Honors Chemistry Study Guide
Measurements and Calculations/Lab skills
 convert mass to moles:
grams of the substance = molar mass of the substance (grams)
moles of the substance
one mole
 AMU = grams (they’re the same thing)
 the atomic mass of an element is equal to one mole of the element
 Significant figures
o Measured numbers tell what was measured and the precision it was measured at
o Sig figs show the precision to which something was measured
o They are the digits in a number that you actually know
o Zero- placeholder when before decimal, show precision after decimal
o Zeroes after decimal don’t count as sig figs (think of scientific notation)
o The number with the fewest figures in the decimal point decides how many sig figs you
can have in your answer
 Flame test- burn an element to determine what it is (depending on the flame’s color)
 Bright line spectra- look through spectrum at gas in cathode ray to see what lines are visible,
compare lines to key to determine which gas is being viewed
 Density- mass/volume
 Percent error
o (accepted value - experimental value) \ accepted value x 100%
o Tells the difference between results obtained and accepted results- how accurate the
results of your experiment are
o Can also be used with theoretical values
Atomic Structure
 Dalton, Thomson, Rutherford, Bohr
o Dalton- atom is a circle
o Thomson(cathode ray)- negatively charged electrons distributed randomly in a circular
atom (plum pudding model)
o Rutherford(gold foil)- central positive nucleus with negative electrons floating around,
most of the atom is empty space
o Bohr- all parts are solid orbiting a central nucleus, orbits are energy levels/orbitals
 Electron cloud models of the atom (quantum)
o Impossible to know momentum and position of an electron at one time
o Certain regions where an electron is most likely to be found
o Most recent model
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Protons, neutrons, and electrons
o Protons- positive, located in nucleus, same as atomic number,
o Neutrons- neutral, located in nucleus, mass - # protons = # neutrons,
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Electrons- negative, travel around atom, normally the same number of protons but can
be lost or gained
Average Atomic mass & isotopes
o The atomic mass of an element is the average of the masses of all the isotopes (same
atom with different number of neutrons)
Lewis dot diagram- shows the valence electrons in an atom
o Bonds are shared electrons and represented by dashes
o The goal is to have a full octet
o If the atom has a charge, the diagram is put into square brackets with the charge on the
top right
Wave mechanical molecule of the atom
o Electrons are standing waves produced only at certain energy levels
o Standing wave- energy in an electron wave
Orbital filling rules
o Aufbau- in the ground state of an atom, an electron enters the orbital with lowest
energy first and subsequent electrons are fed in the order of increasing energies
o Pauli exclusion principle- a single orbital, regardless of size or shape, may contain no
more than 2 electrons, and those electrons must have opposite spin
o Hund’s rule- when filling a set of equal energy orbitals each orbital gets one electron
before any orbital gets 2
Orbital notation
o Written as: number of shell, letter of orbital, number of electrons in that orbital
o Ex. For fluorine its 1s1 2s2 2p5
Orbitals- each orbital space holds 2 electrons(lowest enegry)
o Electrons are removed starting at the highest energy level- highest level electrons
contain the most energy and are therefore easiest to remove
o All electrons in outermost shell gone- large amount of energy needed to remove from
second highest level (full octet)
o S orbital- fits 2 electrons in each of 2 parts, first row of periodic table, spherical pattern
of standing waves around the nucleus,
o P orbital- holds 2 electrons in each of 3 parts (6 total), second row
o D orbital- found in 4th row of table, 5 parts for a total of 10 electron spaces, requires lots
of energy
o F orbital- doesn’t need to be filled as badly as s and p
o This chart can be used to determine the number of orbitals used by drawing a diagonal
line from right to left until all electrons are used (can be used for almost all elements)
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5p 5d 5f
6s 6p 6d
7s 7p
Orbital Hybridization
o Combination of orbitals- interact with each other to be as far away as possible, merge
into identical hybridizations of the two
o 1 s orbital + 3 p orbitals  sp3 hybridization (4 identical sp combinations)
Periodic Table
 Properties and location of metals, nonmetals, and metalloids
Place in periodic
Activity
table
Alkali
First and second
Very reactiveMetals and left columns
radioactivity
Alkaline
increases from top
Earth
to bottom (group 1
Metals
is more reactive
than group 2)
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Transition
Metals
Halogens
Middle
Fairly unreactive
Second right
Highly reactive
(with both types of
alkali metals)
Metalloids
Between
transition and
halogens
Noble Gases
Far right
Unreactive
Properties
Easily lose electrons,
always found naturally
in compounds, achieve
stability by losing
electrons, good
conductors
Malleable, high melting
and boiling points
Almost all brittle solids
at room temp, have
both metal and
nonmetal properties
Complete valence
electron shell, low
boiling points, gas at
room temp
Periodic trends
o Atomic radius
 Half the distance between the nuclei of two atoms bonded together- size of the
atom
 Across the table- radius gets smaller (valence elecrons have the same energy
levels, but # of protons attracting them increases, causing a greater pull to the
center)
 Down a group- radius increases (more inner level electrons, which shield the
attractive forces of the nucleus and allow valence electrons to be more
energetic)
o Ionization energy
 amount of energy needed to remove most loosely bound atom
 across the table- increasing ionization energies (number of protons increases so
the charge/pull of the nucleus increases, electrons are more strongly attracted
to the nucleus and harder to take off)
 down a group- ionization energy decreases (valence electrons are at a higher
electron level- farther from the nucleus- less pull on the electrons)
o Electronegativity
 Measure of its attraction for electrons when bonded to another atom (opposite
of ionization energy)
 Across the table- electronegativity increases
 Down a group- electronegativity decreases
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Ionic radius
Metallic character
Bonding
 Bonds are made to get a full octet
 Opposite forces attract, but if they get too close the nuclei repel and they move to a safe
distance again
 Bond length- the distance between two nuclei at the point of minimum energy (forces cancelled
out)
 Octet rule- full electron shell (8 electrons) is the most stable number, all atoms try to become
stable
 Ionic bond- bond formed by transferring electrons from one atom to another (between a
positive one and a negative one)
 The two atoms in an ionic bond have a high electronegativity difference
 Lewis dot structures for ionic bonding
o Unbonded valence electrons are dots
o Remember to add charges by putting structure in square brackets and adding charge to
top right
 Naming ionic compounds
o First chemical (cation) is normal, second (anion) has –ide added to end
 Sodium chloride
o Roman numerals show the charge of an ion (usually a transition metal)
 Iron(II) chloride
 Properties of ionic, covalent, and metallic compounds
o Ionic compounds are often hard crystalline (crystal lattice) solids, water soluble, can
conduct electricity when dissolved in water
o Covalent compounds- more often softer solids, liquids, or gases, mostly not water
soluble, but even when they are, solutions do not conduct electricity
 Polar covalent bond- uneven sharing of electrons between nonmetals or metalloids
 Lewis dot structures for covalent molecules
o Find total amount of electrons in structure
o Create bonds to share electrons
o Find out how they can share to create as many octets as possible
o Not enough electrons to share normally- share more! (form a double or even triple
bond)
 Nonpolar covalent bond- two atoms have identical/very similar electronegativities (evenly
distributed charges)
 Covalent Bond polarity
o If the two atoms in a bond have very different electronegativities, the electrons in the
bond are attracted to the atom with the higher electronegativity (polar covalent bond)
o That atom becomes slightly negative and the other becomes slightly positive- polarity
(separation of charges)
 Covalent Molecule polarity
o Polar and nonpolar molecules do not mix
o polar molecules have a partial negative charge on the higher EN side and a partical
positive charge on the other
o nonpolar molecules are symmetrical while polar molecules are not- if it has partial
positive/negative charges but still has symmetry, it is nonpolar
o
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a molecule is polar if it has a dipole moment (separated into negative and positive sides)
represents the positive side and represents the negative side, with an arrow
across pointing to the negative side with a plus on the positive side
o Many polar molecules (such as water) line up positive to negative sides
o Mixing polar with nonpolar (water and oil) creates a separation & heterogeneous
mixture
 Lewis diagrams that include atoms that break the octet when bonding
 Predict shapes of molecules formed for up to 4 electron domains
Matter and Change
 Types of matter
o Pure substances
 Only homogenous mixtures- composition is the same throughout the sample
(constant composition)
o Mixtures
 Can be heterogeneous or homogenous
 Combination of two or more pure substances
o Elements
 Cannot be broken down or decomposed into simpler substances
o Compounds
 Combination of two or more elements
o Heterogonous
 A mixture that is not even distributed
o Homogenous
 A material that is evenly mixed throughout (uniform distribution)
 Separation of mixtures
o Filtration
 Use a filter to trap small particles while letting liquid/dissolved particles through
o Distillation
 Boil off the liquid and collect the resulting gas so you can separate a dissolved
solution
o Evaporation
 Boil off the liquid- lose the liquid but keep the thing that was inside it
o Chromatography
 Used to separate ink- water moves up the paper to separate the colors in the
ink
 Chemical vs. physical changes and properties
o physical changes are in appearance only
 boiling and meting are physical changes
o chemical changes affect the chemical composition of the substance
 chemical changes include burning and chemical reactions
 Types of reactions
o Single replacement
 One element replaces another element in a reaction
 A + BX  B + AX
o Double replacement
 Two soluble ionic compounds react and make a precipitate/gas/molecular
compound
 AB + CD  AD + CB
Synthesis
 Two or more reactants combine to form a single product
 A + B  AB
o Decomposition
 Opposite of synthesis, single product broken down into two or more simpler
substances
 AB  A + B
o Combustion
 Oxygen will combine to create carbon monoxide/dioxide and water
Chemical quantities/stoichiometry
 interpreting chemical formulas
 gram to mole to atoms conversions
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calculating GFM (gram formula mass) for compounds
o GFM = mass of one mole of an element
o To find GFM, multiply the number of atoms by the element’s mass
percent composition to empirical formula to molecular formula conversions
1. Convert percents to numbers out of 100g.
2. Divide each number by its atomic mass.
3. Divide all numbers by the smallest numbers.
4. Round to whole numbers.
5. Create a ratio to find the empirical formula.
6. Find the mass of the empirical formula.
7. Divide the given weight by mass of the empirical formula.
8. Multiply the subscripts by this answer.
Hydrates
o Crystals that contain attached water molecules
balancing chemical equations
o sum of the products must equal the sum of the reactants
stoichiometry
o Measuring chemicals that go into and come out of any given reaction
o Allows atoms and molecules to be counted by weighing them
o Atomic mass is measured in amu
o Mole = atomic mass of something in terms of grams
o 1 amu = 6.02 x 10^23 moles
o The mass of the mole depends on the mass of the element
o Both C and O contain 6.02x10^23 moles, but C weighs 12 g and O weighs 16g because it
is more massive
o 1 mole of C = 12.01 g
o Convert amu to moles  1 mole = whatever the mass of the element is
o
Find the mass of 1 mole of a substance: multiply amount of moles by mass, add up all
masses
o Equation with reactants and products- must be balanced
1. Convert balanced equation to molar masses.
2. Create ratio of reactants and compare to ratio of products
342.296/384 = 5/x
3. Calculate!
o Limiting reactant- one that runs out first and causes reaction to stop
o Excess reactant- left over stuff
o What is the max I can make? Which product will cause it to stop?
Phases and intermolecular forces
 Intermolecular forces- substances with higher IMF have lower boiling points, while substances
with weaker IMF have more gas in the air above the, and more particles vaporizing
 IMF are the reasons why liquids and solids exist- hold the parts of a solid together
o Hydrogen bonding (h-boding)
 Special type of dipole dipole force between hydrogen and strongly
electronegative elements(like N, O, F) in polar molecules
 Very strong dipole attraction (strongest to form between molecules)
 Temperature of water is changed by breaking/forming hydrogen bonds
o Dipole-dipole
 Attractions between the opposite partial charges in polar molecules
 Molecules orient themselves so the attraction is maximized, repulsion is
minimized
o London dispersion
 Attractions between noble gases/nonpolar molecules
 Weakest, based on temporary clustering of electrons
 Sometimes electrons become clustered together while moving randomly around
their atom, this place becomes slightly negative and is strong enough to attract
a neighboring molecule
 Very weak and don’t last a long time because the electron cluster only lasts a
short time
 Only thing that lets nonpolar substances stick together well enough to condense
from gas to liquid
 Vapor pressure & boiling point
o Boiling point- vapor pressure of liquid equals atmospheric pressure
o Vapor pressure- pressure created by gas particles hitting the surface of a liquid,
increases with temperature
 Heating/cooling curves
o Heat added at a regular rate (joules) causes steady heat increase of ice
o Reach melting point- solid begins to melt & curve becomes straight
o The larger the mass of something, the longer it will take to melt
o Solid and liquid together at the same time, amount of solid goes down and liquid goes
up as temp increases
o Bubbles in liquid phase are squashed down by atmospheric pressure
o As heat increases, liquids vapor pressure increases and bubbles are more free
o Temperature remains constant during a phase change (even though you’re still adding
heat)
o Heat of fusion- amount of heat you add to a solid to make it liquid
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o Heat of vaporization- heat added to boil off the liquid
o diagonal parts of curve- while temp is changing, kinetic energy goes up
o kinetic energy- related to movement of particles, measured in degrees Celsius
o substances at the same temperature have the same kinetic energy
o Potential energy is the straight part- joules is a unit of potential energy
o Potential energy is related to a substances structure
o Cooling curve is the reverse of heating curve
o Heat flows from hot place to low place- this is why water freezes in the freezer
o More vaporization creates a cooler final temperature, causing the liquid to feel cooler
Formulas
o Q = mCΔT: used when the temperature of a substance is changing
o When 25g of water are cooled from 20 to 10 degrees, the number of joules of heat
energy released is: 1050j  25 (grams of water) * 4.18 (specific heat of water) * 10
(change in degrees of heat)
o Q = mHf : used when there is a change of state between solid and liquid
o Q = mHv : used when there is a change of state between liquid and gas
Videos!
Unit conversion and sig figs
http://www.youtube.com/watch?v=hQpQ0hxVNTg&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&ind
ex=2
periodic table
http://www.youtube.com/watch?v=0RRVV4Diomg&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&ind
ex=4
the electron
http://www.youtube.com/watch?v=rcKilE9CdaA&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&index
=5
stoichiometry
http://www.youtube.com/watch?v=UL1jmJaUkaQ&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&ind
ex=6
bonding
http://www.youtube.com/watch?v=QXT4OVM4vXI&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&ind
ex=22
polar/non polar molecules
http://www.youtube.com/watch?v=PVL24HAesnc&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&inde
x=23
lewis dot diagrams
http://www.youtube.com/watch?v=a8LF7JEb0IA&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&index
=24
orbitals
http://www.youtube.com/watch?v=cPDptc0wUYI&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&inde
x=25
liquids (London dispersion, dipole dipole, hydrogen bonding, heating curves)
http://www.youtube.com/watch?v=BqQJPCdmIp8&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&ind
ex=26
atom models
http://www.youtube.com/watch?v=thnDxFdkzZs&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&index
=38