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Things you need to review
Before IB Chem -11
Concepts and notes are in bold. Problems you should try and solve are in normal
font. Special directions are in italics.
If there are significant issues, please e-mail me: [email protected]
Have a good summer,
Mr. Capello
I. Prefixes
Prefixes
a)
b)
c)
d)
e)
f)
g)
h)
Mega, 106, M
Kilo, 103, k
Deci, 10-1, d
Centi, 10-2, c
Milli, 10-3, m
Micro, 10-6, μ (mew)
Nano, 10-9, n
Pico, 10-12, p
1 mL = 1cc = 1 cm3
Convert the following:
6.7 cm → dm
0.54 kg → g
5.95*108 µL → kL
1.323 Mg/L → mg/cm3
1.79*1032 pm/ns → km/h
II. Scientific Precision
A. Accuracy – refers to the closeness of a measurement to the true or
accepted value of the quantity measured (how close to right)
B. Precision – refers to the agreement among the numerical values of a
set of measurements of the same quantity made in the same way (how
close to others).
C. Significant Figures
1. Rules
a. If the number is less than one, the first nonzero number to the
right of the decimal is significant and all numbers following are
significant.
b. If the number is greater than one, every number is significant
2. Calculations
a. Addition and subtraction – give as many sig figs in the answer as
there are in the measurement with the least number of decimal
places.
b. Multiplication and division – give as many sig figs in the answer
as there are in a measurement with the least # of sig figs.
3. Exact Numbers
a. A number that arises from counting or conversion factor ex// how
many students in the classroom
b. Ignore exact numbers when doing sig fig calculations
D. Percent Error
1. a judgment of accuracy based on a comparison with the accepted
value
2. %Error = | Accepted – Experimental | /Accepted * 100
How many sig figs?
Calculate with the correct sig figs?
0.005 =
0.0050 =
123. =
10. =
5.00 cm + 7.113 cm =
27.8252 L - 5.1 L =
3.975 kg * 2.2 kg=
315.29754 m /5.746 m =
The melting point of a chemical is 53.0°C. Two students try to determine this value
experimentally doing 5 trials. Student A gets 54.9°C, 54.7°C, 55.0°C, 54.8°C, and 55.1°C.
Student B finds 52.3°C, 53.2°C, 54.0°C, 52.5°C, and 53.5°C. Calculate the average value
for each student. Using the average value, determine each student’s percent error. Which
student is more precise, which is more accurate?
Memorize:
III. Polyatomic Ions
Ammonium – NH4+
Hydroxide – OHPermanganate – MnO4Nitrite – NO2-
Nitrate – NO3-
Sulfite – SO32-
Sulfate – SO42-
Carbonate – CO32Bicarbonate or hydrogen carbonate HCO3Phosphate – PO43Hypochlorite – ClO-
Chlorite – ClO2-
Chlorate – ClO3-
Perchlorate – ClO4-
Cyanide – CNChromate – CrO42Thiosulfate – S2O32-
IV. Nomenclature
A. Monatomic ion
1. Review charges on periodic table
2. If cation, name stays the same
3. If anion, name adds -ide to end
B. Binary ionic compounds
1. Write down charges
2. Find LCD
3. Use factor as subscript
4. Take name of cation and name of anion
5. Examples: LiF; BaO; Al2O3; Na2O; AlI3; MgCl2; K3N
C. Multivalent compounds
1. Look at charge of anion
2. Discover charge of transition metal
3. Find LCD
4. Use factor as subscript
5. Take name of metal (Roman numeral) anion name
6. Examples: CuCl; CuCl2; FeO; Fe2O3; ZnS; Zn2S
D. Polyatomic Ion Compounds
1. Write down charges
2. Find LCD
3. Use factor as subscript with polyatomics in () if subscript is
other than one
4. Take name of cation and name of anion
5. Ex: NaNO3; MgCO3; Ca(BrO3)2; K2SO4; (NH4)3PO4
E. Covalent Compounds
1. Prefixes
a. 1 = mono
f. 6 = hexa
b. 2 = di
g. 7 = hepta
c. 3 = tri
h. 8 = octa
d. 4 = tetra
i. 9 = nona
e. 5 = penta
j. 10 - deca
2. Provide less electron negative element first
3. Add a prefix to that element if the subscript is other than 1
4. Give more e-negative element with prefix, name, & ide
5. Ex: NO, N2O; NO2; N2O3, N2O4; N2O5; P4O10; OF2; SO3
F. Acids
1. Binary acid – H plus some single element
a. Add the prefix hydro and change the anion to –ic
b. HF, HCl, HBr
2. Oxyacids – H plus a polyatomic ion
a. If the polyatomic ends in -ite change to –ous
b. If the polyatomic ends in –ate change to -ic
Name these examples
B. Binary ionic compounds
LiF
AlI3
BaO
MgCl2
Al2O3
K3N
Na2O
C. Multivalent compounds
CuCl
Fe2O3
CuCl2
FeO
D. Polyatomic Ion Compounds
NaNO3
K2SO4
MgCO3;
(NH4)3PO4
Ca(BrO3)2
E. Covalent Compounds
NO
N2O5
N2O
P4O10
NO2
OF2
N2O3
SO3
N2O4
F. Acids
1. Binary acid
HF
HCl
HBr
2. Oxyacids
H2SO3
H2CO3
H2SO4
HClO
H3PO4
HClO2
HNO2
HClO3
HNO3
HClO4
NO – name
cation first
Add –ide
to anion
YES – use
Roman
Numerals to
explain
charge
Add –ide
to anion
If ionic, is it a
transitional
metal?
Then, name
more
electroneg.
element
ALWAYS use
prefix and
add -ide
First, name
less
electroneg.
element
Add prefix if
subscript is
other than
one
Covalent
Is it an ionic
bond, covalent
bond, or acid?
Nomenclature
Add
hydro +
halogen
+ic +acid
Is it
binary,
H+
halogen
Balance
charge of
polyatomic
with H’s
Is it oxy, H+
polyatomic
ite = ous+acid
ate = ic+acid
Acid
V. Measurement and Units
A. Base Units
1. Groups
a) Length – meter, m
b) Mass - kilogram, kg
c) Time – second, s
d) Temperature - kelvin, K
i. K = ˚C + 273
ii. ˚F = (1.8 * ˚C) + 32
iii. ˚C = (˚F – 32) / 1.8
(e) Amount of substance - mole, mol
(f) Electric current – ampere, A
(g) Luminous intensity – candela, cd
B. Derived Units – combinations of SI base units
1. Volume, cubic meter, m3, l*w*h
a) 1 mL = 1 cc = 1 cm3
b) 1 dm3 = 1 L = 1000 mL = 1000 cm3
2. Density, mass per unit volume, kg/m3, g/cm3
a) d = m/v
3. Velocity, distance traveled per unit time, m/s
4. Acceleration, velocity changed per unit time, m/s2
5. Pressure, force per unit area, pascal, Pa = kg/(m*s2)
6. Energy, force times length = joule, J = kg*m2/s2
a) 1 calorie (cal) = 4.184 J
b) 1 Calorie = 1 kcal = 1000 cal = 4184 J
8. Concentration, amount per volume, mol/dm3 or moldm-3
VI. Ionic Bonding
A. Definition – bonding that results from electrical attraction between (+)
ions (cations) and (-) ions (anions)
1. Cations – 1+, 2+, 3+ - P.T. Trends – show valence e- and dots
2. Anions – 3-, 2-, 1- - Same as above
B. Difference in e-neg >1.7 ex:// Na=.9 and Cl=3.0
C. Typically the joining of a metal (left of staircase) and a nonmetal (right
of staircase)
D. Ionic Compound – a substance where the total combination of cations
and anions equal to zero
E. The How-To of Ionic Formation
1. Draw dot diagram of each atom participating in bond (Reactants)
2. Determine if it is easier to go to zero or eight and how best to get
there
3. Set up reactants with dots
4. Set up products with dots and charges as necessary.
5. Do the charges add up to zero?
F. Crystal Lattice
1. A 3-D arrangement so that each (+) charged ion is surrounded by
(-) ion and vice versa
2. Lattice energy – the energy required to separate 1 mol of ions
from an ionic compound – the greater the E, the greater the force
of attraction, the hard to break
3. Lattice energy is based on ion size and charge
a. Smaller ions – greater E – KF vs RbF
b. Greater charge – greater E – NaI vs MgO
G.
Polyatomic ions – a group of atoms with a charge
H.
Transitional Metals – multivalent charges
VII. Metallic bonding
A.
B.
Metals form lattice structures and are based on electrical attraction
Electron sea model
1. Valence e- are not attached to specific metal cations instead e- can move
freely about – as if the cations were like buoys surrounded by a sea of
electrons
2. Because the electrons do not have a specific cation they are attached to, they
are called delocalized electrons
C.
Metallic bond – attraction of a metallic cation to delocalized e- which helps
explain most of the properties of metals
D.
Alloys
1. A mixture of elements that have metallic properties
2. Substitutional alloy – different metal cations of similar size are combined
3. Interstitial alloy – smaller atoms fill the holes in the metal lattice
4. Examples
a. Brass – Cu and Zn
b. Bronze – Cu, Zn, and Sn
c. Gold – only 24 carat is pure
d. Pewter – Sn, Sb, and Pb
e. Stainless Steel – Fe, Cr, and Ni
VIII. Covalent Bonding
A. Bonding that results from the sharing of e- between atoms usu.
between two nonmetals
B. Types
1. Polar
a. e- are unevenly shared resulting in a slight charge
b. 0.3 < Difference < 1.7
c. H = 2.1 and Cl = 3.0
2. Nonpolar
a. e- equally shared resulting in a balanced distribution of
charge
b. e-neg < 0.3
c. H2
C. Basic Definitions
1. Molecule – a neutral group of atoms held together by covalent
bonds
2. Chemical Formula – indicates the relative numbers of each
kind of atom in a chemical compound by using atomic symbols
and numerical subscripts
3. Diatomic Molecule – a molecule containing only two atoms
4. Bond Length
a. the average distance between two bonded atoms at their
minimum potential energy
b. Single bonds are longer; triple bonds are shorter
5. Bond Energy
a. The energy required to break a bond and form neutral
isolated atoms.
b. Single bonds have the least E; triple the most
D. Dot Diagrams
1. Lewis Definitions
a. Lewis dot diagram – a positioning of atomic symbols using
e-pairs or dashes to represent bonds between atoms and
adjacent pairs to represent nonbonding eb. Shared pair or bonding e- - pairs of e- that act as a bond
between atoms
c. Unshared, lone, or nonbonding e- pair – e- necessary to fill
octets but not used in the bonding of atoms
2. Octet Rule
a. Chemical compounds tend to form so that each atom has 8
e- in its valence shell
b. Exceptions: H, Ne, B, among others
3. Steps
a. Add up all of the valence eb. Choose central atom
c. Position other atoms around central
d. Bind each atom singly to central and subtract
e. Fill octet of surrounding atoms and subtract
f. Attempt to fill octet of central atom
g. Check for multiple bonds or ion formation
h. Recount to 8
4. Single Bonds – a covalent bond produced by the sharing of
one pair of e- between two atoms: F2; HF; H2O; NH3; CF4
5. Multiple Bonds
a. A covalent bond produced by the sharing of two (double) or
three (triple) pairs of e- between atoms that usually only
occurs with C, N, S and O
b. e- count used up but an atom still has an unfilled octet and
the combination is not recognized as a polyatomic ion
c. Examples: CO2 O2 N2
IX. Reactions Types
A. Synthesis
1. Definition – two or more substances that combine to form a
new compound
2. Generic formula – A + X -> AX
B. Decomposition
1. Definition – a single compound undergoes a rxn to produce
two or more simpler substances
2. Generic – AX -> A + X
C. Single-Replacement
1. Definition – when one element takes the place of a similar
element in a compound
2. Generic Formula – A + BX  AX + B
D. Double-Replacement
1. Definition – the ions of two compounds exchange places to
form two new compounds
2. Generic Formula – AX + BY  AY + BX
E. Combustion
1. Hydrocarbons – a compound that has only C’s and H’s in it
2. General formula –
Hydrocarbon + Oxygen  Carbon dioxide + Water
F. Neutralization - An acid and base combine to form salt and water
X. Activity Series
A. List of elements organized according to the ease in which they
react
B. Helps determine whether a reaction will work or not
Images taken from Advanced
Chemistry Clugston and
Flemming (226)
Will these reactions occur or not? If so, predict the products and balance.
K + BaO 
Na + CaO 
F2 + 2NaCl 
Br2 + KCl 
XI.VSEPR
A. Valence shell electron pair repulsion (VSEPR) explains the shapes of
molecules
B. The shape will help indicate the properties of molecules
C. Based on the idea that bonding pairs take up less space than lone
pairs and the negative charges spread out in space as much as
possible
D. To predict shape
1. Draw covalent structure
2. Count shared vs. unshared pairs on the CENTRAL ATOM
3. Double and triple bonds count as one shared pair
4. Consult table
Solve these probems:
Density
What is the density of sunflower oil if it has a mass of 116.0 g and a volume of 125.0 mL?
If the known density of sunflower oil is 0.919 g/mL, what is the percent error of this
experiment?
What is the volume of a sample that has a mass of 20.0 g and a density of 4.0 g/mL?
A piece of aluminum is placed in a graduated cylinder that contains 10.5 mL of water. After
placing the aluminum in the cylinder the new measurement of the water in the cylinder is
13.5 mL. The density of aluminum was looked up and found to be 2.7 g/mL. What was the
mass of the aluminum placed in the cylinder?
Draw and name the following. Also determine the electron domain and molecular
geometries:
F2
NH3
CF4
CO2
N2
NH4+
PO43-
Balance and understand the following reaction types:
Synthesis Subtypes
a. Metal
i.
ii.
iii.
+ Oxygen
Mg +
Fe +
Fe +
yields a metal oxide
O2 
O2 
O2 
b. Metal + Sulfur yields a metal sulfide
i.
Rb +
S8 
ii.
Ba +
S8 
c. Metals + halogens yields a salt
i.
Na +
Cl2 
ii.
Mg +
F2 
d. Metal Oxides + water yields a metal hydroxide
i.
Na2O + H2O 
ii.
CaO +
H2O 
e. Nonmetal Oxides + water yields oxyacids
i.
SO2 +
H2O 
ii.
P4O10 + H2O 
Decomposition Subtypes
a. Binary Compounds to individual elements
i.
H2O 
ii.
HgO 
b. Metal Carbonates yield metal oxide and carbon dioxide
i.
CaCO3 
ii.
Na2CO3 
c. Metal Hydroxides yield metal oxide and water
i.
Ca(OH)2 
ii.
NaOH 
d. Metal Chlorates yield metal chloride and oxygen
i.
KClO3 
ii.
Mg(ClO3)2 
e. Oxyacids yield nonmetal oxides and water
i.
H2CO3 
ii.
H2SO4 
Single Replacement Subtypes
a. A metal replaces a metal
i.
Al +
Pb(NO3)2 
b. A metal replaces aqueous H to form a metal hydroxide and
hydrogen gas
i.
Na +
H2O 
c. A metal replaces acidic H to form a salt and hydrogen gas
i.
Mg +
HCl 
d. Halogen replacement
i.
Cl2 +
KBr 
Double Replacement Subtypes
a. Formation of a precipitate
i.
KI +
Pb(NO3)2 
b. Formation of a gas
i.
FeS +
HCl 
c. Formation of water
i.
HCl +
NaOH 
Combustion Examples
a. CH4 +
b. C3H8 +
c. C8H18 +
O2 
O2 
O2 
Neutralization Examples
a. HCl +
NaOH

b. Mg(OH)2 +
HNO3 
These are the very basic of concepts you should understand. Other topics we will be
covering in the first month which should be somewhat familiar to you are
stoichiometry and limiting reactant problems.
Here are some websites to look over if those concepts are foreign to you:
http://www4.ncsu.edu/~franzen/public_html/CH201/prac/Limiting_Reagent_Worksheet.pdf
https://www.everettcc.edu/files/students/rainier-learning-center/tutoringcenter/chemistry/w324-limiting-reagent-worksheet.pdf
https://smccd.mrooms.net/pluginfile.php/584957/mod_resource/content/1/StoichiometryWo
rksheet1.pdf