Study Guide for General
Chemistry I*Honors | CHEM
121
Chemistry
Liberty University
(LU)
7 pag.
Chemistry study guide
Chapter 1
1. Hypothesis- tentative explination
a. Theory- explination of general causes of certain phenomena, with
considerable evidence to support it
b. Law- statement or equation that summarizes a broad variety of
observations and experiences
•
• Element- cannot be decomposed into simpler sumbstances
compound- substance composed of two or more elements
mixtures- combonation of two or more substances in which each substance
retains its own identity
• Match 54 elements with symbols. Tell if element or not.
• Physical properties- can be observed w/o changing the identity and composition
of the substance.
-Color, odor, density, melting point, boiling point, hardness
Chemical properties- describe the way a substance may change or react to form
other substances
• Flammability
oPhysical changes- changes the physical appearance
oChemical change- transformed into a chemically different substance
▪Mass kg
▪Length m
▪Time s
▪Temp K
▪Amount of substance mol
▪Electric current A (ampere)
▪Luminous intensity cd (candela)
oGiga 10^9
oMega 10^6
oKilo 10^3
oDeci 10^-1
oMilli 10^-3
oMicro 10^-6
oNano 10^-9
oPico 10^-12
oFemto 10^-15
• Density- the amount of mass in a unit volume of the substance: D=m/v
• 1mL=1cm^3=1cc
• 1 g of water = 1mL @ 4 degress C
• Dimensional analysis- given unit X (desired unit/given unit) =desired unit
Chemistry study guide
• Precision- measure of how closely individual measurements agree with one
another
• Accuracy- how close the measurements agree with the “true” value
• Uncertainties always exist in measured quantities
• Zeros at beginning are never significant.
• Zeros between nonzeros are
• Zeros at the end of a number are ONLY if they are after a decimal
• Addition and subtraction: answer has same amount of sig figs as the number
with fewest decimal places
• Mult. and Div. : answer has same amount as the number with fewest sig figs
• Fahrenheit - C=(F - 32) * 5/9 K=(F - 32) * 5/9 + 273.15
• Celcius- F=(C * 9/5) + 32 K= C+273.15
Chapter 2
• Daltons postulates
oElements composed of atoms
oAtoms of one element are exactly identical but atoms from other
elements are not
oAtoms in one element cant be changed into atoms of other elements by
chemical reaction; atoms are neither created nor destroyed in chemical
reactions
oCompounds are formed when atoms of more than one element combine
• Radioactivity- spontaneous emission of radiation.
• Rutherford- discovered alpha and beta rays
oBeta- negaive…. 1-
oAlpha- positive….2+
• Atomic theory philosophers
oDemocritus
oIssac Newton-
oJohn Dalton
• Sizes
oProton= 1.0073 amu
oNeutron= 1.0087
oElectron= 5..486 X 10^-4
• Isotope- the different possible versions of each element
oKnow how to write the symbols for isotopes
▪Atomic number top left of symbol and mass number bottom right
• Molecular compound- contain more than one type of compound.
• Ionic compound- compound that contains both positively and negatively charged
ions
• Diatomic elements
oHelp Our Needy Class Find Brains Immediately
• Molecular formula- indicates the actual numbers and types of atoms in a
molecule
Chemistry study guide
• Empirical formula- gives only relative number of atoms if each type
Chapter 3
• Combination reactions- 2 or more substances combine to form one
• Decomposition reactions- one substance undergoes reaction to create 2 or more
substances
• Combustion reactions- rapid reactions that produce a flame.
• Percent composition- % element= [(#atoms)(atomic weight)/formula weight] X
100%
• Formula weight- sum of atomic weights
oAccount for subscripts
• Molar mass=formula weight
• To get empirical or molecular formulas:
oConvert percents to grams
oConvert grams to moles of each element
oDivide larger number of moles by smaller number to get ratio
oRound ratio to nearest whole number to get subscripts
• Read 3.6
• Limiting reactant- reactant that is completely consumed in a reaction
oRead top of 103
oFind # of moles of one element needed to react with given amount other
oSubtract # found from # given in problem of same element
oIf more than zero, its not the limiting reactant
• Percent yield: (actual yield/ theoretical yield) X 100%
oLook at problem on 106
o
Theoretical yield= Grams of A1⟶ mol A1⟶ mol A2 grams⟶ A2
oActual yield is given in problem
Chapter 4
• Solvent- Substance in greatest quantity in solution
• Solute- other substances in solution
• Electrolyte- substance whose aqueous solution contains ions
• Nonelectrolyte- substance that does not for ions in solution
• Strong electrolytes- solutes that exist in solutions as complete or nearly
complete ions
• Weak electrolytes-solutes that exist in solutions mostly in the form of molecules
with only a small fraction in the form of ions.
• Table on 125
• Precipitation reactions- result in insoluble product
oPrecipitate- insoluble product
• Acid-base reactions-
oAcids- ionize in aqueous solutions to form hydrogen ions
▪H+ becomes by itself
Chemistry study guide
oBases- accept(react with) H+ ions
▪Dissociate and add OH- ions
• Spectator ions- ions that appear in identical forms among the reactants and
products of a complete ionic equation(same on both sides)
• Complete ionic equation- all soluble strong electrolytes shown as ions
oEverything is shown in equation
• Net ionic equation- when spectator ions are omitted from the equation
oOnly things that change are shown
• Redox reactions- electrons are transferred between reactants (charges change)
oOxidation rxn- loss of electrons by a substance molecule becomes
positively charges)
oReduction rxn- when molecules become more negatively charged
• Oxidation number- actual charge for monatomic a ion
oAtom in elemental for, Ox #=0 (P, H)
oMonatomic ions (K+, S2-) Ox # is the charge of the ion (1, -2)
oNonmeatal- usually negative Ox #
▪Oxygen usually -2
▪Hydrogen- +! When bonded to nonmetals and -1 when bonded
to metals
▪Fluorine=-1. Halogens=-1
oThe sum of the OX #’s in a polyatomic ion equals the charge of the ion
• The higher the OX #, the higher the reactivity
• Molarity- number of moles of solute per liter
• Dilution- process of adding water
oMoles of solute before dilution=moles of solute after dilution\
oMolarity(original) X Volume (orginal)= Molarity (dilute) X volume
(dilute)
• Page 149-150
Chapter 5
• Energy- capacity to do work or transfer heat
• Energy used to cause an object with mass to move against force
• Joule- unit for energy
• Calorie- amount of energy need to raise the temp of 1 gram of water from 14.5 C
to 15.5 C
o1 cal =1.184 J
o1 Cal =1000 cal=1 kcal
• Work- w
• Force- F
• Distance- d
• Heat- q
• Potential energy- energy in position relative to other objects
• Kinetic energy- energy when in motion
• First law of thermodynamics: ENGERGY IS CONSERVED
Chemistry study guide
•
E= q +w
owhen heat is transferred TO the system q is positive
onegative when heat is transferred FROM the system
oWhen work is done on system by surroundings, w is positive
oWhen work done on surroundings by system, w is neg
oq:
ow
o
E
▪pos, gains heat. Neg, loses heat
▪pos- work done on system neg- work done by system
▪pos- net gain energy neg- net loss energy
• State function- property of a system that is determined by specifying the systems
condition or state
oInternal energy and enthalpy
▪E and H
• Enthalpy- thermodynamic function that accounts for heat flw in processes
occurring at constant pressure when no forms of work are performed besides P-
V work.
oH= E + PV
o
H= (E + PV)= E + PV=(q + w) - w= q
•
H= H(final) – H(initial)
• Calorimetry- measurement of heat flow
• Specific Heat (c)
o
q= m x c x T
• Hess’s law- if steps of enthalpy, H will equal the sum of the steps’ H
Chapter 6
•
Frequency()- s^-1
•
Wavelength ()- m
oc=v
▪c- speed of light
• Electromagnetic Spectrum (shortest to longest)
oGamma, x-rays, ultraviolet, infrared, microwave, radio
• Planck’s constant:
oE=hv
▪E- energy
▪h- plancks constant (6.626 X 10^-34)
• Photoelectric effect- light shining on a clean metal surface causes the surface to
omit electrons
• Transition energies with Bohr model:
oE=(-HcR)(1/n^2)
• Principal quantum number (n)
oPositive integral numbers
Chemistry study guide
oAs n increases to orbital becomes larger
oAs n increases so does energy
• Angular momentum quantum number ( l)
oIntegral values from 0 to (n-1)
oDefines shape of the orbital
oS,p,d,f… corresponds to values of 0,1,2,3
• Magnetic quantum number (m(l))
oIntegral values between – l and l
oDescribes orientation in space
• m(s)- spin
o+1/2 or -1/2
• Chart page 227
• Degenerate orbitals- orbitals with the same energy
• Pauli Exclusion Principle
oNo two electrons in an atom can have the same set of four quantum
numbers (n,l,m(l), m(s))
• Hunds rule:
oFor degenerate orbitals, the lowest energy is attained when the number
of electrons with the same spin is maximized
• Electron cinfigurations
• Spin makes nmr/mri possible
Chapter 7
• Mendeleev and meyer organized first periodic table by atomic weight
oTrends
oEffective nuclear charge: increases from right to left and going down
oSize of atoms and ions: increases left to right and going down
▪Anions are bigger than parent ions
oIonization energy: in rows-increases left to right and decreases going
down
oGroups
oAlkali metals: 1A
oAlkaline earth metals: 2A
oOxygen group: 6A
oHalogens: 7A
oThe noble gases: 8A
Chapter 8
• Ionic bond- b/w ions of opposite charge
• Covalent bonds- share electrons
• Lattice energy-energy required to completely separate a mole of a solid
• Electronegativity: increases left to right and going up
Chemistry study guide
o=0, nonpolar
o= small number; polar
o=largerer number; ionic
• Exceptions to the octet rule:
oOdd # of electrons
oFewer than octet of valence electrons
oMore than an octet of valence electrons
• Pi bond- one in every bond
• Sigma bond- in multiple bonds
• Formal charge: sum of the formal charges always equals overall charge of ion
• Bond order: ½(# of bonding electrons - # of antibonding electrons
Ch. 9
• Dipole- when a distance separates two electrical charges of equal magnitude but
opposite signs
•
Ch. 10
• Ideal Gas Law: PV=nRT
Name Expressio
n
Constant
Boyle's law PV=PV
Charles's law V/T=V/T
Avogadro's
law
V/n=V/n
• nM/V=PM/RT=below equation
• d=PM/RT
oM=molar mass
• Partial Pressure:
oP= P1+P2+P3….
