Transcript For metals

How To Prepare
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DO NOT CRAM. Get your studying done with by the
night before. Get a good night’s sleep and have breakfast
the morning of the exam.
Use a review book with old exams, answers and
explanations in it. Take the old tests and grade yourself.
The questions you don’t understand why you got wrong
make sure to see your teacher about.
Actively participate in any and all review classes and
activities offered by your teacher.
Study vocabulary. Identify key words and use flash cards to
help you remember what the meaning of those words are
and the concepts behind them.
(c) 2006, Mark Rosengarten
Outline for Review
1) The Atom (Nuclear, Electron Config)
2) Matter (Phases, Types, Changes)
3) Bonding (Periodic Table, Ionic, Covalent)
4) Compounds (Formulas, Reactions, IMAF’s)
5) Math of Chemistry (Formula Mass, Gas Laws,
Neutralization, etc.)
6) Kinetics and Thermodynamics (PE Diagrams, etc.)
7) Acids and Bases (pH, formulas, indicators, etc.)
8) Oxidation and Reduction (Half Reactions, Cells, etc.)
9) Organic Chemistry (Hydrocarbons, Families, Reactions)
(c) 2006, Mark Rosengarten
The Atom
1) Nucleons – click here for website on nucleons
2) Isotopes – click here for website on isotopes
3) Natural Radioactivity
4) Half-Life
5) Nuclear Power
6) Electron Configuration
7) Development of the Atomic Model
(c) 2006, Mark Rosengarten
Nucleons
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Protons: +1 each, determines identity of element, mass of 1
amu, determined using atomic number, nuclear charge
Neutrons: no charge, determines identity of isotope of an
element, 1 amu, determined using mass number - atomic
number (amu = atomic mass unit)
32 S and 33 S are both isotopes of S
16
16
S-32 has 16 protons and 16 neutrons
S-33 has 16 protons and 17 neutrons
All atoms of S have a nuclear charge of +16 due to the 16
protons.
website
(c) 2006, Mark Rosengarten
Isotopes
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Atoms of the same element MUST contain the same number
of protons.
Atoms of the same element can vary in their numbers of
neutrons, therefore many different atomic masses can exist for
any one element. These are called isotopes.
The atomic mass on the Periodic Table is the weight-average
atomic mass, taking into account the different isotope masses
and their relative abundance.
Rounding off the atomic mass on the Periodic Table will tell
you what the most common isotope of that element is.
(c) 2006, Mark Rosengarten
Weight-Average Atomic Mass
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WAM = ((% A of A/100) X Mass of A) + ((% A of B/100) X Mass of B) + …
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What is the WAM of an element if its isotope masses and
abundances are:
 X-200: Mass = 200.0 amu, % abundance = 20.0 %
 X-204: Mass = 204.0 amu, % abundance = 80.0%
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amu = atomic mass unit (1.66 × 10-27 kilograms/amu)
website
(c) 2006, Mark Rosengarten
Most Common Isotope
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The weight-average atomic mass of Zinc is
65.39 amu. What is the most common isotope
of Zinc? Zn-65!
What are the most common isotopes of:
 Co
Ag
 S
Pb
FACT: one atomic mass unit (1.66 × 10-27
kilograms) is defined as 1/12 of the mass of an
atom of C-12.
This method doesn’t always work, but it usually
does. Use it for the Regents exam.
(c) 2006, Mark Rosengarten
Natural Radioactivity
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Alpha Decay
Beta Decay
website
Positron Decay
Gamma Decay
Charges of Decay Particles
Natural decay starts with a parent
nuclide that ejects a decay particle to
form a daughter nuclide which is more
stable than the parent nuclide was.
(c) 2006, Mark Rosengarten
Alpha Decay
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The nucleus ejects two protons and two neutrons. The
atomic mass decreases by 4, the atomic number decreases
by 2.
238 U 
92
(c) 2006, Mark Rosengarten
Beta Decay
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A neutron decays into a proton and an electron. The
electron is ejected from the nucleus as a beta particle. The
atomic mass remains the same, but the atomic number
increases by 1.
14 C 
6
(c) 2006, Mark Rosengarten
Positron Decay
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A proton is converted into a neutron and a positron. The
positron is ejected by the nucleus. The mass remains the
same, but the atomic number decreases by 1.
53 Fe 
26
(c) 2006, Mark Rosengarten
Gamma Decay
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The nucleus has energy levels just like electrons, but the
involve a lot more energy. When the nucleus becomes
more stable, a gamma ray may be released. This is a
photon of high-energy light, and has no mass or charge.
The atomic mass and number do not change with gamma.
Gamma may occur by itself, or in conjunction with any
other decay type.
(c) 2006, Mark Rosengarten
Charges of Decay Particles
(c) 2006, Mark Rosengarten
Half-Life
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Half life is the time it takes for half of
the nuclei in a radioactive sample to
undergo decay.
Problem Types:
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Going forwards in time
Going backwards in time
Radioactive Dating
website
(c) 2006, Mark Rosengarten
Going Forwards in Time
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How many grams of a 10.0 gram sample of I-131 (half-life
of 8 days) will remain in 24 days?
#HL = t/T = 24/8 = 3
Cut 10.0g in half 3 times: 5.00, 2.50, 1.25g
(c) 2006, Mark Rosengarten
Going Backwards in Time
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How many grams of a 10.0 gram sample of I-131 (half-life
of 8 days) would there have been 24 days ago?
#HL = t/T = 24/8 = 3
Double 10.0g 3 times: 20.0, 40.0, 80.0 g
(c) 2006, Mark Rosengarten
Radioactive Dating
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A sample of an ancient scroll contains 50% of the original
steady-state concentration of C-14. How old is the scroll?
50% = 1 HL
1 HL X 5730 y/HL = 5730y
(c) 2006, Mark Rosengarten
Nuclear Power
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Artificial Transmutation
Particle Accelerators
Nuclear Fission
Nuclear Fusion
(c) 2006, Mark Rosengarten
Artificial Transmutation
 4020Ca
 9642Mo
+ _____ ----->
40 K
19
+ 11H
+ 21H -----> 10n + _____
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Nuclide + Bullet --> New Element + Fragment(s)
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The masses and atomic numbers must add
up to be the same on both sides of the arrow.
Website
(c) 2006, Mark Rosengarten
Particle Accelerators
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Devices that use electromagnetic fields to accelerate particle
“bullets” towards target nuclei to make artificial transmutation
possible!
Most of the elements from 93 on up (the “transuranium”
elements) were created using particle accelerators.
Particles with no charge cannot be accelerated by the charged
fields.
website
(c) 2006, Mark Rosengarten
Nuclear Fission
 9236Kr + 14156Ba + 3 10n + energy
The three neutrons given off can be reabsorbed by
other U-235 nuclei to continue fission as a chain
reaction
 23592U
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+
1 n
0
website
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A tiny bit of mass is lost (mass defect) and converted
into a huge amount of energy.
(c) 2006, Mark Rosengarten
Chain Reaction
(c) 2006, Mark Rosengarten
Nuclear Fusion
 21H
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+ 21H  42He + energy
Two small, positively-charged nuclei smash
together at high temperatures and pressures to
form one larger nucleus.
A small bit of mass is destroyed and converted
into a huge amount of energy, more than even
fission.
website
(c) 2006, Mark Rosengarten
Electron Configuration
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Basic Configuration
Valence Electrons
Electron-Dot (Lewis Dot) Diagrams
Excited vs. Ground State
What is Light?
(c) 2006, Mark Rosengarten
Basic Configuration
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The number of electrons is determined from the atomic
number.
Look up the basic configuration below the atomic number on
the periodic table. (PEL: principal energy level = shell)
He: 2 (2 e- in the 1st PEL)
Na: 2-8-1 (2 e- in the 1st PEL, 8 in the 2nd and 1 in the 3rd)
Br: 2-8-18-7 (2 e- in the 1st PEL, 8 in the 2nd, 18 in the 3rd
and 7 in the 4th)
(c) 2006, Mark Rosengarten
Valence Electrons
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The valence electrons are responsible for all chemical bonding.
The valence electrons are the electrons in the outermost PEL
(shell).
He: 2 (2 valence electrons)
Na: 2-8-1 (1 valence electron)
Br: 2-8-18-7 (7 valence electrons)
The maximum number of valence electrons an atom can have
is EIGHT, called a STABLE OCTET.
(c) 2006, Mark Rosengarten
Electron-Dot Diagrams
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The number of dots equals the number of valence electrons.
The number of unpaired valence electrons in a nonmetal tells
you how many covalent bonds that atom can form with other
nonmetals or how many electrons it wants to gain from metals
to form an ion.
The number of valence electrons in a metal tells you how
many electrons the metal will lose to nonmetals to form an
ion. Caution: May not work with transition metals.
EXAMPLE DOT DIAGRAMS
Click here for website on valence electrons and
electron dot diagrams
(c) 2006, Mark Rosengarten
Example Dot Diagrams
Carbon can also have this dot diagram, which it
has when it forms organic compounds.
(c) 2006, Mark Rosengarten
Excited vs. Ground State
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Configurations on the Periodic Table are ground state
configurations.
If electrons are given energy, they rise to higher energy levels
(excited state).
If the total number of electrons matches in the configuration,
but the configuration doesn’t match, the atom is in the excited
state.
Na (ground, on table): 2-8-1
Example of excited states: 2-7-2, 2-8-0-1, 2-6-3
website
(c) 2006, Mark Rosengarten
What Is Light?
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Light is formed when electrons drop from the excited
state to the ground state.
The lines on a bright-line spectrum come from specific
energy level drops and are unique to each element.
EXAMPLE SPECTRUM
(c) 2006, Mark Rosengarten
EXAMPLE SPECTRUM
website
This is the bright-line spectrum of hydrogen. The top
numbers represent the PEL (shell) change that produces the
light with that color and the bottom number is the
wavelength of the light (in nanometers, or 10-9 m).
No other element has the same bright-line spectrum as
hydrogen, so these spectra can be used to identify
elements or mixtures of elements.
(c) 2006, Mark Rosengarten
Development of the Atomic Model
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Thompson Model
Rutherford Gold Foil Experiment and Model
Bohr Model
Quantum-Mechanical Model
(c) 2006, Mark Rosengarten
Thompson Model
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The atom is a positively charged diffuse mass with
negatively charged electrons stuck in it.
website
(c) 2006, Mark Rosengarten
Rutherford Model
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The atom is made of a small, dense, positively charged nucleus
with electrons at a distance, the vast majority of the volume of
the atom is empty space.
website
Alpha particles shot
at a thin sheet of gold
foil: most go through
(empty space). Some
deflect or bounce off
(small + charged
nucleus).
(c) 2006, Mark Rosengarten
Bohr Model
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Electrons orbit around the nucleus in energy levels (shells).
Atomic bright-line spectra was the clue.
Animation
(c) 2006, Mark Rosengarten
Quantum-Mechanical Model
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Electron energy levels are wave functions.
Electrons are found in orbitals, regions of space where an
electron is most likely to be found.
You can’t know both where the electron is and where it is
going at the same time.
Electrons buzz around the nucleus like gnats buzzing around
your head.
(c) 2006, Mark Rosengarten
Matter
1) Properties of Phases
2) Types of Matter
3) Phase Changes
(c) 2006, Mark Rosengarten
Properties of Phases
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Solids: Crystal lattice (regular geometric pattern), vibration
motion only
Liquids: particles flow past each other but are still attracted to
each other.
Gases: particles are small and far apart, they travel in a straight
line until they hit something, they bounce off without losing
any energy, they are so far apart from each other that they
have effectively no attractive forces and their speed is directly
proportional to the Kelvin temperature (Kinetic-Molecular
Theory, Ideal Gas Theory)
(c) 2006, Mark Rosengarten
Solids
The positive and
negative ions
alternate in the
ionic crystal lattice
of NaCl.
(c) 2006, Mark Rosengarten
Liquids
When heated, the ions move
faster and eventually
separate from each other to
form a liquid. The ions are
loosely held together by the
oppositely charged ions, but
the ions are moving too fast
for the crystal lattice to stay
together.
(c) 2006, Mark Rosengarten
Gases
Since all gas molecules spread out
the same way, equal volumes of
gas under equal conditions of
temperature and pressure will
contain equal numbers of
molecules of gas. 22.4 L of any
gas at STP (1.00 atm and 273K)
will contain one mole
(6.02 X 1023) gas molecules.
Since there is space between gas
molecules, gases are affected by
changes in pressure.
(c) 2006, Mark Rosengarten
Types of Matter
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Substances (Homogeneous)
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Elements (cannot be decomposed by chemical change): Al,
Ne, O, Br, H
Compounds (can be decomposed by chemical change): NaCl,
Cu(ClO3)2, KBr, H2O, C2H6
Mixtures
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Homogeneous: Solutions (solvent + solute)
Heterogeneous: soil, Italian dressing, etc.
(c) 2006, Mark Rosengarten
Elements
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A sample of lead atoms (Pb). All
atoms in the sample consist of lead,
so the substance is homogeneous.
website
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A sample of chlorine atoms (Cl). All
atoms in the sample consist of
chlorine, so the substance is
homogeneous.
(c) 2006, Mark Rosengarten
website
Compounds
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Lead has two charges listed, +2
and +4. This is a sample of lead
(II) chloride (PbCl2). Two or more
elements bonded in a wholenumber ratio is a COMPOUND.
This compound is formed from
the +4 version of lead. This is lead
(IV) chloride (PbCl4). Notice how
both samples of lead compounds
have consistent composition
throughout? Compounds are
homogeneous!
(c) 2006, Mark Rosengarten
Mixtures
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A mixture of lead atoms and
chlorine atoms. They exist in no
particular ratio and are not
chemically combined with each
other. They can be separated by
physical means.
website
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A mixture of PbCl2 and PbCl4
formula units. Again, they are in no
particular ratio to each other and
can be separated without chemical
change.
(c) 2006, Mark Rosengarten
Phase Changes
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Phase Change Types
Phase Change Diagrams
Heat of Phase Change
Evaporation
(c) 2006, Mark Rosengarten
Phase Change Types
website
(c) 2006, Mark Rosengarten
Phase Change Diagrams
website
AB: Solid Phase
BC: Melting (S + L)
CD: Liquid Phase
DE: Boiling (L + G)
EF: Gas Phase
Notice how temperature remains constant during a phase
change? That’s because the
PE is
changing,
(c) 2006,
Mark
Rosengarten not the KE.
Heat of Phase Change
website
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How many joules would it take to melt 100. g of H2O (s) at
0oC?
q=mHf = (100. g)(334 J/g) = 33400 J
How many joules would it take to boil 100. g of H2O (l) at
100oC?
q=mHv = (100.g)(2260 J/g) = 226000 J
(c) 2006, Mark Rosengarten
Evaporation
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When the surface molecules of a gas travel upwards at a
great enough speed to escape.
The pressure a vapor exerts when sealed in a container at
equilibrium is called vapor pressure, and can be found on
Table H.
When the liquid is heated, its vapor pressure increases.
When the liquid’s vapor pressure equals the pressure
exerted on it by the outside atmosphere, the liquid can
boil.
If the pressure exerted on a liquid increases, the boiling
point of the liquid increases (pressure cooker). If the
pressure decreases, the boiling point of the liquid
decreases (special cooking directions for high elevations).
(c) 2006, Mark Rosengarten
Reference Table H: Vapor Pressure of Four
Liquids
website
(c) 2006, Mark Rosengarten
Bonding
1) The Periodic Table
2) Ions
3) Ionic Bonding
4) Covalent Bonding
5) Metallic Bonding
(c) 2006, Mark Rosengarten
The Periodic Table
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Metals
Nonmetals
Metalloids
Chemistry of Groups
Electronegativity
Ionization Energy
Video
(c) 2006, Mark Rosengarten
Metals
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Have luster, are malleable and ductile, good
conductors of heat and electricity
Lose electrons to nonmetal atoms to form positively
charged ions in ionic bonds
Large atomic radii compared to nonmetal atoms
Low electronegativity and ionization energy
Left side of the periodic table (except H)
(c) 2006, Mark Rosengarten
Nonmetals
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Are dull and brittle, poor conductors
Gain electrons from metal atoms to form negatively
charged ions in ionic bonds
Share unpaired valence electrons with other
nonmetal atoms to form covalent bonds and
molecules
Small atomic radii compared to metal atoms
High electronegativity and ionization energy
Right side of the periodic table (except Group 18)
(c) 2006, Mark Rosengarten
Metalloids
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Found lying on the jagged line between metals and
nonmetals flatly touching the line (except Al and Po).
Share properties of metals and nonmetals (Si is shiny like a
metal, brittle like a nonmetal and is a semiconductor).
(c) 2006, Mark Rosengarten
Chemistry of Groups
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Group 1: Alkali Metals
Group 2: Alkaline Earth Metals
Groups 3-11: Transition Elements
Group 17: Halogens
Group 18: Noble Gases
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Diatomic Molecules
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(c) 2006, Mark Rosengarten
website
Group 1: Alkali Metals
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Most active metals, only found in compounds in
nature
React violently with water to form hydrogen gas
and a strong base: 2 Na (s) + H2O (l)  2 NaOH
(aq) + H2 (g)
1 valence electron
Form +1 ion by losing that valence electron
Form oxides like Na2O, Li2O, K2O
(c) 2006, Mark Rosengarten
Group 2: Alkaline Earth Metals
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Very active metals, only found in compounds in
nature
React strongly with water to form hydrogen gas
and a base:
 Ca (s) + 2 H2O (l)  Ca(OH)2 (aq) + H2 (g)
2 valence electrons
Form +2 ion by losing those valence electrons
Form oxides like CaO, MgO, BaO
(c) 2006, Mark Rosengarten
Groups 3-11: Transition Metals
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Many can form different possible charges of ions
If there is more than one ion listed, give the charge as a Roman
numeral after the name
Cu+1 = copper (I) Cu+2 = copper (II)
Compounds containing these metals can be colored.
(c) 2006, Mark Rosengarten
Group 17: Halogens
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Most reactive nonmetals
React violently with metal atoms to form halide
compounds: 2 Na + Cl2  2 NaCl
Only found in compounds in nature
Have 7 valence electrons
Gain 1 valence electron from a metal to form -1
ions
Share 1 valence electron with another nonmetal
atom to form one covalent bond.
(c) 2006, Mark Rosengarten
Group 18: Noble Gases
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Are completely nonreactive since they have eight
valence electrons, making a stable octet.
Kr and Xe can be forced, in the laboratory, to give
up some valence electrons to react with fluorine.
Since noble gases do not naturally bond to any
other elements, one atom of noble gas is
considered to be a molecule of noble gas. This is
called a monatomic molecule. Ne represents an
atom of Ne and a molecule of Ne.
(c) 2006, Mark Rosengarten
Diatomic Molecules(elements)
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Br, I, N, Cl, H, O and F are so reactive that they exist in a more
chemically stable state when they covalently bond with
another atom of their own element to make two-atom, or
diatomic molecules.
Br2, I2, N2, Cl2, H2, O2 and F2
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The decomposition of water: 2 H2O  2 H2 + O2
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(c) 2006, Mark Rosengarten
Electronegativity
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website
An atom’s attraction to electrons in a chemical bond.
F has the highest, at 4.0
Fr has the lowest, at 0.7
If two atoms that are different in EN (END) from each other
by 1.7 or more collide and bond (like a metal atom and a
nonmetal atom), the one with the higher electronegativity will
pull the valence electrons away from the atom with the lower
electronegativity to form a (-) ion. The atom that was stripped
of its valence electrons forms a (+) ion.
If the two atoms have an END of less than 1.7, they will share
their unpaired valence electrons…covalent bond!
(c) 2006, Mark Rosengarten
Ionization Energy
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website
The energy required to remove the most loosely held
valence electron from an atom in the gas phase.
High electronegativity means high ionization energy
because if an atom is more attracted to electrons, it will
take more energy to remove those electrons.
Metals have low ionization energy. They lose electrons
easily to form (+) charged ions.
Nonmetals have high ionization energy but high
electronegativity. They gain electrons easily to form (-)
charged ions when reacted with metals, or share unpaired
valence electrons with other nonmetal atoms.
(c) 2006, Mark Rosengarten
Ions
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website
Ions are charged particles formed by the gain or loss of
electrons.
 Metals lose electrons (oxidation) to form (+) charged
cations.
 Nonmetals gain electrons (reduction) to form (-) charged
anions.
Atoms will gain or lose electrons in such a way that they end
up with 8 valence electrons (stable octet).
 The exceptions to this are H, Li, Be and B, which are not
large enough to support 8 valence electrons. They must be
satisfied with 2 (Li, Be, B) or 0 (H).
(c) 2006, Mark Rosengarten
Metal Ions (Cations-positive ion)
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Na: 2-8-1
Na+1: 2-8
Ca: 2-8-8-2
Ca+2: 2-8-8
Al: 2-8-3
Al+3: 2-8
Note that when the atom
loses its valence electron,
the next lower PEL
becomes the valence PEL.
Notice how the dot
diagrams for metal ions
lack dots! Place brackets
around the element symbol
and put the charge on the
upper right outside!
(c) 2006, Mark Rosengarten
Nonmetal Ions (Anions-negative ion)
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F: 2-7
F-1: 2-8
O: 2-6
O-2: 2-8
Note how the ions all have 8
valence electrons. Also note the
gained electrons as red dots.
Nonmetal ion dot diagrams show
8 dots, with brackets around the
dot diagram and the charge of
the ion written to the upper right
side outside the brackets.
N: 2-5
N-3: 2-8
(c) 2006, Mark Rosengarten
Ionic Bonding
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website
If two atoms that are different in ELECTRONEGATIVITY
(END) from each other by 1.7 or more collide and bond (like
a metal atom and a nonmetal atom), the one with the higher
electronegativity will pull the valence electrons away from the
atom with the lower electronegativity to form a (-) ion. The
atom that was stripped of its valence electrons forms a (+)
ion.
The oppositely charged ions attract to form the bond. It is a
surface bond that can be broken by melting or dissolving in
water.
Ionic bonding forms ionic crystal lattices, not molecules.
(c) 2006, Mark Rosengarten
Example of Ionic Bonding
(c) 2006, Mark Rosengarten
Covalent Bonding
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Video
If two nonmetal atoms have an END of 1.7 or less, they
will share their unpaired valence electrons to form a
covalent bond.
A particle made of covalently bonded nonmetal atoms is
called a molecule.
If the END is between 0 and 0.4, the sharing of electrons is
equal, so there are no charged ends. This is NONPOLAR
covalent bonding.
If the END is between 0.5 and 1.7, the sharing of electrons
is unequal. The atom with the higher EN will be d- and the
one with the lower EN will be d+ charged. This is a
POLAR covalent bonding. (d means “partial”)
(c) 2006, Mark Rosengarten
website
Examples of Covalent Bonding
(c) 2006, Mark Rosengarten
Metallic Bonding
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Metal atoms of the same element bond with each other by
sharing valence electrons that they lose to each other.
This is a lot like an atomic game of “hot potato”, where metal
kernals (the atom inside the valence electrons) sit in a crystal
lattice, passing valence electrons back and forth between each
other).
Since electrons can be forced to travel in a certain direction
within the metal, metals are very good at conducting electricity
in all phases.
website
(c) 2006, Mark Rosengarten
Compounds
1) Types of Compounds
2) Formula Writing
3) Formula Naming
4) Empirical Formulas
5) Molecular Formulas
6) Types of Chemical Reactions
7) Balancing Chemical Reactions
8) Attractive Forces
(c) 2006, Mark Rosengarten
Types of Compounds



website
Ionic: made of metal and nonmetal ions. Form an ionic
crystal lattice when in the solid phase. Ions separate when
melted or dissolved in water, allowing electrical conduction
(electrolytes-video). Examples: NaCl, K2O, CaBr2
Molecular: made of nonmetal atoms bonded to form a
distinct particle called a molecule. Bonds do not break
upon melting or dissolving, so molecular substances do not
conduct electricity. EXCEPTION: Acids [H+A- (aq)] ionize
in water to form H3O+ and A-, so they do conduct.
Network: made up of nonmetal atoms bonded in a
seemingly endless matrix of covalent bonds with no
distinguishable molecules. Very high m.p., don’t conduct.
(c) 2006, Mark Rosengarten
Ionic Compounds
website
Ionic Crystal Structure, then adding heat (or dissolving in water) to break
up the crystal into a liquid composed of free-moving ions.
(c) 2006, Mark Rosengarten
Molecular Compounds
website
(c) 2006, Mark Rosengarten
Network Solids
Network solids are made of nonmetal atoms covalently
bonded together to form large crystal lattices. No individual
molecules can be distinguished. Examples include C
(diamond) and SiO2 (quartz). Corundum (Al2O3) also forms
these, even though Al is considered a metal. Network solids
are among the hardest materials known. They have
extremely high melting points and do not conduct electricity.
(c) 2006, Mark Rosengarten
Formula Writing






website
The charge of the (+) ion and the charge of the (-) ion must
cancel out to make the formula. Use subscripts to indicate
how many atoms of each element there are in the compound,
no subscript if there is only one atom of that element.
Na+1 and Cl-1 = NaCl
Ca+2 and Br-1 = CaBr2
Al+3 and O-2 = Al2O3
Zn+2 and PO4-3 = Zn3(PO4)2
Try these problems!
(c) 2006, Mark Rosengarten
Formulas to Write

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

Ba+2 and N-3
NH4+1 and SO4-2
Li+1 and S-2
Cu+2 and NO3-1
Al+3 and CO3-2
Fe+3 and Cl-1
Pb+4 and O-2
Pb+2 and O-2
(c) 2006, Mark Rosengarten
Formula Naming






Compounds are named from the elements or
polyatomic ions that form them.
KCl = potassium chloride
Na2SO4 = sodium sulfate
(NH4)2S = ammonium sulfide
AgNO3 = silver nitrate
Notice all the metals listed here only have one
charge listed? So what do you do if a metal has
more than one charge listed? Take a peek!
(c) 2006, Mark Rosengarten
website
The Stock System


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

CrCl2 = chromium (II) chloride
CrCl3 = chromium (III) chloride
CrCl6 = chromium (VI) chloride
Try
Co(NO3)2 and
Co(NO3)3
FeO = iron (II) oxide
MnS = manganese (II) sulfide
Fe2O3 = iron (III) oxide
MnS2 = manganese (IV) sulfide
The Roman numeral is the charge of the metal ion!
(c) 2006, Mark Rosengarten
Empirical Formulas




website
Ionic formulas: represent the simplest whole number mole
ratio of elements in a compound.
Ca3N2 means a 3:2 ratio of Ca ions to N ions in the
compound.
Many molecular formulas can be simplified to empirical
formulas
 Ethane (C2H6) can be simplified to CH3. This is the empirical
formula…the ratio of C to H in the molecule.
All ionic compounds have empirical formulas.
(c) 2006, Mark Rosengarten
Molecular Formulas





The count of the actual number of atoms of each element
in a molecule.
H2O: a molecule made of two H atoms and one O atom
covalently bonded together.
C2H6O: A molecule made of two C atoms, six H atoms and
one O atom covalently bonded together.
Molecular formulas are whole-number multiples of
empirical formulas:
 H2O = 1 X (H2O)
 C8H16 = 8 X (CH2)
Calculating Molecular Formulas
(c) 2006, Mark Rosengarten
Types of Chemical Reactions

Redox Reactions: driven by the loss (oxidation) and gain
(reduction) of electrons. Any species that does not change
charge is called the spectator ion.
website
 Synthesis
 Decomposition
 Single Replacement
website

Ion Exchange Reaction: driven by the formation of an
insoluble precipitate. The ions that remain dissolved
throughout are the spectator ions.
 Double Replacement
(c) 2006, Mark Rosengarten
Synthesis

Two elements combine to form a compound
2 Na + O2  Na2O
Same reaction, with charges added in:
 2 Na0 + O20  Na2+1O-2
Na0 is oxidized (loses electrons), is the reducing agent
O20 is reduced (gains electrons), is the oxidizing agent

Electrons are transferred from the Na0 to the O20.

No spectator ions, there are only two elements here.

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
(c) 2006, Mark Rosengarten
Decomposition

A compound breaks down into its original elements.
Na2O  2 Na + O2
Same reaction, with charges added in:
 Na2+1O-2  2 Na0 + O20
O-2 is oxidized (loses electrons), is the reducing agent
Na+1 is reduced (gains electrons), is the oxidizing agent

Electrons are transferred from the O-2 to the Na+1.

No spectator ions, there are only two elements here.

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
(c) 2006, Mark Rosengarten
Single Replacement

An element replaces the same type of element in a compound.
Ca + 2 KCl  CaCl2 + 2 K
Same reaction, with charges added in:
 Ca0 + 2 K+1Cl-1  Ca+2Cl2-1 + 2 K0
Ca0 is oxidized (loses electrons), is the reducing agent
K+1 is reduced (gains electrons), is the oxidizing agent

Electrons are transferred from the Ca0 to the K+1.

Cl-1 is the spectator ion, since it’s charge doesn’t change.

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

(c) 2006, Mark Rosengarten
Double Replacement

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
The (+) ion of one compound bonds to the (-) ion of another
compound to make an insoluble precipitate. The compounds
must both be dissolved in water to break the ionic bonds first.
NaCl (aq) + AgNO3 (aq)  NaNO3 (aq) + AgCl (s)
The Cl-1 and Ag+1 come together to make the insoluble
precipitate, which looks like snow in the test tube.
No species change charge, so this is not a redox reaction.
Since the Na+1 and NO3-1 ions remain dissolved throughout
the reaction, they are the spectator ions.
How do identify the precipitate?
(c) 2006, Mark Rosengarten
Identifying the Precipitate

website
The precipitate is the compound that is insoluble. AgCl is
a precipitate because Cl- is a halide. Halides are soluble,
except when combined with Ag+ and others.
(c) 2006, Mark Rosengarten
Balancing Chemical Reactions

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
Balance one element or ion at a time
Use a pencil
Use coefficients only, never change subscripts(formulas)
Revise if necessary
The coefficient multiplies everything in the formula by that
amount
 2 Ca(NO3)2 means that you have 2 Ca, 4 N and 12 O.
Examples for you to try!
website
website
(c) 2006, Mark Rosengarten
Reactions to Balance

___NaCl  ___Na + ___Cl2

___Al + ___O2  ___Al2O3

___SO3  ___SO2 + ___O2

___Ca + ___HNO3  ___Ca(NO3)2 + ___H2

__FeCl3 + __Pb(NO3)2  __Fe(NO3)3 + __PbCl2
(c) 2006, Mark Rosengarten
Attractive Forces

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
Molecules have partially charged ends. The d+ end of one
molecule attracts to the d- end of another molecule.
Ions are charged (+) or (-). Positively charged ions attract
other to form ionic bonds, a type of attractive force.
Since partially charged ends result in weaker attractions than
fully charged ends, ionic compounds generally have much
higher melting points than molecular compounds.
Determining Polarity of Molecules
Hydrogen Bond Attractions
(c) 2006, Mark Rosengarten
Determining Polarity of
Molecules
website
-----------------------------------------------------------------------------
website
(c) 2006, Mark Rosengarten
Hydrogen Bond
Attractions
website
A hydrogen bond attraction is a
very strong attractive force
between the H end of one polar
molecule and the N, O or F end
of another polar molecule. This
attraction is so strong that water
is a liquid at a temperature
where most compounds that are
much heavier than water (like
propane, C3H8) are gases. This
also gives water its surface
tension and its ability to form a
meniscus in a narrow glass tube.
(c) 2006, Mark Rosengarten
Math of Chemistry
1) Formula Mass
2) Percent Composition
3) Mole Problems
4) Gas Laws
5) Neutralization
6) Concentration
7) Significant Figures and Rounding
8) Metric Conversions
9) Calorimetry
(c) 2006, Mark Rosengarten
Formula Mass





Gram Formula Mass = sum of atomic masses of all elements in
the compound
Round given atomic masses to the nearest tenth
H2O: (2 X 1.0) + (1 X 16.0) = 18.0 grams/mole
Na2SO4: (2 X 23.0)+(1 X 32.1)+(4 X 16.0) = 142.1 g/mole
Now you try:
 BaBr2
 CaSO4
 Al2(CO3)3
Video
website
(c) 2006, Mark Rosengarten
Percent Composition
website
What is the % composition, by mass,
of each element in SiO2?
%Si =
(28.1/60.1) X 100 = 46.8%
%O = (2 X 16.0 = 32.0), (32.0/60.1) X 100 = 53.2%
The mass of part is the number of atoms of that element in the
compound. The mass of whole is the formula mass of the
compound. Don’t forget to take atomic mass to the nearest
tenth! This is a problem for you to try.
(c) 2006, Mark Rosengarten
Practice Percent
Composition Problem

What is the percent by mass of each element in Li2SO4?
(c) 2006, Mark Rosengarten
Mole Problems



Grams <=> Moles
Molecular Formula
Stoichiometry
(c) 2006, Mark Rosengarten
Grams <=> Moles




How many grams will 3.00 moles of NaOH (40.0 g/mol)
weigh?
3.00 moles X 40.0 g/mol = 120. g
How many moles of NaOH (40.0 g/mol) are represented
by 10.0 grams?
(10.0 g) / (40.0 g/mol) = 0.250 mol
Video
(c) 2006, Mark Rosengarten
Molecular Formula

Molecular Formula = (Molecular Mass/Empirical Mass) X Empirical Formula

What is the molecular formula of a compound with an
empirical formula of CH2 and a molecular mass of 70.0
grams/mole?
1) Find the Empirical Formula Mass: CH2 = 14.0
2) Divide the MM/EM: 70.0/14.0 = 5
3) Multiply the molecular formula by the result:



5 (CH2) = C5H10
(c) 2006, Mark Rosengarten
Stoichiometry

Moles of Target = Moles of Given X (Coefficent of Target/Coefficient of given)

Given the balanced equation N2 + 3 H2  2 NH3,
How many moles of H2 need to be completely
reacted with N2 to yield 20.0 moles of NH3?

20.0 moles NH3 X (3 H2 / 2 NH3) = 30.0 moles H2
website
(c) 2006, Mark Rosengarten
Gas Laws





Make a data table to put the numbers so you can eliminate the
words.
Make sure that any Celsius temperatures are converted to
Kelvin (add 273).
Rearrange the equation before substituting in numbers. If you
are trying to solve for T2, get it out of the denominator first by
cross-multiplying.
If one of the variables is constant, then eliminate it.
Try these problems!
website
website
website
(c) 2006, Mark Rosengarten
Gas Law Problem 1




A 2.00 L sample of N2 gas at
STP is compressed to 4.00
atm at constant temp-erature.
What is the new volume of
the gas?
V2 = P1V1 / P2
= (1.00 atm)(2.00 L) / (4.00
atm)
= 0.500 L
(c) 2006, Mark Rosengarten
Gas Law Problem 2



To what temperature must a 3.000 L sample of O2 gas at 300.0
K be heated to raise the volume to 10.00 L?
T2 = V2T1/V1
= (10.00 L)(300.0 K) / (3.000 L) = 1000. K
(c) 2006, Mark Rosengarten
Gas Law Problem 3

A 3.00 L sample of NH3 gas at 100.0 kPa is cooled from 500.0
K to 300.0 K and its pressure is reduced to 80.0 kPa. What is
the new volume of the gas?

V2 = P1V1T2 / P2T1
= (100.0 kPa)(3.00 L)(300. K) / (80.0 kPa)(500. K)
= 2.25 L


(c) 2006, Mark Rosengarten
Neutralization
website

10.0 mL of 0.20 M HCl is neutralized by 40.0 mL of NaOH.
What is the concentration of the NaOH?

#H MaVa = #OH MbVb, so Mb = #H MaVa / #OH Vb

= (1)(0.20 M)(10.0 mL) / (1) (40.0 mL) = 0.050 M

How many mL of 2.00 M H2SO4 are needed to completely
neutralize 30.0 mL of 0.500 M KOH?
(c) 2006, Mark Rosengarten
Concentration

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
Molarity
Parts per Million
Percent by Mass
Percent by Volume
(c) 2006, Mark Rosengarten
Molarity

What is the molarity of a 500.0 mL solution of NaOH (FM
= 40.0) with 60.0 g of NaOH (aq)?



website
Convert g to moles and mL to L first!
M = moles / L = 1.50 moles / 0.5000 L = 3.00 M
How many grams of NaOH does it take to make 2.0 L of a
0.100 M solution of NaOH (aq)?


Moles = M X L = 0.100 M X 2.0 L = 0.200 moles
Convert moles to grams: 0.200 moles X 40.0 g/mol = 8.00 g
(c) 2006, Mark Rosengarten
Parts Per Million

100.0 grams of water is evaporated and analyzed for lead.
0.00010 grams of lead ions are found. What is the
concentration of the lead, in parts per million?

ppm = (0.00010 g) / (100.0 g) X 1 000 000 = 1.0 ppm
If the legal limit for lead in the water is 3.0 ppm, then the
water sample is within the legal limits (it’s OK!)

(c) 2006, Mark Rosengarten
Percent by Mass

A 50.0 gram sample of a solution is evaporated and found
to contain 0.100 grams of sodium chloride. What is the
percent by mass of sodium chloride in the solution?

% Comp = (0.100 g) / (50.0 g) X 100 = 0.200%
(c) 2006, Mark Rosengarten
Percent By Volume

Substitute “volume” for “mass” in the above equation.

What is the percent by volume of hexane if 20.0 mL of
hexane are dissolved in benzene to a total volume of 80.0
mL?

% Comp = (20.0 mL) / (80.0 mL) X100 = 25.0%
(c) 2006, Mark Rosengarten
Sig Figs and Rounding

How many Significant Figures does a number have?

What is the precision of my measurement?

How do I round off answers to addition and subtraction
problems?

How do I round off answers to multiplication and division
problems?
(c) 2006, Mark Rosengarten
website
How many Sig Figs?


Start counting sig figs at the first non-zero.
All digits except place-holding zeroes are sig figs.
Measurement
# of Sig Figs
Measurement
# of Sig Figs
0.115 cm
3
234 cm
3
0.00034 cm
2
67000 cm
2
0.00304 cm
3
_
45000 cm
4
0.0560 cm
3
560. cm
3
0.00070700 cm
5
560.00 cm
5
(c) 2006, Mark Rosengarten
What Precision?

A number’s precision is determined by the furthest (smallest)
place the number is recorded to.

6000 mL : thousands place
6000. mL : ones place
6000.0 mL : tenths place
5.30 mL : hundredths place
8.7 mL : tenths place
23.740 mL : thousandths place

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
(c) 2006, Mark Rosengarten
Rounding with addition and
subtraction

Answers are rounded to the least precise place.
1) 4.732 cm
16.8
cm
+ 0.781 cm
---------22.313 cm
22.3 cm
2)
17.440 mL
3.895 mL
+ 16.77 mL
-------------38.105 mL
38.11 mL
(c) 2006, Mark Rosengarten
3)
32.0
MW
+ 0.0059 MW
--------------32.0059 MW
32.0 MW
Rounding with multiplication
and division

Answers are rounded to the fewest number of significant
figures.
1)
37.66 KW
x 2.2 h
---------82.852 KWh
83 KWh
2)
14.922 cm
x 2.0 cm
----------2
29.844 cm
2
30. cm
website
(c) 2006, Mark Rosengarten
3) 98.11 kg
x 200 m
---------19 622 kgm
20 000 kgm
Metric Conversions

Determine how many powers of ten
difference there are between the two
units (no prefix = 100) and create a
conversion factor. Multiply or divide
the given by the conversion factor.
How many kg are in 38.2 cg?
(38.2 cg) /(100000 cg/kg) = 0.000382 km
How many mL in 0.988 dL?
(0.988 dg) X (100 mL/dL) = 98.8 mL
(c) 2006, Mark Rosengarten
Calorimetry




This equation can be used to determine any of the
variables here. You will not have to solve for C, since we
will always assume that the energy transfer is being
absorbed by or released by a measured quantity of water,
whose specific heat is given above.
Solving for q
website
Solving for m
Solving for DT
(c) 2006, Mark Rosengarten
Solving for q

How many joules are absorbed by 100.0 grams of water in a
calorimeter if the temperature of the water increases from
20.0oC to 50.0oC?

q = mCDT = (100.0 g)(4.18 J/goC)(30.0oC) = 12500 J
(c) 2006, Mark Rosengarten
Solving for m

A sample of water in a calorimeter cup increases from 25oC
to 50.oC by the addition of 500.0 joules of energy. What is the
mass of water in the calorimeter cup?

q = mCDT, so m = q / CDT = (500.0 J) / (4.18 J/goC)(25oC) = 4.8 g
(c) 2006, Mark Rosengarten
Solving for DT

If a 50.0 gram sample of water in a calorimeter cup absorbs
1000.0 joules of energy, how much will the temperature rise
by?

q = mCDT, so DT = q / mC = (1000.0 J)/(50.0 g)(4.18 J/goC) = 4.8oC

If the water started at 20.0oC, what will the final temperature
be?

Since the water ABSORBS the energy, its temperature will INCREASE by
the DT: 20.0oC + 4.8oC = 24.8oC
(c) 2006, Mark Rosengarten
Kinetics and Thermodynamics
1) Reaction Rate
2) Heat of Reaction
3) Potential Energy Diagrams
4) Equilibrium
5) Le Châtelier’s Principle
6) Solubility Curves
(c) 2006, Mark Rosengarten
Reaction Rate



website
Reactions happen when reacting particles collide with
sufficient energy (activation energy) and at the proper angle.
Anything that makes more collisions in a given time will make
the reaction rate increase.
 Increasing temperature
 Increasing concentration (pressure for gases)
 Increasing surface area (solids)
Adding a catalyst makes a reaction go faster by removing steps
from the mechanism and lowering the activation energy
without getting used up in the process.
website
(c) 2006, Mark Rosengarten
Heat of Reaction

Reactions either absorb PE (endothermic, +DH) or release PE
(exothermic, -DH)
Exothermic, PEKE, Temp
Endothermic, KEPE, Temp
Rewriting the equation with heat included:
4 Al(s) + 3 O2(g)  2 Al2O3(s) + 3351 kJ
N2(g) + O2(g) +182.6 kJ  2 NO(g)
(c) 2006, Mark Rosengarten
Potential Energy Diagrams

Video
website
Steps of a reactions:
 Reactants have a certain amount of PE stored in their
bonds (Heat of Reactants)
 The reactants are given enough energy to collide and
react (Activation Energy)
 The resulting intermediate has the highest energy that
the reaction can make (Heat of Activated Complex)
 The activated complex breaks down and forms the
products, which have a certain amount of PE stored in
their bonds (Heat of Products)
 Hproducts - Hreactants = DH
EXAMPLES
(c) 2006, Mark Rosengarten
Making a PE Diagram






X axis: Reaction Coordinate (time, no units)
Y axis: PE (kJ)
Three lines representing energy (Hreactants, Hactivated complex,
Hproducts)
Two arrows representing energy changes:
 From Hreactants to Hactivated complex: Activation Energy
 From Hreactants to Hproducts : DH
ENDOTHERMIC PE DIAGRAM
EXOTHERMIC PE DIAGRAM
(c) 2006, Mark Rosengarten
Endothermic PE Diagram
If a catalyst is added?
(c) 2006, Mark Rosengarten
Endothermic with Catalyst
The red line represents the catalyzed reaction.
(c) 2006, Mark Rosengarten
Exothermic PE Diagram
What does it look like with
catalyst?
(c) a
2006,
Mark Rosengarten
Exothermic with a Catalyst
The red line represents the catalyzed reaction. Lower
A.E. and faster reaction time!
(c) 2006, Mark Rosengarten
Equilibrium
When the rate of the forward reaction equals the rate of the
reverse reaction.
(c) 2006, Mark Rosengarten
Examples of Equilibrium

Solution Equilibrium: when a solution is saturated, the rate of
dissolving equals the rate of precipitating.


Vapor-Liquid Equilibrium: when a liquid is trapped with air in a
container, the liquid evaporates until the rate of evaporation
equals the rate of condensation.


NaCl (s)  Na+1 (aq) + Cl-1 (aq)
H2O (l)  H2O (g)
Phase equilibrium: At the melting point, the rate of solid
turning to liquid equals the rate of liquid turning back to solid.

H2O (s)  H2O (l)
(c) 2006, Mark Rosengarten
Le Châtelier’s Principle





If a system at equilibrium is stressed, the equilibrium
will shift in a direction that relieves that stress.
A stress is a factor that affects reaction rate. Since catalysts
affect both reaction rates equally, catalysts have no effect on a
system already at equilibrium.
Equilibrium will shift AWAY from what is added
Equilibrium will shift TOWARDS what is removed.
This is because the shift will even out the change in reaction
rate and bring the system back to equilibrium
NEXT
website
Video
(c) 2006, Mark Rosengarten
Steps to Relieving Stress



1) Equilibrium is subjected to a STRESS.
2) System SHIFTS towards what is removed from the system
or away from what is added.
The shift results in a CHANGE OF CONCENTRATION for
both the products and the reactants.
 If the shift is towards the products, the concentration of the
products will increase and the concentration of the
reactants will decrease.
 If the shift is towards the reactants, the concentration of the
reactants will increase and the concentration of the
products will decrease.
 NEXT
(c) 2006, Mark Rosengarten
Examples

For the reaction N2(g) + 3H2(g)  2 NH3(g) +
heat





Adding N2 will cause the equilibrium to shift RIGHT, resulting in an
increase in the concentration of NH3 and a decrease in the
concentration of N2 and H2.
Removing H2 will cause a shift to the LEFT, resulting in a decrease in
the concentration of NH3 and an increase in the concentration of
N2 and H2.
Increasing the temperature will cause a shift to the LEFT, same
results as the one above.
Decreasing the pressure will cause a shift to the LEFT, because there
is more gas on the left side, and making more gas will bring the
pressure back up to its equilibrium amount.
Adding a catalyst will have no effect, so no shift will happen.
(c) 2006, Mark Rosengarten
Solubility Curves





website
Solubility: the maximum quantity of solute that can be
dissolved in a given quantity of solvent at a given temperature
to make a saturated solution.
Saturated: a solution containing the maximum quantity of
solute that the solvent can hold. The limit of solubility.
Supersaturated: the solution is holding more than it can
theoretically hold OR there is excess solute which precipitates
out. True supersaturation is rare.
Unsaturated: There are still solvent molecules available to
dissolve more solute, so more can dissolve.
How ionic solutes dissolve in water: polar water molecules
attach to the ions and tear them off the crystal.
 LIKE DISSOLVES LIKE
(c) 2006, Mark Rosengarten
Solubility
Video
Solubility: go to the temperature
and up to the desired line, then
across to the Y-axis. This is how
many g of solute are needed to
make a saturated solution of that
solute in 100g of H2O at that
particular temperature.
At 40oC, the solubility of KNO3 in
100g of water is 64 g. In 200g of
water, double that amount. In 50g
of water, cut it in half.
(c) 2006, Mark Rosengarten
Supersaturated
If 120 g of NaNO3 are added to
100g of water at 30oC:
1) The solution would be
SUPERSATURATED, because
there is more solute dissolved
than the solubility allows
2) The extra 25g would
precipitate out
3) If you heated the solution up
by 24oC (to 54oC), the excess
solute would dissolve.
(c) 2006, Mark Rosengarten
Unsaturated
If 80 g of KNO3 are added to
100g of water at 60oC:
1) The solution would be
UNSATURATED, because there
is less solute dissolved than the
solubility allows
2) 26g more can be added to
make a saturated solution
3) If you cooled the solution
down by 12oC (to 48oC), the
solution would become saturated
(c) 2006, Mark Rosengarten
How Ionic Solutes Dissolve in Water
Water solvent molecules attach to the
ions (H end to the Cl-, O end to the Na+)
Water solvent holds the ions apart and
keeps the ions from coming back together
(c) 2006, Mark Rosengarten
Acids and Bases
1) Formulas, Naming and Properties of Acids
2) Formulas, Naming and Properties of Bases
3) Neutralization
4) pH
5) Indicators
6) Alternate Theories
(c) 2006, Mark Rosengarten
Formulas, Naming and
Properties of Acids

Arrhenius Definition of Acids: molecules that dissolve in water
to produce H3O+ (hydronium) as the only positively charged
ion in solution.




HCl (g) + H2O (l)  H3O+ (aq) + Cl-
Properties of Acids
Naming of Acids
Formula Writing of Acids
(c) 2006, Mark Rosengarten
Properties of Acids
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Acids react with metals above H2 on Table J to
form H2(g) and a salt.
Acids have a pH of less than 7.
Dilute solutions of acids taste sour.
Acids turn phenolphthalein CLEAR, litmus RED
and bromthymol blue YELLOW.
Acids neutralize bases.
Acids are formed when acid anhydrides (NO2,
SO2, CO2) react with water for form acids. This is
how acid rain forms from auto and industrial
emissions.
(c) 2006, Mark Rosengarten
Naming of Acids

Binary Acids (H+ and a nonmetal)
 hydro (nonmetal) -ide + ic acid


HCl (aq) = hydrochloric acid
Ternary Acids (H+ and a polyatomic ion)
 (polyatomic ion) -ate +ic acid


(polyatomic ion) -ide +ic acid


HNO3 (aq) = nitric acid
HCN (aq) = cyanic acid
(polyatomic ion) -ite +ous acid

HNO2 (aq) = nitrous acid
(c) 2006, Mark Rosengarten
Formula Writing of Acids
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Acids formulas get written like any other. Write the H+1
first, then figure out what the negative ion is based on the
name. Cancel out the charges to write the formula. Don’t
forget the (aq) after it…it’s only an acid if it’s in water!
Hydrosulfuric acid: H+1 and S-2 = H2S (aq)
Carbonic acid: H+1 and CO3-2 = H2CO3 (aq)
Chlorous acid: H+1 and ClO2-1 = HClO2 (aq)
Hydrobromic acid: H+1 and Br-1 = HBr (aq)
Hydronitric acid:
Hypochlorous acid:
Perchloric acid:
(c) 2006, Mark Rosengarten
Formulas, Naming and
Properties of Bases

Arrhenius Definition of Bases: ionic compounds that dissolve
in water to produce OH- (hydroxide) as the only negatively
charged ion in solution.
 NaOH (s)  Na+1 (aq) + OH-1 (aq)

Properties of Bases
Naming of Bases
Formula Writing of Bases


(c) 2006, Mark Rosengarten
Properties of Bases
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Bases react with fats to form soap and glycerol. This process is
called saponification.
Bases have a pH of more than 7.
Dilute solutions of bases taste bitter.
Bases turn phenolphthalein PINK, litmus BLUE and
bromthymol blue BLUE.
Bases neutralize acids.
Bases are formed when alkali metals or alkaline earth metals
react with water. The words “alkali” and “alkaline” mean
“basic”, as opposed to “acidic”.
(c) 2006, Mark Rosengarten
Naming of Bases

Bases are named like any ionic
compound, the name of the metal ion
first (with a Roman numeral if
necessary) followed by “hydroxide”.
Fe(OH)2 (aq) = iron (II) hydroxide
Fe(OH)3 (aq) = iron (III) hydroxide
Al(OH)3 (aq) = aluminum hydroxide
NH3 (aq) is the same thing as NH4OH:
NH3 + H2O  NH4OH
Also called
(c) 2006,ammonium
Mark Rosengarten hydroxide.
Formula Writing of Bases

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Formula writing of bases is the same as for any ionic formula
writing. The charges of the ions have to cancel out.
Calcium hydroxide = Ca+2 and OH-1 = Ca(OH)2 (aq)
Potassium hydroxide = K+1 and OH-1 = KOH (aq)
Lead (II) hydroxide = Pb+2 and OH-1 = Pb(OH)2 (aq)
Lead (IV) hydroxide = Pb+4 and OH-1 = Pb(OH)4 (aq)
Lithium hydroxide =
Copper (II) hydroxide =
Magnesium hydroxide =
(c) 2006, Mark Rosengarten
website
Neutralization

H+1 + OH-1  HOH

Acid + Base  Water + Salt (double replacement)

HCl (aq) + NaOH (aq)  HOH (l) + NaCl (aq)

H2SO4 (aq) + KOH (aq)  2 HOH (l) + K2SO4 (aq)

HBr (aq) + LiOH (aq) 

H2CrO4 (aq) + NaOH (aq) 

HNO3 (aq) + Ca(OH)2 (aq) 

H3PO4 (aq) + Mg(OH)2 (aq) 
(c) 2006, Mark Rosengarten
pH



website
A change of 1 in pH is a tenfold increase in acid or base
strength.
A pH of 4 is 10 times more acidic than a pH of 5.
A pH of 12 is 100 times more basic than a pH of 10.
(c) 2006, Mark Rosengarten
website
Indicators
At a pH of 2:
Methyl Orange = red
Bromthymol Blue = yellow
Phenolphthalein = colorless
Litmus = red
Methyl orange is red at a pH of
3.2 and below and yellow at a pH
of 4.4 and higher. In between the
two numbers, it is an intermediate
color that is not listed on this
table.
Bromcresol Green = yellow
Thymol Blue = yellow
(c) 2006, Mark Rosengarten
Alternate Theories



Website-video
Arrhenius Theory: acids and bases must be in aqueous
solution.
Alternate Theory: Not necessarily so!
 Acid: proton (H+1) donor…gives up H+1 in a reaction.
 Base: proton (H+1) acceptor…gains H+1 in a reaction.
HNO3 + H2O  H3O+1 + NO3-1
 Since HNO3 lost an H+1 during the reaction, it is an acid.
 Since H2O gained the H+1 that HNO3 lost, it is a base.
(c) 2006, Mark Rosengarten
Oxidation and Reduction
1) Oxidation Numbers
2) Identifying OX, RD and SI Species
3) Agents
4) Writing Half-Reactions
5) Balancing Half-Reactions
6) Activity Series
7) Voltaic Cells
8) Electrolytic Cells
9) Electroplating
(c) 2006, Mark Rosengarten
Oxidation Numbers





Video
Elements have no charge until they bond to other elements.
 Na0, Li0, H20. S0, N20, C600
The formula of a compound is such that the charges of the
elements making up the compound all add up to zero.
The symbol and charge of an element or polyatomic ion is
called a SPECIES.
Determine the charge of each species in the following
compounds:
NaCl
KNO3
CuSO4
Fe2(CO3)3
website
(c) 2006, Mark Rosengarten
Identifying OX, RD, SI
Species




website
website
Ca0 + 2 H+1Cl-1  Ca+2Cl-12 + H20
Oxidation = loss of electrons. The species becomes more
positive in charge. For example, Ca0  Ca+2, so Ca0 is the
species that is oxidized.
Reduction = gain of electrons. The species becomes more
negative in charge. For example, H+1  H0, so the H+1 is
the species that is reduced.
Spectator Ion = no change in charge. The species does not
gain or lose any electrons. For example, Cl-1  Cl-1, so the
Cl-1 is the spectator ion.
(c) 2006, Mark Rosengarten
Agents
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


Ca0 + 2 H+1Cl-1  Ca+2Cl-12 + H20
Since Ca0 is being oxidized and H+1 is being reduced, the
electrons must be going from the Ca0 to the H+1.
Since Ca0 would not lose electrons (be oxidized) if H+1
weren’t there to gain them, H+1 is the cause, or agent, of Ca0’s
oxidation. H+1 is the oxidizing agent.
Since H+1 would not gain electrons (be reduced) if Ca0 weren’t
there to lose them, Ca0 is the cause, or agent, of H+1’s
reduction. Ca0 is the reducing agent.
(c) 2006, Mark Rosengarten
Video
Writing Half-Reactions



Ca0 + 2 H+1Cl-1  Ca+2Cl-12 + H20
Oxidation: Ca0  Ca+2 + 2eReduction: 2H+1 + 2e-  H20
The two electrons lost
by Ca0 are gained by
the two H+1 (each H+1
picks up an electron).
PRACTICE SOME!
(c) 2006, Mark Rosengarten
Practice Half-Reactions

Don’t forget to determine the charge of each species first!

4 Li + O2  2 Li2O
Oxidation Half-Reaction:
Reduction Half-Reaction:


Zn + Na2SO4  ZnSO4 + 2 Na
Oxidation Half-Reaction:
Reduction Half-Reaction:



(c) 2006, Mark Rosengarten
Balancing Half-Reactions

Ca0 + Fe+3  Ca+2 + Fe0

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Ca’s charge changes by 2, so double the Fe.
Fe’s charge changes by 3, so triple the Ca.



3 Ca0 + 2 Fe+3  3 Ca+2 + 2 Fe0
Try these:
__Na0 + __H+1  __Na+1 + __H20


website
(hint: balance the H and H2 first!)
__Al0 + __Cu+2  __Al+3 + __Cu0
(c) 2006, Mark Rosengarten
Activity Series


For metals, the higher up the chart the
element is, the more likely it is to be
oxidized. This is because metals like to
lose electrons, and the more active a
metallic element is, the more easily it can
lose them.
For nonmetals, the higher up the chart the
element is, the more likely it is to be
reduced. This is because nonmetals like to
gain electrons, and the more active a
nonmetallic element is, the more easily it
can gain them.
(c) 2006, Mark Rosengarten
Metal Activity
3 K0 + Fe+3Cl-13
REACTION


Fe0 + 3 K+1Cl-1
NO REACTION


Metallic elements start out with a charge
of ZERO, so they can only be oxidized to
form (+) ions.
The higher of two metals MUST undergo
oxidation in the reaction, or no reaction
will happen.
The reaction 3 K + FeCl3  3 KCl + Fe
WILL happen, because K is being
oxidized, and that is what Table J says
should happen.
The reaction Fe + 3 KCl  FeCl3 + 3 K
will NOT happen.
(c) 2006, Mark Rosengarten
Voltaic Cells




Animation
Produce electrical current using a spontaneous redox reaction
Used to make batteries!
Animation
Materials needed: two beakers, piece of the oxidized metal
(anode, - electrode), solution of the oxidized metal, piece of
the reduced metal (cathode, + electrode), solution of the
reduced metal, porous material (salt bridge), solution of a salt
that does not contain either metal in the reaction, wire and a
load to make use of the generated current!
Use Reference Table J to determine the metals to use
 Higher = (-) anode
Lower = (+) cathode
(c) 2006, Mark Rosengarten
Making Voltaic Cells
More
Info!!!
Create
Your
Own
(c) 2006, Mark Rosengarten
Cell!!!!
How It Works
Since Zn is listed above Cu, Zn0 will be
oxidized when it reacts with Cu+2. The
reaction: Zn + CuSO4  ZnSO4 + Cu

The Zn0 anode loses 2 e-, which go up the wire and
through the load. The Zn0 electrode gets smaller as the
Zn0 becomes Zn+2 and dissolves into solution. The e- go
into the Cu0, where they sit on the outside surface of the
Cu0 cathode and wait for Cu+2 from the solution to come
over so that the e- can jump on to the Cu+2 and reduce it
to Cu0. The size of the Cu0 electrode increases. The
negative ions in solution go over the salt bridge to the
anode side to complete the circuit.
(c) 2006, Mark Rosengarten
You Start At The Anode
Vital to make a battery
Is this electrochemistry
You take two half-cells
And connect them up so well
With a load to power in between
You need to have electrodes you see
Full of that metallicity
Let electrons flow
Across the salt bridge we go!
Allowing us to make electricity
We start the anode
Electrons are lost there
And go through the wire
And through the load on fire
They get to the cathode
And reduce the cations
And the anions go through the salt bridge
Back to where…
WHERE?
(c) 2006, Mark Rosengarten
Make Your Own Cell!!!
(c) 2006, Mark Rosengarten
Electrolytic Cells



Website & Video
Use electricity to force a nonspontaneous redox reaction
to take place.
Uses for Electrolytic Cells:
 Decomposition of Alkali Metal Compounds
 Decomposition of Water into Hydrogen and Oxygen
 Electroplating
Differences between Voltaic and Electrolytic Cells:
 ANODE:
Voltaic (-)
Electrolytic (+)
 CATHODE:
Voltaic (+)
Electrolytic (-)
 Voltaic: 2 half-cells, a salt bridge and a load
Video
 Electrolytic: 1 cell, no salt bridge, IS the load
(c) 2006, Mark Rosengarten
Decomposing Alkali
Metal Compounds
2 NaCl  2 Na + Cl2
The Na+1 is reduced at
the (-) cathode,
picking up an e- from
the battery
The Cl-1 is oxidized at
the (+) anode, the ebeing pulled off by the
battery (DC)
(c) 2006, Mark Rosengarten
Decomposing Water
2 H2O  2 H2 + O2
The H+ is reduced at
the (-) cathode,
yielding H2 (g), which
is trapped in the tube.
The O-2 is oxidized at
the (+) anode, yielding
O2 (g), which is
trapped in the tube.
(c) 2006, Mark Rosengarten
Electroplating
The Ag0 is oxidized to Ag+1
when the (+) end of the
battery strips its electrons
off.
The Ag+1 migrates through
the solution towards the (-)
charged cathode (ring),
where it picks up an electron
from the battery and forms
Ag0, which coats on to the
ring.
(c) 2006, Mark Rosengarten
Organic Chemistry
1) Hydrocarbons
2) Substituted Hydrocarbons
3) Organic Families
4) Organic Reactions
(c) 2006, Mark Rosengarten
Hydrocarbons

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
Molecules made of Hydrogen and Carbon
Carbon forms four bonds, hydrogen forms one bond
Hydrocarbons come in three different homologous series:
 Alkanes (single bond between C’s, saturated)
 Alkenes (1 double bond between 2 C’s, unsaturated)
 Alkynes (1 triple bond between 2 C’s, unsaturated)
These are called aliphatic, or open-chain, hydrocarbons.
Count the number of carbons and add the appropriate suffix!
website
(c) 2006, Mark Rosengarten
Vocabulary
(c) 2006, Mark Rosengarten
Nucleon – particle found in the nucleus of an atom – includes the proton and
neutron only – equal to the mass number of an atom
Isotope – atoms of the same element which have the same atomic number but
different mass number
Atomic Number - equal to the number of protons in the nucleus of an atom
Mass Number - equal to the sum of the protons and neutrons in the nucleus of an
atom.
Nuclear charge - equal to the number of protons in the nucleus of an atom.
Alpha Particle – A radioactive particle equivalent to a helium nucleus (2 protons,
2 neutrons) - Mass of 4 and a +2 charge
Beta Particle – A radioactive particle equivalent to an electron. Has no mass and
-1 charge
Positron – A radioactive particle equivalent to an positively charged electron.
Has no mass and +1 charge
Gamma Rays – High energy light given off during a nuclear process – Have no
mass or charge
Fission – A nuclear reaction where a large nucleus breaks up into smaller ones.
This is what happens in nuclear power plants
Fusion - process where two or more small nuclei combine to form a larger
nucleus. Fusion is the reverse process of nuclear fission
(c) 2006, Mark Rosengarten
Valence Electron – Electrons in the outermost energy level (furthest
away from the nucleus) – Generally the only electrons involved in
chemical reactions.
Electron Dot Diagram (EDD) – Symbol of an element surrounded
by dots which represent valence elctrons
Stable Octet - The octet rule is a rule that states that atoms tend to
combine in such a way that they each have eight electrons in their
valence shells, giving them the same electronic configuration as a
noble gas.
Bright Line Spectrum - When electrons jump from the excited state
to the ground state, the electrons emit energy in the form of light,
producing a bright-line spectrum. Each element has its own unique
bright-line spectrum.
Orbital- Regions of the most probable electron location in the
wave-mechanical model of the atom
Solid – phase of matter with a definite shape and volume and low
entropy – particles arranged in a(c)regular
geometric pattern
2006, Mark Rosengarten
Liquid – phase of matter that has a definite volume but takes shape
of its container
Gas – phase of matter that takes the shape of & fills its entire
container – has high entropy.
Element - substances that are composed of atoms that have the
same atomic number. Elements cannot be broken down by
chemical change.
Compound - 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.
Mixture - composed of two or more different substances that can be
separated by physical means.When different substances are mixed
together, a homogeneous or heterogeneous mixture is formed.
Homogeneous Mixture – Components are evenly distributed – Also
called solutions.
Heterogeneous Mixture – Components are unevenly distributed
(c) 2006, Mark Rosengarten
Solution - a homogeneous mixture of a solute dissolved in a solvent
Melting – a phase change in which a solid changes to a liquid
Boiling – the process of rapidly converting a liquid to its gaseous
(vapor) state, typically by heating the liquid to a temperature called
its boiling point.
Boiling Point - temperature at which the vapor pressure is equal to
the pressure of the gas above it.
Freezing - a phase change in which a liquid cools and changes to a
solid
Condensation - a phase change in which a gas cools and changes to
a liquid
Sublimation - When a solid can change directly into a gas skipping
the liquid phase
Evaporation – a phase change in which a liquid changes to a gas
Exothermic – process which releases energy causing the
temperature of its surroundings to increase
(c) 2006, Mark Rosengarten
Endothermic – process which absorbs energy causing the
temperature of its surroundings to decrease
Heat of Fusion – Amount of heat in Joules or KF required to melt 1
gram of ice to water
Heat of Vaporization - Amount of heat in Joules or KF required to
vaporize 1 gram of water to vapor
Metals – Found on the left side of the Periodic table – Are
malleable, ductile, lustrous, good conductors and form positive ions
Malleable – can be pounded into thin sheets
Ductile – can be stretched into wire
Luster - shiny
Nonmetals – Found on the right side of the Periodic table – Are
brittle, dull, poor conductors and form negative ions
Metalloids – have the properties of both metals and nonmetals –
found along the stair-step line
Ionization energy – amount of energy required to remove the most
loosely held electron in an atom(c)–2006,
Values
found on table S
Mark Rosengarten
Electronegativity – The attraction a nucleus has for electrons in a
bond – Values found on Table S – Fluorine has highest
Alkali metals – Group 1 metals - Most active metals, only found in
compounds in nature – Form +1 ions
Alkaline Earth Metals – Group 2 metals - Very active metals, only
found in compounds in nature – Form +2 ions
Transition Metals – Groups 3-11 - Many can form different
possible charges of ions - Compounds containing these metals can
be colored.
Halogens – Group 17 nonmetals – Most reactive nonmetals –
Fluorine most active
Noble Gases - Are completely nonreactive since they have eight
valence electrons, making a stable octet.
Diatomic Elements - Br2, I2, N2, Cl2, H2, O2 and F2
Ions - charged particles formed by the gain or loss of electrons.
(c) 2006, Mark Rosengarten
Positive Ion – Formed when an atom, usually a metal, loses 1 or
more electrons
Negative Ion – Formed when an atom, usually a nonmetal, gains 1
or more electrons.
Ionic bond – bond that forms when a metal transfers valence
electrons to a nonmetal.
Covalent bond – bond that forms when nonmetals share valence
electrons
Metallic Bond – Bond that forms between metal atoms such as in
copper wire – Described as “positve ions in a sea of mobile
electrons”
Ionic Compound - made of metal and nonmetal ions.
Molecular Compound - made of nonmetal atoms bonded to form a
distinct particle called a molecule.
REDOX Reaction – Short for oxidation-reduction - driven by the
loss (oxidation) and gain (reduction) of electrons.
(c) 2006, Mark Rosengarten
Oxidation – loss of electrons – oxidation number increases
Reduction – gain of electrons – oxidation number decreases
Precipitate – compound that forms as a result of a double
replacement reaction which is insoluble in water
Intermolecular Attractive Forces(IMAF) – force of attraction
between molecules such as hydrogen bonding, dipole-dipole, etc…
Hydrogen Bond – A special type of dipole-dipole attraction that
occurs when hydrogen is bonded to N, O or F.
Gram Formula Mass - sum of atomic masses of all elements in the
compound – equal to the mass of one mole of a compound
Catalyst – speeds up a chemical reaction by lowering the activation
energy.
Activation Energy – amount of energy needed to start a reaction
Heat of Reaction(DH) – amount of heat absorbed or released during
a chemical reaction
(c) 2006, Mark Rosengarten
Chemical Equilibrium – When the rate of the forward and reverse
reactions are equal
Solubility - the maximum quantity of solute that can be dissolved
in a given quantity of solvent at a given temperature
Arrhenius Acid - molecules that dissolve in water to produce H+ or
H3O+ (hydronium) as the only positively charged ion in solution.
Arrhenius Base - molecules that dissolve in water to produce OH(hydroxide) as the only negatively charged ion in solution.
Bronsted-Lowry Acid – proton (H+) donor
Bronsted-Lowry Base – proton (H+) acceptor
Voltaic cell - Produce electrical current using a spontaneous redox
reaction – used to make batteries
Electrolytic cell - Use electricity to force a nonspontaneous redox
reaction to take place.
Anode – electrode at which oxidation occurs
Cathode – electrode at which reduction occurs
(c) 2006, Mark Rosengarten
Salt bridge – allows for the movement of ions
Hydrocarbons - Molecules made of Hydrogen and Carbon
Alkanes – saturated hydrocarbons with only single bonds between
carbon atoms
Alkenes – unsaturated hydrocarbons with at least one double bond
between carbon atoms
Alkynes – unsaturated hydrocarbons with at least one triple bond
between carbon atoms
Esterification - reaction between an alcohol and organic acid
which produces an ester and water
Fermentation – reaction of a sugar with an enzyme that produces
alcohol and CO2
Polymerization – process of joining many small
molecules(monomers) to make a large molecule(polymer).
Saponification – A fat or oil reacts with a strong base and produces
a soap
(c) 2006, Mark Rosengarten
Isomer – compounds that have the same chemical formula but
different structures
(c) 2006, Mark Rosengarten