Chem 112 Final Exam Notes

Chapter 1 – Matter, Measurement, and Problem Solving1.1 Atoms and Molecules

Chemistry – the science that seeks to understand the behaviour of matter by studying the behaviour of atoms & molecules Properties of matter are determined by the properties of molecules & atoms

o Atoms & molecules determine how matter behaves If they are different, matter would be different Ex. properties of water molecules determine how water behaves

Atoms – submicroscopic particles that constitute the fundamental building blocks of ordinary matter Molecules – two or more atoms joined in a specific geometrical arrangement Properties of the substances around us depend on the atoms and molecules that compose them

o Details of how specific atoms bond to form a molecule & type of atoms in the moleculeo Ex. Properties of CO (g) depend on the properties of CO moleculeso CO (g) have the right properties that fit into cavities within hemoglobin molecules that are normally reserved for O

molecules Too much CO (g) can lead to unconsciousness and even death

CO molecules can take place of oxygen Reducing the amount of oxygen reaching the body’s tissues

Small amounts can result in headache, dizziness, weakness, and confused thinking Hemoglobin – a large protein molecule

Oxygen carrier in red blood cells Each subunit contains an iron atom to which oxygen binds

o Any changes in molecules are likely to result in large changes in the properties of the substances they compose CO2(g) doesn’t kill us extra oxygen

1.2 The Scientific Approach to Knowledge Scientific knowledge is based on observation & experiment

o Some observations & experiments are qualitative (noting/describing how a process happens), but many are quantitative (measuring/quantifying something about the process)

o Observations often lead scientists to formulate a hypothesis (a tentative interpretation/explanation of the observations)

A good hypothesis is falsifiable can be confirmed or refuted by further observations Test hypotheses by experiments (highly controlled procedure designed to generate observations)

Results may support a hypothesis or prove it wrongo Scientists must modify or discard the hypothesis

Series of similar observations can lead to the development of a scientific law Brief statement that summarizes past observations & predicts future ones Ex. Law of conservation of mass

o “In a chemical reaction, matter is neither created nor destroyed” Scientific laws describe how nature behaves often refer to them as principles

o Generalizations of what nature does One or more well-established hypotheses may form the basis for a scientific theory

A model for the way nature is & tries to explain why the nature behaves the way it does Often predicting behaviour far beyond the observations/laws from which they were

developed Ex. Atomic theory

o “matter is composed of small, indestructible particles called atoms”o Over time, scientific community eliminates or corrects poor theories & laws, and good theories & laws remain

1.3 The Classification of Matter Matter – anything that occupies space & has mass

o Classified according to its state (its physical form) and its composition (the basic components that make it up)o Substance – specific instance of matter (ex. air, water, or sand)

Can exist in 3 different states:o Solid – atoms/molecules pack close to each other in fixed locations

They vibrate Has a fixed volume & rigid shape May be crystalline – atoms/molecules are in patterns with long-range, repeating order (ex. diamond

& table salt) Or they may be amorphous – atoms/molecules do not have long-range order (ex. glass & plastic)

o Liquid – atoms/molecules pack about as closely as they do in solid matter

Free to move relative to each other Fixed volume but not a fixed shape Takes shape of their container (ex. water, alcohol, & gasoline at room temperature)

o Gas – atoms/molecules have a lot of space between Compressible Free to move relative to one another Takes the shape & volume of their container (ex. helium, nitrogen, and CO2 at room temperature)

Classifying according to its compositiono First classification

Pure substance – made up of only 1 component Components can be individual atoms, or groups of atoms joined together

o Must always be exactly the sameo Does not vary from one sample to another

Elements – substances that cannot be chemically broken down into simple substanceso Ex. helium

Compounds – substances composed of 2 or more elements in fixed & definite proportiono More common than pure elements on Eartho Water, and table salt (NaCl)

Mixture – composed of 2 or more components in proportions that can vary from one sample to another

Heterogeneous – composition varies from one region to anothero Atoms/molecules that compose them separateo Ex. wet sand

Homogeneous – same composition throughouto Atoms/molecules that compose them mix uniformlyo Ex. sugar water

In general, mixtures are separable b/c different components have different physical/chemical properties o Decanting – carefully pouring off

Ex. pouring the water into another container for a mixture of sand and watero Distillation – process in which the mixture is heated to boil off the more volatile (easily vaporizable) liquid

Volatile liquid is then recondensed in a condenser & collected in a separate flask Component with the lowest boiling point vaporizes first

o Filtration – mixture is poured through filter paper in a funnel designed to pass only the liquid Mixture is composed of an insoluble solid & a liquid

1.4 Physical and Chemical Changes & Physical and Chemical Properties Physical changes – alter only state or appearance, but not composition

o Atoms/molecules that compose a substance do not change their identity o Ex. water boiling changes its state from a liquid to a gaso Physical property – one that a substance displays wo/ changing its composition

Odour, taste, colour, appearance, MP, BP, and density Chemical changes – alter the composition of matter

o Atoms rearrange, transforming the original substances into different substanceso Ex. rusting of iron

Iron atoms combine with oxygen molecules from air to form iron oxide (rust)o Chemical property – one that a substance displaces only by changing its composition via a chemical change

Corrosiveness, flammability, acidity, toxicity1.5 Energy: A Fundamental Part of Physical & Chemical Change

Physical & chemical changes are usually accompanied by energy changeso Ex. when water evaporates from your skin, water molecules absorb energy from your body, making you feel

cooler Physical change

o Ex. when you burn natural gas on the stove, energy is released, heating the food you are cooking Chemical change

Energy – capacity to do worko Total energy of an object is a sum of its kinetic energy & potential energyo Kinetic energy – energy associated with its motion

Thermal energy – energy associated with the temperature of an objecto Potential energy – energy associated with its position or compositiono 1st principle

Law of conservation of energy Energy is neither created nor destroyed It can only change from one form to another Total initial energy = total final energy

o 2nd principle The tendency of systems with high potential energy to change in a way that lowers their potential

energy Objects/systems with high PE tend to be unstable

o Ex. weight will naturally fall (lowering its PE), unless restrained Can harness the high PE to do work lowers the PE become stable

Work – the action of a force through a distance1.6 The Units of Measurement

Units – standard quantities used to specify measurementso Critical in chemistryo 2 common unit systems:

English system – used in United States Metric system – used in most of the rest of the world

o Scientists use the International System of Units (S) – based on the metric system Mass of an object is a measure of the quantity of matter within it Weight of an object is a measure of the gravitational pull on its matter

Temperature is a measure of the amount of average kinetic energy of the atoms/molecules Also determines the direction of thermal energy transfer heat Transfer from hot objects to cold ones °F = English system

o Allows negative temperatures °C = metric system

o Allows negative temperatures K = SI unit

o Avoids negative temperatureso 0K = absolute zero

Which molecular motion virtually stops

Prefix multipliers – multipliers that change the value of the unit by powers of 10o When reporting a measurement, choose a prefix multiplier close to the size of the quantity being measured

Derived unit – combination of other unitso Ex. unit for speed = m/s

Volume – measure of spaceo Any unit of length when cubed

Density – characteristic physical property of materials o Differs from 1 substance to anothero Depends on temperature

Density of most substances as temperature Increasing the temperature will expand the object Mass remains constant so density decreases

o Intensive property – independent of the amount of the substance Often used to identify substances

o Mass is extensive property – depends on the amount of the substance1.7 The Reliability Measurement

More digits = more certainty Fewer digits = less certainty Scientific measurements are reported so that every digit is certain except the last, which is estimated

o Assumed to be 1 in the last digit unless otherwise indicatedo Ex. reporting as 15.0 ppm, the scientists mean 15.00.1 ppm between 14.9 & 15.1 ppmo Ex. reporting as 15 ppm, this would be 151 between 14 & 16 ppm

# of digits reported in a measurement depends on the measuring device

Significant figures – non-place-holding digitso Those that are not simply marking the decimal place

Ex. 0.0008 = 1 SD Ex. 0.000800 = 3 SDs

o The greater the # of SF, the greater the certainty of the measurement Exact numbers have no uncertainty

o Do not limit the # of SD in any calculation Regard as having an unlimited # of SDs

o Ex. 3 atoms = 3.00000… atoms1. In multiplication/division, the result carries the same # of significant figures as the factor with the fewest significant figures

2. In addition/subtraction, the result carries the same # of decimal places as the quantity with the fewest decimal places

3. When rounding to the correct # of SDs, round down if the last digit dropped is ≤4; round up if the last digit dropped is ≥5

4. To avoid rounding errors in multistep calculations round only to the final answer – do not round intermediate steps. If you write down intermediate answers, keep track of SDs by underlining the least significant digit

Scientists often repeat measurements several times to confidence in the result 2 different kinds of certainty:

o Accuracy – refers to how close the measured value is to the actual valueo Precision – refers to how close a series of measurements are to one another or how reproducible they are

Series of measurements can be precise, but not accurateo

Random error – inconsistency inaccurate & impreciseo Has equal probability of being too high/lowo Almost all measurement have some degree of random erroro Can average itself out with enough trials

Systematic error – inaccurate & preciseo tends toward being either too high/lowo does not average out with repeated trials

1.8 Solving Chemical Problems Dimensional analysis – using units as a guide to solving problems

Conversion factor – a fractional quantity with the units we are converting from on the bottom & the units we are converting to on top

o Constructed from any 2 equivalent quantities

Chapter 2 – Atoms and Elements2.1 Imaging and Moving Individual Atoms

Scanning tunnelling microscopy (STM) – a technique that can image & move individual atoms & moleculeso Gerd Binnig & Heinrich Rohrer, March 16, 1981o Works by moving an extremely sharp electrode (an electrical conductor) over a surface & measuring the resulting

tunnelling currento Electrical current flows between the tip of the electrode & the surface even though the 2 are not in physical

contacto Tunnelling current is extremely sensitive to distance

Possible to maintain a precise separation of approximately 2 atomic diameters between tip & surface Current is constant as long as the distance is constant If current starts to a bit, the tip is moved up, away from the surface to the current If current starts to a bit, the tip is moved down towards the surface to the current as the tip goes over an atom, the tip must move up to maintain constant current

Atom – smallest identifiable unit of an elemento 91 different naturally occurring elementso Over 20 synthetic elements (not found in nature) made my scientists

2.2 Early Ideas about the Building Blocks of Matter Leucippus & Democritus proposed that:

o “Matter were composed of small, indestructible particles(atoms)”o “Many different kinds of atoms existed, each different in shape & size, and that they moved randomly through

empty space” Plato & Aristotle

o “Matter had no smallest parts & that different substances were composed of various proportions of fire, air, earth, & water”

o No experimental way to test the idea Nicolaus Copernicus

o Proposed that the sun, not Earth, was at the center of the universe Marks the beginning of the scientific revolution

2.3 Modern Atomic Theory and the Laws That Led to It 3 most important laws that led to the development & acceptance of the atomic theory:

o Law of Conservation of Mass Formulated by Antoine Lavoisier in 1789 “in a chemical reaction, matter is neither created nor destroyed”

Total mass of the substances involved in the reaction does not change Mass of reactants = mass of product(s)

* particles rearrange during chemical rxn, but the amount of matter is conserved b/c the particles themselves are indestructible

o Law of Definite Proportions Proposed by Joseph Proust (1754 – 1826) “all samples of a given compound, regardless of their source or how they were prepared, have the same

proportions of their constituent elements” Applies to every compound Applies to 2 or more samples of the sample compound

* compounds have definite proportions of their constituent elements b/c the atoms that compose them, each with its own specific mass, occur in a definite ratio

o Law of Multiple Proportions Published by John Dalton in 1804 “when 2 elements (call them A & B) form 2 different compounds, the masses of element B that combine

with 1g of element A can be expressed as a ratio of small whole numbers” Applies to 2 different compounds containing the same 2 elements (A & B)

Atomic Theoryo Explained the laws by John Dalton in 1808, which included the following concepts:

1. Each element is composed of tiny, indestructible particles called atoms2. All atoms of a given element have the same mass & other properties that distinguish them from the

atoms of other elements3. Atoms combine in simple, whole-number ratios to form compounds

4. Atoms of 1 element cannot change into atoms of another element. In a chemical rxn, atoms only change the way that they are bound together with other atoms

o Matter is indeed composed of atoms2.4 The Discovery of the Electron

Further experiments revealed that the atom itself was composed of even smaller, more fundamental particles Electrical charge – fundamental property of some of the particles that compose atoms

o Results in attractive & repulsive forces – electrostatic forces - between those particles1. + & - electrical charges attract one another2. + charges repel one another3. - charges repel one another4. + & - charges of exactly the same magnitude sum to zero when combined

o Electric field – area around a charged particle where these forces exist Millikan’s oil drop experiment Charge = -1.60x10-19 C Mass = 9.10x10-28 g Charge of the electron is significant b/c it determines how strongly an atom holds its electrons

o If e- had a much smaller charge atoms will be held more looselyo If e- had a greater charge atoms will be held more tightly

Could result in fewer compounds or maybe even none2.5 The Structure of the Atom

Since atoms are neutrally charged, they must contain + charge that neutralizes the e- J.J. Thomson e- were small particles held within a + charged sphere

o The raisin bread model Raisins = e- Bread = positively charged sphere

Radioactivity – emission of small energetic particles from the core of certain unstable atoms1. Alpha () particles2. Beta () particles3. Gamma () rays

Rutherford proposed the nuclear theory of the atom, with 3 basic parts1. Most of the atom’s mass & all of its + charge are contained in a small, dense core called the nucleus2. Most of the volume of the atom is empty space, throughout which tiny, - charged e- are dispersed3. There are as many – charged e- outside the nucleus as there are + charged particles (protons) within the nucleus,

so that the atom is electrically neutral H atoms contain 1 p+, and He atoms contain 2, yet H atom has only ¼ the mass of a helium atom

o Because the atom contains the previously unaccounted mass due to neutrons – neutral particles within the nucleus

Mass similar to that of a p+, but has no electrical charge Ex. we can think of the e- that surround the nucleus in analogy to the water droplets that make up a cloud

o Although their mass is almost negligibly small, they are dispersed over a very large volumeo Consequently, an atom, like a cloud, is mostly empty space

Matter appears solid b/c the variation in its density is on such a small scale that our eyes cannot see it2.6 Subatomic Particles: Protons, Neutrons, and Electrons in Atoms

All atoms are composed of the same subatomic particles protons, neutrons, and electrons Atomic mass unit (amu) – common unit to express the masses of subatomic particles

o 1/12 the mass of a carbon atom containing 6 p+ & 6 n0

o Mass of p+ & n0 is approximately 1 amu

# of protons defines the element Isotopes – elements with the same # of protons but have different # of neutrons During chemical changes, atoms can lose/gain e- & become charged particles = ions

o Cations – positively chargedo Anions – negatively chargedo Act very differently than the atoms from which they are formed

2.7 Finding Patterns: The Periodic Law and the Periodic Table Periodic law

o When the elements are arranged in order of increasing mass, certain sets of properties recur periodically Listed in order of increasing atomic number Metals – good conductors of heat & electricity

o can be pounded into flat sheets (malleability)o can be drawn into wires (ductility)o are often shinyo tend to lose electrons when they undergo chemical changes

Non-metals – some are solids at room temperature, other are liquids or gaseso Poor conductors of heat and electricityo Tend to gain e- when undergo chemical changes

Metalloids – elements that lie along the zigzag diagonal line that divides metals & non-metalso Exhibit mixed propertieso Semiconductors

Periodic table can also be divided into main-group elements & transition elements/metalso Family/group – each column within the main-group regions

Usually have similar properties Alkali metals – group 1A reactive metals Alkaline earth metals – group 2A fairly reactive metals Halogens group 7A very reactive non-metals

2.8 Atomic Mass: The Average Mass of an Element’s Atoms Atomic mass – average mass of the isotopes that compose that element

o Weighted according to the natural abundance of each isotopeo Atomic mass = (fraction of isotope n)(mass of isotope n)

Ex. Chlorine atomic mass = (0.7577)(34.97 amu) + (0.2423)(36.97) = 35.45 amu Mass spectrometry – measuring the masses of atoms & molecules & percent abundances of isotopes of elements

1. Sample is injected into the instrument & vaporized2. Vaporized atoms are then ionized by an e- beam3. E- in the beam collide w/ the atoms, removing e- & creating + charged ions4. Charged plates w/ slits in them accelerate the + charged ions into a magnetic field, which deflects them

Amount of deflection depends on the mass of the ions Lighter ions are deflected more than heavier ones

5. Ions strike a detector & produce an electrical signal that is recorded6. Result is the separation of the atoms in the sample according to their mass

Position of each peak on the x-axis indicates the mass of the isotope that was ionized

Intensity (indicated by the height of the peak) indicates the relative abundance of that isotope

Peaks are specific to the molecule Can be used to identify an unknown molecule & to determine how much of it is present in a particular sample

2.9 Molar Mass: Counting Atoms by Weighing Them Atoms are counted by weighing them

o Far too small to count by any ordinary means Mole – the amount of material containing 6.02214x1023 particles

o 1 mol of particles = 6.02214x1023 particles

o Avogadro’s numbero ** the value of the mole is equal to the # of atoms in exactly 12g of pure C-12

12g C = 1 mol C atom = 6.02214x1023 C atomso Gives a relationship between mass & # of atoms

Molar mass – mass of 1 mol of atoms of an elemento Equal to the element’s atomic mass in amuo g/mol

Chapter 3 Molecules, Compounds, and Chemical Equations

Elements combine w/ each other to form compoundso Each with its own specific properties

Life could not exist with just 91 different elements3.1 Hydrogen, Oxygen, and Water

H2 is an explosive gas used as a fuel in the space shuttle O2 is a inflammable gas as a natural component of the air on Earth However, when H2 and O2 combine to form H2O, a dramatically different substance results

An entirely new substance results when 2 elements combine to form a compoundo Properties of compounds are generally very different from the properties of the elements that compose them

Although some of the substances that we encounter in everyday life are elements, most are compounds Free atoms are rare on Earth In a compound, elements combine in fixed, definite proportions In a mixture, elements can mix in any proportions

3.2 Chemical Bonds Compounds are composed of atoms held together by chemical bonds

o Result of interactions between the charged particles (e- & p+) that compose atomso Electrostatic forces are responsible for chemical bonding

Like charges repel & opposite charges attract 2 types of chemical bonds:

o Ionic bonds – occur between metal & non-metals Involve the transfer of e- from 1 atom to another

Metals have the tendency to lose e-o Becomes a cation (+ charged ion)

Non-metals have a tendency to gain e-o Becomes an anion (- charged ion)

Oppositely charged ions are then attracted to 1 another by electrostatic forces Solid phase is composed of a lattice (regular 3D array) of alternating cations & anions

o Covalent bonds – occur between 2 or more non-metals Involve the sharing of e- between 2 atoms Shared e- interact with the nuclei of both atoms

Lowering the potential energy of the system through electrostatic interactions Result is a molecular compound

Composed of individual covalently-bonded molecules3.3 Representing Compounds: Chemical Formulas and Molecular Models

Chemical formula – indicates the elements present in the compound & the relative # of atoms/ions of eacho Quickest & easiest way to represent a compound

o Normally list the more metallic (or more positively charged) elements first, followed by the less metallic elementso 3 types:

Empirical formula – gives the relative # of atoms of each element in a compound Lowest whole-number ratio Ex. hydrogen peroxide = HO

Molecular formula – gives the actual # of atoms of each element in a molecule of a compound Ex. hydrogen peroxide = H2O2

Whole-number multiple of the empirical formula Indicates the number & type of each atom in he molecule

Structural formula – uses lines to represent covalent bonds & shows how atoms in a molecule are connected/bonded to each other

Ex. hydrogen peroxide = H–O–O–H Gives a sense of the molecule’s shape Shows the different types of bonds that occur between molecules how atoms are connected

o Single bond 1 shared e- pair

o Double bond 2 shared e- pair Generally stronger & shorter than a single bond

o Triple bond 3 shared e- pair Stronger & shorter than double bond

o Structural formula communicates the most informationo Empirical formula communicates the least

Molecular models o Ball-and-stick models – atoms = balls & chemical bonds = sticks

Balls are typically colour-coded to specific elements Portrays the geometry of the molecule

Ex. the C atom sits in the center of a tetrahedron formed by the 4 H atomso Space-filling molecular models – atoms fill the space between each other to more closely represent best estimates

for how a molecule might appear if scaled to a visible size Gives the best sense of the relative sizes of the atoms & how they merge together in bonding

The atoms that composed a molecule, the lengths of the bonds between atoms, the angles of the bonds between atoms, & its overall shape determine the properties of the substance that the molecule composes

o Change any of these details changes the propertieso

3.4 An Atomic-Level View of Elements and Compounds

Pure substances are either elements or compoundso Subcategorize elements & compounds according to the basic units that compose themo Elements may be:

Atomic elements – exist in nature with single atoms as their basic units Ex. helium is composed of helium atoms & iron is composed of iron atoms

Molecular elements – exist as molecules – 2 or more atoms of the element bonded together Most molecular elements exist as diatomic molecules

o Ex. hydrogen is composed of H2 molecules Few exist as polyatomic molecules

o Ex. phosphorus exists as P4 & sulphur exists as S8

o Compounds may be: Molecular compounds – usually composed of 2 or more covalently bonded non-metals

Basic units are molecules composed of the constituent atoms Ionic compounds – composed of cations & anions bound together by ionic bonds

Basic unit is the formula unit o Smallest, electrically neutral collection of ionso Different from molecules in that they do not exist as discrete entities, but only as part of

a larger latticeo Ex. formula unit of sodium chloride is NaCl

Many common ionic compounds contain ions that are themselves composed of a group of covalently bonded atoms w/ an overall charge

o Polyatomic ion – an ion composed of 2 or more atoms Ex. ClO-

3.5 Ionic Compounds: Formulas and Names Ionic compounds are generally very stable b/c the attractions between cations & anions within ionic compounds are strong,

& b/c each ion interacts with several oppositely charged ions in the crystalline lattice Criss-crossing method for formulaSummarizing Ionic Compound Formulas:1. Ionic compounds always contain + & - ions2. In a chemical formula, the sum of the charges of the cations must equal the sum of the charges of the anions3. A formula reflects the smallest whole-number ratio of ions Some ionic compounds have common names (Ex. sodium chloride = table salt) However, chemists have developed systematic names for different types

of compoundso Determined from its chemical formula

Ionic compounds can be categorized into 2 types, depending on the metal in the compound

o 1st type contains a metal whose charge is invariant from 1 compound to another

No need to be specified in the name of the compound Charges of these metals can be inferred from their

group # in the periodic table

o 2nd type contains a metal w/ a charge that can be differ in different compounds Are often transition metals Can determine the charge of the metal cation by inference from the sum of the charges of the non-metal

anions Sum of all the charges must be zero

ionic compounds that contain a polyatomic ion are named in the same way as other ionic compounds, except that the name of the polyatomic ion is used whenever it occurs

o most polyatomic ions are oxyanions anions containing oxygen & another element one with more oxygen atom has the ending –ate one with few has the ending –ite

hydrates – ionic compounds containing a specific # of water molecules associated w/ each formula unito Ex. MgSO4 7H2O = magnesium sulphate heptahydrateo Waters of hydration can usually be removed by heating the compound

3.6 Molecular Compounds: Formulas and Names Formula for a molecular compound cannot readily be determined from its constituent elements b/c the same combination

of elements may form many different molecular compoundso Ex. nitrogen and oxygen form all of the following unique molecular compounds: NO, NO2, N2O, N2O3, N2O4, and

N2O5

Like ionic compounds, many molecular compounds have common names (H2O and NH3 = water & ammonia)o Requires a systematic approach to naming them

First element is the more metal-like one (toward the left & bottom of the periodic table)

Acids – molecular compounds that release hydrogen ions (H+) when dissolved in watero Composed of hydrogen (usually written first in their formula) & one or more non-metals (written second)o Characterized by their sour taste & ability to dissolve many metalso Binary acids – composed of hydrogen & a non-metal

Ex. HCl (aq) = hydrochloric acido Oxyacids – contain hydrogen & an oxyanion/polyatomic ions (anion containing a non-metal & oxygen)

# of H+ ions depends on the charge of the oxyanion Formula is always charge-neutral

Name of oxyacids depend on the ending of he oxyanion

3.7 Formula Mass and the Mole Concept for Compounds Formula mass – average mass of a molecule (or a formula unit) of a compound

o Molecular mass/weighto Sum of the atomic masses of all the atoms in its chemical formula

o Equivalent to molar mass of a compound3.8 Composition of Compounds

A chemical formula, in combination w/ the molar masses of its constituent elements, indicates the relative quantities of each element in a compound

Mass percent composition – method to express how much an element is in a given compound in percentage

o Mass percent of element X = massof element X∈1mol of compound

massof 1molof the compound X 100%

o Conversion factor between mass of the element and mass of the compound Ex. 58.64% of Cl in CCl2F2 = 58.64g Cl : 100g CCl2F2

o Conversion factor between moles of atoms and moles of molecules Ex. 1 mol CCl2F2 : 2 mol Cl

3.9 Determining a Chemical Formula from Experimental Data Calculate a chemical formula from mass percent composition

o Obtains an empirical formula only

o To get a molecular formula, molar mass of the compound is required Molecular formula = empirical formula x n, where n = 1, 2, 3, …

n = molar mass

empirical formulamolarmass combustion analysis – another common way of obtaining empirical formulas for unknown compounds

o especially those containing C and H1. unknown compound undergoes combustion in the presence of pure oxygen2. when the sample is burned, all of the C in the sample is converted to CO2, and all of the H is converted to H2O3. CO2 and H2O produced are weighed4. Determine the amounts of C and H in the original sample using numerical relationships between moles 5. Any other elemental constituents (O, Cl, or N) can be determined by subtracting the original mass of the sample

from the sum of the masses of C & H3.10 Writing and Balancing Chemical Equations

Chemical reaction – 1 or more substances are converted into 1 or more different oneso Compounds form & changeo Combustion reaction – particular type of chemical reaction

A substance combines w/ oxygen to form 1 or more oxygen-containing compounds Emit heat

o Represented by a chemical equation Reactants – substances on the left side of the equation Products – substances on the right side Add coefficients to balance the chemical equation

Mass conservation law # of atoms on the right = # of atoms on the left

3.11 Organic Compounds 2 types of compounds:

o Organic – originate from living things Composed of C and H and a few other elements (N, O, S)

C is the key element Ex. Sugar – from sugarcane/sugar beet Easily decomposed Major components of living organisms Hydrocarbons – organic compounds that contain only C & H

Main components of most fuels (gasoline, oil, & natural gas) Alkanes – hydrocarbons containing only single bonds Alkenes – containing only double bonds Alkynes containing only triple bonds

Functionalized Hydrocarbons – hydrocarbons containing a functional group (a characteristic atom or group of atoms)

Ex. alcohol – functional group = -OH Hydrocarbon portion of the molecule is represented as “R”

o R – OH Family – group of organic compounds with the same functional group

o Inorganic – originate from the Earth Ex. salt – mined from the ground/ocean Typically more difficult to decompose

o Have different properties

Chapter 4 Chemical Quantities and Aqueous Reactions The amount of product formed in a chemical reaction is related to the amount of reactant that reacts

4.1 Global warming and the Combustion of Fossil Fuels Greenhouse gas – allows sunlight to enter the atmosphere & warm Earth’s surface

o Prevent some of the heat generated by the sunlight from escaping Balancing between incoming & outgoing energy from the sun determines Earth’s average temperature If greenhouse gases were not present, more heat energy would escape, & Earth’s average temperature would be colder

than it is now If the concentration of greenhouse gases in the atmosphere were to , Earth’s average temperature would rise

4.2 Reaction Stoichiometry: How Much Carbon Dioxide? Stoichiometry – numerical relationships between chemical amounts in a balanced chemical equation

o Allows to predict the amounts of products that will form in a chemical rxn based on the amounts of reactants that react

o Allows to determine the amount of reactions necessary to form a given amount of product

o 16 moles of CO2 are produced for every 2 moles of octane burnedo Ex. how many moles of CO2 will form if we burn 22.0 moles of C8H18?

4.3 Limiting Reactant, Theoretical Yield, and Percent Yield

Limiting reactant – reactant that limits the amount of product in a chemical reactiono Produces the least amount of producto Completely consumed in a chemical rxn

Reactant in excess - reactant that do not limit the amount of product o Occurs in a greater quantity than is required to completely react with the limiting reactant

Theoretical yield – amount of product that can be made in a chemical reaction based on the amount of limiting reactant Actual yield – amount of product actually produced by a chemical reaction

o Always equal or less than the theoretical yield Small amount of product is usually lost to other rxns or does not form during a rxn

Percent yield - percentage of the theoretical yield that was actually attained

o % yield = actual yield

theoretical yield X 100%

Convert masses to amounts in moles4.4 Solution Concentration and Solution Stoichiometry

Solution – a homogeneous mixture of 2 substanceso Ex. salt & watero Aqueous solution – water acts as the solvento Dilute solution – contains small amount of solute relative to the solvento Concentrated solution – large amount of solute relative to the solvent

Solvent – majority component of the mixture Solute – minority component Molarity (M) – common way to express solution concentration

o M = amount of solute (¿mol)volume of solution(¿ L)

Laboratories often store solutions in concentrated formed called stock solutionso To save space in storerooms

o Concentration = molevolume

M1V1 = M2V2

For solving dilution problems # of moles of solute does not change when diluting a solution

4.5 Types of Aqueous Solutions and Solubility When a solid is put into a liquid solvent, the attractive forces that hold the solid together (the solute-solute interactions)

come into competition w/ the attractive forces between the solvent molecules & the particles that compose the solid (the solvent-solute interactions)

Electrolytes – substances that dissolve in water to form solutions that conduct electricityo Ex. salt – ionic compoundo Dissociate into their component ions when dissolved in watero Dissolved ions act as charge carriers

allowing the solution to conduct electricityo Strong electrolytes – substances that completely dissociate

into ions Resulting solutions are strong electrolyte solutions Strong acid – completely ionizes in solution

HCl (aq) – hydrochloric acid Represent complete ionization w/ a single rxn arrow

between the acid & its ionized formo Weak electrolytes – substances that dissociate partially into

ions Weak acid – do not completely ionize in water

CH3COOH (aq) – acetic acid Composed mostly of the non-ionized acid Only a small % of acid molecules ionize Represent w/ reverse arrows

Conduct electricity only weakly Non-electrolyte – compounds that do not dissociate into ions when dissolved in water

o Ex. sugar – molecular compound Sugar solution is composed of intact C12H22O11 molecules homogeneously mixed w/ the water molecules

o Most molecular compounds (except acids) dissolve in water as intact moleculeso Resulting solutions (non-electrolyte solutions) do not conduct electricity

Not all ionic compounds dissolve in watero Soluble – if a compound dissolves in watero Insoluble – if it does not dissolve

Forms precipitation4.6 Precipitation Reactions

Precipitation reactions – forms precipitate (solid) when mix 2 solutions Only insoluble compounds form precipitates

4.7 Representing Aqueous Reactions: Molecular, Ionic, and Complete Ionic Equations Molecular equation – shows the complete neutral formulas for each compound in the reaction as if they existed as

moleculeso Ex. Pb(NO3)2 (aq) + 2KCl (aq) PbCl2 (s) + 2KNO3 (aq)o However, in actual solutions of soluble ionic compounds, dissolved substances are present as ions

Complete ionic equations – list all the ions present as either reactants/products in a chemical reactiono Ex. Pb2

+ (aq) + 2NO3- (aq) + 2K+ (aq) + 2Cl- (aq) PbCl2 (s) + 2K+ (aq) + 2NO3

- (aq)

o Spectator ions – ions in solution appear unchanged on both sides of the equation Net ionic equations – show only the species that actually change during the reaction

o Without spectator ionso Ex. Pb2

+ (aq) + 2Cl- (aq) PbCl2 (s) 4.8 Acid–Base and Gas-Evolution Reactions

Acid-base reaction – an acid reacts w/ a base & neutralize each other, producing water (or in some cases a weak electrolyte)

o AKA neutralization reactiono Arrhenius definitions

Acid – substance that produces H+ ions in (aq) solution H+ ions normally associate w/ water molecule to form hydronium ions

o H+ (aq) + H2O H3O+ (aq) Polyprotic acids – contain more than 1 ionizable proton & release them sequentially

o Ex. H2SO4 is a diprotic acid

Base – substance that produce OH- ions in (aq) solution Some bases such as Sr(OH)2 produce 2 moles of OH- per mole of the base

o Antacids employ different bases as neutralizing agentso When an acid & a base is mixed, the H+ (aq) combines with the OH- (aq) to form H2O (l)o Acid-base rxns generally form water and salt that usually remains dissolved in the solution o Titration – a substance in a solution of known concentration is reacted w/ another substance in a solution of

unknown concentration Equivalent point – the point when the # of moles of OH- equals the # of moles of H+ in solution

Typically signalled by an indicator (a dye whose colour depends on the acidity/basicity of the solution)

Gas-evolution reaction – 2 aq solutions mix to form a gaseous product that bubbles out of solutiono Many are also acid-base rxns

When cation of one reactant combines with the anion of the othero Often form an intermediate product that then decomposes to form a gas

Ex. 4.9 Oxidation-Reduction Reactions

Oxidation-reduction reactions – e- transfer from one reactant to anothero AKA redox reactions

Combustion rxn is one type of redox rxno Ex. metal (loses e-) reacts with a non-metal (gains e-)o Oxidation – loss of electrons (LEO)

Increase in oxidation stateo Reduction – gain of electrons (GER)

Decrease in oxidation stateo Transfer of e- need not to be a complete transfer (as occurs in the formation of an ionic compound) for the rxn to

qualify as redox rxno Oxidation state/number – a “charge” given to each atom based on the electron assignment

To keep tract of e- before & after a rxn Ex. H has an oxidation state of +1 & Cl has an oxidation state of -1 Different from ionic charges (1+ & 1-)

Use to identify redox reactions

C changes from an oxidation state of 0 to an oxidation state of +4

o C loses e- & is oxidized S changes from an oxidation state of 0 to -2

o S gains e- & is reducedo Oxidation & reduction must occur togethero Oxidizing agent – substance that causes the oxidation of another substance

Oxidizing agent is always reduced Ex. oxygen

o Reducing agent – substance that cause the reduction of another substance Reducing agent is always oxideized Ex. hydrogen

Chapter 5 Gases5.1 Breathing: Putting Pressure to Work

Pressure – force exerted per unit area by gas molecules as they strike the surfaces around themo Body’s ability to create pressure differences to move air into & out of the lungso Total pressure exerted by a gas depends on several factors

The higher the concentration = the greater the pressure The greater the volume = the lower the pressure

Concentration How the lungs work:

o When exhale, the chest cavity muscles relax which the lung volume = the pressure & forcing air back outo When inhale, the chest cavity muscle expands which lung volume = the pressure & gaseous molecules flow into lungs

5.2 Pressure: The Result of Molecular Collisions Variation in pressure in Earth’s atmosphere creates wind Changes in pressure help us to predict weather

Pressure = forcearea

= FA

o The fewer the gas particles = the lower the pressure Pressure with altitude

Imbalance pressure within ear cavities when at a higher altitude o external pressure drops while internal pressure remains the sameo greater internal pressure forces eardrum to bulge outward causing pain

pressure can be measured in several different units:o mmHg (millimetre of mercury) – measurement originates from a barometer

unit of mmHg is often called torr 1 mmHg = 1 torr

o Atm (atmosphere) – average pressure at sea level 1 atm = 760 mmHg

o Pa (pascal) – SI unit of pressure 1 Pa = 1 N/m2

1 atm = 101,325 Pao in Hg (inches of mercury)

1 atm = 29.92 in Hgo psi (pounds per square inch)

1 atm = 14.7 psi Manometer – U-shaped tube containing a dense liquid (usually mercury) to

measure the pressure of a gas sample in the laboratoryo One end of the tube is open to atmospheric pressureo Other end is attached to a flask containing the gas sampleo Mercury levels on both sides of the tube will be the same if the pressure

of the gas sample = atmospheric pressureo Mercury level is higher on opened side if the pressure of the sample >

atmospheric pressureo Mercury level is higher on the closed side if the pressure of sample <

atmospheric pressure5.3 The Simple Gas Laws: Boyle’s Law, Charles’s Law, & Avogadro’s Law

4 basic properties of a gas sample:o Pressure (P)o Volume (V)o Temperature (T)o Amount in moles (n)o These properties are interrelated

When one changes, it affects the others Boyle’s Law – inverse relationship between volume & pressure

o An of 1 results in a of the other If the volume of gas sample , the same # of gas particles is crowded into a smaller volume, resulting in a more

collisions w/ the walls & in the pressure

o V 1p

(constant T & n)

V = (constant) x 1p

PV = constant P1V1 = constant = P2V2

P1V1 = P2V2

As long as the temperature & the amount of gas are constant Charles’s Law – linear proportional relationship between volume & temperature

o An of 1 results in a of the other with a same factor When the temperature of a gas sample , the gas particles move faster; collisions w/ the walls are more frequent, &

the force exerted w/ each collision is greater Only way for the pressure to remain constant is for the gas to occupy a larger volume, so that collisions become less

frequent & occur over a larger area Ex. when air is heated, its volume , resulting in a lower density

2nd floor of a house is usually warmer than the ground floor Hot-air balloon floats in the colder, denser surrounding air

o V T (constant P and n) V = (constant) x T

VT

= constant

V 1T 1

= V 2T 2

As long as the pressure & the amount of gas are constant o If linearly related lines between volume & temperature on a graph is extended/extrapolated backwards from the lowest

measured temperature, it shows absolute zero Gas will have a zero volume at -273.15°C

Coldest possible temperature Gas with negative volume is impossible

Avogadro’s Law – direct proportional relationship between volume & amount of gaso in amount of gas results in a of volumeo V n (constant T and P)

v 1n1

= v 2n2

As long as the pressure & temperature of the gas are constant Gay-Lussac’s Law – direct relationship between pressure & temperature

o As the temperature , the pressure o P T

As long as the amount of gas and volume is fixed 5.4 The Ideal Gas Law

Combination of Boyle’s law, Charles’s law, & Avogadro’s law PV = nRT

o V 1p

, V T , V n

o V nTP

o V = nRTP

Ideal gas constant – R= 0.08206 Latmmol K

o Quantities in the ideal gas law must be expressed in the units within R Pressure (p) in atm Volume (V) in L Moles (n) in mol Temperature (T) in K

5.5 Applications of the Ideal Gas Law: Molar Volume, Density, and Molar Mass of a Gas Molecular volume – volume occupied by 1 mole of a substance

o Molar volume of an ideal gas under standard temperature (273K) and pressure (1 atm) = 22.4L

Density = molar massmolar volume

o Under standard conditionso Density of a gas is directly proportional to its molar mass

Greater the molar mass of a gas = more dense the gas molar mass lower than that of air tends to rise

o n = mM

d = mV

o PV = nRT nV

= PRT

d = PMRT

Derived Density as molar mass

Molar mass – can be calculated to determine an unknown gas

o PV = nRT PV = mM

RT M = mRTPV

5.6 Mixtures of Gases & Partial Pressures Many gas samples exist as mixtures of gases

o Ex. air is a mixture of nitrogen, oxygen, argon, carbon dioxide & a few other gases in trace amounto Each type of gas is independent of the other gases in the mixture

Partial pressure (Pn) – pressure of individual component in a gas mixtureo Can be calculated using ideal gas law by assuming that each gas component acts independently

Ex. Pa = naRTV

; Pb = nbRTV

; Pc = ncRTV

; …

Dalton’s law of partial pressures - sum of the partial pressures of the components in a gas mixture must equal the total pressureo Ptotal = Pa + Pb + Pc + …

= naRTV

+ nbRTV

+ ncRTV

; …

= (na + nb + nc + …) RTV

= (ntotal)RTV

oPaPtotal

=

na(RTV

)

ntotal (RTV

) = nantotal

o Pa = aPtotal

Mole fraction (a) = nantotal

o Equivalent to its percent by volume divided by 100% When the desired product of a chemical rxn is a gas, it’s often collected by the displacement of water

5.7 Gases in Chemical Reactions: Stoichiometry Revisited Ex. use the conversion factors to determine the mass of product obtained in a chemical rxn based on a given mass of reactant

o Conversion factor between amounts (in moles) of each comes from the stoichiometric coefficients in the balanced chemical equation

Use the ideal gas law to determine the amounts in moles from the volumes, or vice versa5.8 Kinetic Molecular Theory: A Model for Gases

Kinetic molecular theory – models gas as a collection of particles (either molecules/atoms, depending on the gas) in constant motiono A single particle moves in a straight line until it collides with another particles (or with the wall of the container)o Basic assumptions of kinetic molecular theory are

1. The size of a particle is negligibly small Assumes that the particles themselves occupy no volume, even though they have mass

o Space between atoms/molecules in a gas is very large compared to the size of an atom/molecule itself under normal pressures

2. The average kinetic energy of a particle is proportional to the temperature in Kelvins Motion of atoms/molecules in a gas is due to thermal energy (distributed among the particles in the gas) At any given moment, some particles are moving faster than others (there is a distribution of velocities) The higher the temperature, the faster the overall motion the greater the average KE KE T

o The atoms in a sample of He & a sample of argon at the same temperature have the same average kinetic energy

Not the same average velocity Since He atoms are lighter, they must move faster to have the same KE as

argon atom 3. The collision of 1 particle with another (or with the walls of its container) is completely elastic

When 2 particles collide, they may exchange energy, but there is no overall loss of energyo Any KE lost by one particle is completely gained by the other

Completely elastic = particles have no “stickiness” & they are not deformed by the collision o Particles do not exert an forces on one another between collisionso Ex. billiard balls

Ideal gas law can be mathematically derived from kinetic molecular theoryo ideal gas law follows directly from KE theory

Boyle’s Law – force the gas to occupy a smaller space if volume is # of collisions with the surrounding surfaces must necessarily , resulting in a greater pressure As long as the temperature remains the same

V 1P

Charles’s Law – average speed (average KE) as temperature Greater KE results in more frequent collision P

o Only way for pressure to remain constant = volume V T

Avogardro’s Law – # of collisions with the surrounding surfaces as the # of particles Greater # of collisions = greater the pressure

o For pressure to remain constant, volume must V n

Dalton’s Law – total pressure of a gas mixture is the sum of the partial pressures of its components Particles have negligible size & they do not interact Mass is the only property that would distinguish one type of particle from another Particles with different masses have the same average kinetic energy at a given temperature

o exerting the same amount of force upon collision with a surfaceo same amount of pressure

Adding different kinds of gases has the same effect as simply adding more particleso ** PAGE 208 for derivation **

5.9 Mean Free Path, Diffusion, & Effusion of Gases Although gases particles travel at tremendous speeds, they also travel in haphazard paths

o Molecule travels only a short distance before it collides with another molecule, changes direction, only to collide again, etc A molecule in the air experiences several billion collisions per second Mean free path – average distance that a molecule travels between collisions

Increases with decreasing pressure Diffusion – process by which gas molecules spread out in response to a concentration gradient

o Root mean square velocity influences the rate of diffusion Heavier molecules diffuse slowly than lighter ones

first molecules we smell are the lighter ones Effusion – process by which gas escapes from a container into a vacuum through a small hole

o rate of effusion also depends on root mean square velocity heavier molecules effuse more slowly than lighter ones

o rate of effusion is inversely proportional to the square root of the molar mass of the gas

rate 1

√Mo Graham’s law of effusion – ratio of effusion rates of 2 different gases

ratearateb

= √M b

M a

Ex. explains why He balloons only float for a day or so He escapes from the balloon quite quickly since it has a low molar mass Air-filled balloon will remain inflated longer

5.10 Real Gases: The Effects of Size & Intermolecular Forces Gases behave ideally when:

1. The volume of the gas particles is small compared to the space between them Size of gas particles becomes important at high pressure at low pressure, molar volume of a gas is nearly identical to that of an ideal gas at high pressure, molar volume becomes greater than that of an ideal gas

volume of the particles themselves occupy a significant portion of the total gas volume2. The forces between the gas particles are not significant

Intermolecular forces – attractions between the atoms/molecules that compose any substance Typically small in gases

o do not matter how much at low pressure b/c the molecules are too far apart to “feel” the forceo Do not matter much at high temperatures b/c he molecules have a lot of KE

When 2 particles with high KE collide, a weak attraction between them doesn’t affect the collision much

Ex. if 2 billiard balls collide at high velocity, the “stickiness” will not have much effect Ex. if the 2 billiard balls might stick together and not bounce off one another at low

velocity At high temperature, pressure of a gas is nearly identical to that of an ideal gas At low temperature, the pressure of the gas is less than that of an ideal gas

Gas atoms will spend more time interacting w/ each other & less time colliding with the wallso These assumptions are valid for most common gases at STP

But assumptions break down at higher pressure or lower temperatures Van der Waals equation – combines the effects of particle volume & particle intermolecular forces to describe non-ideal gas behaviour

o [P + a(nv

)2] x [V – nb] = nRT

1 mol of an ideal gas

Chapter 9 – Chemical Bonding I: Lewis Theory Chemical bonding is at the heart of chemistry 3 theories:

o Lewis Theory – simple model of chemical bonding Used dots, dashes, and chemical symbols

o Valence bond theory – a chemical bond is the overlap of 2 orbitals that together contain two electrons treats e- in a more quantum-mechanical manner

o Molecular orbital theory – essentially a full quantum-mechanical treatment of the molecule & its e- as a whole Has great predictive power, but at the expense of great complexity & intensive computational

requirements9.1 Bonding Models & AIDS Drugs

X-ray crystallography – a technique in which X-rays are scattered from crystals of the molecule of interesto To determine the structure of HIV-protease (protein synthesized by HIV)

This particular protein is crucial to the virus’s ability to multiply & cause AIDS Wo/ HIV-protease, HIV cannot spread in the human body b/c the virus cannot replicate AIDS can’t

develop pharmaceutical companies create a molecule that would disable HIV-protease by inhibiting its active

site Bonding theories – models to predict how atoms bond together to form molecules

Explain why some combinations of atoms are stable & others are not Predict the shape of molecules

o Determine many of the physical & chemical properties of compounds In this case, it’s used to stimulate the shape of potential molecules & how they would interact

with the protease molecule Lewis theory – valence e- are represented as dots

o Lewis structure –predict moleculeso Simplest model for making quick, everyday predictions about most moleculeso Predict whether a particular set of atoms will form a stable molecule or noto Predict what the molecule might look like

9.2 Types of Chemical Bonds

Wo/ bonding, there would be just 91 different kinds of naturally occurring elementso Life would be impossible

Chemical bonds form b/c they lower the potential energy between the charged particles that compose atoms 3 types of chemical bonds:

o Ionic – transfer of an e- from metal to non-metal Metal becomes a cation Non-metal becomes an anion Lowering their overall PE when attrated

o Covalent – sharing of e- between non-metals b/c non-metals have high IE (their e- are relatively difficult to remove) shared e- interact with the nuclei of both of the bonding atoms

lowering their PEo Metallic – occurs in metals

Metals have low IE (tend to lose e- easily) Electron sea – all of the atoms in a metal lattice pool their valence e-

Pooled e- are no longer localized on a single atom, but delocalized over the entire metal + charged metal atoms are then attracted to the sea of e-, holding the metal together

9.3 Representing Valence Electrons with Dots Lewis structure represents valence e- as dots surrounding the symbol for that element

o With a max. of 2 dots per side

oo Octet rule – stable configuration of 8 e- in the outermost shello Predicts compounds

9.4 Ionic Bonding: Lewis Structures & Lattice Energies

o Transfer of e- gives Cl an octet & leaves K wo/ any valence e- but w/ an octet in the previous principal energy levelo K becomes + charged (cation)o Cl becomes – charged (anion)

Lewis structure of an anion is usually written within brackets with charge in the upper right-hand corner, outside

o 2 Na atoms each lose their 1 valence e-o S atom gains 2 e- & forms an octet

Formation of an ionic compound from its constituent elements is usually quite exothermico Energy comes from the crystal latticeo Lattice energy – the energy associated with forming a crystalline lattice of alternating cations & anions from the

gaseous ions Ex. Na & Cl bond together to lower the PE

Energy is emitted as heat when the lattice forms Born-Haber cycle – hypothetical series of steps that represents the formation of ionic compound from its

constituent elements Easiest way to calculate lattice energy Enthalpy of each step is known except for the last one, which is the lattice energy Change in enthalpy for the overall process is also known Enthalpy change for the unknown last step can be determined using Hess’s law H°f = H°step 1 + H°step 2 + H°step 3 + H°step 4 + H°step 5

H°lattice = H°step 5 = H°f – (H°step 1 + H°step 2 + H°step 3 + H°step 4 ) Lattice energies become less exothermic (less negative) with ionic radius

As the ionic radii as we move down the column, the ions cannot get as close to each other & therefore do not release as much energy when the lattice forms

Lattice energies become more exothermic (more negative) with magnitude of ionic charge9.5 Covalent Bonding: Lewis Structures

Sharing some (or all) of their valence e- in order to attain octets

o Bonding pair – a shared pair of e-

Often represented by a dash to emphasize that it constitutes a chemical bondo Lone pair – e- not involved in bonding

AKA nonbonding electrons Double bond – when 2 e- pairs are shared between 2 atoms

o Shorter & stronger than single bonds Triple bond – sharing of 3 e- pairs between 2 atoms

o Shorter & stronger than double bonds Intermolecular forces – bonding interactions between molecules

o Weakero Overcome this relatively weak force for a substance to boil

Intramolecular forces – bonding interactions within a moleculeso stronger

9.6 Electronegativity & Bond Polarity 1 limitation to the Lewis structure of representing e- as dots & covalent bonds as 2 dots shared between 2 atoms is that the

shared e- always appear to be equally sharedo Not always trueo Ex. hydrogen fluoride H side of the molecule have a slight + charge & F side has a slight – charge

E- density is greater on the F atom than on the H atom e- in HF is unequally shared in reality

Electronegativity – ability of an atom to attract e- to itself in a chemical bond (which results in polar & ionic bonds)o Ex. F is more electronegative than H b/c it takes a greater share of the e- density in HF1. Generally across a period in the periodic table2. Generally down a column3. Fluorine is the most electronegative element 4. Francium is the least electronegative element (sometimes called the most electropositive)o Inversely related to atomic size

The larger the atom, the less ability it has to attract e- to itself in a chemical bond Degree of polarity in a chemical bond depends on the electronegativity difference (EN) between 2 bonding elements

o The greater the EN, the more polar the bondo Purely covalent – equally shared e-

If 2 elements with identical electronegativities form a covalent bond

Molecules = nonpolar o Polar covalent bond – between a pure covalent bond & an ionic

bond - & + poles

o ionic bond – complete transfer of e-9.7 Lewis Structures of Molecular Compounds & Polyatomic Ions

writing Lewis structures for molecular compounds1. write the correct skeletal structure for the molecule

hydrogen atoms are always terminal (at the end) central atoms must form at least 2 bonds

put the more electronegative elements in the terminal position & the less electronegative (other than H) in the centre

2. calculate the total number of e- for the Lewis structure by summing the valence e- of each atom in the molecule3. distribute the e- among the atoms, giving octets (or duets in the case of H) to as many atoms as possible

begin by placing 2 e- between every 2 atoms then distribute the remaining e- as lone pairs

first to the terminal atoms then to the central atom

4. if any atoms lack an octet, form double/triple bonds as necessary to give them octets writing Lewis structure for polyatomic ions

o follows the same procedure charge of the ion must be considered when calculating the total # of e- for a polyatomic ion

add 1 e- for each – charge & subtract 1 e- for each + charge

o usually written within brackets with the charge of the ion in the upper right-hand corner, outside the bracket9.8 Resonance & Formal Charge

2 additional concepts to write the best possible Lewis structures for a large # of compounds Resonance – when 2 or more valid Lewis structures can be drawn for the same compound

o Molecule exists as an average of the 2 Lewis structureso Bonds are equivalent & each is intermediate in strength & length o Resonance structures – method of representing the molecule with both structures, with a double-headed arrow

between them Same skeletal formula, but different e- arrangement

o Resonance hybrid – actual structure of the molecule intermediated

between 2 or more resonance structures Only structure that actually exists

E- are localized either on the atom (lone pair) or between atoms (bonding pair) in Lewis structure E- are delocalized over several atoms/bonds in nature

o Resonance stabilization – stabilizes the structure by lowering their PE Formal charge – fictitious charge assigned to each atom in a Lewis structure that helps to distinguish among competing

Lewis structureso The charge it would have if all bonding e- were shared equally between the bonded atoms

Calculated charge for an atom if we completely ignore the effects of electronegativityo Formal charge = # of valence electrons – (# of nonbonding electrons + ½ # of bonding electrons)

Ex. formal charge for H in HF = 1 – [0 + ½(2)] = 0 Ex. formal charge for F in HF = 7 – [6 + ½(2)] = 0

1. Sum of all formal charges in a neutral molecule must be zero2. Sum of all formal charges in an ion must equal to the charge of the ion3. Small (or zero) formal charges on individual atoms are better than large ones4. When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom

9.9 Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets Octet rule in Lewis theory has some exceptions

1. Odd-electron species Molecules/ions with an odd # of electrons called free radicals (simply radicals) Tend to be somewhat unstable & reactive Ex. nitrogen monoxide has 11 e-

React with oxygen in the air to form NO2 (17 e-)o React with water to form nitric acid

2. Incomplete octets Molecules/ions with fewer than 8 e- around an atom

3. Expanded octets Molecules/ions with more than 8 e- around an atom Often for elements in the 3rd row of the periodic table & beyond of up

to 12 (& occasionally 14) e- b/c the d orbitals in these elements are energetically

accessible & can accommodate the extra electrons never occur in 2nd period elements

9.10 Bond Energies & Bond Lengths bond energy – energy required to break 1 mole of the bond in the gas phase

o used when standard enthalpies of formation are not available for all of the reactants & products of a rxn estimate enthalpy changes of rxn

o Always positive b/c it always takes energy to break a bond

o compounds with stronger bonds tend to be more chemically stable than compounds with weak bonds less chemically reactive

o Particular type of bond can have different bond energies in different molecules Calculate an average bond energy for a chemical bond

Found in a table Depend on the kind of atom & type of bond (single/double/triple)

o Hrxn = (H’s bonds broken) + (H’s bonds formed)

A rxn is exothermic when weak bonds break & strong bonds form A rxn is endothermic when strong bonds break & weak bonds form

o Breaking a chemical bonds always require energyo Forming a chemical bond always releases energy

Bond length – average length of a bond between 2 particular atoms in a large # of compoundso Depend on the kind of atoms involved in the bond & type of bond (single/double/triple)o Triple bond is shorter than double bondo Double bond is shorter than single bondo As the bond gets longer, it also becomes weaker

Triple bond is stronger than double bond Double bond is stronger than single bond

9.11 Bonding in Metals: The Electron Sea Model Metallic bonding – occur between metals

o Each metal atom donates 1 or more e- to an electron sea when bond together to form a solid Ex. each sodium atom donates its 1 valence e- to the “sea” and becomes a sodium ion

Held together by their attraction to the sea of e- metals conduct electricity

E- in a metal are free to move Movement/flow of e- in response to an electric potential (voltage) is an electric current

Excellent conductors of heat Mobile e- help to disperse thermal energy throughout the metal

Accounts for the malleability of metals Capacity to be pounded into sheets No localized or specific “bonds” in a metal it can be deformed relatively easily by forcing the

metal ions to slide past one another Accounts for the ductility of metals

Capacity to be drawn into wires

25% short answer75% 20-25 exam 120-25 exam 230-60 exam 3 + new

Exam 1 Building block of atomsProtonNewtonElectron istopesMolesEmpiralnformula Percent compostionMass dataCombustionCarbon always co2 gHydrogen h20 lBalancing equationStoichemistryNimenculture

YieldActual in gramsTheoretical

MolarityLewis structure

Group numberValence elctronValence shell expansion More than 4 atomsNomenclature t 3-6 atomElectronic Molecular geometry

Exam 2 Chemical reactionSolubility problemsAcid baseAntihydrztesRedoxOxidation numberCovalent bondingAssign Proportion same species oxidizated and reduced

Vander waalsPhase

Exam 3 uiliMolarity Molarity