Chapter 1:
Metric System:
1. internationaly accepted system of measurement
(SI)
2. based on 10
3. categories of measurement
1. mass vs. weight
1. mass-remains
constant, amount matter something has, kilogram
2. weight-changes
with change in gravity, newtons (kg/m/sec2)
2. length--meter
3. volume-cubic (cm3), space inside
container, liter (L)
4. Metric Prefixes
mega-M, 106
kilo-K, 103 deka-D, 101 deci-d, 10-1
centi-c,10-2
milli-m, 10-3
micro-u(mew), 10-6 nano-n, 10-9 pico-p,
10-12
Measurement and Significant Figures:
1. 1 place beyon calibration
2. rules for significant figures apply only to measured
#s
1. exact #-constant (ex.5,280ft)
2. measured #-produced by instruments
that measure
a. all non-zero digits are significant
b. zeros between non-zerp #s are significant
c. zeros to the right of non-zero but to the left
of an understood decimal
point are not significant(if decimal
point is after#, then significant)
d. in #s <1 & >-1 zeros to right of decimal
point & left of non-zero, not sig.
e. in #s <1 & >-1 the zero to the left of
the decimal point is not significant
f. all zero to right of decimal point & right
of non-zero are significant
3. Rules of Rounding
1. # to right >5 round up
2. # to right <5 stay same
3. #=5 (nothing after) round to
closest even (ex. 3.5=4 and 4.5=4)
4. #=5 (non-zero after) round
up
4. Math operations w/significant figures
1. addition & subtraction=
left most uncertain figure determines where
to round (ex.
123.4+1.45=124.9)-not necessarily lowest # of sig. fig
2.multiply & divide= smallest
number of significant figures in problem
determine
# in answer
Dimensional Analysis (Factor Label Method):
1. problem solving method
2. uses unit factors (any 2 terms describing same
amt of something)
Chapter Summary:
chemistry-study of matter, property, & changes
mass vs weight
chem properties-matter changes
phys properties-changes w/o changes in composition
extensive property-depends on amount
intensive peoperty-doesn't depend on amount
chemical change-change type of material had
physical change-change description
energy-ability to cause change(anything not matter)
kinetic energy-motion
potential energy-stationary
mixture-homogenous-solution
heterogenous-different(can see differences)
substance-can not be purified any mote by physical
means
compound-divided into simler @ fixed preportions
element-purest form, simplest matter
Chapter 2: Chemical Formulas & Composition Stoichiometry
400 BC Democritus-atoms make up matter, tiny indivisible
spheres
1800s John Dalton:Atomic Theory (pg.42)
1.elements made of
indivisible atoms (false-protons, neutrons, electrons
2.atoms of element
same (false-isotopes)
3.not created, destroyed,
changed (false-nuclear reactions, synthetic elements)
4.compounds-combinations
in small, whole # ratios (true)
5. relative #/kind
in compounds constant (true)
atoms-smallest part of element keeping chemical
& physical properties
made of 1. protons: in nucleus, positive
charge (atomic number)
2. neutrons: in nucleus, no charge
3. electrons: around nucleus, negative charge
molecule-smallest part of element/compound with
stable, independent esitence
diatomic- O2,N2,H2,F2,Cl2,Br2,I2
polyatomic-P4, S8
Chemical Formulas-composition & ratios of elements
Law of Constant Composition(Def.Prop)-diff
pure samples of a comp always contain same
elements in proportion by mass fixed
molecluar compounds-formed from
molecules that exist individualy
Ionic compound-made of charged
ions(+&-), can exist seperately, formula units
ion-(table2-3) group of sigle atom w/charge (want
to be stable)
cation:+ charge (metals lose -e)
anions:-charge (non-metals gain -e)
Forming Compound: Na+ + Cl- = NaCl (no charge)
1. cation 1st 2. anion
2nd 3. ions in ratio (no net charge)
Bionary Ionic- 2 elements
Complex Ionic-Polyatomic ion in it
Naming:1. ID ions 2. cation 1st 3. binary-ending
-ide, complex-ending Polyatomic
Atomic Weights:
1. atomic mass units (amu): 1/12 of mass of Carbon
12 atom
2. standard for masses on periodic table
3. decimal mass->weighted average of the isotopes
Mole:
1. amt of substance that contains as many items
as there are in exactly 12 grams C12 atom
2. avogardo's #: 1 mole=6.0221367 x 1023 particles
3. mass of 1 mole of atoms of pure element in grams
= atomic mass, molar mass(g/mol)
Molar Mass: the amount of substance that contains
the mass (g) numerically equal to the
formula weight of substance, Avogadro's # of formula units
Percent Composition:express % of each element in
compound
Empirical Formula:smallest whole # ratio of atoms
present in compound
1. relative mass of element
2. relative # of atoms (divide
by atomic weight)-convert to moles
3. divide by smallest # of atoms
4. convert fractions to integers
5. write formula w/whole # subscripts
Molecular Formula:the true formula of a compound
that is either the same as or an integer
muliple
of the simplest formula
n=molecluar wt divided by simplest
formula weight
Law of Mulitiple Proportions:when two elements A&B
form more than 1 copmound, the ratio of the masses
of B that combine with a given mass A can be
expressed by small whole #
Chapter 3: Chemical Equations & Reaction Stoichiometry
3-1 Chemical Equations:
Reactants----> Products (Look
at states of matter, and balancing equations)
Law of Conservation of Matter: no matter is created
or destroyed, only changes form
coefficients- placed in front to balance(preserve
Law of Conservation of Matter)
can mean moles, molecules, or atoms, simplify if possible
3-2 Calculations & Chemical Reactions:
Reaction Stochiometry-relationship based on conversion
factors, uses mole ratio from
balanced equation,4 types=mole-mole,mole-mass,mass-mass,mass-mole
3-3 Limiting Reactant:
limiting reactant-substance that gets used up first,
determines amt of product
A. Determine Limiting Reactant-convert all to moles(use
all of 1, how much other is needed)
compare the one solved for w/
amt have...one that runs out is limiting
B. determine amt of product produced based on limiting
3-4 Percent Yield:
% yield = actual product yield x 100
theoretical yield
3-5 Sequential Reactions:
more than one reaction required to obtain the product
amt of product from 1st reaction is starting material
for 2nd reaction
3-6 Concentrations of Solutions:
Solution-solute & solvent
% by mass..... % solute=mass of solute / mass of
solution x 100
molarity-tells concentration of solutes in soln...
M=mol solute/L soln
3-7 Dilution of Solution:
M x V=amt of solute(mol)......M1 x V1 = M2 x V2
(for dilution)
3-9 Titrations:
controlled addition of one substance to another
*measure V of substance required to react w/given
V of another soln w/known concentration
calculate concentration of 1st soln
need: standard soln-know exact amt
buret-glass
long test tube w/measurements & stopper to add drop by drop
indicator-let
know when stuff reacts (phenothalynine)
looking for equivalence point:stochiometrically
equivalent amts of acid and base
visualized by
end point:color change of indicator
Chapter 4: The Periodic Table
4-1 The Periodic Table
Dimitri Mendeleev-made periodic table
Myer-made other periodic table
Mosley-Properties of elements are functions of atomic
number (Periodic Law)
columns-groups/families--18
rows-periods--7
Alkali & Alkaline Earth (groups 1 & 2)-most
active
transition elements-heavy metals
metalloids-properties in between metals & nonmetals
metals->metalloids->nonmetals
metalic character increases down & left of periodic
table
fewer electrons in outer shell(1
or 2)
form cations
formation of metallic bonds (shared
electrons)-- metal + metal
formation of ionic bonds-- metal
+ nonmetal
nonmetalic character increases up & right of
periodic table
more electrons in outer shell(4+)
form anions
formation of covalent bonds(molecular
compounds)--nonmetal + nonmetal
formation of ionic bonds--
nonmetal + metal
4-2 Intro to Aqueous Solutions:
solution-solute (being dissolved) & solvent
(dissolving- usually H20)
electrolytes-substances aqueous solutions conduct
a current (ions), current conducted by
movements of ions: strength = ions present + ions charge
dissassociation-breaking
apart ionic compound into ions
ionization-breaking
apart molecular compound into ions
strong-conduct a current, ionize
almost completely in aqueous soln
1. strong
acids-acids produce H+ ions is aqueous soln
2. stron
soluable bases-bases produce OH- ions in aqueous soln
3. soluable
salts-salts have cation other than H+ and anion other than OH-
weak-conduct slightly
non-don't conduct
strong & weak acids-common weak acids ionize
slightly
reversible reactions-occur in both directions -->
& <--, usually weak acids
strong soluable bases, insoluable bases, weak bases-
strong-dissociate completely,
group 1 or 2 metal + OH- ion
insoluable-transition metals,
not soluable in water
weak-ionize only slightly but
soluable
solubility rules-pg 124-125
4-3 Reactions in Aqueous Solutions:
3 ways to write reactions:
1. Formula Unit-show complete formula
2. Total Ionic Equation-show predominan form of
each substance in aqueous soln
strong
acid, strong soluable base, soluable salt written in ionic form
3. Net Ionic Equation-show only species that react,
eliminate spectator ions
4-4 Percipitation Reactions:
reaction that results in formation of insoluable
solid, replacement reaction--cations switch
4-5 Acid Base Reactions:
neutralization reactions: Acid + Base = Salt + H2O
strong acid + strong soluable base: H+(aq) + OH-(aq)
--> H2O(l)
weak monoprotic acid + strong soluable base: HA(aq)
+ OH- --> A-(aq) + H2O(l)
4-6 Oxidation Numbers:
keep track of electrons in redox reactions
trasfer of electrons from 1 species to another (charges)
Binary Compounds:
ionic: oxidation # equal to charge
molecular: ox # doesnt carry same
significance, figure our which is more metallic in
character, some nonmetals have more than one charge
ox # of free uncombined element
= 0
ox # of monatomic ion= its charge
sum of all ox # in compound =
0
sum of all ox # in polyatomic
ion = polyatomic ions charge
4-7 Redox Reaction:
oxidation-algebraic increase in ox# corresponds
to loss (apparent loss) of electrons
reduction-algebraic decrease in ox# corresponds
to gain(apparent gain) of electrons
redox reaction-general term for reactions in which
1 substance oxidized while other reduced
oxidizing agent-oxidize other substance,is reduced,
gain electron, decrease ox #
reducing agent-reduce other substance, is oxidized,
lose elctrons, increase ox #
disproportion reaction-redox reaction where same
element is reduced & oxidized
4-8 Displacement Reactions:
displacement-something moving/switiching places
1. double displacement--acid-base reactions,precipitate
reactions
--two elements displace each other fromm compound
2. single displacement--one element displaces another
from a compound
--redox reactions
-- A + BC ---> AC + B (metal +ionic= ioonic +metal)
--more active metal replaces less activve metal
a.more active metal + salt(less
active metal)-->less active metal+salt(more active metal)
b.active metal + nonoxidizing
agent--> hydrogen + salt of acid
c.active nonmetal + salt(less
active metal)-->less active nonmetal + salt(more active)
F2 > Cl2 > Br2 > I2 X2 + YZ = YX + Z2
4-9 Binary Compounds (2 Elements):
more metalic element + unambiguous stem of less
metalic and -ide
1. metals w/one oxidation #
2. metals w/more than one oxidation #
more metallic
(Roman Numeral) then stem of less metalic + "ide"
3. old system: suffix-ous (lower #), -ic (higher
#)...use latin name for metal
4-10 Tenery Acids & their Salts
tenery acid: H, O (diff #), Nonmetal
original # of O---"ic" acid
ionic compound ends in: "ate"
1 more than "ic"---"perstemic" acid
"per...ate"
1 less than "ic"---"ous" acid
"ite"
2 less than "ic"---"hypostemic" acid
"hypo...ite"
ionic compound ends in "ide"---"hydrostemic" acid
Chapter 5: The Atom
A. Democtritus-things made of atoms-concept
B. Dalton-1800s,atom is indivisible, uniform
C. Thompson-current(Fareday),cathode ray tube experiment-electron
exists
1.)beam negative charge 2.)straight
line-shadow 3.)charge plates-toward postive
4.)magnet-opposites attract
5.)has mass-paddle wheel
D. Millikan-oil drop experiment(suspend oil drop-determine
mass, volume,density),mass elec
E. Goldstein-proton exists(cathode ray tube),positive
beam,equal,opposite charge(1836xmass)
F. Rutherford-gold foil experiment,alpha source(Po),scintillation
screen(bombardment)
most go straight through(expected),some
moderated difflections(some surprise), 1 in
8000 bonce back(surprised)-wide
angle difflections,there's something postive,dense,small
nucleus-tiny,positive charged,dense...w/o
proof said there are neutrons
G. Moseley-wavelength and atomic #,as +charge increase,
at.# increase
H. Chadwick-neutrons-neutral, mass in nucleus
mass # = decimal number on periodic table (p+ +
no)
isotope- H(protium,deuterium,tritium), same element,#protons,#electrons,diferent#
neutrons
nueclide terms
atomic mass=1mole=molar mass=avogadros#atoms=weighted
ave of isotopes by %
Electromagnetic Radiation-atomic emmision spectra
-ground state(neutral atom)
-excited state(excited state-short
time-give off radiation-goes back to ground)
wavelength-distance between equal
points on two waves,measured in nm or anstrom
frequency-# crests that pass a
point in 1 amt of time, measured in crests/sec
relationship speed of light=wavelength(frequency)
Visual Spectrum:
violet:short wavelength,high frequency
red:long wavelength,low frequency
Light as Particles:
-each photon has a quantum(measured
amt) of energy
-energy of photon proportional
to frequency
-energy decreases, frequency decreases,
increases wavelength
-Energy=Planck's constant(frequency)
Photoelectric effect-light stikes cathode, remove
electron(stream of electrons=current)
1.) electron ejects when short
wavelength
2.) current
Old Theory-if energy increases, current increase--photoelectric
doesnt support
Einstein:(1905)Nobel Prize1921
1.) extended Planck's particle
theory
2.) each photon transfers energy
to 1 electron during a collision
3.) # photons per unit time= intensity
Neils Bohr:
1.) certain lines in the emmision
spectra of atoms, therefore electrons must be able to
jump to
certain places(energylevels), closer to nucleus=lowerevergy orbit
2.) electrons in certain orbits,
absorb a certain amt of energy
3.) when electron moves to a higher
energy, it cant stay there and returns to ground
emits
a certain amount of energy (photon)
4.) model only worked for atoms
with one electron in outer level
Louis de Broglie:
1.) light has both wave and particle
properties
2.) particles w/mass and velocity
have a wavelength associated w/it
wavelength
= Planck's constant divided by mass times velocity
Section 5-14: Quantum Mechanics-electrons as particles
with wave like properties
"where fo the electrons go?"--explain
behavior(making a whirlpool)->destructive interference
quantized
energy
Basic Quantum Mechanic Ideas:
1.) atoms exist in certain energy states
2.) radiation emitted= change in energy
3.) quantum #s define energy states
Section 5-15: Quantum Numbers:
atomic orbitals-probable electron locations
Werner Heisenburg Uncertainty Principle-impossible
to determine the electrons momentum and
possition at same time
1.) Principle QN- (n), energy level of electron
(ex. 1,2,3 etc)
2.) Subsidiary QN- (l), sublevel (ex. 0->s, 1->p,
2->d, 3->f), max value= n-1
3.) Magnetic ON- (ml), orientation of sublevel ml=-l
-> +l,
4.) Spin QN- (ms), spin of electron-values +1/2
or -1/2 therefore no more than e electrons per
orientation
Section 5-16:Atomic Orbitals
s-> 1 orientation, p->3, d->5, f->7
Section 5-17: Electron Configuration
3 Rules:1.)Aufbau Principle-"building up", electron
fill lowest energy orbitals first
2.)Pauli's Exclsio Principle-exclusive, no 2 electrons can have same 4
QN
(difference=spin)
3.)Hund's Rule-electron's must occupy all orient of a given sublevel equal
before pairing
1.) Orbital Notation-- has lines for each orientation
and arrows representing electrons
with spins
2.) Electon Conf Notation-- energy level, sublevel,
electons...ex. 1s2
3.) Nobel Gas Config-- [preceeding nobel gas] electron
config for rest
Definitions:paramagnetism-element w/unpaired electrons,weak
attraction to magnetic fields
dimagnetism-element w/no unpaired electrons,weak repulsion from magnetic
field
Section 5-18:
period on table= n
1st 2 columns- "s" block,group#= how many electrons
in last level, column 1: ns1, 2:[ ]ns2
transitional- "d" block, group#= sum of electron
in s & d, [ ](n-1)d1-10ns1-2
group 13-18- "p" block, group#= sum of electron
in s & p +10, [ ]ns2np1-6
# electrons in outer energy level = group number
or group #-10 (group # > than 10)
d block exceptions to the trend (n-1)d1-10,ns2:
group 6: (n-1)d5ns1 and group 11: (n-1)d10ns1........due
to stability
Chapter 6: More with Periodic Table
Trends:
group 1--s block--ns1
\
group 2--s block--ns2
- representative (main) elements
group 13-18--p block--ns2np1-6
/
group 3-12--d block--ns2(n-1)d1-10--trannsition
elements
Lanthanide&Actinide Series--f-block
group 18--p block--ns2np6--Nobel Gas-->filled outer
energy level->inert(stable)
Periodicity-elements in similar gorups have similar
properties (function of atm. #)
Section 6-2 Atomic Radii
-radius=1/2 distance between nuclei of adjacent
atoms
-Effective Nuclear Charge--affects atomic radius
size (Zeff)
-attraction of electrons for nucleus
-degree determined by inner electrons
(outer levels shielded by inner)
-Trends: group-radius size increase down group(electrons
in highest energy level)
period-radius size decreases across period(increase Eff.Charge,ncule more+)
Section 6-3 Ionization Energy:
-formation of cation (minimum amt of energy required
to remove electron)
-1st I.E. needed to remove 1st electron...2nd I.E.
to remove second electron
-1st I.E. larger than second I.E. because of more
positive character in nucleus
-Trends: group-IE decreases down group (e- furthest
from nucl, less effect of nucl)
period-increases across period(easier to gain e-,nucleus larger, larger
eff charge)
-Group 2 > Group 1 (increase positive character,
outer e- held tightly)
-Group 13- low in comparison(ns2np1-easy to remove
therefore low IE)
-Group 15> Group 16 (half filled sublevel=more stable..higher
IE)
Section 6-4 Electron Affinity:
-amt of energy required to gain e- and create in
w/ -1 charge(anion)
-positive EA=heat aborbed, forced to gain electrons
-negative EA=heat release, want to gain electrons
-Trends: group-no real trend (except group 2 larger
and group 15 less than expected)
period-increases in negative value across period(gain e- to become nobel
gas)
Section 6-5 Ionic Radii:
-half distance between nuclei of adjacent ions
-isoelectronic different elements w/ same # of e-
-Trends: group-size increases down group (more e-,
higher energy levels)
period-size decreases across period (lose e-, increase positive character)
Section 6-6 Electronegativity:
-Linus Pauling--researched trends of electronegativity
-the relevant measure of the tendency of an atom
to attract e- to itself when chemicaly
combined with other atoms
-elements w/ high EN gain e- therefore form anions
-elements w/ low EN lose e- therefore form cations
-Trends: group-decreases down group (more shielding,
electrons easier to remove)
period-increases across period (want to gain e-)
-Pauling Scale--arbitrary values, Fluorine=4.0 (most
electronegativity)
-Predictions about Bonding:
-greater difference in EN values
for 2 elements, more ionic bond is, higher EN
element
accepts e- from less EN element
-2 nonmetals w/similar ENs form
covalent bonds, share e-
Section 6-7 Hydrogen and Hydrides:
-Hydrogen
-Cavendish 1766-steam through
gun barrel, acid & active metal, decompostion H20
-Reactions of Hydrogens and hydrides
-hydrides--binary compounds of
H+ and a metal
-ionic--active
metal + H-, group 1 or 2 metal, form basic soln in H20
-molecular--nonmetal
+ H+, form halides (group 17), group 16, acidic soln
-industrial
use nitrogen + hydrogen = ammonia (Haber Process)
Section 6-8 Oxygens & Oxides:
-oxygen
-Priestly, 1774
-convert pig iron to steel
-allotrope(same element in unstable
form in same state of matter), ozone
-Reactions of Oxygen and Oxides
-oxygen forms oxides through combinations
with all other elements except nobel gas
-oxides=binary compounds that
contain oxygen
-metal + O2 -->metallic oxide
(ionic solid)
-group
1 metals + O2-->oxides, peroxides, superoxides
-trend--tendency to form oxygen rich coompounds incease down group
-peroxides (O22-), superoxides (O2-)
-group
2 metals + oxygen = metallic oxides
-metals
w/variable oxid.states--limited oxygen=lower state, excess O2=higher
-reactions of metal oxides w/
water
-basic
anhydrides=oxides of metals
-combine
w/H2O to form bases w/no change in oxidation state of metal
-metal
oxide is hydroxide base w/ H2 removed
-reactions of oxygen w/nonmetals
-forms
molecular oxides (covalent)
-trend:nonmental(w/more
ox.states)lower oxstate=limit O2,higher oxstate=excess
-reactions of nonmetal oxides
w/H2O
-acid
anhydrides=nonmetal oxides
-combine
w/H2O to form acids w/no change in oxidation state of metal
-ternary
acids (3 elements: H,O, ?)
-reactions of metal oxides w/nonmetal
oxides
-basic
anhydrides + acid anhydrides-->salt(no change in ox.state of either)
-Combustion Reactions (complete combustion)
- hydrocarbon + O2 -->carbon dioxide
+ H2O
- limited oxygen = carbon oxide
+ H2O
- extremely limited oxygen = carbon
+ H2O
-Combustion of fossil fuels and the problems of
air pollution
Chapter 7: Chemical Bonding
Chemical Bonding is the attractive forces that hold
atoms together in compound
Ionic-metal + nonmetal (electrostatic attraction
of ions), electron exchange, solids have
high melting
point, soluable in polar solvents, conduct current
Covalent-nonmetal + nonmental, share electrons,
solids have low melting point, soluable
in nonpolar
solvents, do not conduct current
Lewis Structure-structures drawn to represent bonding
among atoms in a compound
-element represents nucleus and inner sshell electrons
-dots represent valence (outer energy llevel) electrons
-8 because highest energy level can onlly hold 8 electrons
Section 7-2 Ionic Bonding
-Attraction of anions and cations in large numbers
tend to form solids
-when diff in EN betw 2 elem is high->ionic bond
will form (diff sides of table)
-Coulombs Law-the forece of attraction (F) between
2 opp charged particles is directly
proportional to the product of their charges & inversly proportional
to the
square of the distance
-strong force equals large charge and small ions
-formula unit-simplest ratio in ionic compound
-Redox: loss of e- equals oxidation, gaion of e-
equals reduction
Section 7-3 Covelant Bonding
-2 atoms share one or more pairs of e-
-diff in EN very small or zero ( 2 nonmetals)
-lower melting and boiling pt
--strong intramolecular forces
--weak intermolecular forces
-atoms too close together=repulsions(notgoing to
bond)--higher energy
-H2 molecules-low (neg) energy--attract-most stable
-atoms to far apart-no bond--almost no energy
Section 7-4 Lewis Structure:Molecules & polyatomic
ions
-valence e-, # & types of bonds, no 3D shapes
Section 7-5 Octet Rule
-all elements want an octet (except H--wants 2)
-Lewis structure-boning e- shared
-unshared e- are only w/ 1 atom
-loan pair e- are pair of e- in same orrbital
- S(total shared)= N(valence e- needed) - A(# e-
available)
-Guidlines for Lewis structure
1. skeleton structure-central
atom least EN, H never central, O atoms dont bond to
each other (except O2,O3, peroxides), ternary acids: H bonds
to O (not central), if more than one central-use symetrical form
2. calculate N
3. calculate A (neg ion add charge,
pos ion subtract charge)
4. calculate S
5. place S e- into skeleton stucture
as bonding e-(double & triple as necessary)
6. form octets where needed
Section 7-6 Resonance
-three equivalent structures equals resonance structures
-true structure average of three
-four shared pairs are shared equally--delocalization
Section 7-7 Limitations to Octet Rule
-Beryllium-even ionic bonds have some covalent character,
N=4, 2 valence e- = 2 bonds
-Group 13 (Boron) 3 valence e- equals 3 covalent
bonds, N=6
-Compounds or ions w/ odd number e-
-Compounds or ions where central element needs more
than 8e- to hold all available e-
-if S is less than # needed to
bond all atoms to central atom than increase S
-if u satisfy all octets before
all A e- are used, put extra on cental atom
Section 7-8 Types of Covalent Bonds
Polar covalent-unequal sharing, diff in EN is significant
(slightly pos end-lowerEN,slightly
neg end-high EN==>dipole), crossed arrow points to more EN, heteronuclear
Nonpolar covalent-equal sharing, diff in EN equals
zero, homonuclear
Section 7-9 Dipoles
-dipole movement=distance seperating charges that
are equal in magnitude,opp in sign
times magnitude of charge
-electrically charged plates used to determine
Section 7-10 Continous Range of Bonding Types:
Nonpolar Covalent:EN diff=zero
Polar Cov:EN diff intermediate Ionic:EN diff large
-partial ionic character-HCl (polar cov.)-17%ionic
charact(due to uneual sharing)
-partial covalent charac-LiCl (ionic)-some degree
of e- sharing(due to charge density)
Classification:
1.)all compounds have ionic and covalent character
2.)classification must be consistent with w/physical
properties
3.)some comp have both characters, but one classificaion
is needed
Chapter 8: Molecular Sructure and Covalent Bonding
Theories
1.) shapes
2.) VSEPR (Valence Shell Electron Pair Repulsion
Theory)
-spatial
arragenment of atoms (where bonding occurs)
3.) VB (Valence Bond Theory)
-overlapping
of orbitals explains how bond occurs
-hybridization=blending
of orbitals (5 types)
Section 8-2 VSEPR
-valence e- on central atom repel each other
-decrease repulsion=stable bond
-regions of high e- density (each
bonded atom = 1, each unshared pair= 1)
-Electronic geometry-geometry resulting from arrangement
of e- density around
central atom (stable molecule)
-Moecular geometry-actual shape of molecule
Section 8-3 Polar Molecules
-Polar Molecules have EN differences greater than
zero(look at dipole arrangement)
-must be one polar bond or one pair unshared e-
AND either
the polar bonds or unshared pairs
of e- will not symmetrically cancel
Section 8-4 Valence Bond Theory
hybridization-blending of orbitals to form new,low
energy boning orbitals
1.) sp-2 regions high e- denisty on central atom
2.) sp2-3 regions high e- denisty on central atom
3.) sp3-4 regions high e- denisty on central atom
4.) sp3d-5 regions high e- denisty on central atom
5.) sp3d2-6 regions high e- denisty on central atom
Sections 8-5 thru 8-12 in book
Section 8-13 Compounds containg Double Bonds
sigma bond-bond resulting from head on overlab or
orbitals
pi bond-bond resulting from side to side overlab
of orbitals, exist only w/sigma bond
double bond = 1 sigma bond, 1 pi bond
Section 8-14 Triple Bonds
triple bond= 1 sigma bond, 2 pi bonds
Chapter 10 & 11: Acid-Base Chemistry
Section 10-1 Properties
Acids: sour, lithmus paper(B->R), neutralization,
RXNs w/salt form acid,new salt,current
Bases: biter, lithmus paper(R->B), neutralization,
current
Section 10-2 Arrhenius Theory
Acid produces H+ in aqueous soln
Base produces OH- in aqueous soln
limits Acid/Base classification because of need
for aqueous soln
Section 10-3 Hydronium Ion
H3O+ = hydrated hydrogen ion [H+(H2O)n] <-hydrated
proton
-concentration determines pH
Section 10-4 Bronstead-Lowery Theory
Acid is H+ donor
\
Base is H+ acceptor /aqueous soln not
neccessary
conjugate acid base pairs = species that differ
by 1 proton
conjugate base of strong acid is weak base
conjugate base of weak acid is strong base
conjugate acid of weak base is strong acid
water is H+ acceptor when it reacts with strong
acid
water is H+ donor when it reacts with a weak base
autoionization of water: pure water autoionizes
slightly to form H3O+ and OH-
water is amphiprotic
Section 10-5 Strength of Acid
1.) bond strength (halogens in group order F>Cl>Br>I)--major
part
2.) stability of resulting soln
3.) leveling solvent (water)--role is measureing
strength
a.) H3O+ is strongest acid
found in aq. soln
b.) OH- is strongest base found
in aq. soln
Ternery Acids:
aren't bases because central atom is nonmetal
1.) increase oxygen content = increase acid strength
2.) increase EN of central atom = increase acid
strength
Section 10-6 Reactions of Acids/Bases (Net Ionic)
Strong acid + strong base: H+ + OH- ---> H2O
weak acid + strong base: HA + OH- <==>
A- + H2O
Section 10-7 Acids and Basic Salts
acidic salt-less that stochiometric amount of base
reacts with polyprotic acid
basic salt-polyhydroxy bases that react with less
tha stochiometric amount of acid
Section 10-8 Amphoterism
the ability of a substance to act as an acid or
base
Section 10-9 Preperation --in book
Section 10-10 Lewis Theory
acid-any species that can accept a share in an e-
pair
base-any species that can donate a share in an e-
pair
Section 11-1 Calculation involving Molarity
molarity = moles/L
titration-adding one species to another, how much
of one will react with other
indicator-changes color at end pt.
equivalence pt- stochiometrically equiv. amts of
acid/base
standardization-to know accurate concentration (
[ ] ) of soln
-accuracy-volume,moles
-primary standard-high molec weight,sollubility(most soln), safety, not
very reactive
reactions known, pure
-KHP for base and Na2CO3 for acid
-secondary standard is standarized solnn
Section 11-2 Standarization of Acid and Base
Solns
KHP-abb for organic comp: potassium hydrogen phtalate(M.W.
204.2g) K6H4(COO)(COOH)
Section 11-3 Equivalent Weights and Normality
normality = equivalent weights (eq) of solute per
liter of soln
1 eq of acid-the mass of acid that furnishes 1 mole
hydrogen ions or reacts with 1 mole hydroxide
1 eq of base-the mass of basethat furnishes 1 mole
hydroxide ions or reacts with 1 mole hydrogen
eq weight = formula weight / # of acidic hydrogens
or hydroxide ions
normality always greater than molarity
volume x normality = volume x normality
product of volume and concentration equals amt of
solute
Section 11-6 Balancing Redox Reaction--Half Reaction
Method
1. write as much of overall reaction as possible
(omit spectator ions)
2. write unbalanced half reactions for oxidation
and reduction
3. balance all elements in each half reaction except
hydrogen and oxygen
4. balance hydrogen and oxygen
Balance O
Balance H
acidic:
add water
then
add H+
basic:
for every O:
for every H:
add 2 OH- to side needing O then
add 1 water to side needing H
add 1 water to other side
add 1 OH- to other side
5. balance charge by adding e-
6. balance e- transfer by multiplying by appopriate
integer
7. add half-reactions and eliminate common terms
===>net ionic equation
Chapter 12: Gases and Kinetic Molecular Theory
Section 12-1/12-2: Aspects of Gas Behavior
1. can be compressed
2. exerts pressure on surroundings (collisions of
molecules and container)
3. expand without limits
4. diffuse into each other (mix completely)
5. properties dependent on pressure, volume, temp,
moles
Section 12-3 Pressure
1. collisions
2. force per unit area (P=F/A) units-- torr, atm,
mm Hg, psi, inches, SI unit--pascal (1 N/m2)
3. measuring pressure--barometer, manometer
Section 12-4 Boyle's Law
1. volume vs. pressure (increase pressure= decrease
volume)
2. V1P1 = V2P2 (amt, temp constant)
Section 12-5 Charles' Law
1. volume occupied by certain mass is directly proportional
to absolute temp (K)
2. V1/T1 = V2/T2 (pressure, amt constant)
Section 12-6 STP (Standard Temp and Pressure)
1. reference for dealing with gases
2. temp= 0 Celcius or 273 K
Section 12-7 The Combined Gas Law
1. combined Boyle and Charles' Law
2. P1V1/T1 = P2V2/T2 (amt constant)
3. P1/T1 = P2/T2 (volume, amt constant)
Section 12-8 Avogadro's Law
1. at constant temp and pressure, equal volumes
of gas contain same # of molecules
2. V1/n1 = V2/n2 (temp, pressure constant)
3. 22.4L @ STP = 1 mole of gas = molar mass = Avogadro's
# particles
Section 12-9 Ideal Gas Law
1. ideal gas-gas that obeys laws exactly
2. PV = nRT (R= universal gas constant .0821
Latm/mol k or 8.315 dm3 kPA/mol K
Section 12-10 Molecular Formula and Empirical Formula
1. Dumas method-heat substance to force out air
to only have gas vapor to figure out how
much gas
produced
Section 12-11 Dalton's Law of Partial Pressure
1. mixture of gases-partial pressure of each gas
(total moles = sum of all moles present)
total pressure = sum of pressures
of all gases (P1 + P2 + P3 + .....)
2. mole fraction and partial pressures
XA = the number of moles of A
out of total # of moles in sample (nA / ntotal)
PA = XA x Ptotal (mole
fraction x total pressure = partial pressure)
3. Collection of Gas over water
Ptotal = Pgas + Pwater
(must acount for pressure of water vapor) Patm = Pgas + Pwater
Section 12-12 Kinetic Molecular Theory
1. gases composed of molecules (small, far apart,
compressable, low density)
2. molecules in constant, straight line morion w/
varying velocities
3. collisions are elastic (energy transfered)
4. there are no attractive forces between molecules
Kinetic Energy of Gas KE= 1/2 m v2
1. KE and T ==> average KE of molecules proportional
to T
2. KE of molecules of diff gases are equal at same
temp
Ideal Gas Laws with KMT
1. Boyle=>pressure caused by force of collisions
(1/2 V = 2 P)
2. Dalton=>molecule spaced out (each gas exerts
a pressure)
3. Charles=>increase V, more collisions, increase
temp (more energy)
4. Avogadro=>increase amt, increase volume to keep
constant P
Section 12-13 Diffusion and Effusion of Gases
1. diffusion--gas molecules expand to fill the volume
of container contained in
2. effusion--gas molecules move through tiny openings
in porous materials
Section 12-14 Real Gases -- Deviations from ideality
1. under most conditions real gases behave ideally
2. nonideal behavior most significant at high pressure/low
temp (liquifies)
3. volume--ideal=gas molecule move w/o interuptions
--real=volume that gas molec has availaable is not measure volume
-- "nb" correction factor
--larger molec have greater "b"
--larger # molec have greater "n"
--larger the product of nb, the larger the colume correction
4. pressure--ideal=attractive forces btw molec not
significant
--real=molec attract one another (less energy collisions)
--"a" correction for attraction
--large value for "a" equals strong atttractive forces
--large values for "n" equal more molecc present
Van der Waals Equation:
n2a
( P + V
) (V - nb) = nRT
(a & b are constants must be given)
Ex. He low "a" b/c noble
gas (few attractions)
low "b" b/c monatomic, very small
Section 12-15 Volume-Mass Relationship in RXNs involving
gases
stochiometry
Chapter 13: Liquids and Solids
13-1 Kinetic-Molecular Description of Liquids &
Solids
Liquids-disorganized clusters, close, difficult
to compress
Solids-ordered arrangement, vibrational motion about
fixed position, nearly incompressible
13-2 Intermolecular Attractions & Phase Changes
Properties:
Liquids-boiling pt. vapor pressure, viscosity, heat
of vaporizaiton
Solids-melting pt., heat of fusion
intermolecular-bwt particles intramolecular-withion
compds (bonds)
Some intermolecular attractions
ion-ion --- columbs law
dipole-dipole --- polar covalent
hydrogen bonds --- not really bonds H, O/N/F
some Cl/C
London forces
13-3 Viscosity- measure of liquids resistance to
flow
13-4 Surface Tension- measure of inward forces that
must be overcome to increase surf area
13-5 Capillary Action --cohesive and adhesive forces
13-6 Evaporation -- condensation liquid/gas phase
LeChatelier's Principle
13-7 Vapor pressure-- partial pressure of vapor
molecules above surface
13-8 Boiling Point & Distillation BR when
VP= external pressure
13-9 Transfer involving Liquids-- specific heat,
molar heat capacity, molar heat of vap/cond
13-10 Melting pt/freezing pt-- temp solid / liquid
phase in equilib
13-11 Heat Transfer involving solids-- molar heat
of fusion, heat of solidification
13-12 Sublimation
13-13 Phase Diagram
13-15 Crystal Structure-- unit cell, 7 simple, complex,
13-16 Bonding in Solids-- metallic, ionic, molecular,
covalents
13-17 Band Theory -- properties due to electron
structures