June 2011 Final Review Accelerated Chemistry 1
June 2011 Final
A
CCELERATED CHEMISTRY
Note: Much of this was covered during the first semester, and can be found in more detail in my review sheet for the January final. Some topics will be less detailed to save time and space.
Matter
PHASES
Matter is found in three phases: Solids
Liquids Gases
PURE SUBSTANCES VS.MIXTURES
A pure substance is a molecule (or a group of molecules) that have definite ratios of atoms. For example, water (H2O) is a pure substance, because it will always have 2 atoms of hydrogen for every one atom of oxygen. Bear
in mind that drinking water is a mixture.
A mixture does not necessarily have things in set proportions. There are two types of mixtures:
A homogenous mixture has uniform texture and color throughout. You cannot tell the different parts of the mixture
A heterogeneous mixture is a mixture where the different parts can be easily discerned
Physical and Chemical Properties and Changes
A physical property is a property that is apparent without changing the substance. An example of a physical property is the boiling or freezing point of a substance. Similarly, physical changes are changes that take place that do not change the substance on a chemical level, but instead change the state (solid to liquid, for example) of the matter
A chemical property is a property that can be seen when the substance reacts chemically with another, such as pH or pOH level and reactivity with water. A chemical change is a change that is not reversible and changes the substance on a chemical level.
Methods for Separating Mixtures
There are three ways to separate mixtures:
1) DISTILLATION: You distil, or boil off, all of the water in a mixture, leaving everything that has a boiling
June 2011 Final Review Accelerated Chemistry 2 2) FILTRATION: You can filter out the different components of the mixtures using a filter, such as a
semi-permeable membrane.
3) DECANTATION: A process by which all of the layers of a mixture are poured off to leave just the precipitate or certain parts of the mixture. For example, let’s say you have a mixture of water and oil. Oil, as you know, won’t dissolve in the water, and will remain in its own layer on top of the water because it is less dense. Thus, decanting the mixture would be pouring off either the oil or the water to leave the other component.
Double Displacement
Reactions and Precipitates
A double displacement reaction is a unique reaction where to substances (usually liquids or aqueous solutions) chemically react with each other to form two new compounds, including one precipitate, or a solid created as a result of two liquids. A double displacement reaction will only occur if there is a precipitate; otherwise, there was no reaction. Let’s look at an example. Let’s see what happens when we combine lead (II) nitrate (Pb(NO3)2) with potassium iodide (KI):
Pb(NO3)2 (aq) + KI (aq)
In a double displacement reaction, the anions (with negative charges; here, the nitrate and iodine molecules) switch to form two new compounds:
PbI2 + KNO3
When we look at the solubility chart1, we see that iodine will always be soluble, meaning that it will always dissolve, unless it is part of a molecule with Ag+, Hg2+, or Pb2+, like it is here. Thus, we can write the following
balanced equation:
Pb(NO3)2 (aq) + 2KI (aq) → PbI2(s) + 2KNO3(aq)
In this equation, PbI2 is a solid because the molecule is insoluble. KNO3 is still in its aqueous form because K+ is
always soluble.
Solutions
A solution is a homogenous solution (see above) which has both a solute and a solvent, such as mixing salt in water. A solute is the substance that is getting dissolved, going back to our example, the salt would be the
1
June 2011 Final Review Accelerated Chemistry 3 solute, because it is being dissolved in the water. The solvent is the substance that the solute is being dissolved into, in our example, the water would be the solvent.
DIFFERENT TYPES OF SOLUTIONS
There are three different types of solutions:
a) A saturated solution is a solution where there is the maximum amount of solute dissolved into the solvent. For example, a saturated solution of water and sodium nitrate (NaNO3) at 10:C would have 100g of water and approximately 80g
of NaNO3. (See below for how to read a solubility curve).
b) An unsaturated solution is a solution where yet more solute can be added to the solvent. Going back to the example above, an example of an unsaturated solution would have only, say, 70g of NaNO3.
c) Finally, a supersaturated solution holds more solute than possible, for example, 90g of NaNO3, and adding any more would
result in the creation of a precipitate.
READING A SOLUBILITY CURVE
A solubility curve is a visual representation of how much solute can be added to a solvent at a given temperature. For example, only 60 grams of KNO3 can be added to 100 grams of water at
40:C and be a saturated solution, but 80 grams of KNO3 would
make a saturated solution at 50:C2. As you raise the temperature, the solubility of a solid goes up — more solid can be dissolved in the same amount of water at a higher temperature. The same is true for gaseous solutes. If, say, you have less or more than 100g of water, assuming the temperature remains constant, you can set up a proportion to figure out the solubility of the solute.
DISSOCIATION
HCl→H+ + Cl
-Weaker acids (which will be given to us on the final) will not completely ionize in water, whereas stronger acids will.
BOILING POINT ELEVATION/FREEZING POINT DEPRESSION
By creating a solution of a salt (which is any ionic compound) and water, we can raise the boiling point and lower the freezing point, so a water/salt solution will not freeze as quickly, which is why salt is added to ice to make it melt. To find the freezing point depression, use the following equation:
In this equation:
= the change in freezing point
2
June 2011 Final Review Accelerated Chemistry 4 Kf= a constant equal to 1.86:C
m = molality of the solute3
i = number of ions created by the solute when it’s dissolved
When calculating the boiling point elevation, use the following equation:
Here, everything is exactly the same, except that the constant “Kb” is equal to .51:C
Percent Composition by Mass
Percent composition by mass is a way of measuring how much of a certain element is present in a molecule. It can be calculated as follows:
This can also be used for finding the percent concentration of how much solute is in a solution in terms of percentages.
Basic Structure of the Atom
Each atom has a nucleus, depicted in blue in the diagram on the right, and electrons, which can be found in orbitals and are represented by yellow dots circling around the nucleus. Each electron has an atomic mass of 1/1840 amu (atomic mass unit), and we regard it as weightless Each nucleus has two other types of particles in it:
Protons, which are positively-charged particles within the nucleus. Each proton has an atomic mass of 1 amu.
Neutrons, which have no charge. Each neutron has an atomic mass of 1 amu.
ISOTOPES
Each element on the periodic table has an atomic number, which is the number of protons in it. Isotopes, however, are versions of an element that have different numbers of neutrons, raising their atomic mass. Isotopes are noted by referencing their atomic masses (through which we can discern how many extra neutrons are in the isotope. For example, Carbon-14, a common isotope used for finding out how old things are, can be written as follows:
14C
3
Molality is equal to the moles of solute per kilogram of solution.
June 2011 Final Review Accelerated Chemistry 5 Carbon-14
Some people also include the atomic number of the element, although it is not necessary. Since the atomic number of carbon is 6, we know that in carbon-14, there are 8 neutrons.
ELECTRON CONFIGURATIONS AND QUANTUM MECHANICS
Due to the very in-depth explanations in my review sheet for the January final, this part will be abridged. For a more in-depth explanation, see my other review sheet.
There are four types of orbitals:
An S orbital, which can only hold 2 electrons
A P orbital can hold a total of 6 electrons (3 sets of 2 electrons). A D orbital, which can hold 10 electrons (5 sets of 2 electrons). An F orbital, which is can hold 14 electrons (7 sets of 2 electrons).
The first energy level will only hold an S orbital, whereas the second energy level will hold the second energy level can hold both an S orbital and a P orbital, for a total of eight electrons. The third energy level will hold an S orbital, a P orbital, and a D orbital, and so on and so forth. As you get higher and higher, some energy levels actually overlap. Use this diagram to write out electron configurations:
QUANTUM NUMBERS
Each electron in an atom has its own specific number so that it can be identified easily, called a quantum number. Each quantum number itself is made up of three integers and one fraction, and can be determined as follows:
1. The first number tells you what energy level theelectron is in: a. First energy level = 1
b. Second energy level = 2, and so on.
2. The second number tells you what type of orbital the electron is in:
a. S orbital = 0 b. P orbital = 1 c. D orbital = 2 d. F orbital = 3
3. Since each type of orbital is divided into groups of two electrons, the third number tells you which group of two the electron is in:
June 2011 Final Review Accelerated Chemistry 6 a. Since an S orbital only has one set of two electrons, an electron in the S orbital would be 0 b. In a P orbital, which has three sets of two electrons:
i. First group = 0 ii. Second group = 1 iii. Third group = 2
c. In a D orbital, which has five sets of two electrons: i. First group = -2
ii. Second group = -1 iii. Third group = 0 iv. Fourth group = 1
v. Fifth group = 2
d. In an F orbital, which has seven sets of two electrons: i. First group = -3
ii. Second group = -2 iii. Third group = -1 iv. Fourth group 0
v. Fifth group = 1 vi. Sixth group = 2 vii. Seventh group = 3
4. The fourth and final number is always a fraction. In each set of two electrons, one electron is pointing up, and the other is pointing down. This number tells us if the electron is pointing up or down:
a. If the electron is facing up = +½ b. If the electron is facing down = -½
See below for how to determine the quantum number of an electron when given an orbital energy diagram
ORBITAL ENERGY DIAGRAMS
Orbital energy diagrams are visual representations of how the electrons are shaped in an atom. Take a look at the orbital energy diagram for a standard atom of cobalt:
June 2011 Final Review Accelerated Chemistry 7 o Since it is in the second energy level, the first number is 2
o Since it’s in a P orbital, the second number is 1
o Since it’s in the third group of the P orbital, the second number is 2 o Since the electron is facing down, the final number is –½
Thus, the quantum number for the circled electron is: 2,1,2,-½ LEWIS STRUCTURES
Lewis Structures are diagrams that show the configuration if valence electrons (which are used for bonding with other atoms to form compounds) around an atom. Here are some examples of Lewis Structures:
Zoning in on the nitrogen diagram, we can see that you first add electrons per orbital, so the 2 electrons on top belong to the S orbital, and the three around the sides belong the following P orbital. Carbon is the exception because it has a special SP orbital, which has a total of four electrons.
Additionally, Lewis Structures can also be used to show how individual atoms in a molecule fit together:
Writing Chemical Formulas
CHEMICAL FORMULAS FOR IONIC COMPOUNDS
The first coloumn, called the alakali metals, will always have a charge of +1. These elements are hydrogen (H), lithium (Li), sodium (Na), potassium(K), rubidium (Rb), cesium (Cs), and francium (Fr). The second coloumn, called the alkaline eart metals, will always have a charge of +2. These elements include beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The second couloumn from the right will have a charge of -1, the third from the left will have a charge of -2. The middle section, however,
hydrogen nitrogen chlorine carbon
June 2011 Final Review Accelerated Chemistry 8 changes depending on the element and what it is boding with. There are some elements that do stay the same, however. Here are a few of the common ones:
1. Silver (Ag) will always have a charge of +1
2. Copper is usually found with either a charge of +1 or +2.
When naming ionic compounds that do not include transition metals (even ones that do not change their charge), you write the cation (the positive ion)’s name, and then the root of the name of the anion (the negative ion) and add __+ide to the end of it. For example:
NaCl
Sodium = Na Chlorine = Cl
Sodium chloride
When using a polyatomic ion, you simply just use the name of the polyatomic ion. For example, NaNO3 is
sodium (Na+) bonded with NO3-, which is nitrate. Thus, NaNO3 is sodium nitrate.
When using transition metals, things become more complicated. Because transition metals change their charges depending on which anion they are bonded to, we have to express that in the name. Thus, we must look at the anion to see what charge the cation is, since it must always add up to 0, or neutral, if it is not an ion. For example, let’s look at CuCl2. Since copper (Cu) is a transition metal, it can either have a charge of +2 or
+4. Looking at chlorine (Cl), we see that it has a charge of -1, and there are two molecules of chlorine, for a total charge of -2. Since we need to add up to zero, we can conclude that the charge of copper in CuCl2 is +2.
We write the name of the compound like this:
copper (II) chloride
Remember, you only include the charge when you are dealing with a transition metal. Metals that are not a transition metal are not written with the charge.
CHEMICAL FORMULAS FOR MOLECULAR COMPOUNDS
Naming molecular compounds is different from naming ionic compounds. You use the name of the first ion, then a prefix to describe how much of the second element you have, the root of the second element, and then add the suffix __ide.
Mono- 1 Di- 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6
Root of anion + ide Charge of
cation Name of
June 2011 Final Review Accelerated Chemistry 9 Septa- 7
Octa- 8 Nona- 9 Deca- 10
For example, CO2 is called carbon dioxide, since there are two carbons (when there is only one of the first
element, you do not need a prefix. Compounds like H2O2, or rubbing alcohol, has the name dihydrogen
dioxide.) Naming a molecular compound is much easier than naming an ionic compound, because you are given all of the information that you need.
Remember: Unless otherwise indicated, you can assume that a compound is neutral, or equal to 0. If the compound in question is an ion, then you will be told.
The Periodic Table
METALS VS.NONMETALS
Metals are generally:
Conductors of electricity Malleable (can be flattened) Ductile (can be made into a wire) Have higher boiling points Very high melting points
Give up electrons to form compounds with nonmetals to form salts
All metals are to the left of the step ladder on the periodic table.
Nonmetals are generally:
Not conductors of electricity Not malleable
Not ductile
Very brittle and shatter easily Have very low boiling points Have very, very low freezing points
All nonmetals are to the right of the step ladder on the periodic table. PERIODIC TABLE TRENDS
Atomic radius is the distance from the nucleus to the outermost orbital. As you go across the period table from left to right, atomic radii become smaller, and as you go down the period table, atomic radius increases. This is because the smaller the atomic radius, the more nuclear force, or power the nucleus has, increases
June 2011 Final Review Accelerated Chemistry 10
The Mole
If a dozen is twelve of something, then a mole is 6.02x1023 of something. A mole is generally used to identify the number of molecules. Additionally, molar mass is used to express the number of moles per gram of an atom, and is equal to the atomic mass of each of the elements in a compound. For example, the molar mass of water (H2O) is twice the atomic mass of hydrogen added to the atomic mass of oxygen. Since the atomic
mass of one molecule of hydrogen is 1.008amu (atomic mass unites), two molecules of hydrogen is 2.016amu. The atomic mass of oxygen is 15.999amu. Thus, if we were to add the atomic masses of the two molecules of hydrogen and the one molecule of oxygen, we would get a molar mass of 18.015g/mol (read: grams per mole). We can use the molar mass to convert between moles and grams. For example, if I had 3 moles of water and I wanted to find out how many grams that would be, I would set up the following equation:
(
)
In the above equation, all I did was essentially multiply by the molar mass. If I wanted to convert from grams to moles, I would divide by the molar mass:
(
)
We can also use moles to find the quantity of atoms in a given amount of moles. For that, we use Avogadro’s Number, or 6.02x1023. For example, if I had 5 moles of something — for this type of example, it does not matter what element or compound it is — and I wanted to find out how many molecules are in 5 moles, I would set up the following equation:
(
)
All I am doing is essentially multiplying by Avogadro’s Number. If I wanted to convert from molecules to moles, I would divide by Avogadro’s Number:
(
)
Empirical and Molecular Formulas
A molecular formula is the formula that tells us the actual not-reduced ratio of atoms per element in a compound. For example, rubbing alcohol has a molecular formula of H2O2. However, an empirical formula tells
us the reduced ratio of atoms per element in a compound. Therefore, referring back to the above example, since both numbers can be reduced, the empirical formula for rubbing alcohol is HO.
Let’s say were given the following information about a compound. The percent composition is as follows: o 40% carbon
o 6.67% hydrogen o 53.33% oxygen
We can now assume that the % sign really means “out of 100 grams”. Thus, we can say that: o 40% carbon = 40 grams of carbon
o 6.67% hydrogen = 6.67 grams of hydrogen o 53.33% oxygen = 53.33 grams of oxygen
June 2011 Final Review Accelerated Chemistry 11 o (40 grams of carbon / 12.01 grams) = 3.33 moles carbon
o (6.67 grams of hydrogen / 1.008 grams) = 6.60 moles of hydrogen o (53.33 grams of oxygen / 15.99 grams) = 3.33 moles of oxygen
Because the subscripts in an empirical or molecular formula can also refer to the number of moles of each element, we can write the following formula:
C3.33H6.60O3.33
However, it is impossible to have fractions of an element in a compound, thus, we divide each of the numbers by the smallest number in the formula:
CH2O
Because the number 6.60/3.33 is so close to a whole number (approximately 2/10 away from the number two), we may round. This is the empirical formula. If we are given no further information, then we stop here. If we are given the molar mass of the compound, however, we can continue to find the molecular formula. For example, let’s say that we are given that the molar mass of 180 g/mol for the molecular formula. We can find the molar mass of the empirical formula — which in this case is 30 g/mol — and divide the two. We would get the number 6, and we would then multiply each of the subscripts by six and get:
C6H12O6
or glucose.
Stoichiometry
Stoichiometry is the term used for determining how many grams I can make of a product, given the grams of reactants, or vice versa. For example, let’s refer back to the equation for the formation of water:
2H2 + O2→ 2H2O
If we were given, for example, 10 grams of oxygen gas, we must now find how many grams of water can be produced. However, we cannot use grams alone to calculate this — we must first convert to moles:
(
)
We can now set up the following proportion:
Using this proportion, we can now solve for x:
( )
We can now see that, since x is equal to 1.25, 1.25 moles of water is produced for every .625 moles of oxygen gas. If the question is asking for moles, then this is our answer. However, the question is asking for the number of grams produced, not moles. Thus, we must convert 1.25 moles of H2O to grams:
(
)
Thus, we can now say that, for every 10 grams of oxygen gas, 22.5 grams of water is produced. If we were to just set up the proportion using grams, we would have been of by 2.5 grams.
June 2011 Final Review Accelerated Chemistry 12 Not all chemical equations require just one of each of the reactant to form the product. Observe the following reaction that is used to make water:
H2 + O2→ H2O
There is an evident problem with this equation, however. The Law of Conservation of Mass states that no matter can be created or destroyed. However, we see that there are two atoms of hydrogen and two atoms of oxygen in the reactants (both of these elements are diatomic), but there is only one atom of oxygen in the product. That last atom of oxygen cannot exist as oxygen gas, because oxygen is diatomic. Thus, perhaps the equation yields two molecules of water:
H2 + O2→ 2H2O
This equation also does not make sense. Although now we both atoms of oxygen are reacting to form water, there are now four atoms of hydrogen in the product, and only two in the reactant. Thus, perhaps this reacting to molecules of hydrogen gas (H2) to react:
2H2 + O2→ 2H2O
At last, we now have an equation that has the same number of reactants as it does products. There are four molecules of hydrogen and two molecules of oxygen in the reactants, there are four molecules of hydrogen and two molecules of oxygen in the product. This equation is now balanced.
Oxidation-Reduction Reactions
If in a double displacement reaction, two anions switch, then in an Oxidation-Reduction Reaction, a compound and either a neutral atom or an ion switch to form a new compound and a neutral atom. In an oxidation-reduction reaction, the oxidized part loses electrons, and the reduced part gains electrons. Observe the following reaction:
3Ca (s) + 3CrCl3(aq) → 3CaCl2 (aq) + Cr (s)
In the reaction above, calcium (Ca) is oxidized, since it must give electrons to chlorine to form CaCl2.
Meanwhile, chromium (Cr) gains electrons in the process and become neutral — it reduces. Thus, because calcium is being oxidized, it is the reducing agent, and chromium is the oxidizing agent.
Heating and Cooling Curves
A
2.1j/g:C
C
4.2j/g:C
B
6 kJ/mol
E
1.8j/g:C
D
4.1kJ/mol
June 2011 Final Review Accelerated Chemistry 13 The above diagram is called a heating curve, and is used for showing how substances — in this case, water — change from solid to liquid to gas. A cooling curve is the opposite, and goes in the opposite direction.
HOW TO READ A HEATING CURVE
There are two types of energy: potential energy and kinetic energy. Potential energy is stored energy. Kinetic energy is the energy made my motion.
“A” is an example of kinetic energy. In it, the ice is slowly being heated up and nearing its melting point. Here, potential energy remains constant.
“B” is potential energy. As the ice begins to change from solid to liquid, the temperature remains constant, and kinetic energy remains constant.
“C” is another example of where kinetic energy is released. After completely melting into liquid water, the temperature begins to rise again, and potential energy remains constant. This will continue until the temperature of the water reaches 100:C, the boiling point.
“D” is the change from liquid to gas. As the water reaches its boiling point, the temperature levels out and the water begins to boil off. Here, potential energy is released, and kinetic energy remains constant.
“E” is the continuation after the water becomes water vapor. After turning completely into gas, the temperature again begins to shoot up. This will continue on forever.
Specific heat is the amount of energy required to raise 1 gram of a substance by 1:C, and is measured in j/g:C. On the heating chart above, you can find the specific heats for ice, water, and water vapor.
The molar heat of fusion and molar heat of vaporization tell you how much energy is required to turn one mole of a substance from either solid to liquid or liquid to gas respectively, and is measured in kJ/mol.
CALCULATING THE ENERGY RELEASED OR ABSORBED WHEN HEATING OR COOLING RESPECTIVELY
We can use the following equation for finding out how much energy is released or absorbed when a substance, in this case, water, is cooled or heated respectively:
( ) ( ) ( )
When a substance is being heated up, energy is released, resulting in an exothermic reaction. When a substance is being cooled down, energy is absorbed, resulting in an endothermic reaction. If the enthalpy change of a reaction is negative, then the reaction is exothermic.
CALCULATING ENERGY ABSORBED OR RELEASED DURING A PHASE CHANGE
When calculating how much energy is absorbed or released during a phase change, you don’t use the specific heat, but the molar heat of fusion (when going from solid → liquid or vice versa) or the molar heat of vaporization (when going from liquid → gas or vice versa). Thus, let’s say that you’re calculating how much energy is released when 100g of water is changed from ice to liquid. First, we have to change 100 grams to moles (see above for more information about that):
(
June 2011 Final Review Accelerated Chemistry 14 Then, we can set up the following proportion:
And then solve for x, which is 30.36 kilojoules of energy released.
Remember that when calculating energy released during a phase change you are working with kilojoules, while you are working with joules when not. In problems where you are required to use both equations, make sure that you are working with the same units when calculating the total amount of energy released.
Thermodynamics and Calorimetry
As noted above, if the enthalpy changes of a reaction is negative, then the reaction is exothermic; if the enthalpy change is positive, than the reaction is endothermic.
The heat of combustion is the heat released when a substance is completely burned, and is measured in kilojoules per gram. The molar heat of combustion is the same concept, but is measured in kilojoules per mole, not gram. For example, let’s say we are asked to calculate the molar heat of combustion of wax4 (in kJ/mol) when 10 grams of wax is burned and 10,000 joules of energy are released. First, we must convert joules into kilojoules:
(
)
Now that we have the number of kilojoules, we can set up the following proportion to find the heat of combustion:
We can now see that the heat of combustion is 1 kJ/gram.
If this is all we were asked to calculate, then we stop here. However, we were also asked to find the molar heat of combustion. First, we must convert 10 grams of wax, which has a chemical formula of C25H52 and has a molar
mass of 352g/mol:
(
)
Now, we can set up the same proportion as before, using moles as opposed to grams:
And we can solve for x, which is equal to 357.143 kJ/mol.
Hydrocarbons
4
June 2011 Final Review Accelerated Chemistry 15 A hydrocarbon is a molecular compound that is made up of only carbon and hydrogen atoms.
THE DIFFERENT TYPES OF HYDROCARBONS
There are three different kinds of hydrocarbons:
An alkane has only single bonds and end with the suffix –ane. All alkanes will have a molecular formula of CnH2n+2, where n is equal to a positive whole number.
An alkene is a hydrocarbon with both single and double bounds; all alkenes end with the suffix –ene. All alkenes will a molecular formula of CnHn, where n is equal to a positive whole number.
An alkyne is a hydrocarbon that has a mixture of triple, double bonds; all alkynes end with the suffix; all alkynes end with the suffix –yne. They will have a chemical formula of CnH2n-2, where n is equal to a
positive whole number.
ISOMERS
Look at the diagram on the left. Although they look like completely different hydrocarbons, you’ll see that they all have 5 carbon atoms and 12 hydrogen atoms — they are all molecules of pentane. That is because they are all isomers of the same molecule. Due to the fact they they are all different isomers of the same molecule, their flexibility and their shape differ. Being isomers, these all have different boiling points, too. The reason for is that hydrocarbons tend to tangle up because of intermolecular forces called hydrogen dispersion forces that cause attractions between the hydrocarbons. Since isomer (a) is the longest, it is most likely to become tangled and would be harder to untangle, giving it the highest boiling point, since it takes the most energy to move. Isomer (c) would have the lowest boiling point, since it is the most compact and requires the least amount of energy to move or untangle.
This is also the reason that larger hydrocarbons also have higher boiling points. Intermolecular forces, such as hydrogen bonds, which cause water molecules to be attracted to each other, cause molecules to have higher boiling points since more energy is required to separate them. This is why water has a boiling point of 100:C, which is relatively high for such a small molecule.
Intra- and Intermolecular Forces
An intramolecular force is a force that holds two atoms together in a molecule. Here is a list of intramolecular forces from weakest to strongest:
June 2011 Final Review Accelerated Chemistry 16 א. Ionic bonds, which are bonds that hold a salt together
ב. Polar covalent bonds, which bond some covalent molecules together, and whose valence electrons are attracted more to a particular atom or atoms in the molecule
ג. Covalent bonds, which bond the atoms in a molecular compound together and whose electrons are equally shared between the atoms.
Covalent bonds are the strongest because they have an equal sharing of the electrons, whereas ionic bonds are the weakest because they involve giving one electron away.
Intermolecular forces are forces of attraction between two molecules. Here is a list of intermolecular attractions from weakest to strongest:
א. No attraction
ב. Dipole-dipole forces, which are forces between any two generic molecular compounds. ג. Hydrogen bonds, which are bonds between water molecules, H2N, and H2F molecules.
ד. Ion-ion bonds, which are bonds between different ions between the cations and anions in different ionic compounds, and whose strength are equivalent to that of ionic compounds
ELECTRONEGATIVITY
Electronegativity is a measure of how tightly the electrons are held together by an atom in a bond, and ranges from 0-4.0 (non-electronegative → very electronegative). The electronegativity of a molecule is calculated by subtracting the electronegativity of the anion from that of the cation:
If the electronegativity is between 0.0 and 0.4 → covalent bond If the electronegativity is between 0.4 and 2.0 → polar covalent bond If the electronegativity is between 2.0 and 4.0 → ionic bond
MOLECULAR GEOMETRY
Molecular Geometry is the study of how the atoms within a molecule fit into each other and the different shapes of the molecules. The idea behind molecular geometry is that the electrons want to be as separate from the positive parts of the molecule as possible, and the molecule will shape itself to accommodate this. Take a look at the following diagrams:
As you can see here, the electrons are putting as much distance between themselves and the positive parts of the molecule. This is why the following shapes result. This is also why all of these molecules are polar molecules.
Pyramidal geometry
June 2011 Final Review Accelerated Chemistry 17
Gases
The different gas laws show the various relationships between pressure, temperature, and volume:
Boyles’ Law says that pressure and volume are inversely proportional: when one goes up, the other goes down and vice versa. Temperature remains constant.
Charles’ Law says that temperature and volume are directly proportional: when one goes up, so does the other, and vice versa. Pressure remains constant.
Since, here, temperature values can be negative in this equation, you must convert all temperature values into Kelvin before using this equation. See my notes about Absolute Zero for more notes about converting between Celsius and Kelvin.
Gay-Lussack’s Law says that temperature and pressure are also directly proportional when volume remains constant. Here, temperature must also be measured in Kelvin.
The Ideal Gas Law (nicknamed “Pivnert”) combines all of these laws together and allows you to calculate the pressure, volume, moles of gas, or temperature given all of the others:
Like in the above equations, temperature must be measured in Kelvin (or converted to it), and “R” is a constant equal to .0821atmL/Kmol
FINDING THE DENSITY OF A GAS AT STP
STP, or standard temperature and pressure, is equal to 22.4 moles per liter. When given the molar mass and the number of moles of a specific gas and how many liters of gas there are, we can also find the density of the gas in question. For example, let’s say that we have 20 moles of hydrogen gas (H2(g)) that has a volume of
approximately 5 liters. First, since density is measured in grams per liter (g/L), we must convert the moles to grams:
(
)
Then, we can calculate the density:
ABSOLUTE ZERO
June 2011 Final Review Accelerated Chemistry 18 example, 0:C is equal 273K (no need for the :). You must always use the Kelvin system when using temperature in the gas laws. Otherwise, your answer will be WRONG. This is key.
SPECIAL PROPERTIES OF GASES
What is perhaps most important about gases is that they spread to the form of their container, and thus we cannot talk about mass and gases, since their mass can change using any of the factors above. Additionally, pressure has become a variable here (whereas it is not with liquids and solids), since changing the volume or temperature can and does affect the pressure exerted on the gas.
Acids & Bases
Acids are things that give off protons, or hydrogen ions when dissolved in water or otherwise ionized. For example, HCl, when dissolved in water, breaks down into H+ and Cl- ions. Because hydrogen only has 1 proton and one electron, when it is broken down, it’s like a lone proton is roaming free in the world. All bases have a pH between 1.0-6.99999… and a pOH between 7.000…1 and 14.0.
Bases are the opposite of acids. Instead of giving off lone protons, they give of hydroxide, or OH- ions. For example, when sodium hydroxide is ionized, it breaks down into Na+ and OH- ions. Ammonia is the only base that, as a solid, does not have any hydroxide ions. All bases have a pH of 7.000…1-14.0 and a pOH between 1.0-6.99999….
When an acid and a base react chemically with each other, they form water and a salt5 of some sort:
HCl + NaOH → H2O + NaCl
Water (H2O) has a pH of exactly seven, and is considered to be amphoteric, or both an acid and a base, since it
has a hydrogen ion and a hydroxide ion.
CALCULATING PH OR POH
pH is literally the power of hydronium, or the concentration of hydrogen ions in a solution to do this, we use the following formula:
-log[H3O+]
Concentration is here calculated in molarity, or moles of solute per liter of solution (assuming, unless otherwise told, that the solute has a negligible to the volume of the solution). Let’s say that you have .1 grams of HCl in a 5 liter solution of water and want to find the pH. First, you’d need to find how many moles, which we can calculate to .0027 moles, and then calculate the molarity:
.0027 moles/5 liters = 5.5x10-4
Then, we use the logarithm function on our calculators and input the following information: -log(5.5E-4
You would then calculate the pH to be 1.46, which means that this acid is very acidic.
STRONG VS.WEAK ACIDS
5
June 2011 Final Review Accelerated Chemistry 19 A weak acid has a very low Ka constant, which means that it has very, very few hydrogen ions. All Ka constants
will be given to us on the test. Additionally, this also means that the equation given with the Ka constant is
always going to be at equilibrium. Weak acids will never ionize fully. Strong acids are the complete opposite of weak acids: they have no Ka constant, completely ionize, and will never be at equilibrium.
MOLARITY
The following formula is used for calculating molarity:
TITRATIONS
A titration is the process by which one can calculate the molarity of an acid or a base by neutralizing it with its opposite. We can then use the following equation for calculating the missing date (give, of course, the other three):
Equilibrium
Equilibrium is the state where there everything is at balance and everyone is in a state of peace. A sign of equilibrium is that there is both a forward and backward reaction. Look at the two glasses of water below:
H2O(l) H2O (g)
Both will be able to achieve equilibrium, even though there is more water in glass “A” than glass “B”. In fact, if here were to be added to these glasses, which we can call equilibrium systems, the water vapor and liquid water would work to remain at equilibrium. This is called La Châtlier’s Principle. La Châtlier’s Principle states
An example of two different equilibrium
June 2011 Final Review Accelerated Chemistry 20 that when the equilibrium system is troubled, it will work to stay at equilibrium. For example, look at the following equilibrium expression:
2SO2 (g) + O2 (g) 2SO3 (g) + heat
If we were to increase the pressure, then we would favor the forward reaction, because an increased pressure would favor an environment with fewer molecules (there are 3 moles of molecules on the left side, but only 2 on the right)
If we were to increase the concentration of reactants, than we would favor the forward reaction because there would be more molecules to react.
If we were to lower the temperature, we would favor the backward reaction, because the left side of the equation favors an environment with a lower temperature
If were to decrease the concentration of SO3, then we would favor the forward reaction because the
reactants want to stay at equilibrium with the products. This would also increase the heat.
USING EQUILIBRIUM CONSTANTS
Equilibrium concentration = molecular solubility
Let’s say we are asked to molecular solubility of copper (II) iodide, or CuI when the Keq is equal to 1.27x10-12.
We can set up the following equation:
[ ][ ]
Since the copper and the iodine are in a ratio of 1:1, we can assume that they are equal. Thus, we can set up the following equation:
√[ ][ ] √
[ ]
Thus, we can say that both have an equilibrium concentration of 1.13x10-16. NO UNITS NEEDED FOR EQUILIBRIUM CONSTANTS!!!!!
Another example with an ICE chart: Let’s say we’re given the following equilibrium expression:
I2 (g) + H2 (g) 2HI (g)
We are told that 3 moles of both iodine gas and hydrogen gas react, and are told that, at 490:C and in a 1 liter closed box, the system is at equilibrium and that the Keq is 4.59. Now, we must find the concentration of each
substance at equilibrium. We can set up the following chart:
Hydrogen gas Iodine gas Hydroiodic acid gas
Initial 3 moles 3 moles 0 moles
Change -x -x +x
Equilibrium 3-x 3-x +x
June 2011 Final Review Accelerated Chemistry 22
Electrochemistry
As the zinc is broken down and is oxidized (Zn → Zn2+ + 2e-), the electrons travel up the solid zinc and flow into the flask with the copper sulfate and ionized copper ions. The electrons bond with the sulfur ions, reducing them to neutral atoms (Cu2+ +2e- → Cu). Meanwhile, since copper would rather become stable than become part of a compound, excess SO4 flies across the salt bridge into the
flask where there is excess Zn2+ ions floating around. This will continue until all of there is no more left to ionize and reduce. Meanwhile, energy is released because of the reactions occurring.
This is a battery, also called a voltaic or galvanic cell. Since the zinc is being ionized, it is the anode, and since the copper is being reduced from its ionized form, it is the cathode. This battery can be written out as follows:
Zn|Zn2+||Cu2+|Cu
The cell potential of a battery is the amount of energy in volts released when the anode and cathode react. For example, let’s say that we have a lithium-flourine cell, Li|Li+||F2+|F:
2Li → 2Li+ +2e- +3.05V positive because chart shows reduction potential; Li is oxidize F2+ + 2e- → F +2.87 V positive because reduction potential is positive
5.92 V potential of battery
If something with a negative reduction potential was used as a cathode, then it would have e negative reduction potential; if something with a positive reduction potential was used as an anode, then it would have a negative oxidation potential.
Nuclear Chemistry
Radiation has two forms:
Ionizing radiation, which ionizes everything it comes in contact with. This radiation is harmful. Non-ionizing radiation increases the motion of molecules and is less dangerous.
There are three types of radiation:
Alpha radiation, symbolized by the Greek letter α. A specific isotope of helium, is omitted. This creates a transmutation of the radioactive material. All forms of radiation create transmutations.
June 2011 Final Review Accelerated Chemistry 23 Beta radiation, symbolized by the Greek letter β. A proton turns into a neutron and a high-energy electron is ejected from the atom, . Beta radiation will always yield an element with a higher atomic
number, but with the same atomic mass. This type of radiation occurs naturally, and is usually relatively harmless.
Gamma radiation, symbolized by the Greek letter γ. This is the emission of charge-less and mass-less rays.
Nuclear radiation is generally harder to control. The source of nuclear radiation is high-energy particles being ejected from the nucleus of an atom, hence the name. This usually happens naturally; particles do not need to be excited to be ejected from the nucleus. Nuclear radiation can occur artificially, however, as is the case of nuclear energy.
HALF-LIVES
Radiation is measured in half-lives, after one half-life, half of the radioactive material is gone, after the second half-life, half of the radiation is left. This results in the graph on the right. The amount of radioactive material is cut in half from the previous, and will never reach zero, resulting in the asymptotic graph shown left.
NUCLEAR FISSION
Nuclear fission is the process by which is split, causing the release of tremendous amounts of energy. This is how nuclear energy is produced, and how the atom bomb was created. This is how the atom is split:
A neutron is rammed into an atom of uranium, causing it to break into isotopes of barium and krypton. This is the part of the reaction that is referred to as “splitting the atom”. The resulting reaction creates another neutron, which can, in turn, react with more uranium
A graph of the progression of half-lives of radioactive substances
The process of splitting the atom
