Chemistry Final Study Guide

Math Review:

· Scientific Method

1. Observe

2. Hypothesize

3. Test by experiment

4. If the date does not support the hypothesis, change the hypothesis then re-do the experiment

Accuracy vs. Precision

Accuracy – how close the measurement is to the accepted value
Precision – how close all the measurements are to each other

Percent Error
Scientific Notation

Chapter One:

Gases in a breath

Nitrogen- 78%

Most abundant gas in the air
Less reactive than O

Oxygen- 21%
Argon, CO₂, and others- make up 1% of all air we breathe

Argon- very unreactive (considered chemically inert)

Pollutants in a breath

Carbon Monoxide (CO)

Odorless
Disrupts the delivery of oxygen throughout the body
Generated by charcoal grills, kerosene heaters, propane stoves

Ozone (O₃)

Sharp odor (around copy machines, electric motors)
Very toxic to lungs, even in low concentrations
Pollutant at earth’s surface

Sulfur Oxides and Nitrogen Oxides

Particularly affect elderly and young children, people with breathing diseases like emphysema and asthma
Contribute to acid precipitation

PM (particulate matter)

PM₁₀à particles = 1x10^-1 mm in diameter
PM_2.5 à particles = 2.5 mm in diameter
A mix of many chemicals
Visible as smoke and soot
Affects cardiovascular system and lungs

Conversions

PPH (%) – parts per hundred
PPM – one part out of a million and is 10,000 times less concentrated than 1 PPH
PPB – one part out of one billion, or 1000 times less concentrated than 1 PPM

AQI (air quality index, calculated for: tropospheric ozone, particulate matter, CO, SO₂, NO2)

Required by law to report this in newspaper
Tells the amount of pollutants in air for that day
Limit is decided by how toxic it is

Toxicity = the health hazard of a substance

Levels of the atmosphere

Mesosphere = highest
Stratosphere = middles

Includes the ozone layer

Troposphere = where we are

Elements

Substances that cannot be broken down into simpler ones by any chemical means
90 of the 110 occur naturally on earth
Named by chemical symbols on the periodic table

Naming compounds (non-ionic)

Covalent compounds: shared electrons between two non-metals

Prefix- element name (prefix root+ide, ex. Oxygen = oxide)

1. Mono-

2. Di-

3. Tri-

4. Tetra-

5. Penta-

6. Hexa-

7. Hepta-

8. Octa-

9. Nona-

10. Deca-

o Ex. H₂O = Dihydrogen monoxide

Other rules:

1. Write down more metallic element first (the element that is closer to the left side of the periodic table)

2. Suffix –ide is added to second element

a. Ex. Oxygen = oxide, sulfur = sulfide, fluorine = fluoride

3. Mono- is omitted for the first element (the element that is closer to the left side of the periodic table

a. Ex. NH₃ = nitrogen tri-hydride, N₂O = di-hydrogen monoxide

Ionic bonds: transferred electrons between a metal and a non-metal

First element: element name
Second element: root+ide
Ex. CaCl₂ = calcium chloride

Hydrocarbons:

Mother – Methane (1 carbon atom)
Eats – Ethane (2 carbon atoms)
Peanut – Propane (3 carbon atoms)
Butter – Butane (4 carbon atoms)

Reactants and products in chemical reactions
Balancing equations

Laws of conservation of mass

Mass reactants = mass products

Same number of atoms always conserved
Identity of atoms remains the same
What may change:

Physical states of elements
Number of molecules

Other rules:

1. If an element is present in just one compound on each side, balance it first

2. Balance anything that exists as a free element last

3. Check when done: same number of atoms and same total charge (if any) on both sides

Periodic Table

Elements go in order of atomic number
Columns

Groups/families

Group 1 = very reactive

All elements in each column has the same number of valence electrons

All have similar chemical properties

Properties are a result of how electrons are distributed within the atom

Rows

Periods
Periodicity determined by rows

Metals

Shiny
Conduct electricity and heat well
Ex. Iron, gold, copper

Nonmetals

Varied appearance
Don’t conduct well
Ex. Chlorine, sulfur, oxygen

Metalloids

Elements that fall in between metals and nonmetals because they don’t fit entirely into either group
Ex. Silicon, germanium (semiconductors)

Noble gases

Elements that are inert (do not readily undergo chemical reactions)
Group 8 of the periodic table (8 valence e^- à why they are unreactive)

Classifying Substances

Atom

The fundamental units of matter (the smallest unit of an element that can exist independently)

Molecule

More than one type of atom
Diatomic molecules always exist in pairs of atoms

Ex. H₂O, O₂, N₂ (all halogens)

Element

Only one type of atom (O, N, Cl)

Compound

Can be broken down into something simpler by chemical or physical means
Two or more types of elements

In fixed combinations (Ex. Carbon dioxide, cocaine)

Mixture- a physical combination of two or more substances present in variable amounts

More than one element and/or compound (variable amounts- not fixed)
Can be separated by a physical process
Ex. Air, soda

Sources of pollutants

Indoor:

CO (furnace)
Smoke from cigarettes, incense, candles, etc.
Asbestos (in ceilings)
Radon from basements
Acetone, paint thinner
Aerosols
Ventilation makes a huge difference

Outdoor:

Primary pollutants:

CO- comes from cars/auto emissions
PM- smog
SO₂- coal plants
NO/NO₂- auto emissions (ozone formation)

Secondary pollutants (formation of tropospheric ozone):

NO₂ (sunlight) à NO + O
O + O₂ à O₃
When sunlight decreases, ozone production decreases

Chapter 2

Ozone

Allotrope of O

Energy + 3O₂ à 2O₃ (energy must be absorbed for this reaction to occur)
Resonance Lewis Structure:

Very distinct smell

Where is it?

Bad in troposphere, good in stratosphere (acts as a screen)

Allotropes- two or more forms of the same element that differ in their chemical structure and therefore in their properties

Have different physical properties
Ex. O = O₂, O₃, C = graphite, diamonds, bucky balls

DU (Dobson units)

Corresponds to one ozone molecule for every billion molecules and atoms present in the air
Used to measure the concentration of ozone in the atmosphere

Ozone layer

When the ozone later gets below 220 DU, it is considered to be a hole

Lowest recorded amount = 88 DU

Atoms

Structure

Nucleus- made up of protons (positive charge) and neutrons (neutral charge), both have almost the same mass
Electrons (negative charge)- surround the nucleus like a cloud, mass is so small it is insignificant to atomic mass

Valence electrons = outermost layer of e^-, can be determined by group number in periodic table

Periodicity of elements

Elements are arranged by atomic number- elements with similar properties fall into the same groups (columns), result of how e^- are distributed within the atom
Group number tells the number of valence e^- for groups 1-8 (skips middle of periodic table)

Total e^- can change but only valence e^- determine chemical properties because they are available for bonding
Ex. Noble gases (group 8) are very stable (full valence shell)

Mass number

Sum of the number of protons and neutrons in the nucleus

Atomic number

Number of protons in an atom

Isotopes- two or more atoms of the same element that differ in the number of neutrons and mass

Covalent bonds (see above)

Represented by Lewis structures

Count the number of valence electrons contributed by each atom then add them together
Arrange the rest of electrons in pairs so that each atom has a filled valence shell

Octect rule

Each atom needs 8 electrons in its outer shell (valence electrons) to be stable (except for H)

Resonance

Structures that represent hypothetical extremes of electron arrangements in a molecule (drawn both ways with an arrow in between)
Draw resonance if the bonds are different on each side but the elements are the same
Ex. O₃, SO₂

Different types of bond sizes/strengths

Single bond – one pair of shared electrons

Represented by one line or two dots

Double bond – two pairs of shared electrons

Shorter, stronger and harder to break than single bonds
Represented by two lines or four dots

Triple bond – three pairs of shared electrons

Shortest and strongest
Harder to break than double bonds of the same atom
Represented by three lines or six dots

Light waves

Wavelength

The distance between successive peaks
Represented by l

Frequency

The number of waves passing a fixed point in one second
Represented by n
Frequency (n) = speed of light (c) / wavelength (l)

c = 3.0 x 10⁸ ^m/_s

EM spectrum

The ranges from short, high-energy X-rays and gamma rays to long, low-energy radio waves
From right: IR (10^-4) à Visible (10^-6) à UV (10^-8)

From right: Visible spectrum = ROYGBIV red à violet (low frequency to high frequency, long wavelength to short wavelength)

Photons- individual bundles of energy that excite bonding e^- and cause bonds to break apart

Energy (E) = ^hc/_l
Quantization- photons are a fixed amount of light (the energy distribution is fixed at certain levels, like stairs)

Oxygen/Ozone screen

O₂takes a shorter wavelength to break than O₃

Need to know which one is bigger (more energy is needed to break an O₂than an O₃bond)

O₂ l £ 242 nm
O₃ l £ 320 nm

Ozone is good in stratosphere, bad in troposphere
Chapman Cycle

O₂molecules are broken apart by UV photons (must be a certain wavelength to break bonds) à 2O + a new O that’s fed into the cycle à O + O₂ à O₃ (subcycle) à O + O₂ (collisions) à 2O₂
Unbalanced: O₂ + photon à O, O + O₂ (collisions) à O₃, O₃ + O (collisions) à O₂

· Stratospheric ozone depletion

o Ozone hole (see above)

o Free radicals

§ Lone elements that break apart the ozone bonds

o Most common natural cause of ozone destruction = water vapor

§ Evaporates from oceans

o Most common man-made cause of ozone destruction = chlorine (from CFCs)

o How does ozone destruction work?

§ A free radical ends up reacting with an oxygen atom in an ozone molecule then you are left with O₂ (doesn’t necessarily return to O₃)

o Montreal Protocol

§ Made in response to ozone layer destruction

§ Agreed to ban/limit CFCs

o CFCs (trichlorofluoromethane) and their replacements (what makes them better?)

§ Refrigerants/coolants

§ Extremely stable – lasts forever in troposphere, goes up to stratosphere

§ Freon 11 (CCl₃F)

§ Freon 12(CCl₂F₂)

§ HCFCs – don’t destroy ozone/make it to stratosphere (decompose more readily in the atmosphere, do not accumulate to same extent in stratosphere)

Chapter 3

· Energy balance

o The natural energy balance process between the earth, sun and atmosphere

o Global warming = an increase in average global temperatures (natural, fluctuates)

o Greenhouse effect = what causes global warming, atmospheric gases trapping and returning a large portion of the heat radiated by the earth (normal warming of the earth)

§ When greenhouse gases trap and return more than 80% of the heat energy radiated from the earth

o Energy balance of radiation to earth/away from earth

o Enhanced greenhouse effect = process by which atmospheric gases trap more than usual heat radiated by earth

· Molecular geometry

o Draw Lewis structure à look at bonding on central atom à determine geometry

§ If a molecule has > 2 atoms, the most stable arrangement is where the electron pairs are as far apart as possible

o To determine geometry:

§ Draw Lewis structure

§ Determine how electron pairs can be as far apart as possible

§ Valence Shell Electron Pair Repulsion (VSEPR) theory

§ Don’t need to know angles of bonds

o Geometry:

§ Linear – one pair of bonding electrons = 180

§ Trigonal planar – 3 electron pairs (all bonding) = 120

§ Bent – three electron pairs (two bonding, one non-bonding) ~ 117 OR four electron pairs (two bonding, two non-bonding) ~ 105

§ Tetrahedral – 4 electron pairs (all bonding) ~ 109.5

§ Trigonal pyramidal – four electron pairs (three bonding, one non-bonding) ~ 107

o IR released by the earth gets absorbed in the atmosphere by vibrating greenhouse gases

§ Through the vibrations, many molecules knock into each other which creates heat (radiates in all directions, whole atmosphere is vibrating)

o Vibrating molecules

§ When the IR range radiation can’t break bonds between molecules, it vibrates them

· Different energies (different types of radiation) = different modes of vibration

o Why are some gases greenhouse gases?

§ Some gases are greenhouse gases because they can’t be broken by IR so they vibrate which causes other molecules to vibrate and create heat

o Percent transmittance = the amount of radiation that makes it through the atmosphere

o Different frequencies of light affect the molecules differently (Ex. microwave doesn’t have enough energy to vibrate molecules, just spin them)

o Vibrating states are quantized (see above)

· Carbon cycle

o The carbon cycle is the natural cycle of carbon in the form of CO₂ on earth

§ Most of the CO₂ is in the deep ocean and fossil fuels

o Humans disrupted the cycle by burning coal and fossil fuels (thanks to the Industrial Revolution)

· Average atomic mass

· The average of the masses of all the naturally occurring isotopes of an element relative to C^-12 (12 AMU)

· Ex. (abundance in decimal of isotope A) x (isotope A mass) + abundance in decimal of isotope B) x (isotope B mass) +…

· Avogadro’s number

· 6.022 x 10²² things (fixed quantity – like a dozen) particles / mol

· Unit = mole

· Molar mass – the mass in grams of one mol of something

· Ex. One mol ¹²C has a mass of 12 g \ molar mass ¹²C is 12g/mol

· Mass percent (mass composition) – the percent amount of an element in a compound

· Ex. % composition by mass = mass of element in compound (molar mass) / total mass of compound (molar mass) x 100

· Empirical Formula – the lowest whole number ratio of the atoms in a molecule

· Used to find molecular formula

· Ex. Given:

1. % of each element by mass of a compound

2. Assume a 100 g sample

3. Grams of each element

4. Use molar mass

5. Moles of each element

6. Divide by the lowest number of moles to get subscripts à 7. Find empirical formula

· Molecular Formula – actual formula of compound

· Find scaling factor (molar mass of molecule / molar mass of empirical formula)

· Multiply each subscript by scaling factor

· More greenhouse gases

o CH₄ (methane)

§ 40% due to natural sources

· Leakage from rock fissures, decaying plants, landfills, rice paddies, agricultural sources (ruminants), termites

§ Emissions are going down but atmospheric levels remain the same (stays in atmosphere for 12 years)

§ Effect is not as intense as CO₂ because there’s not as much of it

o N₂O (nitrous oxide-laughing gas)

§ Bacteria in soil, catalytic converters, ammonia fertilizers, burning biomass, industrial process

§ Lasts 114 years in atmosphere

§ Global levels rising but US levels decreasing slightly

o O₃ (Ozone)

§ Efficiency of it as a greenhouse gas depends on altitude

§ Has a warming effect in upper troposphere but higher it actually cools

· GWP (Global Warming Potential)

· The relative contribution of a molecule of a gas to global warming

· CO₂ = 1, all other gases measured relative to it

· Computer modeling

o Used to predict climate change

§ People use them to try to predict into the future, used to set policies/limits on gases

o Albedo = reflectivity of the surface

§ As albedo decreases, surface temperatures increase

o Forcing = anything that affects the global surface temperature (Ex. Greenhouse gases, aerosol, O₃ changes)

o Humans = most likely cause of global warming

o Warming is already affecting plants and animals (can’t adapt fast enough to changing environments à more extinction)

· What to do about global warming?

· US and Russia = largest carbon emitters

· Must reduce emissions without hurting economy

· Ideas:

§ Reduce dependence on fossil fuel

§ Clean coal technologies

· C capture/sequestration à store under ocean and in rocks

· Kyoto Protocol

o In 1977, scientists gathered in Kyoto to find a way to reduce greenhouse gas concentrations

o Countries split into annexes:

§ Annex 1 – industrialized countries

§ Annex 2 – developed countries à help pay for developing countries to grow (must reduce CO₂ by 1990s)

§ Developing nations (no restrictions on emissions)

o US has still not ratified the Kyoto Protocol

§ Disagree with developing countries not having restrictions

§ Would hurt economy too much

§ Has non-binding proposals to reduce emissions

· Global warming vs. ozone depletion

o Are they the same or is one the result of the other? – Not related

o Ozone depletion not contributing to global warming

§ Has other damaging effects but doesn’t warm earth (not right frequency of radiation)

§ Even if ozone layer was depleted more, the light doesn’t have the ability to warm the atmosphere

Chapter 4

Energy- the ability to do work

Work (W) = force (f) x distance (d)

Movement against a force

Heat- transfer of energy from a warmer object to a colder object

A consequence of motion at the molecular level

Temperature- the average kinetic energy of the molecules in a substance (determines the direction of heat flow)

Measure of the average speed of that motion
Potential energy- the energy that’s stored within a molecule that gets released by heat (chemical potential)
Kinetic energy- motion
Breaking/forming bonds

Calorie

· The amount of energy necessary to raise the temp of something by one °

· 1 cal = 4.184 J

· 1 kcal = 1000 cal = 1 Cal (one dietary cal) = 4.184 kJ

Entropy- things going from a state of organization to a state of disorganization (loses that you can’t control in a power plant- entropy is always increasing)
Thermodynamics

1^st law – conservation of mass (no matter is ever created or destroyed)

Energy is transformed from one type to another but never created or destroyed
There will always be loses in a power plant

2^nd law – entropy is always increasing

Power plants

Inevitably inefficient

Energy is being used/lost at each step to convert to other types of energy

How do we decide what to use as a fuel?

Determined by combustion: fuel combines rapidly with oxygen to release energy (burning)

CO₂and H₂O are always products of combustion

Combustion- rapid combination of fuel and oxygen

Exothermic reactions- release of heat in a chemical reaction
Endothermic reactions- absorption of heat in a chemical reaction
Products = heat, water and CO₂ ALWAYS

How does a calorimeter work?

Something is burned inside a contained environment, water heats and calorimeter measures change in temperature

Measures specific heat

Q (specific heat in J) = m (mass in g) x c (specific heat in ^J/_g C°) x DT (temperature difference in °C)

Heat of combustion = kJ/mol
Specific heat – the heat energy that must be absorbed to increase the temperature of 1 g of a substance 1°C (calorie)

Q= m x c x DT

Bond energies – chemical reactions involve rearranging of atoms

Energy is used to break bonds/form bonds
Bond energy – the amount of energy that must be absorbed to break a chemical bond
Must know how to determine endothermic vs. exothermic reactions

Chapter 5

Where does our drinking water come from?

Surface water – comes from lakes, rivers, reservoirs
Ground water – pumped from wells drilled into underground aquifers

Water is a great solvent

Solvent- the thing that does the dissolving

Substances capable of dissolving other substances- usually present in the greater amount

Solute- the thing that’s dissolved

Substances dissolved in a solvent-usually present in the lesser amount
IMFs must be overcome to dissolve
Solute molecules orient themselves around the solvent particles, encapsulating them (dissolve because IMFs attracting solute particles to each other are overcome by attraction to solvent)
A solute will only dissolve a solvent with similar strength IMFs (like dissolves like)

Solution - a homogeneous mixture of uniform composition
Aqueous solution (when water is the solvent)

Electrolytes – conduct electricity, dissociate in solution (have ions)
Non-electrolytes – do not conduct electricity

Concentration – the ratio of amount of solute to amount of solution

Percent by mass
Molarity – the number of moles of solute present in one liter of solution

Polar molecules- unequal sharing of electrons due to differences in electronegativity (separation of charge)

Electronegativity- the measure of an atom’s attraction that it shares in a covalent bond

Ex. Water is a polar molecule

Hydrogen bonding- electrostatic attractions between a hydrogen atom bearing a partial positive charge in one molecule and an O, N or F atom bearing a partial negative charge in a neighboring molecule

Are only about one-fifteenth as strong as the covalent bonds that connect atoms together within molecules
Characteristics of water
IMF (intermolecular force) – the force that exists within a molecule

Ionic compounds – made up of electrically charged ions that are present in fixed proportions and are arranged in a regular, geometric pattern

Some atoms form more than one ion:

Must memorize: Groups 1, 2, 5, 6 and 7

Polyatomic ions – made of more than one atom or element

Cations – more metallic element (positive charge, always comes first)
Anions – non-metal element (negative charge, second in name)
Common polyatomic ions:

Naming of ionic compounds:

Metals and non-metals (transfer of electrons)

Double displacement reactions

Be able to predict products and balance
Must be able to figure out solubility of products from charge

Usually results in a precipitate, insoluble gas or H₂O

Colligative properties- dependent on the number of dissolved particles (gizmo)

1. Vapor pressure – pressure of a vapor in equilibrium with the liquid (evaporation rate = condensation rate)

a. With the introduction of a solute, water vapor pressure lowers (salt lowers the vapor pressure more than sugar because there are more molecules with less elements that break apart more readily)

2. Boiling point elevation – when the vapor pressure is at atmospheric pressure

3. Freezing point depression – with the introduction of a solute, freezing point decreases

a. Ex. when you pour salt on an icy sidewalk, the freezing point decreases and the ice melts – can use any solute like sugar but salt is most effective \ must be at a colder temperature to freeze

4. Osmosis

a. Solvent flows to equalize concentration on either side of membrane (as solution concentration increases, osmotic pressure increases)

Chapter 6

Acids – a substance that releases H ions

Ex. Citrus juices, vinegar, vitamin e (citric acid)
Forms H⁺ ion in aqueous solution

Since the H ion only has one proton and no electron, it is sometimes referred to as a proton
H⁺ ions are much too reactive to exist alone so they attach to something else, like water molecules

Strong acids dissociate completely, weak acids do not dissociate as much

Ex. Of strong acids: (* = must memorize)

*HCl à H⁺ + Cl^- (hydrochloric acid)
HBr à H⁺ + Br^- (hydrobromic acid)
HI à H⁺ + I- (hydroiodic acid)
*H₂SO₄ à H⁺ + SO₄^- (sulphuric acid)
*HNO₃ à H⁺ + NO₃^- (nitric acid)
HClO₄à H⁺ + ClO₄^- (perchloric acid)

All other acids are weak

Bases – the opposite of acids, any compound that produces a hydroxide ion (OH^-)

Ex. Lye (oven cleaner), soap, ammonia, bleach
Feels slippery on skin, bitter taste
Strong bases dissociate completely in aqueous solution

Ex. Of strong bases: Group I and II metal hydroxides

NaOH
KOH
CaOH, etc.

All other bases are weak

When acids and bases react, it’s called a neutralization reaction

H ions combine with hydroxide ions to form H₂O molecules, the other product is salt:

How to represent quantitatively

Hydrogen ion (hydronium ion concentration)

[H⁺] à Concentration in moles/liter (molarity)
H⁺ = H₃O⁺ (hydronium ion, the hydrogen ion never exists stably by itself)

Hydroxide ion concentration

[OH^-] à concentration in moles/liter (molarity)

Concentrations of H⁺ and OH^- in a neutral solution

H⁺= 1.0 x 10^-7M
OH^- = 1.0 x 10^-7M

Product of [H⁺][OH^-] = 1 x 10^-14 M² for aqueous solution
pH (parts of hydrogen or potential hydrogen) – a number, between 0 and 14, that indicates the acidity of a solution

pH scale

Goes from 1 to 14
7 = neutral

Less than 7 = acidic, greater than 7 = basic

pH of pure water = 7
pH of strong acids and bases

Acids = pH < 7
Bases = pH >7

Why is rain naturally acidic?

CO₂ in the atmosphere dissolves to a slight extent in H₂O and reacts with it to produce a slightly acidic solution of carbonic acid
Recognize equation for rain

Cause of extra acidity:

SO_x and NO_x are the chief culprits of acid rain

SO_x emissions are highest in regions with many coal-fired EPPs, steel mills, etc. that rely on coal
Highest NO_x emissions are generally found in states with large urban areas, high population density and heavy auto traffic

Where is acid rain the worst? – Along the Ohio River Valley

Recognize equations that contribute to extra acidity:

What contributes to acidity over LA?

NO_x from auto emissions causes acid rain over LA (car capital of the country)

Titration (lab experiment)

Know how to read a titration curve – when the graph starts to level out = knee, goes from low pH to high pH, middle = neutral
Be familiar with experiment

Chapter 8

· Redox Reactions

o What is oxidized and what is reduced?

§ Oxidation – shows the reactant that loses electrons

§ Reduction – shows the reactant that gains electrons

o Half reactions – a type of chemical equation that shows the electrons either lost or gained

o Assign oxidation numbers

§ Rules:

1. An action that is free and in its elemental state has an oxidation number of 0

2. For simple ions, the oxidation number is the charge on the ion

3. For molecular compounds and polyatomic ions, the electrons are assigned to the more electronegative ion

4. O always has an oxidation number of -2

o Recognize whether a simple reaction is redox or not

o Calculate cell potentials from reduction potentials

· Galvanic Cells – convert the energy released in a spontaneous chemical reaction into electrical energy

o Ex. Wired together, they make car batteries, etc.

· Single displacement reactions

o Recognize by equations

o Given an activity series, be able to predict if reaction will happen