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7/31/2019 Chemistry Notes Chapters 1-5
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I. Quantitative Chemistry1. Matter
a. Occupies space and has massb. Can be subdivided into mixtures and pure substances
i. Mixtures: different substances not chemically combinedii. Pure substance: one substance, consistent physical and chemical properties
throughout the substance
2. Elementa. A substance that only contains one type of atom, so it cannot be converted into
anything simpler.
b. Isotopes are atoms with different atom masses, because of more neutrons in thenucleus.
c. 92 elements3. Atoms
a. Moles: a quantity measuring the amount of substance (in the unit mol), that isproportional to the number of particles in a sample of substance.
i. Consistent with the amount of atoms in 12 grams of 126C .
b. 1 mol = 236.02 10 particles4. Stoichiometry
a. The study of quantitative aspects of chemical equationsb. Using a balanced equation, one is able to find the amount of substance used and
amount of substance made.
i. Should be equalc.
If given a mass of substance reacted with excess of another substance, one canfind the amount of product made.
i. Calculate the amount of substance whose mass is given (convert fromgrams to mols)
ii. Use the balanced equation to calculate the amount of required substance inthe reaction
iii. Calculate the mass of the required substance from the amount given.d. Limiting reagents
i. The reactant which limits the amount of product madeii. Can be found by calculating the amount of reagent present and dividing it
by the number of mols given in the formula, i.e. the number of mols thatare given and the number of mols that are needed
e. Applied to gasesi. At a constant temperature and pressure, a volume of any gas will always
contain the same number of particles.
ii. Called Avogadros hypothesis.
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iii. This means that the coefficients of the chemical reaction show the ratio ofthe volumes
iv. If temperature and pressure are given, one can calculate the molar mass.The equation
22.4
stpV
n shows this, where Vstp is the volume of gas and
22.4 is the molar mass.
v. This can be rearranged to do calculations in the same way as otherstoichiometric problems
5. Ideal Gas Equationa. An ideal gas is one in which the particles have negligible volume, no attractive
forces between particles, and the kinetic energy of the particles is proportional to
the absolute temperature.
b. Equation: PV nRT , where P is pressure, V is volume, n is number of mols, R isa gas constant (in most cases, 0.08206), and T is temperature (mostly in Kelvin).
c. This equation can be used to find any one of those values, granted that all othervalues are known
6. Gas Equationsa. Used when some values (for example, n and T) are constant
i. Boyle-Mariotte Law: at a constant n and T, 1 1 2 2PV P V
This graph shows the relationship between pressure and volume; the y-axis is volume, and the x-axis is pressure. When volume is high, pressure
is low, and vice versa. (They have a inverse relationship). This is
explained by the fact that more pressure means less space for the particles
of the gas to occupy.
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ii. Charles law: when n and P are constant, 1 21 2
V V
T T (T will be measured in
Kelvin)
This shows that they have a direct relationship; as temperature increases,
so does volume. This is explained by the fact that the higher the
temperature, the more kinetic energy each particle has, which causes the
gas to expand i.e. takes up more space
iii. Gay-Lussacs law: when n and V are constant, 1 21 2
P P
T T
This law states that as temperature increases, so does pressure. This can be
explained, once again, by kinetic energy. Temperature is a measure of
kinetic energy; it shows us how rapid the particles move. As the
temperature increases, the particles are colliding with the container moreand more, making the pressure increase (volume is held at a constant).
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iv. If all three variables are present, combined gas law can be used.Combined gas law: when n is constant, 1 1 2 2
1 2
PV P V
T T (Temperature must be
expressed in Kelvin; others may be expressed in any units).
7.
Solutionsa. Made up of the liquid (solvent) and the thing being dissolved in the liquid(solute).
b. Solubility: the quantity of that substance that will dissolve to form a certainvolume of solution in that solvent.
i. Solubility varies with temperature; in general, the higher the temperatureof the solvent, the more soluble the solute is.
ii. For gases being dissolved in solvents, it decreases with higher temperatureiii. Can be considered dilute (little amount of solute) or concentrated (a lot of
solute)
iv. Saturated solution: when no more solute will dissolve at a giventemperature (it reached its limit for the temperature given)
v. Supersaturated solution: when the solute concentration exceeds the statedlimit; usually occurs when the temperature of the solution goes from high
to low, or if the solution is produced using chemical reactions. This is the
cause of precipitation
c. Concentration, or [ ]: the amount of substance contained within a given volume ofsolution, given by the equation
nM
V , where M is molarity, n is number of
moles, and V is volume in liters.
8. Titrationsa. Titration is a technique which involves measuring the volume of one solution
which just reacts completely with another solution
b. One solution will have a known concentration. This solution is the called thestandard solution. To check the solution, one may titrate it against a primary
standard, or a solution which is prepared by dissolving a precise mass of solute to
make an accurate concentration
c. An indicator is usually added to the standard solution, and the second, unknownsolution is run in from a burette until the indicator changes colors.
d. The amount of solute can be calculated from the volume of the solution of knownconcentration. The amount of unknown may be found using the balancedequation. The concentration of the unknown can be calculated from this and the
volume of the second solution used.
i. Calculate the amount of moles in the solution of known concentrationii. Use a balanced equation to calculate the amount of the unknown (ratios)
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iii. Calculate the concentration of the unknown using nMV
.
e. Back titration: an insoluble solid reagent is reacted with a known excess of onereagent. When the reaction with the sample is complete a titration is then carried
out to determine how much of the reagent in excess remains unreacted.
i. Same basic principle of titrationsII. Atomic Structure
1. Atomic Theorya. John Dalton-1807
i. All matter is made up of a small number of different kinds ofatoms, which are indivisible and indestructible, but which can
combine to make molecules and compounds
ii. This is mostly true; however, atoms are not indivisible, becausethey are composed of many different, smaller parts.
2. Subatomic particlesa. Proton, Neutron, Electroni. Proton and Neutron have an amu (atomic mass unit) of 1 each,
while the electrons amu is negligible.
ii. Protons and neutrons are located in the nucleus, while electrons arelocated in the electron cloud.
iii. Most of the size of the atom is empty space, called the electroncloud
iv. Electrons: -1Neutrons: 0
Protons:+1
b. The difference between elements, in essence, is the number of protons inthe nucleus
i. Isotopes have different number of neutrons, but are still the sameelement, and one can add or take away (to a certain extent)
electrons, creating ions.
ii. Mass number: the sum of the protons and neutrons in the nucleusiii. Atomic number: number of protons in the nucleus
c. Isotopic notation: AZ X , where A is the mass number, Z is the atomicnumber, and X is the element in question
d. Neutrons and protons are usually the same at light elements (such asCarbon), but because of proton repulsion (think magnets, when trying to
put positive pole to positive pole), a greater number of neutrons are
needed for stability when more protons are present in the nucleus
e. Protons and electrons are equal in atoms
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i. Atoms can gain or lose electrons to form ions, which have a netelectrical charge because the number of protons is different from
the number of electrons.
ii. This is judged by number of electrons; electrons are the ones beingchanged
f. Isotopes: atoms that have the same number of protons but differentnumber of neutrons.
i. Same atomic number, different mass numberii. Mass number is not how heavy an atom weighs; rather, it is the
average weight of all isotopes of a certain atom.
3. Mass Spectrometera. An instrument which separates particles according to their masses and
records the relative proportions of these particles
b. Has six stepsi.
Vaporization1. If the sample is either a solid or liquid, the substance will
be heated to produce vapor
ii. Ionization1. The particles are converted from neutral atoms or
molecules into positive ions, usually from the
bombardment with fast moving electrons (can knock an
electron from the atom/molecule off)
iii. Positive ions are accelerated by electrical field1. Positive ions are accelerated by the high electrical potential
difference between the two parallel electrodes with holes intheir centers
iv. Ions deflected by a magnetic field1. The fast moving electrons enter a magnetic field produced
by an electromagnet, which causes them to deflect
v. Detector records ions of a particular mass1. Those with greater masses will not travel as far, while those
with smaller masses will travel very far
vi. Vacuum prevents molecules from colliding4. Electron Configuration
a. Energy levels and sub levelsi. Each energy sub level is divided into orbitals, which can contain
up to two electrons that have opposite spin
1. Pauli exclusion principle-no two electrons can occupy thesame space at the same time
b. s sub level
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i. Energy level closest to the nucleusii. Only contains one sub level and one orbital; spherical in shape
2 electrons, with opposing orbits
c. p sub leveli. figure eight electron distribution
ii. contain three p sub levelsiii. only differ in that one is oriented along the x-axis, a second along
the y axis, and a third along the z axis
iv.
each orbital can hold two electrons, making six p-electrons and atotal of eight in the second level (2s and 2p)
v. because of the increased number of electrons, there is an increasein the amount of electron-electron repulsion
Pz: 2 electrons Py: 2 electrons Px:2 electrons
Crudely-drawn p sub level; total of 6 electrons
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d. d sub levelsi. 5 d orbitals, all of the same energy
ii. Can hold ten electrons, giving the third level 18 electrons (2 from3s. 6 from 3p, and 10 from 3d)
iii. Comes after 4s, because d orbitals occur at a higher energy levelNote: this only has a total of 10 electrons, 2 in each level; even
though 4 spaces are made, only 2 can be filled
e. Electrons in atoms always adopt the lowest energy level possible by tryingto fill one sublevel before attempting to fill the next. This is called the
Aufbau principle
f. Hunds rule: the principle of maximum multiplicity, or that sub levelorbitals will be formed with as many same-spinning electrons as possible
g. In general, the 4s orbital will be filled before the 3d orbital.i. This is because the 3d has a higher energy level
ii. Two exceptions to the rule: Chromium and Copper1. This is because these elements form cations, and not
anions.
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2. Electron configuration: Cr: 3d54s1; Cu: 3d104s1h. Electron waves
i. Schrodinger wave equation1. Describes wave like nature in atoms2. Involves the quantum numbers
a. Principle quantum number (n)i. Dictates the energy level (1,2,3,4 etc.)
b. Azimuthal (subsidiary) quantum number (l)i. Dictates the sublevel (when l=0, the sublevel
is s, when l=1, the sublevel is p, etc.)
c. Magnetic quantum number (m)i. Dictates the orbital in which the electron
resides (for example, px or dxy) The value of
m can be dictated from the value ofl. For
example, when l is 0, m is 0. When l is 1, mcan be -1,0, and 1. When l is 2, m can be -2,-
1,0,1, or 2, and so on.
d. Spin quantum numberi. Differentiates between the two electrons in
the orbital by spin direction. In an electron
configuration model, upward pointing
arrows are equal to +1/2, while
downward point arrows are equal to -1/2.
3. Pauli exclusion principle: no two electrons in a given atomcan have the same four quantum numbers
i. Ionization energiesi. The minimum amount of energy required to remove a mole of
electrons from a mole of gaseous atoms to form a mole of gaseous
ions.
ii. The more electrons that have been removed from an atom, thegreater the energy required to remove the next electron.
iii. Going down a group, the ionization energy decreases. This isbecause the nuclear charge stays constant (the new protons are
being canceled out by the new electrons), but the valence shell is
farther away from the nucleus
iv. Going across a period, the ionization energy increases. This isbecause all of the electrons are in the same energy level going
across a period, but the charge of the nucleus increases because of
increasing protons.
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III. Periodicitya. Physical properties
i. Electronegativity, effective nuclear charge, atomic radii, ionic radii,melting points, electron affinity, ionization energy
1. Electronegativity: a measure of how strongly the atom attracts theelectrons in a covalent bond
2. Effective nuclear charge: the amount of force which the nucleusexerts on the electrons on the valence electrons
3. Atomic radius: the average distance from the nucleus to the edgeof the electron cloud (like the radius of a circle)
4. Ionic radius: same thing as atomic radius, but for ions5. Melting point: the temperature at which a transition from the solid
state to the liquid state occurs
6.
Electron affinity: the amount of energy released when a neutralatom is turned into a negative ion
7. Ionization energy: the amount of energy needed to take an electronfrom a gaseous mole of atoms or ions
Note: Electron affinity is the amount of energy released, while
ionization energy is the amount of energy needed.
ii. Trends1. Effective nuclear charge: relatively the same for each element,
because with each increasing interval on the periodic table, both a
proton and a neutron are added. These charges cancel each other
out.2. Electronegativity: decreases down a group and increases across a
period. Can be explained by effective nuclear charge. Going down
the group, for successive elements, there are more energy levels
filled with electrons, so the outer valences are farther away from
the nucleus. So, when going down a period, the valence electrons
are farther away from the nucleus, which means the attraction is
much less. This makes it harder for the nucleus to attract electrons
for a covalent bond; ergo, as one goes down a period, the
electronegativity decreases. Going across the period, another
proton is added to the nucleus, thus making it strong. This makes
the effective nuclear charge increase, which pulls the valence
electrons in closer. This also means that, in a covalent bond, the
nucleus would exert more force on the shared electron. Thus, as
one goes across a period, the electronegativity increases.
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3. Atomic and ionic radius: increases down a group and decreasesacross a period. As one goes down the group, the electrons are
farther away from the nucleus, so they are loosely attracted, so the
atomic and ionic radii increase. As one goes across the period, the
increasing amount of protons cause the electron cloud to become
more compact; thus, as one goes across the period, the atomic and
ionic radii decrease.
4. Ionization energy: decreases down a group and increases across aperiod. This is because down a group, the electrons are not very
compacted, and the nuclear charge is decreasing; across a period,
the protons cause the electron cloud to become more compacted,
thus making it harder for an electron to be taken away.
5. Electron affinity: electron affinity decreases down a group andincreases across a per
b. Chemical propertiesi. Alkali metals
1. Li, Na, K, Rb, Cs.a. Soft, malleable metals with low melting points and low
densities
i. Low density is the result of the atoms of alkalimetals being the largest atoms in their period
ii. Softness/low melting point a result of their loneelectron
b. Alkali metals are very chemically reactive
Down:
Electronegativity, ionizationenergy, and electron affinity
decrease; Atomic/ionic radii
increase
Across
Electronegativity, ionization
energy, and electron affinity
increase
Atomic/ionic radii decrease
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i. This a result of only having one electron in theirvalence shell
ii. Electron is very easily lostiii. They will always form cationsiv. Tend to react with group 7 elements (O, Cl, Br)
c. Going down the period, atomic radius increases, theionization energy of the elements decreases, and the
reactivity increases
ii. The Halogens1. F,Cl,Br,I.
a. Mostly nonreactiveb. All exist as diatomic molecules ( 2X )
i. Joined by covalent bondingc. F and Cl are gases, Br is liquid, and I is solid
i. This is a result of the increasing strength of the VanDer Waals forces
d. Because they only require one electron to complete theoctet, they are very electronegative
i. Electronegativity decreases down the group,because they valence shell is farther away from the
nucleus
ii. Means reactivity decreases down the groupe. Oxidizing power decreases down the group
i. This means that a higher halogen will displace alower halogen from its salt (e.g. Cl would displace
I-)
iii. Trends in oxides of period 3 elements1. At the left hand side of the periodic table, Na and Mg have
relatively low ionization energies and so they bond to other
elements to form ionic compounds in which they have lost their
electron. The oxides of these therefore are ionic, and have an oxide
ion. The oxide ion can form a bond to hydrogen ions and as a
result, act as bases dissolving water.
2.
Toward the middle, ionization energy increases, which causescovalent bonding. As such, these elements (C and Si) become
slightly acidic
3. At the far right, the elements continue to form covalent bonds bysharing electrons, but taking electrons from metals is an option.
The oxides of these are able to make acidic solutions
c. First row d-block elements
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i. D-block elements are all dense, hard metallic elements.ii. 3d electrons effectively shield the outer 4s electrons
1. This means the first ionization energy remains somewhat constant2. Also indicates that they have similar physical and chemical
properties
iii. Because of the semi-filled 4s orbital, the d block elements have the abilityto form a variety of stable oxidation states, to form complex ions, to form
colored ions, and to be used as a catalyst in some reactions
1. Variable oxidation states: occurs because of how close 3d and 4sare in energy levels. All d block elements have an oxidation state
of +2 except for Scandium. All d block elements also show an
oxidation state of +3 except for Zinc.
2. Complex ions: ions of d block metals and those in the low p-section (such as lead) have a low energy unfilled d and p orbitals.
These orbitals are able to accept a lone pair of electrons fromligands to form a bond between the ligand and the metal ion.
3. Colored ions: complexes including d block elements are usuallycolored
a. Exception: d0 (Sc3+ and Ti4+) d10 (Zn2+) ionsb. These colors can be explained by:
i. Electron transitions of d-electrons within the d subshell.
ii. Electron transitions from the metal ion to the ligandor the ligand to the metal ion, which are known as
charge transfer transitionsiii. Ligands themselves may be colored and this color
colors the comples.
c. The electrons move from one lower energy orbital to ahigher energy orbital, which creates color as light passes
through and absorbs the energy
4. Catalytic activitya. Occurs because d block elements can form complex ions
with ligands donating one lone pair of electrons, and the
fact that they have multiple oxidation states, so they can
readily gain and lose electrons in reduction-oxidation
reactions.
i. Two important mentions: Haber process, whichuses iron to create ammonia; Contact process,
which uses vanadium(V) oxide to create sulfuric
acid.
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IV. Chemical Bondinga. A chemical bond is an interaction between atoms or ions that results in a
reduction in the potential energy of the system which hence becomes more stable.
i. The bond type depends on the extent to which the atoms involved attractvalence (electronegativities)
1. If the elements have very different electronegativities, then ionicbonding results
2. If the elements have very similar electronegativities, covalentbonding results
b. Ionic bondingi. Occurs between elements that have a large difference in electronegativities
1. In ionic bonding, a metal atom with a low electronegativity loseselectrons to form a positively charged cation, while a nonmetal
atom with a high electronegativity gains electrons from the metal
to form an anion.
2. The resulting electrostatic connection between these ions causesionic bonding.
3. This means we can predict ionic bonding.a. In most cases, ionic compounds are isoelectric with the
noble gases.
b. The three dimensional shape of the anions and cations inionic crystals account for the high melting points and
stability of ionic solids.4. In most s block elements, the elements lose all electrons from their
valence.
a. This means that the alkali metals (group 1), which onlyhave one electron, create an X+ cation. The alkaline earth
metals (group 2) for X2+ cations.
5. Outside of the s block, predictions become harder.a. In transition metals, the atoms can make multiple stable
cations
6. Nonmetals usually gain electrons, because they have higherelectronegativities
a. Happens for the same reason metals lose electrons7. The anions and cations have opposite electrical charges and are
attracted to each other into a crystal lattice
a. Each anion is surrounded by cations and vice versa. (shownbelow)
+-+
- + -
+ +-
-+-
+ - +
- -+
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b. In these lattices, the charges add up to a net charge of zeroi. This means that the formulas of the ionic
compounds must equal charges of 0; no positive or
negative compounds
c. This model can be made into three dimensionsd. An ionic substance is held together by strong electrostatic
attractions in all three dimensions. This means that no
molecules are present in the ionic substance
e. Physical propertiesi. Hard, brittle crystalline solids
ii. Relatively high melting and boiling pointsiii. Do not conduct electricity when solid, but do when
molten or in aqueous solution (electrolytes)
iv.
Are more soluble in water than in other solventsc. Covalent bonding
1. Occurs between atoms that have quite high electronegativitiesa. Usually nonmetals
2. In covalent bonding, the two atoms involved share some of theirvalence electrons since neither element loses electrons easily.
a. The attraction of the two positively charged nuclei for theseshared pairs of electrons results in the two atoms bonding
The arrows represent the nuclear force on the shared
electrons, which causes the covalent bonding
b. In most cases, each atom donates one electroni. In some cases, known as dative covalent bonding, one
atom can donate both electrons.
c. When forming covalent bonds, the atoms involved usually filltheir valence shell.
i. The number of bonds formed are equal to the numberof electrons needed to form the valence shell (since, in
general, each atom donates one electron).
1. Good example: carbon needs 4 electrons, and isable to make 4 bonds
+ +Shared e-
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ii. Octet rule: atoms in compounds usually have eightelectrons in their valence shell
d. If two pairs of electrons are shared, a double bond is formedi. Joins the two atoms more tightly and closer together.
ii. Carbon forms four bonds and oxygen forms two bonds;when oxygen and carbon atoms bond, two double
bonds are used
e. Two atoms can also share three pairs of electrons, giving a triplebond
f. Covalent bonding can lead to two different types of structuresi. They can form molecules, which are two atoms bonded
together.
1. Physical properties of being soft in the solidstate, not conducting electricity, and being
more soluble in nonpolar solvents
ii. A lattice can be held together using covalent bondsoccasionally
1. Physical properties: very hard, very high meltingand boiling points
d. VSEPR theory1. Valence Shell Electron Pair Repulsion
a. Determines the shape of a moleculeb. Most molecules have filled valence levels that contain four pairs
of electrons
i. To be as widely separated as possible, these electronsdistribute themselves so that they are pointing toward
the corners of the tetrahedron.
ii. Some molecules have nonbonding pairs of electrons,which increase electron-electron repulsion and affect
the shape of the electron
Number of regions of
high electron density
Number of non-
bonding electron pairs
Example Shape and bond angle
Two None Carbon dioxide Linear; 180 degrees
Three None Boron trifluoride Trigonal Planar; 120
degrees
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Three One Sulfur dioxide V-shaped; 117 degrees
Four None Methane Tetrahedral; 109.5
degrees
Four One Ammonia Trigonal pyramidal; 107
degrees
Four Two Water Bent; 104 degrees
Five None Phosphorus
pentafluoride
Trigonal bipyramidal; 90
and 120
Five One Sulfur tetrafluoride See-saw; 90 and 117
Five Two Iodine trichloride T-shaped; 90
Five Three Xenon difluoride Linear; 180
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Six None Sulfur hexafluoride Octahedral; 90
Six One Bromine pentafluoride Square pyramid; 88
Six Two Xenon tetrafluoride Square planar; 90
2. Molecules with more than four electron pairsa. Happen in the third and lower p-block because the p block
can promote one or more electron from a doubly filled s or
p orbital into an unfilled low d orbital.
i. In essence, taking an electron from one level andbringing it up to a higher level
ii. This causes them to have an expanded valence shell1. A valence shell with more than 8 electrons
in it
iii.
This usually only happens when the atom is able toform very strong covalent bonds
1. This means they are very, veryelectronegative elements; so mostly groups
7 and 8 (esp Fl O Cl)
iv. These atoms attach to small central atoms, so thatthey can fit around without much electron-electron
repulsion
v. Molecules with expanded valence shells can onlyhave octahedral, trigonal bipyramidal, or square
pyramidal1. Trigonal bipyramidal: has 5 electron areas;
two axial (along the y axis) and three
equatorial (sticking out from the origin, to
create that pyramid shape)
2. Square planar=Octahedral shapetwo filledelectron areas
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3. Polaritya. Based off of electronegativities
i. In a polar bond, one is very, very electronegative,the other is weakly electronegative
ii. This causes the electrons of the less electronegativeatom to basically attach to the more electronegative
one, creating a positive (the less electronegative)
and negative (greater electronegative)
iii. The greater the difference in the electronegativity,the greater the polarity of the bond
iv. Dipole: a separation between the positive andnegative ends.
1. This is based off of symmetry as well. Somedipoles can cancel themselves out if they
have a linear symmetry; others have a dipolebecause the shape is not symmetrical
v. Dipole moment: a measure of the polarity of amolecule; the greater the polarity, the higher the
dipole moment is.
e. Hybridization1. When an atom bonds the atomic orbitals involved in forming the
bonds, or accommodating the lone pair of electrons, interact
with each other to form an equal number of directional hybrid
orbitals of equal energy
a. When atoms join together to form molecules, their outeratomic orbitals interact with each other to produce hybrid
orbitals.
b. The shapes of the hybrid orbitals correspond with theshapes of the molecules according to VSEPR.
i. The best way to determine the hybridization of amolecule is to look at the shape of the molecule
Geometry Hybrid Orbitals Number of Orbitals
Atomic Orbitals used
to form Hybrid
Orbitalslinear sp 2 s, pz
trigonal planar sp2 3 s, py, pz
tetrahedral sp3 4 s, px, py, pz
trigonal bipyramidal dsp3 5 s, px, py, pz, dz2
octahedral d2sp3 6s, px, py, pz, dz2,
dx2-y2
square planar dsp2 4 s, px, py, dx2-y2
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f. Multiple bonds1. Double bonds formed between two atoms are not identical
a. The first is formed by either the interaction with the sorbital or a hybrid orbital.
i. When they interact, they produce a bond in whichthe electron density is at its greatest on the inter-
nuclear axis, and symmetrical about it
ii. This is a bondb. The second bond in a double bond is formed by the side-on
interaction of electrons in p orbitals at right angles to the
inter-nuclear axis.
i. This bond has a low electron density on the inter-nuclear axis, but regions of high electron density on
opposite sides of this.
ii. This is called a bond.
c. Single bonds are always bonds, and double bonds arealways made up of one and one bond.
i. Triple bonds are made up of 1 sigma and two pibonds.
ii. Double and triple bonds are much stronger thansingle bonds, so the nuclei involved are closer
together.
1. This means bond energies increase, whilebond lengths decrease
g.Intermolecular forces
i. van der Waals forces, dipole dipole forces, and hydrogen bonding1. van der Waals forces: London forces are relatively weak forces of
attraction that exist between nonpolar molecules and noble gas
atoms.
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a. They are caused by instantaneous dipole formation; in thisprocess, electron distribution in the individual molecules
suddenly become asymmetrical, and the newly formed
dipoles are now attracted to one another.
i. The larger the molecules electron cloud, the strongvan der Waals forces which exist.
2. Dipole-dipole: take place when two or more neutral polarmolecules are oriented such that their positive and negative ends
are close to each other,
a. Because this is an attraction between unlike charges, theytend to be a strong form of bonding
3. Hydrogen bonding: not true bondsa. In essence, just held together by the strong attraction
between hydrogen (lowest electronegativity) to a highly
electronegative atom on a nearby moleculei. This means that it is found constantly in bonds with
oxygen, nitrogen, and fluorine: the three most
electronegative atoms on the periodic table
ii. Explains the unique characteristics of water, such ashigh specific heat and boiling point.
iii. Hydrogen bonds have a high partial positive charge,while the more the other atom has a negative
charge. This results in bonding
iv. Hydrogen bonding is one of the stronger bonds thatcan occur.
h. Metallic bonding
i. Occurs between atoms which all have low electronegativities1. Close packed lattice: when metal atoms are all packed together in a
fashion so that all molecules are as closely packed together as
possible.
2. In this fashion, no valence electrons belong to a specific atom-they are delocalized among all atoms in the lattice.
3. Because the electrons are not with any of the atoms, each of theatoms becomes a cation
4. The attraction between the cations and the free floating electronscausing the force which holds the structure together.
a. A lattice of cations within a sea of electronsii. Because the atoms are attracted to the free floating electrons and not the
ions themselves, this allows the layers of ions to slide past each other
without any bonds breaking.
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1. If another ion is introduced, the ability to slide without any bondsbreaking is lessened, those making alloys harder than pure metals.
2. The delocalized electrons are free to move from one side of thelattice to the other when some sort of potential difference occurs.
a. This makes them good electrical and thermal conductors
i.Covalent bonding: part IIi. Diamond
1. Diamond is the most common example of a substance that has agiant three dimensional covalent structure. Each carbon atom in
diamond is sp3
hybridized, and is joined to four others, arranged in
a tetrahedron.
a. This means that there is strong bonding in all threedimensions.
2. Silicon has an almost identical structure to diamond.a. The sideways overlap between the p orbitals of the larger
atoms is less, so other allotropes that involve pi bonding do
not occur.
b. Silicon dioxide has a similar structure, but each C isreplaced with Si, and the C-C bonds are replaced with
oxygen bridges.
3. Fullerences: recently discovered allotrope of pure carbon.a. They contain approximately spherical molecules made up
of five and six membered carbon rings.
b. C60i. Acts as an electron deficient molecule readily
accepting electrons from reducing agents.
V. Energeticsa.Thermochemistry
i. The study of energy changes associated with chemical reactions.1. Most chemical reactions absorb or evolve energy, usually in the
form of heat
ii.
Enthalpy1. The total energy of a system, some of which is stored as chemical
potential energy in the chemical bonds of the system.
2. In chemical reactions, bonds are broken and bonds are made.a. The energy involved in making new bonds is rarely equal
to the energy absorbed in breaking the old ones.
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b. This means that every chemical reaction has an enthalpychange
i. Given the symbol H .ii. It is equal to the difference in enthalpy between the
reactants and the products
c. Two types of reactions: endothermic and exothermici. Endothermic:positive H
ii. Exothermic: negative H1. This is because endothermic needs energy,
which means that the amount of energy in
the system increases; conversely, an
exothermic reaction releases energy, which
causes the energy of the system to be
negative
iii.
Enthalpy changes1. Temperature: the average kinetic energy of the particles measured
a. Is an intensive propertyb. The absolute (K) temperature is proportional to the mean
kinetic energy and is independent of the amount of the
substance present.
2. Heat: the measure of the total energy in a substance and doesdepend on the amount of substance present
a. Does have an effect on the temperature of a system3. When temperature increases, heat energy is absorbed from the
surroundingsa. This is dependent on the mass, m, of the substance, the
specific heat capacity, c of the substance, and the amount of
increase of temperature T.
b. Gives the equation: Heat energy=m.c.T4. Calorimetry: a technique used to measure the enthalpy associated
with a particular change
a. This technique is dependent on the assumption that no heatis gained or lost from the surroundings.
i. This is why calorimeters tend to be well insulated5. Hesss law
a. First Law of Thermodynamics (conservation of energy):states that energy cannot be created or destroyed
i. This means that the total change in chemicalpotential energy must be equal to the energy lost or
gained by the system
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ii. Also means that the total enthalpy change onconverting a given set of reactants to a particular set
of products is constant, irrespective of the way in
which the change is carried out
iii. This is Hesss law.1. Basically states that the total enthalpy
change of a reaction is equal to the enthalpy
changes of each step.
2. 1 2 3H = H H iv. Bond Enthalpies
a. All chemical reactions involve the breaking and making ofbonds.
b. Bond enthalpies=the measure of the strength of a covalentbond
i. The stronger the bond, the more tightly the atomsare joined together.
c. The breaking of bonds is an exothermic reaction; as such, itreleases energy
d. The formation of chemical bonds is an endothermicreaction i.e. it requires energy from the surroundings to
work.
e. The bond enthalpies are dependent on how the rest of themolecule is bonded. This means that the average bond
enthalpies may be defined as the enthalpy required to break
a particular covalent bond in a range of molecules.
i. o o oreaction formation formationH = BH (products) - B H (rea
ii. B=the coefficient in the formulaNote: if the bonds being made are weaker than
those being made, the reaction will be exothermic
and vice versa.
1. Also: bond enthalpies are for the conversionof a mole of gaseous molecules into gaseous
atoms.
2. This means that bond enthalpies are lessprecise than other methods.
1. The values, in most cases, are within10% of the actual value.
f. Bond strength increases from single bonds, through doublebonds to triple bonds.
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i. As bonds become stronger, the bonds also becomeshorter
1. Recall: electron density in the bondincreases the attraction of the nuclei for this
electrons, pulling the nuclei together.
g. It is not possible to determine bond enthalpies directly, sothey must be determined indirectly.
i. This is done by applying Hesss law.2. Standard Enthalpy changes of a reaction
a. The amount of energy evolved or absorbed in the formationof one mole of the compounds, in its standard state, from
its constituent elements in their standard state.
i. Standard state is the state the element is normallyfound in at room temperature (298 K).
ii.
If there are allotropes, the more stable one isconsidered the standard state.
1. O2(g) is the standard state of oxygen, notO3(g)
b. The sum of the enthalpies of formation of the reactants willgive the total enthalpy change to form the reactants from
the component elements in their standard states.
i. Similarly, the sum of the enthalpies of formation ofthe products will give the total enthalpy to form the
products.
c. oreaction formation formationH = BH (products) - B H (reactants)
d. Standard enthalpy change of combustioni. The enthalpy change when one mole of the
compound undergoes complete combustion in
excess oxygen under standard conditions
ii. Many covalent compounds will undergocombustion and hence it is often easy to determine
the standard enthalpy change of combustion for
molecules.e. Standard enthalpy change of neutralization
i. The enthalpy change when one mole of the acidundergoes complete neutralization with a strong
base (can also start with a base and end with an
acid).
ii. Always exothermic
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v. Born-Haber cyclea. The formation of an ionic compound can be considered as
the sum of a number of individual processes converting the
elements from their standard states into gaseous atoms,
losing and gaining electrons to form the cations and the
anions respectively and finally, these gaseous ions coming
together to form the solid compound
b. Standard enthalpy change of atomization: the enthalpychange required to produce one mole of gaseous atoms of
an element from the element in the standard state.
c. Electron affinity: the enthalpy change when one mole ofgaseous atoms or anions gains electrons to form a mole of
negatively charged gaseous ions.
i. This change is mostly exothermic for the first level;second level is endothermic because of the electronrepulsions
d. Lattice enthalpy: the energy required to convert one moleof the solid compound into gaseous ions.
i. Very very highly endothermic1. Think how much energy it takes to change
solid water into water vapor
ii. Lattice enthalpies depend upon the nature of theions involved
1. The greater the charge on the ions, thegreater the electrostatic attraction and hence,the greater the lattice enthalpy (and vice
versa)
2. The larger the ions, the the greater theseparation of the charges and the lower the
lattice enthalpy (and vice versa)
e. The Born-Haber Cyclei. Born-Haber cycle, if the magnitude of every term
except one is known, then the remaining value may
be calculated
1. The equation for this is: Enthalpies ofatomization + Electron affinities +Ionization
energy = Enthalpy of formation + Lattice
Enthalpy
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ii. The Born-Haber cycle provides a way in whichlattice enthalpies can be indirectly measured
through experimental techniques.
iii. It is also possible to calculate theoretical latticeenthalpies for ionic compounds.
1. This is done by assuming the ionic model,then summing the electrostatic attractive and
repulsive forces between the ions in the
crystal lattice.
f. Enthalpy change of solutioni. The enthalpy change when one mole of the
substance is dissolved in water to form a diluteaqueous solution
1. Uses lattice enthalpy2. Hsol= Lattice enthalpy+ (hydration
enthalpies of the component ions)
g. Enthalpy change of hydration
Elements in
standard statesSolid compound
Enthalpy of formation
Elements in
gaseous ions
Gaseous
anions and
metal atoms
Gaseous anions and
cations
Enthalpies of
atomization
Electron
affinities
Ionization
energies
Lattice enthalpies
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i. The enthalpy change (always exothermic) when onemole of the gaseous ion is added to excess water to
form a dilute solution
vi. Entropya. A measure of the degree of disorder or randomness in a
system
i. Some states are inherently more probable thanothers
ii. In general, the less order there is in a state, thegreater the probability of the state and the greater its
entropy
iii. There is an increase in entropy going from solid toliquid, and from going to liquid to gas.
b. A solid, with a regular arrangement of particles, has a lowentropy
i. When it melts, the particles can move more easilyii. The system has become more disordered, and its
entropy increasesc. Gas molecules move fast and independently of one anothersince inter-particle forces are negligible
i. Gases have high entropies.d. Entropy decreases as gas pressure increases
i. Higher pressure reduces the volume for gasparticles to move in, resulting in less disorder
Solid
compound
Aqueoussolution
Gaseous ions
Lattice
enthalpySum of hydration
enthalpies
Enthalpy of
solution
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e. When a solid or liquid dissolves in a solvent, the entropy ofthe substance generally increases
i. When a gas does the same, the entropy decreasesii. Hard solids (diamond) have less entropy than softer
solids (lead)
1. This is because there is less movement indiamond (thus making it hard) than there is
in lead
f. The entropy of a perfectly ordered crystal at absolute zerois zero
i. There is no randomness in the crystal form aka nomovement from the particles
ii. Unlike enthalpy, absolute values of the entropy of asubstance in a particle state can be measured
relative to thisiii. Real substances always have a higher randomness
than the crystal; therefore, all entropy is positive
1. And measured in J K-1 mol-1g. Changes in entropy
i. Entropy change will likely be positive if: there is adecrease in order through a decrease in the number
of moles of solid, or an increase in the number of
moles of gas (meaning a reactant, which is solid, is
converted into a gas, or a gas reacts and creates
more gases); an increase in temperature and anincrease in the number of particles also increases
entropy.
ii. Entropy change will likely be negative if: thenumber of moles of solid increases, or a gas turns
into a liquid or a solid.
h. Entropy change: S= B S(products) - B S(reactants) vii. Spontaneity
1. Any change may occur spontaneously if the final state is moreprobable than the initial state
a. If, as a result, the system is more stable in the final state,the final entropy of the universe is greater than the initial
entropy of the universe.
2. S measures the change in the entropy of the system.a. The major effect of chemical changes on the entropy of the
surroundings results from the gain and loss of heat energy.
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b. If chemical potential energy is converted to heat energy,which is then transferred to the universe, then this results in
an increase in the entropy of the surroundings
i. The opposite is true for an endothermic change (theabove lost energy)
c. uni verse surroundi ngs system HS =S +S =T+S
i.
H
T+Sis the magnitude of the entropy change
d. Processes will be spontaneous if: the final state has a lowerenthalpy than the initial state and the final state is more
disordered than the initial state
3. Gibbs Free energya. G H T S
i. This equation tells us if a reaction is spontaneous ornot.
ii. To be spontaneous, G must be negativeb. The Gibbs free energy for a change is equal to the amount
of energy from that system that is available to do useful
work
i. Ergo, for any system in equilibrium, Gibbs freeenergy must be exactly zero.
ii. If G is zero, this means that the stoichiometricamounts of both the reactants and the products are
all mixed together, meaning that no further change
will occur.
iii. If G is negative, it was produce a reaction whichwill increase the amount of products and decrease
the amount of reactants until equilibrium is reached
iv. If G is slightly positive, a reaction favoring thereactants will occur; if G is very positive, the
reaction will be very much non-spontaneous and
will not occur.
v. The values of G can be calculated at anytemperature, as long as one has data about the
reactants and the products, specifically data about
enthalpy and entropy.
c. Gibbs free energy of formationi. G under standard conditions (298 K and 101.3
kPa) can be calculated using the standard Gibbs free
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energy of formation data in the same way standard
enthalpy of formation data is used.
ii. f f fG = G (products) - G (reactants)