chemistry notes chapters 1-5

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)

chemistry notes chapters 1-5

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