Unit 1: Matter & Measurement

1. Substance:

a. Homogeneous matter

b. ONLY ELEMENTS & COMPOUNDS ARE SUBSTANCES

2. Element:

a. Atoms of a single type

b. CANNOT be broken down by a chemical change c. Particle diagrams:

element-Ca

elemet-N2

3. Compound:

a. Substances made of 2 or more elements chemically combined in a definite proportion

b. Composition is fixed (homogeneous composition)

c. CAN be broken down by a chemical change

d. Can be represented by specific formulas

e. Particle diagram:

4. Mixture: 2 or more distinct substances physically combined

a. General properties

i. Composition can vary

ii. Retains the properties of their components

iii. Separation methods:

1. Magnet – can take iron out of a mixture

2. Distillation – used to separate substances with different boiling points &

molecular polarities

3. Filtration – used to take an INSOLUBLE substance out of a mixture

4. Chromatography – used to separate different sized molecules in a mixture

5. Centrifuge – used to separate substances by density

b. Types of Mixtures:

i. Homogeneous-uniform composition

1. solution that is "completely dissolved"

ii. Heterogeneous-non-uniform composition

compound-H2O

c. Particle diagrams:

Heterogeneous

mixtures

Homogeneous

mixtures

element -compound

element-compound

element-element

element-element

compound-compound compound-compound

5. Physical Changes: DO NOT result in the formation of a new substance; can be either reversible or

irreversible

a. Examples:

i. Phase change (reversible)

1. Evaporation/Condensation

2. Melting/Freezing

3. Sublimation/Deposition

ii. Dissolving (reversible)

iii. Cutting into smaller pieces (irreversible)

6. Physical Properties: are constant & do not change

a. Examples:

i. Color

ii. Odor

iii. Solubility

iv. Melting Point/Boiling Point

v. Hardness

vi. Conductivity

vii. Phase

viii. Density-Table T

1. in a given volume the greater the number of atoms, the greater the density does

not result in the forming of a new substance

7. Chemical Changes: results in the formation of a new substance & indicated by the words "reacts",

"forms", "combines"

a. Chemical changes are independent of the amount of the substance

i. 300 kg Al2O3 & 400 kg Al2O3 have the same chemical properties

b. Examples:

i. Electrolysis

ii. Rust

iii. Combustion

8. Chemical Properties: behaviors exhibited when substances react

a. Examples:

i. Combustibility

ii. Reactivity

iii. Reaction rate

iv. Ability to oxidize

v. Ability to rust

vi. Reacts with acids/bases

9. Phases of Matter

Solid Liquid Gas

Definite shape & volume, particles

vibrate

No definite shape, definite volume,

particles vibrate & rotate

No definite shape or volume,

particles vibrate, rotate, translate

10. Density

1. calculating volume

a. regular solid=l x w x h

b. irregular solid=displacement method object is put into water

i. volume(object)=v(final)-v(initial) Example: What is the density of a substance with a mass of 900 grams and a volume of 30cm

3? Units:

volumeL, mL, cm3

massg, kg

densityg/mL, g/ cm3

d = 900 = 30 g/cm3

30

* Densities of elements can be found on Table S!

d = m

V (Table T)

11. Percent Error (Table T)

% error=(measured-accepted) x 100

accepted

measured=experimental=determined

accepted=standard=expected=actual

12. TERMS

(s)=solid

(l)=liquid

(g)=gas

(aq)=aqueous refers to a substance in water & therefore is ALWAYS a mixture

Unit 2: Atomic Structure

1. Models

a. Dalton: Hard Sphere

i. All elements made of atoms. Atoms of one element are alike, but different from atoms of

different elements.

b. Thomson: Plum Pudding

ii. electrons have negative charge

c. Rutherford: Gold Foil Experiment

iii. atoms are mostly empty space, nucleus is dense and positively charged

d. Bohr: Planetary Model

iv. electrons are in fixed orbits around nucleus

e. Modern: Wave Mechanical Model (Electron Cloud)

i. Protons & neutrons in nucleus

ii. Electrons in orbitals around nucleus

1. orbitals: areas of high probability where electrons will most likely be found

2. Atom: a. positively charged nucleus surrounded by negatively charged electrons

b. electrically neutral: the number of protons=the number of electrons; (p=e-)

c. subatomic particle:

i. proton

1. the mass of a proton is 1 amu (1 u) & has approx. the same mass as a neutron; the

charge is +1

2. number of protons in an atom is used to identify the element

ii. neutron

1. the mass of a neutron is 1 amu (1 u) & has approx. the same mass as a proton; the

charge is 0

iii. electron (also called a beta particle)

1. the mass of an electron is 1/1836 amu (1/1836 u); the charge is -1

See Table O for symbols

Proton Neutron Electron

Charge +1 0 -1

Mass 1 amu 1 amu 1/1836 amu

Location nucleus nucleus Orbitals/electron

cloud

3. Isotopes:

a. Have the same atomic # (same # of protons) but different mass #s BECAUSE THEY HAVE

DIFFERENT NUMBERS OF NEUTRONS b. have the SAME BRIGHT-LINE SPECTRUM because they are the SAME ELEMENT

i. the isotope that is more abundant will have an atomic mass closer to the atomic mass

listed on the periodic table

ii. Example 1:

a. Boron-10 & Boron-11; since atomic mass of Boron (B) is 10.81, Boron 11

is more abundant

iii. Example 2: H-1, H-2, H-3 (the number indicates the mass #)

4. Atomic Mass (average atomic mass): weighted average mass of the naturally occurring isotopes of an

element

a. Example: The isotopes of chlorine are Cl-35 and Cl-37. In nature 75 % of all Cl atoms have a

mass of 35 amu (Cl-35) and 25 % of all Cl atoms have a mass of 37 amu (Cl-37). Calculate the

average atomic mass of Cl.

M.A.D. = Multiply (the % by the given mass # for each isotope)

Add (the answers from the multiplication problems)

Divide (the sum by 100)

(35 x 75) + (37 x 25)

100

5. Ground State/Excited State a. Bright-Line Spectrum: produced when electrons move from higher energy states to lower energy

states releasing energy in the form of light

i. excited state-when an electron gains enough energy to move from a lower energy level to

a higher energy level

ii. ground state-lowest energy configuration (listed on the periodic table as the electron

configuration)

6. Electron Dots (Lewis Structures) a. Used to show the number of valence electrons an element has

1 2

8 3

5 6

7 4

7. Notation

a. isotopic 235

U

b. argon-40 (40 is the atomic mass)

X

94

Unit 3: Periodic Table

1.

A. Periods go across. Elements in the same period have the same # of principal energy levels (PELs)

B. Groups go down. Also called families. Families of elements share similar chemical properties because

each element in a family has the same # of valence electrons.

2. Metals: to the left of the zig zag line

1. Most active metal: Fr (bottom left)

2. Metals have low ionization energies, low electronegativities, lose electrons, form positive ions, are good

conductors, are malleable & ductile, most are solid (mercury Hg is the only liquid).

3. Nonmetals: to the right of the zig zag line

1. Most active nonmetal: F (top right, group 17, not 18)

2. Nonmetals have high ionization energies, high electronegativities, gain electrons, form negative ions

with metals, share electrons with nonmetals, lack luster, are brittle, are poor conductors, can be solid,

liquid (bromine), or gas (H, He, N, O, F, Cl, group 18)

4. Metalloids: touch the zig zag line (not Al, though)

1. Have some properties of both metals and nonmetals

5. Noble Gases: Group 18

1. Monatomic gases that don’t usually react with other elements. Kr & Xe can form compounds with F or

O.

2. Have complete valence PEL = stable configuration, stable octet

3. Have dispersion forces between molecules (weak forces of attraction). Dispersion forces increase with

larger mass so boiling point of Kr is higher than He.

6. Ions: do not have equal number protons & electrons & therefore they have a charge.

+ ions = less electrons than protons

- ions = more electrons than protons

7. Allotropes: forms of the same element that have different molecular formulas.

1. Ex: carbon can be in many forms, coal, diamond, graphite, buckminsterfullerene.

a. In all forms, atoms are arranged differently, which leads to different physical & chemical

properties.

8. Properties of Substances

Physical Chemical

Color, odor, solubility, MP, BP, hardness,

conductivity, phase, density

Things you can observe without changing the

substance into something else

Combustibility, reactivity, reaction rate

Things that describe how an element behaves in a

chemical reaction

9. Groups:

1. Elements in the same group react similarly because they have the same # of valence electrons.

2. Down a group

a. radius increases due to an increase in the # of PELs.

b. Electronegativity and ionization energy decrease down a group b/c the valence electrons are

farther away from the nucleus.

c. For metals, activity increases down a group because the electrons become more shielded by inner

PELs, making it harder to hold onto the valence electrons.

3. Group 1: alkali metals = most active metals, form +1 ions, lose electrons easiest

4. Group 2: alkaline earth metals = very active metals, form +2 ions

5. Transition metals: groups 3-12: make solutions with colorful ions

6. Group 17: halogens = most active nonmetals (gain electrons easiest)

7. Group 18: noble gases = mostly inactive, stable electron configurations

10. Periods:

1. as you go across a period (LR)

a. atomic radius generally decreases

b. electronegativity and ionization energy increase across a period due to increase in nuclear

charge.

c. More positive protons in the nucleus mean a tighter hold on the negative electrons.

11. Ionization Energy: amount of energy required to remove an atom’s most loosely held electron.

1. Lower IE, easier to take away an electron.

2. IE decreases down a group & increases across a period.

12. Electronegativity: attraction of a nucleus for electrons.

1. Higher electronegativity, tighter hold on e-.

2. Electronegativity decreases down a group & increases across a period.

13. Atomic Radius: half the distance between nuclei that are next to each other (or the distance between the

nucleus & its valence e-)

1. Radius increases down a group & decreases across a period.

14. Ionic Radius

Metals Nonmetals

Since metals lose e- to become positive ions, the ionic

radii for metals are smaller than the radii for metal

atoms

Loses electrons +

Metal atom metal

ion (+)

Since nonmetal atoms gain electrons to become ions,

the ionic radii for nonmetals are larger than the radii

for the nonmetal atoms

Gains electrons Nonmetal atom

nonmetal ion (-)

-

Unit 4: Chemical Bonding

1. Chemical Bond: attraction between the nucleus of one atom and the electrons of another atom.

BARF-Break (bond) Absorb (energy), Release (energy) Form (bond)

Energy Endothermic Exothermic

Definition Bonds are BROKEN & ENERGY

is ABSORBED (reactant)

Bonds are FORMED & ENERGY IS

RELEASED (product)

Is it spontaneous No Yes

General reaction A + energy -->B + C A + B -->C + energy

Examples Cl2 -->Cl +Cl Na(s) + 0.5Cl2(s) --> NaCl(s) +energy

CaCl2 +NaHCO3 +H2O -->CaCO3 +CO2 +energy

(energy is absorbed as bonds are broken & energy is released as bonds are formed)

2. Types of Molecules:

SNAP-Symmetrical (charge distribution) Nonpolar, Asymmetrical (charge distribution) Polar

Covalent

Bonding

Polar Nonpolar

Definition Unequal sharing of electrons Equal sharing of electrons

Bond 0.41-1.99

(electronegativity difference)

0-0.40

(electronegativity difference)

Degree of

polarity

most polar-->largest difference in

electronegativity)

Least polar-->smaller difference in

electronegativity

Bond Examples O-H=3.4-2.2=1.2

F-H=4.0-2.2=1.8 (more polar--

>largest difference in

electronegativity)

C-H=2.6-2.2=0.4

Cl-Cl=3.2-3.2=0.0 (least polar-->smallest

difference in electronegativity)

Molecule 1. Asymmetrical

2. unbonded electrons or LONE

pairs around the central atom

1. Symmetrical

2. No unbonded electrons or No LONE pairs

around the central atom

Molecule

Examples

H2O, HCl Cl2, H2, CO2, CCl4, CH4

3. Types of Bonds

TICS-->Transfer (of e-) Ionic, Covalent Shared (e

-)

Type of

Bonding

(Intramolecular

Force)

Ionic Covalent (Molecular) Metallic

Definition Compounds formed by the

TRANSFER of

ELECTRONS

Compounds formed by the

SHARING of

ELECTRONS

Type of chemical bond

between atoms in a metallic

element, formed by the

valence electrons moving

freely through the metal

lattice ("Sea of Mobile

Valence Electrons")

Types 1. Metal + Nonmetal

2. Metal + Polyatomic Ion

3. Polyatomic ion +

Polyatomic

4. Polyatomic Ion +

Nonmetal

Nonmetal +Nonmetal Metals

Examples NaCl, KI, HCl, NaOH (has

both ionic & covalent), BaSO4

(has both ionic & covalent)

He, CH4, CO2, CCl4, CH4,

NH3

Hg(l), Au(s), Pb(s)

Lewis structures

[K]+[ Br ]

-

(+1-1=0)

**compounds are electrically

neutral (sum of all charges =0)

H – O H - N – H

H H

(H2O) (NH3)

Number of

Shared

Electrons/Shared

pairs of electrons

O = C = O

Between C & O atoms:

1. 4 electrons 2. 2 pairs

Electronegativity

difference

>2.0 0-0.4=nonpolar

0.41-1.99=polar

-----

Ionic Character Greatest

(large electronegativity

Least

(small electronegativity

-------

difference between two elements)

difference between two elements)

Conducts as

Solid

No No Yes

Conducts as

Liquid

Yes No Yes

Conducts in

Solution

Yes No

Hardness Hard Soft Hard

Melting Point /

Boiling Point

High Low Very high

Covalent Properties: SPLash – Soft, Poor conductor, Low MP & BP (due to weak IM forces)

4. Intermolecular Forces: attraction between molecules

a. The stronger the IMF, the higher the melting point & boiling point

b. IMF increases from GasLiquidSolid

i. Hydrogen bonding = strongest intermolecular force ONLY BETWEEN Hydrogen &

Fluorine, Oxygen, or Nitrogen (Hydrogen bonding is FON)

ii. Dipole-Dipole = partial positive end of one polar molecule attracts partial negative end of

another polar molecule. Stronger than dispersion forces.

iii. Dispersion forces = weak forces of attraction between all types of molecules. ONLY

between nonpolar molecules. Dispersion forces are stronger for higher molecular

weights and when molecules are closer together.

iv. Molecule-Ion attraction = attraction between ions of an ionic compound and partial + or –

end of water molecule in solution. Ex. NaCl(aq)

5. Octet Rule: atoms tend to lose/gain or share electrons to complete the outer/valence shell of electrons to

become more stable (like noble gases).

6. Chemical Formulas:

a. Molecular = indicates # of atoms of each element in the molecule

b. Empirical = simplest ratio of atoms (lowest term)

7. Writing Formulas:

Ionic compound –

1. metal comes first, sum of oxidation #s = 0,

2. Roman numeral needed for oxidation # of metal if there is more than one option

Steps: Criss cross oxidation numbers, drop signs, reduce to lowest form

Unit 5: Stoichiometry 1. Compounds:

a. substances made of 2 or more elements chemically combined in a definite proportion

b. can be broken down into new substances by chemical means (change/reaction)

c. can be represented by specific formulas

2. Formulas – describe both which elements are present in compound as well as their ratio (quantity) in

compound

a. Molecular formula – indicate total # of atoms of each element in molecule (actual ratio of atoms)

Ex: NH3 (1N & 3H in molecule of ammonia), C6H12O6 (6C, 12H, 6O in molecule of glucose)

b. Empirical formula – indicate simplest whole number ratio of atoms in compound (fully reduced)

Ex: CH2O (simplest ratio for C6H12O6)

c. Structural formula – shows how atoms are joined or connected in molecule

Ex:

3. Finding Empirical Formula from Molecular Formula

Ex: Molecular is P4O10. What is empirical? (Divide subscripts by greatest common factor; in this case,

2.) = P2O5

4. Formula Mass – sum of atomic masses in a substance (amu)

Gram Formula Mass – (a.k.a. molar mass, gram atomic mass, gram molecular mass) – sum of atomic

masses in grams/mol

Ex: Formula mass of Al(NO3)3

Al 27 x 1 = 27 N 14 x 3 = 42 O 16 x 9 = 144 27 + 42 + 144

= 213amu *gram formula mass would be 213

grams/mole

5. Percent Composition = mass of part / mass of whole x 100 (Table T)

Ex: What is the percent by mass of magnesium in magnesium oxide (MgO)?

1. Find formula mass of the compound Mg 24 x 1 = 24 O 16 x 1 = 16 24 + 16 = 40

2. Substitute into formula 24 / 40 x 100 = 60%

Ex: What is the percent of water in BaCl2 ● 2H2O

1. Find formula mass of the compound Ba 133 x 1 = 133 Cl 35 x 2 = 70 H2O 18 x 2 = 36

133 + 70 + 36 = 239 2. Substitute into formula 36 / 239 x 100 = 15%

6. Finding Empirical Formula from Percent Composition

Formaldehyde consists of 40% carbon, 6.7% hydrogen, and 53.3% oxygen. What is its empirical formula?

1. Drop the % sign & divide each # by the element’s atomic mass C 40/12 = 3.33 H 6.7/1 = 6.7

O 53.3/16 = 3.33 2. Divide each answer by the smallest answer from step 1 3.33/3.33 = 1 6.7/3.33 = 2

3.33/3.33 = 1

3. The answers from #2 are the subscripts for each element in the formula C1H2O1 = CH2O

7. Finding Molecular Formula from Empirical Formula & Molecular Mass

Empirical formula is CH2 and molecular mass is 42. What is the molecular formula?

1. Find mass of empirical formula C 12 x 1 = 12 H 1 x 2 = 2 12 + 2 = 14

2. Divide molecular mass from problem by mass of empirical formula 42 / 14 = 3

3. Multiply subscripts of empirical formula from problem by answer from step 2 C1 x 3H2 x 3 = C3H6

8. Finding Moles from Gram Formula Mass # moles = given mass / gram formula mass (Table T)

How many moles of CH2 are in 56 grams of CH2?

1. Find gram formula mass of substance in problem (if not given): CH2 = 14

2. Substitute into formula: # moles = 56 grams / 14grams/mole = 4 moles

9. Finding Mass from Number of Moles # moles = given mass / gram formula mass (Table T)

How many grams of CH2 are in 2.5 moles of CH2?

1. Find gram formula mass of substance in problem (if not given): CH2 = 14

2. Substitute into formula: 2.5 moles = given mass / 14 = 35 grams

10. Balancing Equations

Coefficients in front of compounds in a chemical reaction indicate molecule or mole ratios for the

reaction. There must be equal numbers of molecules (or moles) on each side of the reaction (law of

conservation of matter).

Balance the equation, using the smallest whole-number coefficients: H2 + O2 H2O

Balance by inspection OR by listing elements/polyatomic ions underneath the equation and multiplying

by coefficients on each side until all parts are equal. 2H2 + O2

2H2O

Balanced equations fulfill the laws of conservation of mass, energy, and charge.

11. Solving Mole-Mole Stoichiometry Problems

Ex: Given Ca + 2H2O Ca(OH)2 + H2

How many moles of calcium are needed to completely react with 6 moles of water according to the

above equation?

1. Be sure the equation is balanced. (This one is balanced.)

2. Write the number(s) from the word problem directly above the substance to which they refer. Write an

“X” above the

substance in question.

3. Set up a simple proportion using the equation’s coefficients & the #s you added and solve. X 6

Ca + 2H2O

Ca(OH)2 + H2

X / 1 = 6 / 2 2X = 6 X = 3 moles

12. Density D = m/v (Table T)

Ex: What is the density of a substance with a mass of 900 grams and a volume of 30cm3? (vol can be in g

or cm3)

D = 900/30 = 30g/cm3

*Remember – the densities of elements can be found on Table S!! (Everyone forgets this!!)

13. Types of Chemical Reactions

a. Synthesis – 2 or more reactants combining to form 1 product: A + B AB 2Na + Cl2 2NaCl

b. Decomposition – 1 reactant breaking apart into 2 or more products: AB A + B 2NaCl 2Na + Cl2

c. Single Replacement – a free element reacts with a compound, & replaces a part of the compound (if the

free element is ABOVE

the corresponding element on Table J – a metal replaces a metal; nonmetal replaces nonmetal)

A + BX B + AX Fe + CuSO4 Cu + FeSO4 X + AY Y + AX F2 + 2LiCl Cl2 +

2LiF

d. Double Replacement – two compounds (in aqueous solutions) react and swap partners

AX + BY AY + BX AgF(aq) + NaCl(aq) NaF(aq) + AgCl(s)

Unit 6: Physical Behavior of Matter

1. Phases of Matter

Solid Liquid Gas

Definite shape & volume, particles

vibrate

No definite shape, definite volume,

particles vibrate & rotate

No definite shape or volume,

particles vibrate, rotate, translate

Phase changes: SL = fusion, melting LS = freezing, crystallization, solidification

LG = boiling, evaporation, vaporization GL = condensation

SG = sublimation GS = deposition

4. Energy

Q = m C ∆T (Tables B & T)

Example: How much heat is needed to raise the temperature of 10g of water from 3˚C to 5˚C?

Q = x, m = 10g, C = 4.18J/g˚C, ∆T = 2˚C

x = 10 x 4.18 x 2

x = 83.6 J

Energy in a phase change:

Q = mHf

Melting, freezing, crystallization, solidification

Equation Q = mHf Q = mHv

Key Words Melting, freezing, crystallization,

solidification

Boiling, condensing, vaporizing, evaporating

Phases Involved S L or L S L G or G L

Example How much heat is absorbed when 2g ice

MELTS?

Q = mHf

Q = x, m = 2g, Hf = 334J/g (Table B)

x = 2 x 334

x = 668J

How much heat is absorbed when 2g of

water BOILS?

Q = mHv

Q = x, m = 2g, Hv = 2260J/g (Table B)

x = 2 x 2260

x = 4520J

5. Temperature: measures average kinetic energy

Higher temperature means faster moving particles

Heat flows from an area of higher temperature to lower temperature

6. Heating Curve (Cooling Curve is opposite)

Remember – there is no

change in Kinetic energy

during a phase change – only

a change in potential energy.

If a question asks for a

melting, freezing, or boiling

point, it wants the

TEMPERATURE.

7. Liquids and Vapor Pressure: the vapor pressure of a liquid is the pressure the liquid exerts on the side of

the container it is in. When the vapor pressure reaches the atmospheric pressure, the liquid boils.

Table H vapor pressure/boiling temperatures for four liquids, as temperature increases, so does vapor pressure.

Use Table H to determine boiling temp of liquids at different pressures. The higher the boiling temp of a liquid

is at the same vapor pressure, the stronger the intermolecular forces among the molecules.

8. Gas Laws

Formula (Table T) (PTV stick!!)

P1V1 = P2V2 P up, V down

T1 T2 T up, V up

T up, P up

Use Kelvin Temps ONLY! (If temps are in Celsius, use formula from Table T to convert.)

Example: A 400mL gas sample is at STP. If the pressure is changed to 50.65kPa and the volume is changed to

551mL, what is the final temperature?

At first glance, you may not think that there is enough information here to solve the problem…but there is! The

question mentions STP, the values for which are found on Table A. Use the Kelvin temperature and the kPa

pressure, since kPa’s are used later in the problem.

P1 = 101.3 kPa P2 = 50.65 kPa (101.3)(400) = (50.65)(551)

V1 = 400 mL V2 = 551 mL 273 X cross multiply to solve

T1 = 273 K T2 = x (101.3)(400)(X) = (273)(50.65)(551) X = 188 K

9. Kinetic Molecular Theory (for ideal gases)

- gas particles are in random, straight-line motion

- there is a complete transfer of energy between colliding particles (elastic collisions)

- volume of individual gas molecules is negligible (they’re very tiny)

- gas particles have no force of attraction for each other

REAL GASES DEVIATE FROM THIS! (lightest ones – H & He – are closest to ideal)

Gas particles do have some volume & they have small attractive forces for each other

Real gases behave most like ideal gases at HIGH TEMP and LOW PRESSURE.

10. Avogadro’s Hypothesis: equal volumes of gases at the same temperature and pressure have the same

number of molecules.

Example: 1L O2 and 1 L Ne each have the same number of molecules at STP

11. Attraction between particles in the three phases

Solid: strongest forces of attraction

Liquid: attract each other less than in solids

Gas: attract each other the least

Unit 7: Solutions

1. Solution – homogeneous mixture composed of solute (what gets dissolved) and solvent (what does the

dissolving – usually water)

The more solute that is dissolved, the more concentrated the solution is. The less solute dissolved,

the more dilute it is.

Solubility – shows the most solute an amount of water can hold at a specific temperature

2. Temperature and Solubility – gases become LESS soluble at higher temperatures, & solids usually become

MORE soluble at higher temperatures (Table G)

Table G shows the maximum number of grams of various solutes that can be dissolved in 100g of

water at various temperatures. If a point for a substance is found to be ON the line for that substance, the

solution will be SATURATED. If the point is BELOW the line, it’s UNSATURATED, and if it’s ABOVE the

line, it’s SUPERSATURATED.

3. Types of Solutions

a. Saturated: contains the most solute a particular solvent can hold at a particular temperature. They’re at

equilibrium; the rate of crystallizing is equal to the rate of dissolving. If you add more solute, it won’t dissolve

– it’ll settle to the bottom, and the concentration of the solution will REMAIN THE SAME.

b. Unsaturated: contains less than the maximum amount of solute in a solvent at a certain temperature. You

can add more solute and it will dissolve, increasing the concentration of the solution.

c. Supersaturated: temporary condition resulting from the slow cooling of a saturated solution. In this

solution, there is more solute dissolved than should be allowed to dissolve. If you add more solute, all of the

excess solute will come out of solution.

4. Pressure and Solubility – Pressure ONLY affects solubility of GASES. HIGH PRESSURE, HIGHER

SOLUBILITY of gas.

5. Like Dissolves Like – Nature of solute & solvent determines solubility

Polar (+ end and – end) substances will determine other polar substances as well as ionic substances.

Nonpolar (symmetrical) substances will only dissolve other nonpolar substances.

Ionic or polar will not dissolve nonpolar.

6. Determining if a Substance is Soluble (Table F)

Left side of table Right side of table

Lists ions that are soluble. If ions present in the formula are Lists ions that are insoluble. If ions present

in the formula are

present in the column of soluble ions, and none of the present in the column of soluble ions, and

none of the

exceptions are present, the substance is soluble. exceptions are present, the substance is insoluble.

7. Determining Electrical Conductivity – All soluble ionic compounds according to Table F are electrolytes;

they will conduct electricity when dissolved in water. All insoluble ionic compounds according to Table F are

NOT electrolytes, as they will not dissolve in water.

8. Concentration of Solutions (Table T)

Molarity = moles solute / Liters solution

How many moles of solute are contained in 200mL of a 1M solution?

M = 1M L = 200mL .200L moles solute = X 1 = X/.2L X = .2 moles

9. % by Volume – use the % composition formula on Table T

part / whole x 100 *Read carefully! And remember, a solution is composed of SOLUTE AND SOLVENT!

What is the percent of juice in a beverage containing 10mL juice in 100mL of SOLUTION? 10 / 100 x 100

= 10%

What is the percent of juice in a beverage containing 10mL juice in 100mL of WATER? 10 / 110 x 100

= 9%

10. Parts per Million (ppm) (Table T)

Ppm = part / whole x 1000000 *Read carefully! (Same reason as above)

If a NaCl solution contains 0.050g NaCl in 900g solution, what is the concentration in parts per million?

0.050 / 900 x 1000000 = 55.6 ppm

11. Boiling Point Elevation

Dissolving a solute in water RAISES its boiling point. Ionic substances raise the boiling point higher than

covalents due to dissociation of ions.

12. Freezing Point Depression

Dissolving a solute in water LOWERS its freezing point. Ionic substances lower the freezing point more than

covalent substance due to dissociation of ions.

Unit 8: Acids, Bases, & Salts 1. Electrolyte: soluble substance whose solution conducts electricity due to presence of mobile ions. Acids,

bases, & salts.

2. Properties of acids & bases:

Acids Bases

Electrolytes, react w/metals above H on Table J to Electrolytes, red litmus turns blue,

phenolphthalein turns pink,

produce H2, blue litmus turns red, phenolphthalein react w/acid to form salt and water in

neutralization, slippery,

turns colorless, react w/bases in neutralization, sour bitter

3. Arrhenius Acids & Bases (see Tables K & L for common acids & bases)

Acids

Have H+ or H3O

+ ions in solution as only

positive ions

H + nonmetal or negative polyatomic ion from

Table E

OR ends in COOH

Ex: HCl, HNO3, CH3COOH

Bases

Have OH- ions in solution as only negative ions

Metal OR NH4+ + OH

Ex: NaOH, Mg(OH)2, NH4OH

4. Alternate Acid-Base Theory (Bronsted-Lowry Theory)

Acids

Proton (H+) donors, give away H

+ in reactions

Bases

Proton acceptors, take on H+ ions in reactions

BAAD – Bases Accept, Acids Donate

5. Amphoteric Substances: can act as both acids and bases in reaction

Ex: H2O….. H2O + H2O H3O+ + OH

-

6. Salt: ionic compound that doesn’t have H+ ions or OH

- ions in solution. Any ionic compound that isn’t an

acid or a base is a salt.

7. Neutralization: occurs when there are equal H+ and OH

- concentrations in solution

Acid + Base Salt + Water

HCl + NaOH NaCl + H2O 2HNO3 + Mg(OH)2 Mg(NO3)2 + 2H2O 3H2SO4 + 2Al(OH)3

Al2(SO4)3 + 6H2O

Net ionic equation is always H+ + OH

- H2O (can also be written as HOH)

8. Titration: laboratory method used to find the concentration of an acid or base using a solution with known

concentration. In

titration, measured volumes of acid and base are used, and an indicator signals the endpoint (when

neutralization has occurred).

Table T: nAMAVA = nBMBVB, where nA = # of H in the acid formula and nB = # of OH in the base formula.

Ex: How many milliliters of 4M NaOH are required to exactly neutralize 50mL of a 2M solution of H2SO4?

(2H)(2M)(50mL) = (1OH)(4M)(XmL)

X = 50mL of NaOH

9. Molarity: remember, regular molarity is M = mol solute/L solution

10. pH: indicates strength of acid or base. 0-6 acid, 7 neutral, 8-14 base.

Acid: H3O+ > OH

-, Base: H3O

+ < OH

-

pH is negative log of concentration of H+ in solution. If hydrogen (or hydronium) ion concentration is 1x10

-

3, pH is 3.

Every change of pH unit is a change in a power of 10 in hydrogen ion concentration.

As pH decreases, H+ concentration increases and OH

- decreases – solution becomes more acidic and less

basic.

As pH increases, H+ concentration decreases and OH

- increases – solution becomes more basic and less

acidic.

11. Indicators: chemicals that change color at various pH levels (Table M)

Ex: Methyl orange is red in a solution with a pH of 3.2 or less and is yellow in a solution with a pH of 4.4 or

more. A pH in

between these two levels will cause the indicator to be orange.

Ex: If a solution is blue in bromcresol green and yellow in bromthymol blue, what is the pH range of the

solution?

Bromcresol green is blue = pH 5.4 or more. Bromthymol blue is yellow = pH 6.0 or less. pH range is from

5.4-6.0.

Unit 9: Kinetics & Equilibrium 1. Kinetics – studies rates and mechanisms of reaction (mechanism = steps reaction takes)

2. Collision Theory – for reaction to happen, particles must collide with enough energy and in the proper

orientation

3. Factors Affecting Reaction Rates

a. Temperature: increasing temperature generally increases reaction rate; causes more rapid, more effective

collisions

b. Concentration: increasing concentration increases rate of reaction by making more particles available for

collision

c. Surface Area: powder reacts faster than solid block; more spots for collision to happen

d. Nature of Reactants: ionic substances usually react faster than covalent substances

e. Catalyst: reduces activation energy required for forward and reverse reactions by altering the pathway a

reaction takes. It has no

effect on concentrations at equilibrium. It has no effect on heat of reaction (∆H).

4. Heat of Reaction: amount of heat given off or absorbed in a reaction

∆H = HP – HR

5. Types of Reactions

a. Exothermic: a reaction that gives off more heat than it takes in. ∆H is negative (Table I) and is written on

the RIGHT side of eq.

b. Endothermic: a reaction that takes in more heat than it gives off. ∆H is positive (Table I) and is written on

the LEFT side of eq.

6. Activation Energy: the smallest amount of energy needed to start a reaction. Every reaction needs

activation energy. A catalyst

lowers activation energy, making the reaction go faster in both the forward & reverse direction.

7. Potential Energy Diagram

8. Spontaneous Reaction: takes place without help, exothermic (less energy at end), more entropy

(randomness) at end

Entropy increases when… solid liquid gas, substance aqueous solution, compound elements

9. Equilibrium: when the rate of the forward reaction is equal to the rate of the reverse reaction. The amounts

(concentrations) of

reactants and products are not necessarily equal to each other, but they remain constant over time.

10. Physical Equilibrium

a. Phase equilibrium: when rate of forward phase change is equal to rate of reverse phase change. It happens

on flat lines of

heating & cooling curves. Ex: H2O(l) H2O(g)

b. Solution equilibrium: when rate of dissolving a solute equals rate of crystallizing. A saturated solution is

in equilibrium (See any

point on any curve on Table G). Ex: NaCl(s) Na+

(aq) + Cl-(aq)

For solution equilibrium, increasing amount of solute has no effect on equilibrium; since the solution is

saturated, no more solid

can dissolve.

11. Chemical Equilibrium: when rates of forward and reverse chemical reactions are equal

12. LeChatelier’s Principle: if a system at equilibrium is subjected to stress, the reaction will shift to favor the

direction that will relieve the stress. If equilibrium shifts to the right, concentrations of substances on the right

will increase, and on the left will decrease (and vice versa). (ADD AWAY; TAKE TOWARD)

a. Change in concentration: If you increase the concentration of a substance, equilibrium shifts AWAY from

the substance that increased. Decreasing concentration of something shifts equilibrium TOWARD the

substance that decreased.

b. Change in temperature: If you increase the heat/temperature of a reaction, equilibrium shifts AWAY from

where the heat is written in the equation. (It shifts in the direction of the ENDOTHERMIC process.) (Vice

versa for decrease in temp.)

c. Change in pressure: Only affects gases. If you increase pressure on a reaction, equilibrium shifts to the

side with LESS MOLES (less total coefficients). (Vice versa for decrease in pressure.)

13. Reactions that go to completion: these cannot reach equilibrium, because they cannot go in the reverse

direction.

This can happen when: one product is a GAS, one product is WATER, or a PRECIPITATE is formed (Table

F - insoluble).

Unit 10: Redox & Electrochemistry 1. Finding oxidation #s

a. free elements 0

b. group 1 in a compound always +1, group 2 in a compound always +2

c. F is -1 in all compounds

d. sum of oxidation #s in compound is 0

e. sum of oxidation #s in polyatomic ion (Table E) is equal to overall charge on the ion

f. O usually -2 EXCEPT is -1 in peroxides (like H2O2) and +2 in OF2

g. H usually +1 EXCEPT is -1 when only following a group 1 or group 2 metal (like NaH or CaH2)

Ex: Find oxidation #s of elements in Na2S2O3

Na 2 x +1 = +2

S 2 x ___ = ___ All #s need to add to 0, so sulfurs must total +4. Therefore, each sulfur is

+2.

O 3 x -2 = -6

2. Oxidation: loss of electrons LEO

Metals tend to lose electrons and become more positive – the oxidation number goes up. Substance is a

reducing agent.

3. Reduction: gain of electrons GER

Nonmetals tend to gain electrons and become more negative – the oxidation number goes down. Substance

is the oxidizing agent.

4. Redox Reactions: involve both processes. One species loses electrons to another (simultaneous transfer).

Can recognize because there will be a change in oxidation # of element(s) indicating the loss and gain of

electrons.

Look for equation with a free element w/no charge on one side of reaction and that same element in a

compound or having a charge

on the other side.

Ex: Na0 + F

0 NaF What was oxidized, and what was reduced? (Pick from reactant side only!)

Na changed from 0 to +1, so Na0 was oxidized. F changed from 0 to -1, so F

0 was reduced.

5. Half Reactions: indicate parts of reaction that are just oxidation or reduction & show amount of electrons

involved.

Ex: NaCl + K KCl + Na

Na+Cl

- + K

0 K

+Cl

- + Na

0

Cl didn’t change, so ignore it.

Oxidation half reaction: K0 K

+ + 1e

-

Reduction half reaction: Na+ + 1e

- Na

0

6. Balancing Redox Reactions: same number of moles of electrons must be lost in oxidation and gained in

reduction – law of

conservation of mass, charge, and energy must be followed in all chemical reactions.

Ex: Given reaction: __Mg + __Cr+3

__Mg+2

+ __Cr, balance equation using smallest whole number

coefficients.

1. Write half reactions for oxidation and reduction

Mg Mg+2

+ 2e- Cr

+3 + 3e

- Cr

2. Multiply each half reaction by a whole number that makes the electrons on each side of the reactions

equal.

3(Mg Mg+2

+ 2e-) 2(Cr

+3 + 3e

- Cr) (makes 6 electrons for each reaction)

3. Insert those numbers in front of the elements in the overall equation.

3Mg + 2Cr+3

3Mg+2

+ 2Cr

7. Table J (again): can be used to determine if a reaction is spontaneous or not.

Metals higher on Table J are MORE easily OXIDIZED, and LESS easily REDUCED (and vice versa). The

ionic form of the element is the opposite.

Nonmetals higher on Table J are MORE easily REDUCED, and LESS easily OXIDIZED (and vice versa).

The ionic form of the element is the opposite.

8. Electrochemical Cells: 2 types – voltaic and electrolytic

Voltaic cell – uses spontaneous redox chemical reaction to convert chemical energy to electrical energy (is a

battery)

Table J – more active metal is oxidized & anode

2 half cells contain electrolyte solutions

Wire allows electrons to flow from anode to cathode

Electrodes are pieces of metal

Salt bridge allows for flow of ions

AN OX – oxidation at the anode (Zn) (- charge)

RED CAT – reduction at the cathode (Cu) (+ charge)

Anode gets smaller, cathode gets larger

Electrolytic cell – uses electricity to produce a nonspontaneous redox reaction (electrical energy converted to

chemical energy)

Ex: electroplating

1 container containing electrolyte solution

Wire allows electrons to flow from anode to cathode

Electrodes are pieces of metal

Battery or power source necessary

AN OX – oxidation at the anode (Ag) (+ charge)

RED CAT – reduction at the cathode (spoon – item to be plated) (- charge)

Positive ions from solution attract to neg charged cathode (where electrons

had traveled through the wire to the cathode.

Anode gets smaller, cathode gets larger

Unit 11: Organic Chemistry

1. Definition & Characteristics (Tables P, Q, R)

Study of carbon & carbon compounds, generally nonpolar, generally insoluble in water, but soluble in

nonpolar solvent,

nonelectrolytes (except for organic acids), low melting points, slow reactions b/c of high activation energy

2. Bonding: each C atom makes total of 4 covalent bonds

3. Hydrocarbons: contain only hydrogen & carbon, 3 types of straight-chain hydrocarbons (Table Q)

Alkanes

CnH(2n+2)

Single bonds – saturated

Ex: ethane C2H6

Alkenes

CnH2n

One double bond – unsaturated

Ex: ethane C2H4

Alkynes

CnH(2n-2)

One triple bond – unsaturated

Ex: ethyne C2H2

Ex: 2-methylbutane – draw main carbon chain first (butane), then add branches. Make sure each carbon has

4 bonds, filling in

available sites with hydrogens.

Ex: 3-hexene – hex means “6”, -ene tells you there’s a double bond, 3 means the double bond is after the 3

rd

carbon (counting from

either end).

4. Boiling Point: the more carbons, the higher the boiling point

Ex: Which of the following has the highest normal boiling point? CH4 C2H6 C3H8 C4H10

5. Isomers: same molecular formula, different structural arrangement. More carbons, more possible isomers.

Ex: Which is an isomer of ?

6. Functional Groups (Table R) atoms other than carbon and hydrogen that can be part of an organic

compound and give the

compound specific physical and chemical properties.

a. Alcohols (OH) – 1 or more hydrogens replaced with OH (hydroxyl) group, named by dropping “ane” and

adding “ol”

Use a # to indicate carbon(s) where OH group(s) attached 2-propanol

Monohydroxy alcohol (1 OH group), dihydroxy alcohol (2 OH groups), trihydroxy alcohol (3 OH groups)

b. Aldehydes – O atom double bonded to an END carbon, named by dropping “ane” and adding “al”

propanal

c. Ketones – O atom double bonded NOT to an end carbon, named by dropping “ane” and adding “one”

Use a # to indicate carbon where O is attached 2-butanone

d. Organic acids – COOH attached to an end, named by dropping “ane” and adding “oic acid”

butanoic acid

e. Ether – O atom connecting 2 carbon atoms, name 1st chain (yl ending), 2

nd chain (yl ending), “ether”

methyl ethyl ether

f. Halide – has halogen (Cl, Br, F, or I) attached somewhere, say # of carbon where halogen is, name

halogen, then name chain

1-chloropropane

g. Amine – has nitrogen somewhere in the chain

h. Amide – has nitrogen and a double bonded O in the chain

i. Amino acid – has an amine group and an organic acid group – just recognize, not name

j. Ester – has COO in middle of carbon chain. Name the part with ONLY Cs and Hs by saying proper chain

prefix with “yl” at end,

then name part with the double bonded O by saying proper prefix with “oate” at end

methyl ethanoate

7. Like Dissolves Like: small alcohols and acids are polar and dissolve in water. Straight chain hydrocarbons

are nonpolar and will not dissolve in water

8. Organic Chemical Reactions

a. substitution – occurs in alkanes when one atom replaces another. 2 reactants, 2 products CH4 + Cl2

CH3Cl + HCl

b. addition – occurs in alkenes and alkynes when a double or triple bond is broken, & 2 atoms from a

diatomic molecule add on to

the carbons that were involved in the double bond

c. fermentation – sugar alcohol + carbon dioxide, C6H12O6 C2H5OH + CO2

d. esterification – alcohol + organic acid ester + water, CH3OH + CH3COOH CH3COOCH3 + H2O

e. saponification – ester + base soap + alcohol (glycerol)

f. combustion – hydrocarbon + oxygen carbon dioxide + water, CH4 + 2O2 + CO2 + 2H2O

g. polymerization – nC2H4 (C2H4)n, where n is a very large number

addition polymerization – polymerizes by breaking double or triple bond

condensation polymerization – polymerizes by removing water (water is always a product)

Unit 12: Nuclear Chemistry 1. Stable & Unstable Nuclei – ratio of neutrons: protons determines stability, closer to 1 is stable, higher is

unstable, unstable nuclei

break apart or decay by releasing particles or energy, n:p ratio differs depending on isotope of element

2. Alpha Decay – nucleus releases alpha particles (Table N, O)

3. Beta Decay – nucleus releases beta particles

4. Positron Decay – nucleus releases positrons

5. Gamma Decay – nucleus decays and gives off high energy gamma rays with no mass or charge

6. Table N tells decay mode of radioisotopes

7. Separating Alpha, Beta, & Gamma Radiation – alpha has positive charge, is attracted to negative

electrode; beta has negative

charge, is attracted to positive electrode; gamma is uncharged, is undeflected with charges

8. Half Life – time it takes for half of the mass of a radioisotope to undergo its decay process, is CONSTANT

regardless of

environmental conditions (Table N, T) RHTOF (Radioisotope, Half life, Total time, Original mass, Final

mass)

Ex. If 100 grams of a radioisotope decayed to 12.5 grams in 90 years, what was the half life of the isotope?

R (Doesn’t matter), H (?), T (90yrs), O (100g), F (12.5g)

If you have 2 times, divide them; if you have 2 masses, do arrows (100502512.5) 3 arrows, 3 half

lives passed

Plug in to Table T #HL = t/T, 3 = 90/T, T = 30yr

Ex. What fraction of a sample of I-131 will be left after 32 days?

FR = (1/2)(t/T)

t = 32 days, T = 8.07 days, so t/T = 4 days

FR = (1/2)4 = ½ x ½ x ½ x ½ = 1/16

9. Transmutation – when an isotope turns into an isotope of a different element by nuclear decay

a. Natural – happens on its own – everything on Table N – only 1 reactant, 2 products

b. Artificial – needs help – nucleus must be bombarded with a particle – 2 reactants, 2 products

10. Nuclear Reactions – some mass is converted to energy, give off much more energy than chemical reactions

a. Fission – larger nucleus splits into 2 smaller nuclei by bombarding with a neutron. Additional neutrons

released as result, leading to chain reaction – atomic bomb (uncontrolled), nuclear reactor (controlled)

b. Fusion – 2 light nuclei (H or He) unite, requires tremendous activation energy because both nuclei

positively charged & want to repel.

11. Benefits & Risks of Radioactivity

Benefits

Provide a lot of energy for electric power

Can be used to trace reactions

Can be used for medical diagnoses (Tc-99, I-131)

Can be used for medical treatment (Co-60)

Can be used to make food stay fresh longer

Radioactive dating (U-238 for ancient geologic

nonliving things, C-14 for organic material)

Risks

Biological damage (mutations, etc)

Long term storage of waste products (long half life

before safe radiation levels are reached)

Accidents – fuel & waste can escape from nuclear

reactor

Pollution – too much radioactive waste is harmful