CHE2A EXAM 1 REVIEW (ch1-4)

Scientific method

A logical, systematic approach to the solution of a scientific problem

Natural law

explains what happens

theory

explains WHY something happens

Density =

mass/volume

Kelvin =

273 + C

Farenheit =

9/5(C) + 32

Celcius =

(F-32) x 5/9

Common Conversions

0°C = 273.15 K
1 in = 2.54 cm
1 mL = 1 cm^3
1 bar = 100 kPa
1 atm = 101.325 kPa
760 torr = 1 atm
1 mol = 6.022 x 10^23 particles

Rules for sig figs (multiplication/division)

use number of sig figs of least sig fig number

ex. *1.4* x 5.376 = 7.5

Rules for sig figs (addition/subtraction)

use number of decimal places as number w/ least decimal places

ex. 1.30 x *2.7* = 3.5

law of conservation of mass

total mass of substances present after a chemical reaction is the same as the total mass of substances before the reaction

law of definite proportions

(law of constant composition)
all samples of a compound have the same composition

law of multiple proportions

If 2 elements combine to form different compounds, the ratio of masses of the second element that react with a fixed mass of the first element will be a simple, whole-number ratio

isotopes

Atoms of the same element that have different numbers of neutrons

number of protons defines ...

the element (atomic number)

number of neutrons defines ...

isotopes, mass of element (protons + neutrons)

number of electrons defines ...

charge of element

molecular formula

shows the types and numbers of atoms combined in a single molecule of a molecular compound

empirical formula

The simplest whole number ratio of atoms of each element in a compound

Where are *alkali metals* located on the periodic table?

Column 1

Where are *alkaline earth metals* located on the periodic table?

Column 2

Where are *transition metals* located on the periodic table?

Columns 3-12

Where are *chalcogens* located on the periodic table?

Column 16

Where are *halogens* located on the periodic table?

Column 17

Where are *noble gases* located on the periodic table?

Column 18

Where are *the main groups* located on the periodic table?

Column 1, 2, 13-18

What are *periods* on the periodic table?

horizontal rows

Where are *nonmetals* located on the periodic table?

Metals = left of staircase
Nonmetals = right of staircase

9 transition metals with uncertain charges

*+2 or +3 charges*
Cr (Chromium)
Fe (Iron)
Co (Cobalt)

*+1 or +2 charges*
Cu (Copper)
Ag (Silver)
Au (Gold)
Hg (Mercury)

*+2 or +4 charges*
Sn (Tin)
Pb (Lead)

Oxygen common oxidation state

-2

Hydrogen common oxidation state

+1

Fluorine common oxidation state

-1

Naming *covalent compounds*

between 2 nonmetals
= PREFIXES

ex. SF10 = Sulfur decafluoride

Covalent compound prefixes

1: mono
2: di
3: tri
4: tetra
5: penta
6: hexa
7: hepta
8: octa
9: nona
10: deca

Naming *binary ionic compounds (type 1)*

with a metal with specified charge
= full name cation, anion +ide

ex. CsF = Cesium Fluoride

Naming *binary ionic compounds (type 2)*

with a metal with UNspecified charge
= ROMAN NUMERALS

ex. FeO3 = Iron (III) Oxide

Naming *acids (with NO oxygen)*

add hydro- and -ic

ex. HF = hydrofluoric acid

Naming *acids (with oxygen)*

*for -ate endings*
= add - ic

ex. HClO3 = chloric acid

*for -ite endings*
= add -ous

ex. HClO2 = chlorous acid

Charges of periodic table

Common polyatomic ions (MEMORIZE)

*NH4+ = ammonium*
NH3+ = ammonia
NO2- = nitrite
*NO3- = nitrate*
SO3 = sulfite (2- charge)
*SO4 = sulfate (2- charge)*
*OH- = hydroxide*
CN- = cyanide
*PO4 = phosphate (3- charge)*
*CO3 = carbonate (2- charge)*
C2H3O2- = acetate
MnO4- = permanganate
*CrO4 = chromate (2- charge)*
O2 = peroxide (2- charge)
ClO2- = chlorite
*ClO3- = chlorate*

Polyatomic ion prefixes

per- = +1
-ate = __
-ite = -1
hypo- = -2

Atomis Mass =

*sum of [(isotopic masses) x (fractional abundance)]*

ex. 99.630% N-14 and .37% N-15 = (.99630)(14)+(.0037)(15) = 14.0037 u

Molar Mass =

*sum of amount of grams per element in compound*

ex. Na2SO4 = 2(Na) + S + 4(O) = 2(22.99) + 32.07 + 4(16) = 142.05 g/mol

Percent Mass =

*(mass of element ÷ mass of compound) x 100*

ex (using formula). CoCl2 x 6H2O find mass % of water
6(2.016+16) ÷ 237.93 g/mol x 100 = 45.44%

ex (using grams). 87.28 g P, 112.72 g O, 200g total compound
87.28 g P ÷ 200 g x 100= 43.64% P
112.72 g O ÷ 200 g x 100 = 56.36% O

Combustion reaction

reaction where an organic compound (contains C, H, and sometimes others) reacts with O2 to create CO2 + H2O

Combination reaction (synthesis)

2 or more species combine to form compound

*A + B --> AB*

Decomposition reaction

compound decomposes into 2 or more species

*AB --> A + B*

Single Replacement reaction

a chemical change in which one element replaces a second element in a compound

*AB + C --> A + BC*

Double Replacement reaction

a chemical change that involves an exchange of positive ions between two compounds

*AB + CD --> AD +BC*

Redox reaction

A chemical reaction involving the transfer of one or more electrons from one reactant to another; also called oxidation-reduction reaction

Limiting Reactant steps

1. Balance equation
2. Calculate moles or mass of reactants
3. Calculate moles/mass of opposite reactant that is produced from using each reactant
4. Smaller production = limiting reagant
5. Use moles/mass of limiting reagant to calculate product
6. To find excess: subtract available moles/mass of excess with used moles/mass of excess

Percent Yield =

*actual yield ÷ theoretical yield x 100*

aka. given amount ÷ calculated amount x 100

Molarity =

*moles of solute ÷ volume in liters*

Empirical Formula steps (from mass %)

1. base mass % out of 100 g of compound (each % = g)
2. Determine moles of each element
3. Divide each mole by smallest mole value
4. Whole number values = subscripts of each element in compound

Molecular Formula steps (from empirical formula)

*molar mass ÷ empirical formula mass* = factor to multiply subscripts by

Dilution =

*Molarity (M1) x Volume (V1) = Molarity (M2) x Volume (V2)*

ex. 1 L of 11.75 M HCl reacts with 2 L NaOH, find molarity NaOH
(11.75 M)(1 L) = (M2)(2 L) = M2 = 5.9 M NaOH

Titration =

(molarity of subject) (volume of subject) = (molarity of titrant) (volume of titrant)

Strong electrolytes

completely dissociate into ions when dissolved in water

- strong acids (aq)
- strong bases (aq)
- ionic compounds (aq)
- salts (aq)

Weak electrolytes

don't completely disassociate in water

- weak acids
- weak bases
- molecular compounds
- organic compounds
- liquids
- solids
- gases

Precipitate reactions

reactions that result in the formation of an insoluble product (precipitate)

molecular formula:
ex. AgNO3 (aq) + KCl (aq) --> AgCl (s) + KNO3 (aq)

Precipitate reactions: Net ionic equation

ionic equations that include only the particles that participate in the reaction (no spectator ions)

ex. HCl(aq) + NaOH(aq) --> NaCl(aq) + H2O (l)
full ionic equation:
H+ + Cl- + Na+ + OH- --> 2H+ + Cl- + Na+ + OH-
net ionic equation:
*H+(aq) + OH-(aq) --> H2O(l)*

What is an acid?

compound that provides H+

What is a base?

compound that provides OH-

- NOT including compounds that do not contain a metal in addition to OH- (not CH3OH, not MeOH)

Strong Acids (MEMORIZE)

- HCl
- HBr
- HI
- HClO4
- HClO3
- HNO3
- H2SO4

Strong Bases (MEMORIZE)

- LiOH
- NaOH
- KOH
- RbOH
- CsOH
- Ca(OH)2
- Mg(OH)2
- Sr(OH)2
- Ba(OH)2

redox reactions (oxidation-reduction reactions)

chemical reactions that transfer electrons between reactants

- oxidized = loses electrons (reducing agent - electron donor)
- reduced = gains electrons (oxidizing agent - electron acceptor)

remember OIL RIG

*O*xidation *I*s *L*osing electrons
*R*eduction *I*s *G*aining electrons

Half reactions

equations that show either oxidation or reduction alone

Half reaction balancing steps

1. elements (not H2O)
2. oxygen (use H2O)
3. hydrogen (use H+)
4. charge (use electrons)

Half reactions: reduction

ex. SO2 + Cr2O7 (-2) ---> SO4 (-2) + Cr (+3) + H2O
1. reduction equation: Cr2O7 (-2) ---> 2Cr(+3)
ox states: Cr went from +6 to +3 (gained 3 electrons - reduced)
2. balance oxygen
Cr2*O7* (-2) ---> 2Cr(+3) + *7H2O*
3. balance hydrogen
Cr2O7 (-2) + *14H+*---> 2Cr(+3) + *7H2*O
4. balance charges
Cr2O7 (-2) + 14H+---> 2Cr(+3) + 7H2O
(Cr2 = +12 to 2Cr(+3) = +6) 12-6 = add 6 electrons to left

Cr2O7 (-2) + 14H+ + *6e* ---> 2Cr(+3) + 7H2O

Half reactions: oxidation

ex. SO2 + Cr2O7 (-2) ---> SO4 (-2) + Cr (+3) + H2O
1. oxidation equation: SO2 ---> SO4(-2)
ox states: S went from +4 to +6 (lost 2 electrons - oxidized)
2. balance oxygen
SO2 + *2H2O* ---> S*O4*(-2)
3. balance H+
SO2 + *2H2*O ---> SO4(-2) + *4H+*
4. balance charge
SO2 + 2H2O ---> SO4(-2) + 4H+
(S=+4 to S=+6) so, 4-6 = add 2 electrons to right

SO2 + 2H2O ---> SO4(-2) + 4H+ + *2e*

Half reactions: adding reduction and oxidation

reduction: Cr2O7 (-2) + 14H+ + 6e ---> 2Cr(+3) + 7H2O
oxidation: SO2 + 2H2O ---> SO4(-2) + 4H+ + 2e

Cr2O7 (-2) + 14H+ + 6e + SO2 + 2H2O ---> 2Cr(+3) + 7H2O + SO4(-2) + 4H+ + 2e

14H+/4H+ cancels, 6e/2e cancels, 7H2O/2H2O cancels

*Cr2O7 (-2) + 10H+ + 4e + SO2 ---> 2Cr(+3) + 5H2O + SO4(-2)*