CHE 121-1 (Gen Chem- Mehne) 8/12/13 8:22 PM

8/30/13

Beginnings of modern chemistry :

count beginning of chemistry with the beginning of quantitative

(measured by quantity rather than quality) studies instead of

qualitative studies.

o 1772 (before American revolution) Lavoisier studied

carbon + oxygen -> carbon dioxide + heat + light

asked how much (quantitative) of each of these

elements was used and produced

“Law of” summary of observations

“theories” or “models” = things we make up to try and organize

and rationalize our observations.

Measurements:

Property = Magnitude Units

Mass = 3.7 Grams

Weight = 152 Pounds

Cost = 48.95 $

Use the SI system (French) “international system of units” (“metric”

is now obsolete)

We are pretending the English system does not exist and are using

the SI system

property SI unit “lab unit”

Mass Kilogram

**base unit**

Gram = kg/1000

Volume Cubic meter

**derived unit**

Liter = m^3/ (1000)

Milliliter = m^3

/(1,000,000)

Temp Kelvin

(Capitalized bc name

of dead scientist)

Degree Celsius

0K=-273.15 degrees

C

Time second second

Pressure -

heat -

In lab normally measure to thousandth of a gram

Kelvin considered absolute bc 0 K means 0 temp

room temp about 298 K… Kelvin is NOT degrees (F and C both are)

Properties:

“Extensive” properties” = ex: mass; depends upon the amount

of matter chosen for measurement

o size, length, mass, volume, *heat*

“Intensive properties” = do NOT depend on the amount of

matter chosen for measurement (memory> doesn’t matter how

much time coach gives you he needs intensity)

shape, color, pressure, *temperature*, density

“In” dependent of mass

heat/ temp

o heat> measure of energy (more mass= more heat)

o temperature> concentration of heat

Ex of charcoal

o mass and volume would both be different than friends, but

dividing mass by volume (m/v) gives you density and both

peoples density would be the same

Don’t observe density directly

We measure extensive properties and use them to find an intensive

property and something that is a characteristic

9/2/13

Measurements continued:

every measurement has 3 parts:

o

o

o

No measurement is completely precise (there are limits)

o we are rarely counting in the lab and normally making

measurements

Truth in Measurement:

o Significant Figure convention

Report only one digit that is uncertain… The last digit of

a number should be the uncertain number.

Don’t use zeros to locate decimal points

leading zeros never count (ex: 0.07090 g)

see notes in ch. 1 reading tab on sig figs

Be careful with ambiguous numbers (ex: 2300 mm); fix

problem by using scientific notation or changing units as

we need to (ex: 2300 mm = 2.30 m)

o Propagation of Imprecision (“error analysis”)

Addition and Subtraction:

A result has the same number of decimal digits

as the fewest in the data

ex: see diagram 1.1

Multiplication and division:

A result has the same number of significant

figures as the fewest in the data.

Balances in lab precision of 3 decimal digits

ex: mass total = 2.361 g; mass container = 2.340

g

fewest given decimals is 3 so 3 in the

results of 0.021 g (*only 2 sigs)

ex: (line of ink on paper)

w= .1124 cm

L= 2.0 cm

Want area so multiply two numbers. from

calc A=0.2248 cm^2 (CALC LIES) Fewest

given sig figs is L that has 2. So answer

should be written A=0.22 cm^2

Another ex: (paint strip)

L= 52.3 cm; W= 5 cm

A= 261.5 cm^2…. W is only 1 significant figure so

you have to round answer to 300 cm but this

answer is ambiguous so write answer as 3x10^2

cm^2 bc this makes it clear there is only 1 sig fig.

If you leave number in the display without rounding

every step in a series of calculations you can round at

end of calculations and be ok.

“Factor-Label Method” aka “Conversion Factor Approach” aka

“Dimensional Analysis”:

(all reference to the same thing which is solving problems by

looking at the units instead of plugging measurement into

equations)… Ask not what the equation is, but how do I get the

units for the quantities I’ve measured into the units of what I want.

Diagram 1.2 and 1.3

Doing this method if you get the right units you probably got the

right answer.

Ch. 1 Reading Notes 8/12/13 8:22 PM

1.2 The Scientific Method = a systematic approach to research

Carefully define problem

Perform experiments and record data

o data obtained may be…

qualitative = general observations about the system.

quantitative = comprising numbers obtained by various

measurements of the system.

Interpretation (attempt to explain observed phenomena)

o Form a hypothesis= tentative explanation for a set of

observations.

Perform more experiments to test validity of hypothesis

Summarize info as a law

o “law” = a concise verbal or mathematical statement of a

relationship between phenomena that is always the same

under the same conditions.

hypothesis after many tests may become theory

o “theory” = a unifying principle that explains a body of facts

and/or those laws that are based on them.

**Scientific progress is seldom, if ever, made in a rigid, step by step

fashion

3 steps:

o observation> representation> interpretation… cycle

that continues!

1.3 Classifications of Matter:

Matter = anything that occupies space and has mass

o In principle, all mater can exist as solid, liquid, or gas

Chemistry = the study of matter and the changes it undergoes

melting point = the temp at which a solid will melt to form a liquid

boiling point= goes from liquid to gas

Water is unique… molecules are more closely packed in liquid state

than solid state

substance = matter that has a definite or constant composition

and distinct properties; substances differ from one another in

composition.

at present over 66 million substances are known

ex: water silver, table salt, carbon dioxide

mixture = a combination of 2 or more substances in which the

substances retain their distinct identities.

o do NOT have constant composition (ex: air from different

cities)

o ex: soft drinks, milk, cement

o homogeneous or heterogeneous

homogeneous= SAME composition throughout

solution (sugar in water)

heterogeneous= Composition is NOT uniform

throughout (sand and nails)

o *ANY mixture can be created and then separated by PHYSICAL

means to where the components of the mixture have the

same composition and properties as they did to start with

Elements and Compounds

Element= a substance that cannot be separated into simpler

substances by chemical means

o 118 identified (periodic table)

o only 5 elements (oxygen, silicon, aluminum, iron, and calcium,

comprise over 90% of the earths crust.

Compound = a substance composed of 2 or more elements

chemically united in fixed proportions

o Can be separated only by chemical means into pure

components.

1.4 Physical and Chemical Properties of matter:

physical property = can be measured and observed without

changing the composition of a substance.

Chemical property = to observe this property we must carry out a

chemical CHANGE (ex: burning hydrogen and oxygen gas… originals

will have vanished and different substance water will be created.

Impossible to recover originals by physical change)

All measurable properties of matter fall into 2 categories: extensive

and intensive properties

o extensive property = depends on how much matter is being

considered (mass, volume)

o intensive property = does NOT depend on the amount of

matter being considered. (temp, melting point, boiling point,

and density)

**Intensive property is IN dependent of mass

1.5 Measurement:

macroscopic properties = can be determined directly

microscopic properties = on the atomic or molecular scale, must

be determined by an indirect method (ex: density)

SI Base Units

Length Meter m

Mass Kilogram kg

Time Second s****** (NOT sec)

Electrical current Ampere A

Temperature Kelvin K

Amount of substance Mole Mol

Luminous intensity Candela cd

Mass and Weight

o Mass= measure of the quantity of matter in an object

o weight= the force that gravity exerts on an object

Volume = length (m) cubed… so SI unit is the cubic meter; non SI

unit is Liter (L)

o Liter= the volume occupied by one cubic decimeter

L and mL used normally for liquid volume. One Liter =

1000 milliliters or 1000 cubic centimeters

1 mL = 1000 cm^3

Density = the mass of an object divided by its volume (m/v)

o unit = kg/m^3 but that’s large so grams per cubic centimeter

and grams per milliliter more commonly used

Temp scales

o **complicated equations on page 11 to convert bw degrees

Celsius and Fahrenheit

o Absolute 0 on the Kelvin scale = -273.15 degrees Celsius

Celsius and Kelvin scales have units of equal magnitude

1.6 Handling Numbers:

Scientific Notation

o To add and subtract with notation first write each quantity

with the same exponent n.

o to multiply… multiply like normal and then add the exponents

together.

Significant figures = the meaningful digits in a measured or

calculated quantity. The last digit is understood to be uncertain. The

amount of uncertainty depends on the measuring device used.

o Guidelines for sig figs on pg. 15 of textbook

any digit that is not 0 is significant (845 cm. = 3 sigs)

zeroes bw nonzero digits are significant (606m= 3 sigs)

zeroes to the left of the first nonzero digit are not sig.

(0.08 contains 1 sig)… leading zeros never sig.

if a number is greater than 1, then all the zeroes written

to the right of the decimal point count as sig (3.040 has

4 sigs; 0.090 has 2 sigs)

For numbers that do not contain decimal points, the

trailing zeros (that is, zeros after the last nonzero

digit) may or may not be sig. (ambiguous) Use scientific

notation to make clear… 400 cm> 4x10^2 or 4.0x10^2

or 4.00x10^2)

Sig figs apply to measurements, but something that is

true by definition is exact and does not influence sig

figs.

o How to apply significant figures in calculations:

In addition and subtraction, the answer cannot have

more digits to the right of the decimal point than either

of the original numbers.

In multiplication and division, the number of significant

figures in the final product or quotient is determined by

the original number that has the smallest number of sig

figs.

exact numbers obtained from definitions (not

measurements) can be considered to have infinite sig

figs!

usually can round intermediate answers as well.

Accuracy and Precision

o accuracy = how close a measurement is to the true value of

the quantity that was measured (3 measurements can be 3

different extremes and very different therefore not precise but

average out to a number that is the value you should be

getting and accurate)

o precision = how closely two or more measurements of the

same quantity agree with one another. (3 measurements can

all be very similar and therefore precise but the reality is

you’re not getting the “accurate number”)

1.7 Dimensional Analysis in Solving Problems:

Dimensional analysis is the procedure we use to convert between

units in solving chemistry problems.

Lectures over chapter 2 9/9/13 8:22 PM

Missed lecture on Atomic Theory: (periodic table, elements)

9/9/13

Chemical Laws:

mass is conserved during chemical reactions

o ex: carbon + oxygen->carbon dioxide + heat + light

total mass -> total mass

Chemical compounds have definite composition

Atomic Theory

matter is composed of atoms (atom is smallest thing, nothing

smaller)

an element’s atoms are identical to each other and unique to that

element

o carbon atom is a carbon atom and an oxygen atom is an

oxygen atom but they are different

chemical reaction is the rearrangement of atoms (nothing created

nor destroyed)

Dalton predicted a chemical law that had not been observed (bc only applies

to certain chemical reactions)> law of “multiple proportions”

Cathode ray (beam of electrons) tube: physicists in 1890s> showed atoms

can come apart; gives us modern ideas about atoms

Structures of an atom:

more details on handout

mass number, A

atomic number, Z

properties of atom determined by number of protons in the nucleus

9/11/13

Common Compounds and Ions:

Allotropes = each of two or more different physical forms in which an

element can exist. (ex: O2 diatomic oxygen and O3 still oxygen but known as

ozone)

Naming Compounds (Nomenclature)

> 70 million compounds to name

common names (trivial names)

o trivial means don’t tell you what compound is made out of

o ex: water is common name for substance with molecular

formula H20; lime for CaO

systematic names:

o tell you something about the composition of the substance

o ex: water = dihydrogen monoxide; lime = calcium oxide

o *see handout

o 9/13/13

first word: cation positive ion

“ous” if it is lower charge

“ic” if it is higher charge

Stock System:

o chromium (II) and chromium (III)

o lead (II) and lead (IV)

could have “lead (II) chloride”

In class use stock system because it is easier but be able to

read and understand the other system with “ous” and “ic”

o naming continued

second word: anion (negative ion)

“ide”

monatomic

anion (negative ion)

“ate”

anion (negative ion)

contains oxygen

Acid…

oxy-acids

o “-ic acid”

other acids

o hydro-…-ic acid

Examples:

carbon monoxide = CO

carbon dioxide = CO2

o *see table 2.4 for prefixes

ICl = iodine monochloride

N2O5= dinitrogen pentoxide

o sometimes prefixes on both first and second word

o Use prefixes if elements are close to each other on the

periodic table and not forming ionic bonds and might be

multiple options.

Ch. 2 Atoms, Molecules, and Ions 8/12/13 9:22 PM

2.1 The Atomic Theory

5th century B.C. Democritus expressed belief that all mater

consists of very small, indivisible particles he named “atomos”

(uncuttable or indivisible)

1808> John Dalton (English scientist and school teacher) defined

atoms which marked the beginning of the modern era of chemistry.

Hypotheses Dalton’s atomic theory is based on:

o Elements are composed of extremely small particles, called

atoms

o All atoms of one element are identical (size, mass, chemical

properties), and the atoms of different elements are different

o Compounds made of atoms of more than one element. Ratio

of the numbers of atoms of any two of the elements present is

either an integer or a simple fraction.

extension of Proust’s law of definite proportions

which says “different samples of the same compound

always contain its constituent elements in the same

proportion by mass”

law of multiple proportions = if 2 elements can

combine to form more than one compound, the masses

of one element that combine with a fixed mass of the

other element are in ratios of small whole numbers.

aka… the compounds differ in the number of

atoms of each kind that combine (ex: carbon

monoxide and carbon dioxide)

o Chemical reaction involves separation combination, or

rearrangement of atoms; it does not result in their creation or

destruction.

law of conservation of mass = matter can be neither

created nor destroyed

2.2 The Structure of the Atom

atom = basic unit of an element that can enter into chemical

combination

Subatomic particles:

o electron

cathode ray experiment showed that electrons are

negative because beam attracted to positive anode

plate and repelled by plate bearing negative charges

very small mass (JJ Thomson and R.A. Millikan found)

o protons

Thompson’s “plum pudding” model was accepted

for years which says that the atom could be thought of

as a uniform, positive sphere of matter in which

electrons are embedded

image of plum pudding

another discovery about atoms:

o in 1910, Rutherford conducted gold foil experiment and

measured the scattering of alpha particles. Most of the alpha

particles passed trough with little or no deflection but a few

were deflected at wide angles or turned back completely.

Conclusion was that instead of being uniform like Thompson

thought all of the charged particles were concentrated in the

nucleus and most of the atom was empty space.

o neutrons

In 1932, James Chadwick provided proof and name for

neutrons (knew Rutherford’s model had a problem

because Hydrogen and Helium should have had mass

ratio of 2:1 according to Rutherford but it was 4:1 in

reality)

Radioactivity:

o spontaneous emission of particles and/or radiation

o 3 types of rays are produced by decay of radioactive

substances:

Alpha rays (positively charged particles)

Beta rays or Beta particles (are electrons)

high energy gamma rays (no charge)

2.3 Atomic Number, Mass Number, and Isotopes:

atomic number (Z)

mass number (A)j

hydrogen in most common form has 1 proton and 0 neutrons but

every other atomic nuclei have both protons and neutrons

Isotope = same atomic number but different mass numbers. aka

different number of neutrons

chemical properties of element determined primarily by number of

protons and electrons. Neutron do not take part in chemical

changes under normal conditions. (isotopes same types of

compounds and reactivities)

2.4 The Periodic Table:

chart in which elements having similar chemical and physical

properties are grouped together.

Horizontal rows = periods= by atomic number

o from left to right physical properties gradually change from

metallic to nonmetallic

vertical columns = groups or families = similarities in chemical

properties

3 categories:

o metal > good conductor of heat and electricity

majority of known elements

o nonmental > poor conductor of heat and electricity

17 elements

o metalloid> intermediate between properties of metals and

nonmetals

8 elements

Special names:

o alkali metals> first column (excluding hydrogen)

o alkaline earth metals> 2nd column

o halogens > column to the left of noble gases beginning with

F

o Noble gases or rare gases> far right column

2.5 Molecules and Ions:

Molecules:

o of all the elements, only the 6 noble gases (He, Ne, Ar, Kr, Xe,

Rn) exist in nature as single atoms “monatomic gases” Most

matter is composed of molecules or ions formed from atoms

o molecule = at least 2 atoms in a definite arrangement held

together by chemical forces (chemical bonds)

o Difference in molecule and compound:

molecule can be atoms of the same element or atoms of

different elements

Hydrogen gas

compound by definition is 2 or more elements

water

o Molecules are electrically neutral

o diatomic molecule = contains only two atoms (nitrogen,

oxygen, fluorine, chlorine, bromine, iodine)

can contain atoms of different elements (HCl and CO)

o polyatomic molecules = molecules containing more than

two atoms (ex: water, NH3)

Ions:

o atom or a group of atoms that has a net positive or negative

charge

o ionic compound = formed from cations and anions

o monatomic ions = contain only one atom

ex: Mg2+, Fe3+, etc

o Polyatomic ions = ions containing more than one atom

ex: NH4+, OH-, etc

o

2.6 Chemical Formulas:

chemical formula = composition of molecules and ionic

compounds in terms of chemical symbols

3 types:

o molecular formula = EXACT NUMBER of atoms of each

element in the smallest unit of a substance; the “true

formulas of molecules”

Can not talk about molecular formula if it does not exist

as a molecule (ex: NaCl) not for ionic compounds

o empirical formula = Simplest whole number ratio of atoms

(not necessarily the actual number of atoms in a given

molecule) Relative number of atoms of each element in any

amount of compound.

reducing the subscripts in the molecular formulas

can be determined for any type compound

o structural formula= shows how atoms are bonded to one

another in a molecule (H---O---H); different types are in use.

2 models:

o Ball and stick models

balls not proportional to the size of atoms, greatly

exaggerate space between atoms in a molecule; show

3d shape

technically still structural formula

o space filling models

MORE ACCURATE; but don’t show 3d shape well and

hard to put together.

Formulas of Ionic Compounds:

o usually the same as their empirical formulas

o arrangement of cations and anions is such that the

compounds are electrically neutral.

o *charges on cation and anion not shown in the formula

2.7 Naming Compounds: (image pg. 49)

3 categories:

o ionic compounds

binary compounds = formed from just 2 elements

metal cation then nonmetallic anion (ex: NaCl…

sodium chloride)

anion named by adding “ide”

ternary compounds = consisting of 3 elements

designate different cations of the same element

with roman numerals (stock system)

o molecular compounds

usually composed of nonmetallic elements

many are binary compounds

name of first element then second element root with

“ide”

ex: HCl > hydrogen chloride

common for one pair of elements to form several

different compounds

CO carbon monoxide

CO2 carbon dioxide

mono (1)

di (2)

tri (3)

tetra (4)

penta (5)

hexa (6)

hepta (7)

octa (8)

nona (9)

deca (10)

guidelines with prefixes:

“mono-“ omitted for first element

for oxides, ending “a” in the prefix

sometimes omitted

o ex: N2O4> dinitrogen tetroxide

o Acids

acid = substance that yields hydrogen ions (H+) when

dissolved in water

HCl

hydrogen chloride > molecular compound in

gaseous or pure liquid state

hydrochloric acid > dissolved in water molecules

break up into ions

Oxoacids = acids that contain hydrogen, oxygen, and

another element (central element)

ex: H2CO3 carbonic acid; H2SO4 sulfuric acid

usually written “H central O”

names all end with “ic”

Rules for naming 2 or more oxoacids with

different number of O atoms

add one O atom to “ic” acid = “per…-ic”

acid

o ex: HClO4 perchloric acid

remove one O atom… “-ous” acid

o ex: nitrous acid HNO2

remove two O atoms… “hypo… -ous” acid

o HBrO hypobromous acid

“ic” acids

H2CO3 carbonic acid

HClO3 chloric acid

HNO3 nitric acid

H3PO4 phosphoric acid

H2SO4 sulfuric acid

Oxoanions = anions of oxoacids

Rules for naming:

when all H ions removed anion ends with

“ate”

o ex: CO3 with neg 2 charge

(carbonate)

When all H ions removed from the “ous”

acid, anion ends with “-ite”

o ex: ClO2 (chlorite)

one or more but not all of the hydrogen ions

have been removed must indicate the

number of H ions present

See naming oxoacids and oxoanions chart on pg. 51

o Bases

substance that yields hydroxide ions when dissolved in

water

name element and then put hydroxide

ex:

NaOH Sodium hydroxide; KOH Potassium

hydroxide

o Hydrates

compounds that have a specific number of water

molecules attached to them

“anhydrous” means the compound no longer has water

associated with it

use prefixes to indicate how many water molecules

2.8 Introduction to Organic Compounds

hydrocarbons = contain only carbon and hydrogen atoms

alkanes = one class of hydrocarbons

all names end with “ane”

Chart on pg. 53 of first 10 straight-Chain Alkanes

Chapter 3 Lecture 8/12/13 9:22 PM

Stoichiometry:

dictionary: the relationship between the relative quantities of

substances taking part in a reaction or forming a compound,

typically a ratio of whole integers.

“atomic mass” = “atomic weight” = means relative mass

is no units… can write amu after as abbreviation

for atomic mass units

periodic table contains mean (average) of atomic

mass of all isotopes of an element

Mass of an atom:

o based on carbon-12 isotope which is assigned a mass of

exactly 12 atomic mass units (amu).

Normally we are not considered with just one atom but a compound or

molecule…

Molecular weight add up to get total:

carbon dioxide = C + 2(O) = (12.0 amu) + 2(16.0 amu) = 44.0 amu

Defining the mole:

NOT “the quantity represented by Avogadro’s number”

“The amount of substance that contains as many entities (whatever

you are measuring out… molecules if molecular, formula weights if

working with empirical formula, atoms if you have atoms) as there

are atoms in exactly 0.012kg of C-12

o SI, 1971

Restated differently… “the numerousness of entities in a system

equal to that of atoms in (exactly) 0.012kg of C-12

Mole is the unit of matter;

Mass of 1 mole of 1 element and 1 mole of another element can be

very different.

9/20/13 (missed class notes from Joseph)

Avogadro’s number = 6.02214129 x10^23

If what you have is only the relative amount of 2 substances than all you can

figure out in the lab is the empirical formula

Memorize 6.02x10^23 for Avogadro’s number = NsubscriptA

1 mole definition not measurement so don’t have to worry about significant

figures

9/27/13

3 types of balanced equations:

overall equations:

complete ionic equation:

net ionic equation: Generalized; you can not do directly in lab. all

reactions between strong acids and strong bases have the same net

ionic equation. (sour + bitter = salty)

CONCEPT: Balanced equations are indicating really a conservation of

mass. Matter is rearranged but conserved.

States (l, s, aq, g)> may be liquid under reaction conditions, but

label according to what it is on its own at room temperature.

chemical equation is…

o 1) counting atoms (ex: 2 formula units of KCLO^3 and that

will give 2 formula units of KCL and 3 molecules of oxygen)

IONS> formula units

non ionic> molecule

o 2) measuring substance in moles (ex: 2 moles, 2 moles, and

three moles)… Advantage can convert to mass using the

molar mass

o 3) measuring mass (plug in molar masses) shows that mass is

conserved!

WARNING: when doing conversions to find an answer be very careful that

you write moles that are equivalent based off the balance equation

(coefficients) and not just 1 mole.

Formula units:

empirical formula (lowest whole # ratio)

Ionic compounds do not exist as individual molecules; a formula unit

thus indicates the lowest reduced ratio of ions in the compound

if don’t know if a substance is ionic or molecular can refer to

everything as “formula unit.” However, calling all ‘molecules’ is

lying.

A unit of the formula listed (ex: NaCl)

Yields:

actual yield = in laboratory; reality; what you find

theoretical yield = equation used; not what you find in lab; mass

you SHOULD get for product if you have ____ mass of limiting

reagent

relative yield = ‘percentage yield’ = (actual yield / theoretical

yield) x 100%

9/30/13

Quantitative Use of “Chemical Equations”

Limiting Reagent = the reactants that determines the max amount

of products that can be formed

If you don’t already know which is the limiting reagent (maybe told

one is in excess) than first thing you need to find out is which one is

limiting

No matter how many reactions, only one limiting reagent;

often have excess of everything besides most expensive reagent

Chapter 3 Stoichiometry (pg. 60) 8/12/13 9:22 PM

3.1 Atomic Mass (pg. 61)

even the smallest speck of dust that eyes cant see contains as

many as 1x10^16 atoms (God you are amazing!)

it is possible to determine the mass of one atom RELATIVE to

another atom experimentally

o Step 1: assign a value to the mass of one atom to use as a

standard

atomic mass = aka. atomic weight = the mass of an

atom in atomic mass units (amu).

one atomic mass unit = a mass exactly equal to

one twelfth the mass of one carbon-12 atom

atomic mass of C-12 is 12 amu

ex: hydrogen 8.40% as massive as carbon 12 atom so atomic must

be .0840 x 12 amu= 1.008 amu

Average Atomic Mass

o periodic table contains the average atomic mass of all

isotopes of that element (not actual atomic mass for any atom

in any of the isotopes)

multiply percent abundance of each isotope times the

atomic mass of that isotope and add up for the average.

MUCH more abundance of C-12 than C-13

3.2 Avogadro’s Number and the Molar Mass of an Element:

bc atoms so small it is convenient to have a special unit to describe

a very large number of atoms. Chemists measure atoms and

molecules in moles

o mole= amount of substance that contains as many

elementary entities (atoms, molecules, or other particles, as

there are atoms in exactly 12 g (or .0012 kg) of the carbon-12

isotope.

actual number of atoms in 12 g of carbon-12 is

determined experimentally and called Avogadro’s

number. (6.022 x 10^23)

o CONCEPT: Because atoms and molecules are so tiny we need

a huge number to study them in manageable quantities.

recap:

1 mole of carbon-12:

o mass of exactly 12 g

molar mass: mass (in grams or kilograms) of 1 mole of

units (such as atoms or molecules) of a substance

molar mass of Carbon-12 (in grams) = atomic

mass in amu

ex: atom mass of Na= 22.99 amu and its molar

mass is 22.99 g (found on periodic table)

CONCEPT: If we know the atomic mass of an

element, we also know its molar mass.

in calculations, molar mass units = g/mol or

kg/mol.

o contains 6.022 x 10^23 atoms

CONCEPT : knowing the molar moss and A’s number we can calculate the

mass of a single atom in grams; A’s number can be used to convert from the

atomic mass units to mass in grams and vice versa

ex: (12 g carbon-12 atoms)/(6.022x10^23 carbon-12 atoms) =

1.993x10^-23 g

*Conversion factor b/w grams and moles is molar mass

On problems try beginning by thinking about what the molar mass is.

3.3 Molecular Mass

Molecular mass= “molecular weight” = in molecules “formula mass” =

sum of the atomic masses (in amu) in the molecule

multiply atomic masses of each element by the # of atoms of that

element and sum all elements

ex: water> 2(1.008amu) + 16.00 amu = 18.02 amu

CONCEPT: The molar mass of a compound (in grams) is numerically equal to

its molecular mass (in amu)

CONCEPT: molar mass> grams; molecular mass> amu

For ionic compounds that do not contain discrete molecular units, we use the

term “formula mass”

formula unit of NaCl is one Na+ ion and one Cl- ion. Formula mass is

mass of one formula unit. (amu)

3.4 The Mass Spectrometer

most direct and most accurate way to determine atomic and

molecular masses

a gaseous sample is bombarded by a stream of high energy

electrons. Collisions bw electrons and gaseous atoms (or molecules)

produce positive ions by dislodging an electron. Positive ions

accelerated by 2 oppositely charged plates as they pass through.

Emerging ions are deflected into a circular path by a magnet.

Radius depends on the charge to mass ratio. mass of each ion is

determined from the magnitude of deflection

o wider curve> smaller charge/mass ratio

o smaller curve> greater charge/mass ratio

first one developed in 1920s

amount of current (detector registers) directly proportional to the

number of ions, enables us to determine the relative abundance of

isotopes.

masses of molecules can also be determined.

3.5 Percent Composition of Compounds:

percent composition = percent by mass of each element in a

compound [(n x molar mass of element / molar mass of the

compound) x 100]

o n= number of moles of element in 1 mole of compound

o ex: water H20 has two moles of hydrogen… Subscript.

Both molecular and empirical formula tell us percent composition by

mass of compound.

can check: do percentages add up to almost exactly 100?

CONCEPT: if given percent composition can find the empirical

formula; but always empirical formula!

o convenient to assume we started with 100 g

o 1) mass percent

convert to grams and divide by molar mass

o 2) moles of each element

divide by the smallest number of moles

o 3) mole ratios of elements

change to integer subscripts

o 4) empirical formula

CONCEPT: in a chemical formula, the subscripts represent the ratio of the

number of moles of each element that combine to form one mole of the

compound.

3.6 Experimental Determination of Empirical Formulas:

1) chemical analysis tells us # of grams of each element present in

a given amount of a compound

2) convert quantities in grams to number of moles of each element

3) find empirical formula (method assuming 100g)

empirical literally means “based only on observation and

measurement”

Determination of Molecular Formulas:

o must know approximate molar mass

o use method assuming 100 g when given percent compositions

to find empirical formula (molecular may be same or may be a

multiple) so find the molar mass of the empirical formula and

see how it compares to the molar mass of the compound. (ex

3.11 pg. 74)

3.7 Chemical Reactions and Chemical Equations:

chemical reaction = process in which a substance (or substances)

is changed into one or more new substances.

chemical equation = uses chemical symbols to show what happens

during a chemical reaction

Writing chemical equations:

o to conform with law of conservation of mass, there must be

the same number of each type of atom on both sides of the

arrow (‘balance equation’)

o CONCEPT: ratio of the number of molecules is equal to the

ratio of the number of moles (coefficient can read to mean

either)

o reactants = starting materials in chemical reaction

o products = substance formed as a result of a chemical

reaction

o aq denotes not any liquid (l) but water environment

o Balancing formulas:

CONCEPT: changing the subscripts would change the

identity of the substance

use the simplest possible set of whole number ratios to

balance

3.8 Amounts of Reactants and Products:

Stoichiometry = quantitative study of reactants and products in a

chemical reaction

whatever unit given use moles to calculate the amount of product

formed in a reaction. (mole method = stoichiometric coefficients in

a chemical equation can be interpreted as the number of moles of

each substance)

means “stoichiometrically equivalent to” or “equivalent to”

o CONCEPT: molar mass is conversion factor between moles

and grams!

o General steps for solving a stoichiometry problem:

1) write a balanced equation for the reaction

convert the given amount of reactant in whatever unit it

is given to moles

use the mole ratio from the balanced equation to

calculate the number of moles of product formed

convert moles of product to grams (can use molar

mass) or other units of product.

o CONCEPT: 2 main useful tools are molar mass and mole ratio

3.9 Limiting Reagents

in real life, reactants are usually not present in the proportions

indicated by the balanced equation but the goal is to produce the

max quantity of a useful compound so there is frequently a large

excess of one reactant to ensure the more expensive reactant is

completely converted.

o limiting reagent = the reactant used up first in a reaction

o excess reagents = reactants present in quantities greater

than necessary to react with the quantity of the limiting

reagent.

image> men and women dancing partners

to determine which is limiting and which is excess when

given a equation take amount of moles initially (info

given to you in prob) and use equivalency of coefficients

to find moles of product. (one that’s less is limiting)

3.10 Reaction Yield:

theoretical yield = amount of product that would result if all the

limiting reagent reacted (max obtainable yield predicted by the

balanced equation)

actual yield = amount of product ACTUALLY obtained from carrying

out a reaction (almost always less than theoretical)

percent yield = proportion of actual yield to theoretical yield

o % yield = (actual yield/ theoretical yield) x 100%

o temp and pressure can affect

Ch. 4 Reactions in Aqueous Solutions 8/12/13 9:22 PM

9/30/13

Another kind of common calculation with chemical equation planning…

not directly involved with stoichiometry

Rxns w/ solutions

o pure substances: masses, atomic wts --> numerousness

o solutions: volumes, concentrations -->

concentration =

amount of solute in a given amount of solution OR

in a given amount of solvent

Units : no one standard unit; could be by mass in

solution (ex: vinegar) or could be confusing like

mass and volume (i.v bags)

chemist like to make concentration mole of solute

in a given volume of solution

unit: Molar, italics M, or underlined capital

M. (amount of solute, in moles/ volume of

solution, in Liters)

o most of the time measuring in mL (a

thousand make up L)

ex: .1 M CuSO4

o means 1000 mL solution is equivalent

to .1 mole CuSO4

o half a mole would be 500 mL (.05)

Definitions:

solution = a homogenous mixture; at least 2

components and one is normally in smaller

amount

solute = substance(s) present in a smaller

amount

may be numerous solutes! They would each

have unique concentrations.

solvent = substance present in a larger amount

10/2/13

4.5 and 4.6

Planning Reactions – Solutions (sans “Titrations”)

“ity” way we make nouns out of adjectives

o molar concentration or molarity

o (amount in moles of solute)/ (volume of solution in Liters)

o put line under M for unit

Dilution Calculations:

amount in moles (of solute) before dilution is always equal to

amount in moles (of solute) after dilution

Amount of solvent is what changes

amount of solute before dilution = amount of solute after dilution

Ch. 4 Reading (4.1 - 4.5) 8/12/13 9:22 PM

4.1- 4.4 Types of Chemical Reactions> see handout instead of text

4.1 General properties of aqueous solutions

solute = substance present in a smaller amount

solvent = substance present in a larger amount

all solutes that dissolve in water fit into one of 2 categories:

o electrolytes > when dissolved in water results in solution

that can conduct electricity

acids and bases electrolytes

strong electrolytes > solute 100% dissociated into

ions in solution (compound breaks up into cations

and anions)

all ionic compounds

weak electrolytes >

o nonelectrolytes > when dissolved in water results in

solution that can NOT conduct electricity

o hydration = process in which an ion is surrounded by water

molecules arranged in a specific manner

o chemical equilibrium = a chemical state in which no

change can be observed (break up as fast as recombine)

4.2 Precipitation Reactions:

formation of insoluble product, or precipitate (solid that separates

from solution)

usually involve ionic compounds

metathesis reaction = double displacement = exchange of parts bw

two compounds

Solubility = max amount of solute that will dissolve in a given

quantity of solvent at a specific temp

Equation types:

o molecular

formulas of compounds written as though all species

existed as molecules or whole units

shows reagents but not what is actually happening in

solution

o ionic

shows dissolved species as free ions

spectator ions = ions that are not involved in the overall

reaction (appear on both sides unchanged)

o net ionic

only what actually takes part in reaction

ions that are the same on both sides removed (take out

spectator ions)

CONCEPT: when ionic compounds dissolve in water they break apart

into their component cations and anions

4.3 Acid-Base Reactions: (proton transfer processes)

Acids:

o donate H+ ions when dissolved in water

o sour taste

o react with certain metals such as zinc, magnesium, and iron

to produce hydrogen gas

o aqueous acid solutions conduct electricity

Bases

o bitter

o slippery

Bronsted’s definitions do not require acid or base to be in aqueous

solution

neutralization reaction> acid and base

generally acid/ base reactions produce a salt

4.4 Oxidation- Reduction Reactions (electron transfer reactions)

redox = a process in which one substance or molecule is reduced

and another oxidized; oxidation and reduction considered together

as complimentary processes: redox reactions involve electron

transfer.

oxidation = loss of electrons

reduction = gain of electrons

o OIL RIG (oxidation is loss reduction is gain)

reducing agent = donates electrons

oxidizing agent = accepts electrons (causing element it accepts

from to be oxidized)

Common Oxidation-Reduction Reactions

o Combination

= 2 or more combine to form a single product

o Decomposition

= breakdown of a compound into components

o Combustion

= substance reacts with oxygen, usually with the

release of heat and light to produce a flame

all redox

o Displacement (replacement)

= ion (or atom) in a compound is replaced by an ion (or

atom) of another element

hydrogen, metal, and halogen

o Exchange

positive and negative ions switch places

4.5 Concentration of Solutions

concentration of solution = amount of solute present in a given

amount of solvent or a given amount of solution

molarity (M) = molar concentration = number of moles of solute

per liter of solution (solution not solvent)

o depends on temp

o unit: moles/ liter

Dilution of Solutions:

o procedure for preparing a less concentrated solution from a

more concentrated one

o (moles of solute b4)(volume of solution initially) = (moles of

solute after dilution)(volume of solution finally)

Energy Relationships in Chemical Reactions(Thermodynamics) 10/2/13 2:24 PM

Started… Thermo (heat) dynamics (motion)… the movement or

motion of heat.

Now… study of movement of all kinds of energy