear
charge makes for smaller radii; it takes more energy to remove and electron.
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Topic 3: Moles/Stoichiometry
1. A compound is a substance composed of two or more different elements that are
chemically combined in a fixed proportion. A chemical compound can be broken
down by chemical means. A chemical compound can be represented by a specific
chemical formula and assigned a name based on the IUPAC system.
There is usually one particle diagram question on each exam. Black and white
circles are used to represent elements. They are used to distinguish elements from
compounds and simulate chemical reactions. Look at your copies of regents
exams for examples.
There is always one naming question on each exam. They often use polyvalent
metals (Fe, Pb, Cu, Sn, etc) in ionic compounds. The non-metal oxidation number
(always the top one on the periodic table) will allow you to determine the charge
on the metal. Beware of endings on the anion: -ide endings generally mean binary
ionic compounds; -ate or –ite endings indicate polyatomic ions (Table E)
2. A chemical formula can be represented as an empirical formula, a structural
formula, or a molecular formula.
Structural formulas are for organic compounds only. They are needed because of
the existence of isomers for larger molecules. Remember, a dash represents a pair
of shared electrons.
3.
The empirical formula of a compound is the simplest whole-number ratio of
atoms of the elements in a compound. It may be different than the molecular
formula, which is the actual ratio of atoms in a molecule of that compound.
Sometimes they will give you molar mass of a compound and its empirical
formula. You must find the “empirical formula mass” and divide the molar mass
by the efm.
4.
The molecular formula is a whole-number multiple of the empirical formula.
5.
In all reactions there is a conservation of mass, energy, and charge.
They can ask questions like “Which equation does/does not show conservation of
mass?” This can mean balanced or unbalanced chemical equations, half-reactions
or faulty half-reactions, net ionic equations (usually in electrochemistry), which
require you to balance electrons.
6.
A balanced chemical equation represents conservation of atoms. The coefficients
in a balanced chemical equation can be used to determine mole ratios in a
reaction.
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It is a certainty that you will have to balance an equation in at least one question
on the exam.
7. The formula mass of a substance is the sum of the masses of its atoms. The gramformula mass of the substance equals 1 mole of that substance.
Remember, the atomic mass of everything in a parenthesis in a chemical formula
must be multiplied by its subscript when calculating its gram formula mass.
Coefficients in chemical reactions are not included in molar mass calculations!
Table T (back page) has the formula for mole calculations.
8. The percent composition by mass of each element in a compound can be
calculated mathematically.
Table T has this formula also
9. There are many types of chemical reactions, e.g., synthesis, decomposition, single
replacement, and double replacement.
All of these reactions except double replacement (and acid-base neutralization)
are redox reactions.
There are also a number of organic reactions you must know how to recognize:
-
-
substitution (by halogens) of saturated hydrocarbons
addition (by hydrogen and halogens) of unsaturated hydrocarbons
esterification (condensation): alcohol + organic acid  ester + H2O
combustion: organic compound + O2  CO2 + H2O
fermentation: C6H12O6  2CH3CH2OH + 2 CO2
saponification: fatty acids + base  soap + glycerol
There is also addition polymerization and condensation polymerization.
These are all redox reactions. Combustion reactions are favorite balancing
problems, because they can have large coefficients.
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Topic 4: Chemical Bonding
1. Compounds differ in composition as well as chemical and physical properties.
Two major categories of compounds are ionic compounds and molecular
(covalent) compounds.
Very Important!
Ionic compounds:
metal + non-metal or polyatomic ion
Covalent compounds: all non-metals
You must be able to discern whether an element is a metal (left and lower parts of
the table) or whether it is a non-metal or polyatomic ion (upper right part of the table)
2.
Chemical bonds are formed when valence electrons are:
 Transferred from one atom to another (ionic)
A metal from the left and a non-metal from the right.
 Shared between atoms (covalent)
Can be polar (unequal sharing), non-polar (equal sharing)
 Mobile within a metal (metallic)
The phrase “mobile valence electrons” is often used.
3. In a multiple covalent bond, more than one pair of electrons is shared between
two atoms. Unsaturated organic compounds contain at least one double or triple
bond.
Single, double, or triple bonds – each bond is a shared pair of electrons
4. Molecular polarity can be determined by the shape and the distribution of charge.
Examples of symmetrical (nonpolar) molecules include CO2, CH4, and the
diatomic elements.
Examples of asymmetrical (polar) molecules include HCl, NH3 and H2O.
These examples are used year after year. Symmetrical molecules with polar bonds
(the C=O double bond is very polar) are non-polar.
You should be able to draw or recognize Lewis structures for all of the molecules
above.
5. When an atom gains one or more electrons, it becomes a negative ion and its
radius increases. When an atom loses one or more electrons, it becomes a
positive ion and its radius decreases.
This fact appears nearly every year in a question or answer.
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6. When a bond is broken, energy is absorbed. When a bond is formed, energy is
released.
Absorbing energy – endothermic
Releasing energy – exothermic
7.
ΔH is (+)
ΔH is (-)
Atoms tend to bond so that a stable valence electron configuration, like that of a
noble gas, is achieved.
The octet rule governs much of chemistry. Sometimes they will ask about
isoelectronic atoms and ions; they have the same number of electrons; for
example, O2- = Ne = Mg2+ = 10 electrons (eight in the second energy level)
8. Electron dot diagrams (lewis structures) can represent the valence electron
arrangement in elements and compounds.
You must be able to distinguish ionic and covalent compounds to write Lewis
structures successfully! Ionic compounds have brackets and charges. Covalent
compounds have shared electron pairs.
..
NaCl right: [Na]+ [: Cl : ]1wrong: Na:Cl :
HCl
right:
H:Cl:
wrong: [H]+ [: Cl : ]1-
9. Electronegativity indicates how strongly an atom of an element attracts electrons
in a chemical bond. Electronegativity values are assigned according to arbitrary
scales.
The definition of electronegativity is a common test question. Sometimes they do
not use the word “electronegativity” in the question at all – you have to recognize
that they are talking about it!.
Electronegativity values for the elements are found in Table S.
10. The electronegativity difference between two bonded atoms is used to assess the
degree of polarity in the bond.
They no longer ask students to calculate electronegativity differences. Little or no
difference: non-polar bond; large difference: polar bond
The electronegativity of O and N are responsible for the different properties of
organic functional groups (Table R). All are polar to some extent.
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Topic 5: Physical Behavior of Matter
1.
Matter is classified as a substance or a mixture of substances.
2.
The three phases of matter, i.e., solids, liquids, and gases, have different
properties.
Drawing and interpreting particle diagrams are favorite activities here.
3.
A substance (element or compound) has a constant composition and constant
properties throughout a given sample, and from sample to sample.
Distinguishing mixtures from pure substances is a common question.
4.
Elements are substances that are composed of atoms that have the same atomic
number. Elements cannot be broken down by chemical change.
Both of these points are tested regularly.
5.
A mixture is not a substance because it is made up of two or more different
elements and/or compounds. The proportions of components in a mixture can be
varied. Each component in a mixture retains its original properties.
6.
Differences in physical properties such as mass, particles size, molecular
polarity, boiling point, and solubility permit physical separation of the
components of the mixture.
Sometimes they ask about separating mixtures. Separations exploit differences in
physical properties: solubility (filtration), boiling point (distillation)
7.
A solution is a homogenous mixture of a solute dissolved in a solvent. The
solubility of a solute in a given amount of solvent is dependent on the
temperature, pressure, and the chemical natures of the solute and solvent. The
concentration of a solution in expressed in molarity (M), percent by volume,
percent by mass, or part per million (ppm).
The formulas for molarity and ppm are found on Table T. Molarity is a hard
calculation at times because it contains moles in the numerator, which is another
calculation. A molarity question is a certainty; ppm questions occur occasionally.
8.
The addition of a nonvolatile solute to a solvent causes the boiling point of the
solvent to increase and the freezing point of the solvent to decrease. The greater
the concentration of particles, the greater the effect.
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For a given number of moles of a substance, the ionic compounds have the
biggest effect on boiling point and freezing point because they dissociate into
ions. The more moles of ions you get, the bigger the effect.
Example:
CH3OH:
one mole of methanol molecules
NaOH:
one mole of Na+ and 1 mole of OH- ions
Ba(OH)2:
one mole of Ba2+ ions and 2 moles of OH- ions.
The barium hydroxide has the biggest effect for a given concentration of solute.
9.
Energy can exist in different forms, e.g., chemical, light, heat, nuclear.
Energy conversion questions are common: chemical potential energy to kinetic
energy (exothermic rxns); chemical potential energy to electrical (voltaic cells);
electrical to chemical potential energy (electrolytic cells) Nuclear energy to
kinetic energy (fission and fusion)
10.
Heat is a transfer of energy (usually thermal energy) from a body of higher
temperature to a body of lower temperature. Thermal energy is the energy
associated with the random motion of atoms and molecules.
Thermal energy = kinetic energy.
11.
Temperature is the measure of the average kinetic energy of the particles in a
sample of material. Temperature is not a form of energy.
The bolded sentence above is tested somewhere on every single exam!!
12.
Kinetic molecular theory (KMT) for an ideal gas states:
a) All particles are in random, constant, straight-line motion.
b) Gas molecules are separated by great distances relative to their size; the
volume of the gas molecules is considered negligible.
c) The molecules have no attractive forces between them.
d) Collisions between gas molecules may result in the transfer of energy between
gas particles, but the total energy of the system remains constant.
A consequence of statement b) is that equal volumes of different gases at the same
temperature and pressure have the same number of molecules.
Real gases have significant attractive forces between them and/or are not
separated by great distances relative to their size. High pressures and low
temperatures will create these conditions for gases. Under conditions of life on
earth, H2O behaves as a real gas, while N2 and O2 behave ideally.
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13.
Particles are in constant motion except at absolute zero (zero Kelvin).
Kinetic molecular theory describes the relationships of pressure, volume,
temperature, velocity, and frequency and force of collisions.
You need to be able to use the ideal gas equation (Table T). Any variable that is a
constant can be dropped from the equation. Free response questions usually only
require setting up the equation, not solving for it.
Remember, Kelvin temperature must be used for all gas law calculations!!!
The formula for Celsius – Kelvin conversion is also on Table T.
14.
The concepts of kinetic and potential energy can be used to explain physical
properties that include: fusion (melting), solidification (freezing), vaporization
(boiling, evaporation), condensation, sublimation, and deposition.
Sublimation questions are common, deposition is rare.
Most questions are variations on heating and cooling curves. The horizontal lines
are the phase changes. Be sure you know which one (melting-freezing or boilingcondensing) is which on the diagram. Review questions in your review book on
this topic – it appears every year.
At the melting point, solid and liquid phases are in equilibrium with each other.
At the normal boiling point, liquid and gas phases are in equilibrium.
15.
A physical change results in the rearrangement of existing particles in a
substance. A chemical change results in the formation of different particles with
changed properties.
Particle diagrams are commonly used on the regents to demonstrate phase
changes.
16.
Chemical and physical changes can be exothermic or endothermic.
Table I shows heats of reaction (ΔH) for a number of combustion reactions, a
bunch of synthesis reactions (often called the heat of formation), and some
solubility changes (a physical change). There is usually one question using this
table on every test. Often they ask for which reaction is endo/exothermic. Three
have to have the opposite sign to the correct answer.
Physical changes of water are given special attention. Information can be found
on Table B. You should know when and how to use the “heat” equations on Table
T. For phase changes, q=mHf and mHv. For liquid water temperature increases
and decreases, q=mcΔT is used.
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17.
The structure and arrangement of particles and their interactions determine the
physical state of a substance at a given temperature and pressure.
18.
Intermolecular forces created by the unequal distribution of electrons result in
varying degrees of attraction between molecules. Hydrogen bonding is an
example of a strong intermolecular force.
Small molecules that have equal or symmetrical charge distributions are nonpolar, and are gases to very low temperatures and pressures (that is, they behave
ideally under most conditions).
Larger molecules, even those with symmetrical charge distributions, have larger
attractive forces and higher boiling points and melting points. The boiling point
(distillation) is used to separate crude oil (a mixture of hydrocarbons) into
fractions based on molecular size.
Unequal distribution of charge is a euphemism for a polar molecule. These
molecules have dipoles [partial (+) and (–) parts of the molecule]. For similar
molecules, the larger the dipole, the higher the melting point and boiling point are.
Hydrogen “bonding” is the strongest dipole force. Only N, O, and F form
hydrogen bonds. Although it may seem obvious, hydrogen bonds are only
between these elements and hydrogen: N – H , O – H, and H – F. Remember,
hydrogen bonding is not a true “bond”; it has 5% of the strength of a covalent
bond.
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Topic 6: Kinetics and Equilibrium
The stability of a compound is dependent on the amount of energy absorbed or released
during the formation of the compound from its elements.
If energy is absorbed – endothermic; the heat term is a reactant
If energy is released – exothermic; the heat term is a product
Collision theory states that a reaction is most likely to occur if reactant particles collide
with the proper energy and orientation.
We call these “effective collisions.” More effective collisions means higher reaction
rates.
The rate of a chemical reaction is the change in concentration of a reactant or product
per unit time.
They only test this in a qualitative way; the rate increases or decreases depending on
reaction conditions.
The rate of a chemical reaction depends on several factors: temperature,
concentration, nature of reactants, surface area, and the presence of a catalyst.
There are always 1-2 questions on every exam on these factors:
Temperature: increased kinetic energy means more collisions.
Concentration: more molecules per unit volume means higher reaction rates.
Surface area: smaller particles, more surface area, higher reaction rates
Catalysts: lower activation energy, making effective collisions more likely.
Some chemical and physical changes can reach equilibrium.
Physical changes: phase changes, solubility, vapor pressure of liquids
Chemical changes: most reactions are reversible; we can speak about the forward and
reverse directions.
At equilibrium the rate of the forward reaction equals the rate of the reverse reaction
while the measurable quantities of reactants and products remain constant.
This definition appears somewhere nearly every year.
LeChatelier’s principle can be used to predict the effects of stress (change in pressure,
volume, concentration, and temperature) on a system at equilibrium.
Increased pressure: favors direction with the least number of moles of gas. If equal moles
of gas are found on the product and reactant side, there is no effect.
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Pressure only affects gases; solutions are unaffected.
Concentration: increasing a reactant or product concentration creates conditions that lead
to higher concentrations in the opposite direction; add reactant, more products, add
product, more reactants.
Temperature: heat can be a reactant (endothermic) or a product (exothermic) The
endothermic direction is favored when the temperature is increased. For an exothermic
reaction, the reverse direction would be favored.
Energy released or absorbed during a chemical reaction is equal to the difference
between the potential energy of the products and the potential energy of the reactants.
This is ΔH, the heat (enthalpy) of reaction. You need to know how ΔH is represented on
an energy diagram. It is tested nearly every year, and many students don’t get it.
Usually there is one problem using Table I every year. Often it is merely an
endothermic/exothermic question, although there have been a few questions where they
ask to calculate the number of kJ of energy absorbed/released for a given number of
moles of reactant used/product formed.
A catalyst provides an alternate reaction mechanism, which has lower activation energy
than an uncatalyzed reaction.
Catalysts have no effect on equilibrium; only the initial rates of reaction (in both
directions) are affected.
Be able to discern the activation energy on a potential energy diagram. Often they like
you to draw a line for a catalyzed reaction. Be sure that your line begins with the
potential energy of the reactants and ends with the potential energy of the production
Entropy is a measure of the randomness or disorder of a system. A system with greater
disorder has greater entropy.
The following processes have increased entropy:
- Solutes dissolving in solvents
- Phase changes that occur with increasing temperature (sublimation,
melting, vaporization). They often use equations to express this:
H2O(s)  H2O(g)
- Reactions that produce more moles of gas in the product side
Systems in nature tend to undergo changes in energy and entropy.
Reactions that release energy (exothermic) and have increased disorder (positive
entropy) are always favored.
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Topic 7: Organic Chemistry
Organic compounds contain carbon atoms that are bonded to one another in chains,
rings, and networks to form a variety of structures (polymers, oils, and other large
molecules).
Carbon has four valence electrons and always makes four covalent bonds with other
atoms.
Functional groups impart distinctive physical and chemical properties to organic
compounds.
The functional groups contain electronegative elements (halogens, O, and N). The more
polar compounds have higher melting and boiling points than their hydrocarbon
counterparts.
Hydrocarbons, organic acids, alcohols, esters, amines, amides, and amino acids are
categories of organic molecules that differ in their structural formulae as a result of
different functional groups.
You need to be able to use Tables P, Q and R with ease to get the organic questions.
Identifying functional groups, naming compounds, recognizing saturated and unsaturated
hydrocarbons, especially in reactions; these are skills you need to have.
Note that you will see full structural formulas and condensed structural formulas. Table R
has partially condensed formulas. You need to be able to discern between these different
ways of describing molecules.
O
||
Table R:
R-C-O-H
often appears as –COOH on the test!
O
||
R-C-O-R’
often appears as –COO-
Esters and acids always have two oxygens; the rest have one.
Aldehydes: -CHO
Ketones: R-CO-R’
Ethers: R-O-R’ there will always be two - CH2 - groups on each side of the oxygen
Hydrocarbons, organic acids, alcohols, and esters are names using the IUPAC system.
The IUPAC system provides a method of distinguishing among isomers of organic
compounds.
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Isomers have the same molecular formula, empirical formula, and % composition. They
test this definition a lot. If you’re not sure about an isomer, count the carbons, hydrogens,
and oxygens; make sure they’re the same as the molecule in question.
Unsaturated organic compounds contain at least one double or triple bond.
Each bond is a pair of shared electrons. How many electrons (or pairs) in double and
triple bonds?
You need to be able to recognize a alkane, alkene, or an alkyne by its chemical formula
(Table Q). Know how to use the general formulas, and know. They often ask this
question as part of a substitution or addition reaction question.
Addition, hydrogenation, substitution, polymerization, esterification, fermentation,
saponification, oxidation, and combustion are examples of organic reactions.
Addition is probably the most popular reaction they test with, followed by esterification.
Addition, hydrogenation, substitution, polymerization, esterification, fermentation,
saponification, oxidation, and combustion are examples of organic reactions.
Here’s that list again:
-
-
substitution (by halogens) of saturated hydrocarbons
addition (by hydrogen and halogens) of unsaturated hydrocarbons
esterification (condensation): alcohol + organic acid  ester + H2O
combustion: organic compound + O2  CO2 + H2O
fermentation: C6H12O6 (glucose) 2CH3CH2OH (ethanol) + 2 CO2
saponification: fatty acids + base  soap + glycerol
addition polymerization: usually polyethylene (ethene)
condensation polymerization: esters, proteins, polysaccharides, fats
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Topic 8: Oxidation-Reduction
1.
An oxidation-reduction reaction involves the transfer of electrons (e-).
Be able to recognize oxidation/reduction reactions under the following
circumstances:
In chemical equations; everything except double replacement and acid base are
redox reactions. If you’re not sure, look for changes in oxidation numbers of
reactants and products.
In half reactions, oxidation always has electrons on the product side, reduction
has them on the reactant side.
In net ionic equations (which are always single replacement reactions on the
regents), the ion on the reactant side is being reduced and the metal is being
oxidized. The metal is always more active than the ion (above the ion on Table J)
2.
Reduction is the gain of electrons.
Reduction results in a reduction in the oxidation number (becomes less positive,
more negative).
3.
A half-reaction can be written to represent reduction.
Generically, Mn+ + ne-  M
Unbalanced half-reactions (oxidation and reduction) are often wrong answers on
regents questions.
4.
Oxidation is a loss of electrons. Historically, oxidation was explained as the
chemical combination with oxygen.
Oxidation raises the oxidation number of the element (more positive); reduction
lowers the oxidation number (more negative)
5.
A half-reaction can be written to represent oxidation.
M  Mn+ + ne-
6.
In a redox reaction, there is a conservation of charge and mass. The number of
electrons lost is equal to the number of electrons gained.
There is usually one question on the law of conservation of mass (and charge, and
energy) on every exam.
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7.
Oxidation numbers can be assigned to atoms and ions. Changes in oxidation
numbers indicate oxidation and reduction.
Elements always have an oxidation number of zero. In spontaneous reactions, the
more active metal (higher on Table J) will be oxidized, and get a (+) oxidation
number (usually) based on the number of valence electrons.
8.
Corrosion of metals, combustion of fuels, and spoilage of foods are examples of
redox reactions.
The most common corrosion reaction involves adding acid (H+) to active metals,
producing hydrogen gas. Often they will ask you for the metal that doesn’t react:
it must be Cu, Ag, or Au, which are below H2 on the activity series.
Combustion reactions are usually used to balance equations. Try balancing the
combustion of methane, ethane, and propane:
Hydrocarbon + O2  CO2 + H2O
9.
In an electrochemical cell, oxidation occurs at the anode and reduction at the
cathode.
The more active metal is the anode (higher on Table J); the less active metal
(lower on Table J) is the cathode. The ion of that metal will be reduced.
When the electrolyte in the reduction half-cell is used up, the battery “dies.”
Electrons flow from the anode to the cathode. They ask questions about this
nearly every year.
Be on the alert for the inevitable question about salt bridges. If they ask you what
they are for, write these two words: “mobile ions.”
10.
A voltaic cell can operate without an outside energy source.
At the heart of every battery is a spontaneous redox reaction. The electrons flow
from the more active metal, whose valence electrons are in higher energy states
than they will be when they reach the ion of the less active metal.
11.
Batteries are practical applications of voltaic cells.
They no longer ask specific questions about batteries.
12.
An electrolytic cell requires an outside energy source.
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The power supply makes a non-spontaneous process happen. It “pulls” electrons
from elements with low oxidation numbers and puts them on the element with the
higher oxidation number. For example, for the non-spontaneous process
2H2O  2H2 + O2
H goes from +1 to 0; O goes from –2 to O
Electrolytic cells often use two identical electrodes; there is also no separation
between the electrodes (no half-cells)
13.
Refining of metals and electroplating are practical applications of electrolytic
cells.
Group 1 and 2 metals are produced this way. Electrolysis of “fused” salts also
produce pure halogens like Cl2 and F2.
The anode is always where oxidation occurs, and the cathode is always where
reduction occurs.
an-ox
red-cat
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Topic 9: Acids and Bases
Behavior of many acids and bases can be explained by the Arrhenius theory. Arrhenius
acids and bases are electrolytes.
Acids and bases are often composed of non-metals. All other organic molecules are not
electrolytes.
An electrolyte is a substance which, when dissolved in water, forms a solution capable of
conducting an electric current. The ability of a solution to conduct an electric current
depends on the number of ions.
Note that CH3OH and CH3CH2OH are not electrolytes, while NaOH and KOH are. Be
clear on why this is so.
Arrhenius acids yield H3O+ (hydronium ion) as the only positive ion in an aqueous
solution.
The common acids are found on Table K. Remember, organic acids are written R-COOH.
H+ and H3O+ are used interchangeably to describe acids.
Arrhenius acids react with active metals to produce hydrogen gas.
All of the metals above H2 on Table J will react with acid. The metals below, Cu, Ag, and
Au are resistant to corrosion.
Arrhenius bases yield OH- (hydroxide ion) as the only negative ion in aqueous solution.
The common bases are found on Table L. All of them contain hydroxide ion except
ammonia. Ammonia occasionally steals a proton from water, creating hydroxide ion.
In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form a
salt and water.
Note that neutralization reactions are a special form of double replacement. They are not
redox reactions. Note that the complete neutralization of a strong acid by a strong base
will give a solution with pH 7.
Titration is a laboratory process in which a volume of solution of known concentrations
is used to determine the concentration of another solution.
Titration problems are very common on the free response section. Typically they will
give you a data table, occasionally it will be a word problem. Lately they have only been
asking you to set up the calculation, not finish it.
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The titration equation is found on Table T (back page). If you get a word problem,
carefully read the question and determine what your knowns are and unknown is. If you
get a table, you will have to do some subtraction to figure out the ml of acid or base used.
An aqueous solution of a salt conducts electricity. The solution can have a pH greater
than, equal to, or less that 7.
Neutralization reactions always produce salts and water. A “salt” in chemistry means any
ionic compound, not just NaCl.
The acidity or alkalinity of a solution can be measured by its pH value. The relative level
of acidity or alkalinity of a solution can be shown by using indicators.
Table M shows some common indicators and their color change region. They will usually
ask what the color of an indicator is at a specific pH.
The mathematical definition of pH is: pH= -log[H3O+]. On the pH scale, each decrease of
one unit of pH represents a ten-fold increase in hydronium concentration.
The 10-fold increase per unit is a common question:
1 unit difference
2 unit difference
3 unit difference
10x
100x
1000x
The questions always involve a dramatic change in the pH of a solution, say from 5 to 8
or from 5 to 3. They are usually accompanied by an indicator (Table M) question: “What
color will thymol blue be at pH 5?
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Topic 10: Nuclear Chemistry
Isotopes of atoms can be stable or unstable. Stability of isotopes is based on the number
of protons and neutrons in its nucleus. Some nuclei are unstable and spontaneously
decay, emitting radiation.
For small nuclides (atomic numbers 20 or less) equal numbers of protons and neutrons
are stable. For larger numbers, more neutrons are needed to insure stability.
Each radioactive isotope has a specific mode and rate of decay (half life).
There are 24 nuclides listed in Table N. Usually one of these appears on a question
somewhere.
There are four types of radioactive decay problems. The equation can be found at the
bottom of Table T. The key is to find the number of half lives. Figuring out the fraction
remaining will usually help you do that:
½ = 1 half life
¼ = 2 half lives
1/8 = 3 half lives
1/16 = 4 half-lives
1/32 = 5 half-lives
Remember these fractions. If they ask you:
“How much is left?” Divide the initial mass in half as many times as you have halflives.
“How long did it take?” They have to give you the fraction remaining. Once you figure
out the number of half lives, multiply that number by the half life (often on Table N)
Less commonly, they ask for the half-life of an unfamiliar radionuclide. They must
give you the fraction remaining, so you can figure out the # half-lives. Take the time
elapsed and divide it by that number.
“How much is did you start with?” The number of half life can be determined by
taking the time elapsed and dividing it by the half-life. Then take the amount remaining
and double it as many times as you have half-lives.
Spontaneous decay can involve the release of alpha particles, beta particles, or gamma
radiation from the nucleus of an unstable isotope. These emissions differ in mass, charge,
and penetrating power.
Be able to recognize and write the symbols for all of the particles and radiation emitted
from radioactive nuclides – they’re all found on Table O.
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Any change in the nucleus of an atom that converts it from one element to another in
called transmutation. This can occur naturally or can be induced by the bombardment of
the nucleus by high-energy particles.
Artificial transmutation equations have two particles on the left side of the equation. One
of them is either a neutron or an alpha particle. Natural transmutation has only one
particle, the radioactive nuclide.
Nuclear transitions can be represented by equations that include symbols that represent
atomic nuclei (with the mass number and atomic number), subatomic particles (with
mass number and charge) and/or emissions such as gamma radiation.
Be able to distinguish nuclear decay equations from chemical equations. You won’t see
mass numbers or atomic numbers in chemical equations.
Be able to find some component X in a transmutation equation. The mass numbers and
the atomic numbers on both sides of the equation have to add up.
For alpha decay, the mass number decreases by 4 and the atomic number decreases by 2.
For beta minus decay, the mass number doesn’t change and the atomic number increases
by one.
For positron decay (beta plus) the mass number is unchanged and the atomic number
decreases.
Energy released in a nuclear reaction (fission/fusion) comes from the fraction amount of
mass converted into energy. Nuclear changes convert matter into energy according to
E=mc2.
Fission: Two small atoms (usually hydrogen) make a larger one (usually helium)
Fusion: A large atom is split into two smaller atoms by a neutron. Several neutrons are
also produced.
The mass  energy question is frequently asked on the regents. The mass converted into
energy is small, but the energy produced is enormous.
Energy released during nuclear reactions is much greater than the energy released
during reactions.
Lately this fact has appeared in some questions.
There are inherent risks associated with radioactivity and its uses, such as long-term
storage and disposal of radioactive isotopes, and nuclear accidents.
Often there are give reading passages followed by questions on the free response section.
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Radioactive isotopes have many beneficial uses. Radioactive isotopes are used in
medicine and industrial chemistry, e.g., radioactive dating, tracing chemical and
biological processes, industrial measurement, detection and treatment of diseases.
This is another common reading assignment on the test. Medicial radioisotopes must
have short half-lives, for obvious reasons. Dating rocks requires isotopes with long halflives.