Chem 2202. Year End Review.

Atomic Number (Nuclear Charge) is equal to the number of protons in an atom. The mass number of an atom is equal to the sum of its protons and neutrons. For example, an atom of helium has two protons and two neutrons, so its mass number is four, for example, helium-4.

Isotopes: It is possible for atoms of the same element to have different numbers of neutrons, like chlorine-35 because its nucleus contains 17 protons and 18 neutrons and chlorine-37 (17 protons & 20 neutrons). The atomic mass of this element cannot be one value or the other - it has to be a combination of the two.

Ex: The relative abundance of chlorine atoms in nature is 75% chlorine-35 and 25% chlorine-37. If you do a straight average of the two atomic masses, you get 36 amu, but this doesn’t reflect the percent abundances of the two isotopes. Average atomic mass is the weighted average of the atomic masses of the isotopes of an element. It calculated by multiplying the percent abundance by the mass number for each isotope and then adding the products.

Ex

The Mole (Avogadro’s Number): Atoms are so small that it takes an incredibly huge number of them to make up a visible amount. An iron spike for example might contain more than 6.02 x 1023 atoms of iron and has a mass of 55.85 grams. To find the molar mass of a compound, list the number of atoms of each element, and multiply this number by the molar mass of the element. For example water’s molar mass = 18.02 g/mol. Ba(OH)2·8H2O = 315.51 g/mol.

Molar Volume (Measuring Gases): Another way of measuring amounts of substances is by volume. Amounts of gases like hydrogen, nitrogen and oxygen can be measured in litres if the samples are at the same temperature and pressure. Chemists compare gas volumes at 0°C and 101.3 kPa. These conditions are known as Standard Temperature and Pressure or STP: 1 mol of any gas at STP has a volume of 22.4 L.

Ex: Calculate the number of moles of neon in a 6.81 L sample at STP

Mass to Mole Conversions : The mass of a substance can be found using a balance. If you know the molar mass of the substance, then you can convert the mass to a mole amount. Ex. Calculate the mass of 2.40 mol of aluminum nitrate, Al(NO3)3.

Mole - Particle Conversions: to get Molecules Multipy.

Example : How many particles are in 0.25 mol of helium

Here is a handy dandy helper!!!!

Percentage Composition : A good way to get an understanding of what percentage composition is all about is to calculate it for a known substance first. Let's use AlBr3 as an example. The sum of the masses is the molar mass: one mole of AlBr3 has a mass of 266.68 g.

Finding an Empirical Formula: An empirical formula is the lowest whole number ratio of the atoms or ions in a compound. Example A prospector finds a rock that may contain rutile - an ore of titanium. A crystal from the sample is analyzed and found to be 59.94% titanium and 40.06% oxygen by mass. Determine the empirical formula of the crystalline substance.

Molecular formulas are multiples of the E.F.

Stoichiometry is making predictions about the amounts of chemicals involved in a chemical reaction. What is a Mole Ratio? A mole ratio is a mathematical expression that shows the relative amounts of two species involved in a chemical change. Stoichiometry problems involve predicting the mole amount of one substance given the mole amount of another substance in a chemical reaction. Example: Calculate the mass of oxygen gas that should react with 6.49 g of aluminum metal.

   4 Al (s)  +  3 O2 (g) ---- > 2 Al2O3 (s)                            

Gas Stoichiometry.  Example : The fuel for the space shuttle is hydrogen. It reacts with oxygen gas to produce water vapour. Calculate the volume of oxygen gas required to completely burn 250.0 L of hydrogen gas at STP conditions.

v = 250.0 L   v = ? L

Limiting Reagents: In a chemical reaction, the thing that "runs out" is called the limiting species and the unused or leftover species are called excess species.

Ex: Identify the excess and limiting species in the reaction between 5.0 mol of aluminum and 7.0 mol of bromine.

2 Al (s)  +  3 Br2 (l)  >  2 AlBr3 (s)

5.0 mol    7.0 mol              

Since the amount of Al required (4.7 mol) is less than the amount available to react (5.0 mol), aluminum is in excess. Since the amount of bromine required (7.5 mol) is greater than the amount available (7.0 mol), bromine is limiting.

Percent Yield: In chemistry, an expected amount of product in a chemical reaction is called the theoretical yield. A measured amount is called the actual yield. Percent yield is a ratio of the actual to the theoretical yield expressed as a percentage. Example : Solid iron(II) hydroxide is the product of a double replacement reaction between iron(II) chloride and sodium hydroxide. When this reaction was carried out in a laboratory, the following data were recorded:

1) mass of NaOH reacted = 25.0 g mass 2) mass of filter paper = 1.61 g 3) mass of the filter paper and the dry precipitate = 28.91 g.

Ex: Calculate the theoretical yield, actual yield, and percent yield of iron(II) hydroxide.

Solutions: A solution is a homogeneous mixture. A solution forms when two substances are mixed together so evenly that they appear to be a single phase. The components of a solution are the solute - the substance that dissolves, and the solvent - the substance in which the solute is dissolved. Any solid, liquid, or gas which is evenly distributed throughout another solid, liquid or gas is said to be in solution. There are nine possible combinations of these states of matter.

Solid in solide.g. brass

solid in liquide.g. sugar water

solid in gase.g. mothball in air

Liquid in solide.g. dental amalgam

liquid in liquide.g. ethanol in water

liquid in gase.g. water in air

gas in solide.g. hydrogen in palladium

gas in liquide.g. O2 in water

gas in gase.g. oxygen in nitrogen

Solubility and Miscibility: Have you ever noticed that certain substances like table salt dissolve in water easily while other like caulk do not? Different substances have different solubilities. Solubility refers to the maximum amount of a solute that can be dissolved in an amount of solvent under specific

temperature and pressure conditions. A substance that cannot be dissolved in another (or does so to a very limited extent) is said to be insoluble.

Dissociation and ionization are processes that do involve changes in the structure of the solute. Dissociation is the separation of cations (positive ions) and anions (negative ions) in an ionic solid.

Electrolytes and Non-Electrolytes: A solution that conducts electrical current is said to be electrolytic and the solute is called an electrolyte. The sodium chloride solution is an electrolytic solution. The solute in a solution that does not conduct electrical current is a non-electrolyte. Sucrose is a non-electrolyte. Generally, dissociated (soluble) ionic compounds are electrolytes whereas dissolved molecular compounds are non-electrolytes. The exceptions to this rule are the molecular acids (the weak ones).  Acids form electrically conductive solutions; therefore they are electrolytes. However, only a few of them are strong electrolytes (the 6 strong ones). We fully ionize only the 6 strong acids into their ions

Dissocation Equations: Break soluble (see chart) ionics into their ions. Ex: magnesium chloride

Concentration: Generally, concentrated means a high amount of solute relative to the amount of solvent and dilute means a low amount of solute relative to the amount of solvent.  We use the term molar concentration to distinguish it from other ways of expressing concentration. The unit for molar concentration is mol/L or mol·L-1 (e.g. 1.0 mol/L). The symbol M is sometimes used to represent this unit (e.g. 1.0 M). Calculate the molar concentration of 4.60 g of ammonium nitrate in enough water to make a 500.0 mL solution.

Preparing Solutions: Example : Calculate the mass of barium nitrate required to produce 100.0 mL of 0.100 mol/L solution.

Preparing a Solution Using a Pure Solid Once you determine the mass of solute to be dissolved, you are ready to carry out a solution preparation procedure. Preparation of a solution from a mass of solid solute, requires these materials:

1) a source container of the solute2) a beaker e.g. 250 mL3) a scoopula4) a balance5) a full wash bottle (water)6) a stirring rod7) a funnel8) a volumetric flask (they come in various sizes) and cap

To prepare a solution, follow these steps carefully.

A) Obtain the desired mass of solute in the beaker by: finding the mass of the beaker.adding the mass of the beaker and the mass of solute to be obtained together and presetting the balance to that amount.transferring solute from its container to the beaker until the balance beam levels.

B) Add less than half of the final volume of water (i.e. about 50 mL) to the solute in the beaker and stir to dissolve.C) Place the funnel in the neck of the volumetric flask and, using the stirring rod as a guide, pour the solution into the flask.D) Use the wash bottle to rinse leftover solution on the stirring rod and walls of the beaker into the flask, repeat until you are convinced that all the solute

residue has been rinsed into the flask. Rinse walls of the funnel into the flask and remove.E) Carefully add water to the flask until the water level reaches the etched mark on the neck of the flask. You may want to use a medicine dropper to add

the last few drops so that the bottom of the dip in the water level (known as the meniscus) is even with the etched mark on the neck of the flask.F) Stopper the flask and invert several times to ensure the mixture is homogeneous.

Dilution: When a solution is diluted, the number of moles of solute does not change. In other words, the number of moles of solute (ni) in the concentrated volume is equal to the number of moles (nf) in the diluted volume:

 Ex: Calculate the volume of 11.6 mol/L hydrochloric acid required to prepare 250.0 mL of 1.00 mol/L.

Preparation of a Solution by DilutionThese are the steps involved in diluting small quantities of concentrated solutions.

1) Use a pipette or a graduated cylinder to deliver a specific volume of concentrated solution into a volumetric flask .A pipette is used for volumes such as 5.00 mL, 10.00 mL, or 25.00 mL.

2) Use a wash bottle to bring the volume of the diluted solution to the desired volume.3) Stopper and invert the volumetric flask/container several times.

Calculating Ion Concentrations:Report the concentrations of individual ions instead of their compounds. You can calculate the concentration of a specific ion in solution using the molar concentration of the solution and a dissociation or ionization equation.

ExampleCalculate the nitrate ion concentration in a 0.200 mol/L magnesium nitrate solution.

Solution Stoichiometry. Example : A student uses 32.5 mL of 1.02 mol/L sodium hydroxide solution to neutralize 250.0 mL of nitric acid. Calculate the molar concentration of the acid.

Bonding:The valence electrons of an atom are labelled based on whether they are single or paired. Single or unpaired electrons are called bonding electrons. The paired electrons are called lone pairs. The number of chemical bonds an atom can form is a function of its number of unpaired electrons.

It is possible to use the electronegativity values of the elements in a binary compound to make simple predictions about the type of bonding in the compound. Here are some general guidelines you can use:

1) if both elements in a compound have low electronegativity (<1.7), then the bonding that occurs between them may be metallic bonding.2) if both elements in a compound have high electronegativity (>1.7), then the bonding that occurs between them may be covalent bonding.3) if one element has low electronegativity and the other element has high electronegativity, then the bonding between them may be ionic bonding.

Properties of MetalsMetals are excellent conductors of electrical current. Since electricity is essentially the movement of electrons, the electron-sea model can be used to explain the movement of electrons into, through, and out of a metal sample. The sea of electrons is due to the low electronegativity of metals enabling the electrons to “float” around the atoms in a metal. A recurring point in this unit is the connection between electron distribution and chemical bonding. Here we focus on ionic bonds.

Why do Atoms Become Ions? Ions are atoms or groups of atoms that have gained or lost valence electrons. The atoms or groups of atoms that lose electrons are called cations while those that gain electrons are called anions. The noble gases are the key to understanding why atoms become ions. You should notice that each atom has a group of 8 valence electrons, this is called the octet rule.There is a relationship between the position of an element in the periodic table and the charge possessed by its ions. This table provides some general rules you can follow when determining ion charges. Group Number 1 2 13 14 15 16 17 18Ion Charge 1+ 2+ 3+ 4+ 3- 2- 1- 0

Properties of IonicsRecall that an ionic compound always consists of cations and anions.  Reactions between metals and nonmetals can occur when metals lose their loosely held electrons to high electronegativity nonmetals. The result is the formation of ions of opposite charge. It is these oppositely charged ions that mutually attract each other to form ionic compounds. The attraction between these oppositely charged ions is called an ionic bond. Ionic compounds are solids at room temperature.

1) The regular or repeating three-dimensional distribution of cations and anions in an ionic compound is called an ionic crystal lattice. If you could reduce a crystal to the lowest number of cations and anions forming a neutral unit, you would have a formula unit. Formula units however, do not exist as separate independent things like molecules do. 

2) Ionic compounds are hard crystalline solids with high melting points and boiling points.

3) Ionic compounds are brittle. If a force is applied to the crystal, the cations and anions may be forced towards “like” charged ions. This shift may cause repulsions that cause the crystal structure to shatter. 

Ionic compounds conduct electricity as liquids and as aqueous solutions. Heating an ionic compound so that it melts or placing it in water so that it separates into anions and cations allows the anions and cations to move around freely. Ionic solids are nonconductors of electricity and poor conductors of heat. In an ionic crystal lattice, the ions are in fixed positions. The lack of mobility means that the ions are not free to carry electrical current

Network Covalent CompoundsDiamond (pure crystalline carbon) and silicon carbide are two of the hardest and highest melting point substances known. What makes these molecular substances so different from other molecular substances like water or methane?

PropertiesNetwork covalent solids are insoluble in most substances. They are not malleable or ductile. Any theory about bonding in network covalent solids must take these properties into account. 

The properties of network covalent solids may be a function of the high bonding capacities of their atoms. Carbon and silicon atoms can form four covalent bonds. Carbon atoms easily bond to each other to form branched chains or rings of atoms. Carbon is wayyyy special!!!! Diamond, if a carbon atom bonds to four more carbons, then these four carbons can form bonds with more carbons and so on, and so on, and so on. The result is a macromolecule of uncountable numbers of carbon atoms in a rigid network. In other words, when you look at a diamond, you are looking at one huge molecule. The idea of a network of covalently bonded atoms might explain why network solids are poor conductors of electricity. Silicon carbide is a network covalent solid: each carbon is bonded to four silicon atoms and vice versa. The network of covalent bonds in SiC results in a hard, high melting point substance.

Valence Shell Electron Pair Repulsion Theory (VSEPR)We live in a three-dimensional world. The objects that surround us are three-dimensional. Is it reasonable to assume that individual molecules can also have distinctive shapes? Electrons are negatively charged particles that repel each other in the same way that the “like” poles of two magnets repel each other. Valence Shell Electron Pair Repulsion Theory (VSEPR) theory is based on the simple idea that electrons repel each other. Groups of valence electrons spread out as far apart as possible over the surface of a central atom in order to minimize the repulsive forces between them.

Applying VSEPR Theory You can predict the shape of a simple molecule using its Lewis diagram and VSEPR theory. A simple molecule is defined as one that has a single central atom. Here are the steps to follow:

1) draw the Lewis diagram for the molecule2) count the number of lone pairs and bonding groups around the central atom2) predict the shape of the molecule based on the number of lone pairs and bonding groups.

You will explore five simple (ON THE EXAM) shapes: tetrahedral, pyramidal, bent, linear, and trigonal planar. 

Table 1: Summary of VSEPR theoryNumber of Lone Pairs

Number ofBonding Groups

Shape AroundCentral Atom

Bond Angles Example

0 4 Tetrahedral 109.5° CH4

1 3 Pyramidal 107° NH3

2 2 Bent 105° H2O0 3 trigonal planar 120° H2CO0 2 Linear 180° CO2

Electronegativity and Polar Covalent BondsIf one atom has a greater attraction for valence electrons than the other, it is reasonable to conclude that the electrons will be pulled closer to the more electronegative atom. The sharing of the electron pair will be unequal. When this occurs, the bonding is said to be polar covalent, or a bond dipole.

Nonpolar Covalent BondsWhen two atoms with the same electronegativity share valence electrons, the bond between them is a nonpolar covalent bond (sometimes called pure covalent). There is no charge separation - the electrons are equally shared by the two atoms. All the diatomic elements exhibit nonpolar covalent bonding.

The Bond ContinuumAt one extreme there are the pure covalent or nonpolar covalent bonds, and at the other extreme there are the highly polar bonds that border on being ionic. This range from nonpolar to highly polar is known as the bond continuum.The polarity of a covalent bond is a function of the electronegativity difference between two bonded atoms. The greater the difference, the more polar.

When valence electrons are pulled closer to one end of a molecule, a molecular dipole is created and the molecule becomes polar. You can predict whether or not a polyatomic molecule is polar by using your knowledge of bond dipoles and molecular shapes. The steps you perform to make these predictions are:

1) draw a Lewis diagram2) draw a shape diagram for the molecule and use electronegativity values to identify bond dipoles3) analyse the bond dipole distribution to determine molecular polarity

a. if all the bond dipoles cancel each other, the molecule does not possess a molecular dipole and is nonpolarb. if bond dipoles do not cancel each other, the molecule has a molecular dipole and is polar

Although each bond is polar, the poles cancel, it is a non-polar molecule overall.

Intramolecular versus Intermolecular Forces: A covalent bond is an intramolecular attraction within a molecule. An intermolecular force is an attraction between two molecules. They are the result of attractions between positively and negatively charged regions of molecules. They are not as strong as chemical bonds. There are two types of intermolecular forces, van der Waals forces and hydrogen bonds. Van der Waals are broken into dispersion and dipole forces.

Dipole - Dipole ForcesIn a polar molecular substance, the intermolecular attractions involving permanent dipoles are appropriately called dipole-dipole force. Permanent δ+ and δ- regions are important in the formation of attractions that result in condensation or solidification.

London Dispersion ForcesAn instantaneous dipole in one atom may cause a redistribution of electrons in a neighbouring atom or an induced dipole. When this happens, the two helium atoms are momentarily attracted to each other. This force of attraction is called London dispersion force.

Factors Influencing the Strength of London Dispersion ForcesA substance’s boiling point is the temperature at which it changes from a liquid to a gas. Boiling point is used as an indicator of the strength of the forces holding particles together as a liquid. Molecules with more electrons have stronger London dispersion forces, and therefore higher melting and boiling points.

More Than One Force?

Substances that possess both dipole-dipole and London dispersion forces will have higher melting and boiling points than substances that possess London dispersion forces only. Highest will be those who possess all 3 forces. Consider the forces present in these three substances: C3H8, CH3Cl, and C2H5OH.

C3H8 is a nonpolar molecular substance with a tetrahedral shape around the central carbon atoms. It has LD force onlyCH3Cl is a polar molecular substance with a tetrahedral shape around the central carbon atom. It has LD and DD forceC2H5OH is a polar molecular substance with a tetrahedral shape around the central carbon atoms and a highly polar O-H bond at one end. All 3 forces here.

Solubility A general rule about solubility or miscibility is that "like dissolves like", where the term "like" where polar solutes dissolve in polar solvents and nonpolar solutes dissolve in nonpolar solvents.

Relative Strengths of Ionic, Metallic and Covalent Bonds: Recall that network covalent substances have some of the highest melting and boiling points known. Since melting a network solid requires the breaking of covalent bonds, we can infer that covalent bonds are the strongest of the forces of attraction Diamond and silicon carbide have the highest BP and MP of all. Ionic compounds tend to have melting points well above 300°C. They are all solids in pure form at room temperature. Metals are solids at room temperature with relatively high melting points.

OrganicIn 1828, Frederich Wöhler discovered that urea - an organic compound - could be made by heating ammonium cyanate (an inorganic compound). Today over 98% of all known chemicals are organic. Millions of years ago, the organisms that inhabited earth were quite different than those we find here today. Plants were fast growing with broccoli-like stems died, and decayed to form rich organic soils upon which more and more plants grew. Eventually, thick layers of decomposing organic matter accumulated in much the same way that peat bogs do today. Over time these massive organic layers were buried under tremendous pressures and transformed into various types of coal or oil. Every living organism is a source of organic compounds. Each species is capable of producing a wide range of compounds, some of which are unique to that single species. The scent of a rose or the taste of a strawberry. Humans have extracted and purified thousands of useful compounds from plants and animals. For example, the penicillin used to fight bacterial infections is extracted from a naturally occurring mold. Acetylsalicylic acid, commonly known as aspirin, comes from the bark of a willow tree.

Stability of Carbon to Carbon Bonds: Carbon atoms form stable covalent bonds with other carbon atoms. Carbon to carbon bonds are very strong.The tremendous diversity of organic compounds is due mainly to the ability of carbon atoms to form stable chains, branched chains, rings, branched rings, multiple rings, and multiple bonds (double and triple bonds). Add to this the ability to bond to many other nonmetal atoms, and you can certainly see why organic compounds outnumber all other Cl . Consider these two structures:

n-butane                       methylpropane

Structures that have the same molecular formula but different structural formulas are called structural isomers. 

Hydrocarbons: There are two main classes of hydrocarbons: aliphatic and aromatic hydrocarbons. Aliphatic hydrocarbons consist of carbon atoms bonded together in straight chains, or branched chains, or rings. Aromatic hydrocarbons are distinguished by the presence of a special group of six carbons known as the

benzene ring. Three classes of aliphatic hydrocarbons: alkanes, alkenes and alkynes.

General Formula CnH2n+2 CnH2n CnH2n-2

Class of Hydrocarbon alkanes alkenes (one double bond)cycloalkanes

alkynes (one triple bond)cycloalkenes (one double bond)

Alkanes: are hydrocarbons in which the carbon atoms have single bonds to other atoms. A series consisting of a group of compounds in which the compounds differ by a constant increment is called a homologous series. The methane, ethane, propane and butane are an example of a homologous series.

Nomenclature (Naming) of Alkanes: The nomenclature system for organic compounds is based on sets of prefixes and suffixes. Table 1: IUPAC prefixes.

Meth eth Prop But Pent hex Hept oct non dec1 2 3 4 5 6 7 8 9 10

Naming Simple Alkanes: To name continuous-chain (simple) alkanes from either a chemical or structural formula: 1) count the number of carbon atoms and indicate this number using the appropriate prefix and add the -ane ending. Ex: pentane is C5H12. Structural formula

or the skeletal diagram

The condensed formula shows each carbon atom with the number of hydrogen atoms bonded to it.

     or           

Alkyl groups have have the general formula CnH2n+1. They have one less hydrogen atom than a corresponding alkane. For example the methyl group, -CH3, has one less hydrogen than methane, CH4, or ethyl:  -C2H5 . Alkyl groups are examples of substituents: atoms that replace a hydrogen a chain carbon atoms.

Naming Branched Alkanes

1) First, the name of a molecule is based on the longest continuous chain of carbon atoms containing a functional group. 2) Second, lowest possible numbers are used to indicate the location of substituents or functional groups on the continuous chain.3) If an alkyl group occurs more than once, use a Latin prefix to indicate the number present. The Latin prefixes are di = 2, tri = 3, tetra = 4, penta = 5, and so on.

-   e.g. two methyl groups would be represented as dimethyl4) Use proper punctuation: commas are used to separate numbers, and hyphens are used to separate numbers and letters. 

Example : Write a IUPAC name to represent this structural formula.

Answer: 4-ethyl-3-methylheptane

Notice that the alkyl groups are listed in alphabetical order and their locations on the parent chain are indicated using numbers. Hyphens separate the numbers from the letters. Notice that the "methyl" and "heptane" become one name. In organic chemistry, the term saturated refers to organic compounds which contain single carbon to carbon bonds or which have the maximum number of hydrogen atoms bonded to carbon atoms. Alkanes are saturated hydrocarbons.

Unsaturated (enes and ynes): Hydrocarbons whose molecules contain double or triple carbon to carbon bonds (multiple bonds) are said to be unsaturated. When naming alkenes and alkynes, a number is used to designate the location of the multiple bond. In fact, priority in the numbering of the longest continuous chain in unsaturated hydrocarbons is given to the location of the multiple bond.

Cylco’s: A ring of three or more carbons connected by single bonds is called a cycloalkane. Cyclic alkanes have two less hydrogen atoms than their corresponding continuous-chain alkanes, it is thus an ENE ( CnH2n) which is the same as the general formula for an alkene that has one double bond. Cyclic alkenes are rings that possess a double carbon to carbon bond. They are sometimes referred to as cycloalkenes. Ex: of cyclohexene.

Carbon to Carbon Bonds in Benzene

The benzene ring consists of six carbon atoms, each bonded to a hydrogen atom. Resonance means that there are two or more possible distributions of bonding electrons for a compound.

Naming Monosubstituted Alkyl Benzenes

A benzene compound in which one hydrogen is replaced by an alkyl group is called a monosubstituted alkyl benzene. Consider these examples:

Methylbenzene Propylbenzene

The benzene ring is the parent and the alkyl group is the substituent.

When Benzene is the Substituent

There are instances when a benzene ring is bonded to a non-end carbon, of an alkyl group. The alkyl groups become the parents and the benzene rings become the branches. As a branch, the benzene ring is called a phenyl group. 

2-phenylpropane

Disubstituted Alkyl Benzenes: When two hydrogen atoms on the benzene ring are replaced by alkyl groups, the result is a disubstituted alkyl benzene. The two alkyl groups may be the same or different. Consider these examples:

1,3 – dimethylbenzene or m-dimethylbenzene

Oil Refining and the Properties of Hydrocarbons: As you get larger and larger molecules of alkanes, the boiling point goes up (due to increased electrons or LD force) Crude oil is a mix of many different organic compounds. The individual compounds or groups of compounds are called fractions. When crude oil is refined, these fractions are separated from each other. Separation is achieved by heating crude oil so that the various fractions boil off and condense at different heights in a distillation tower. The separation of crude oil on the basis of the different boiling points of its fractions is called fractional distillation.

Combustion Reactions: A typical example is the combustion of propane:

Cracking and Reforming

Cracking involves converting large alkanes to smaller alkanes, alkenes, and hydrogen. Two important types of cracking are thermal cracking and catalytic cracking. Thermal cracking involves heating large hydrocarbons in the absence of air until the carbon to carbon bonds break. Catalytic involves a the use of heat and a special chemical substance called a catalyst to break the bonds.

Polymerization Addition: is the joining of thousands of unsaturated hydrocarbons (the monomer) to form one very long molecule (the polymer). 

Functional Groups: The term alkyl halide is often used to represent organic halides derived from hydrocarbons. The general formula of an alkyl halide is R-X where R represents an alkyl group and X represents a halogen substituent. Naming alkyl halides is a lot like naming branched alkanes. Here are the steps to follow:Identify and name the longest continuous chain of carbon atoms and name the halogen substituent(s). Assign lowest possible numbers to the substituents. List the substituents in alphabetical order using appropriate prefixes.

1-chloro-1,2-difluoroethane.

Production of Organic Halides: Substitution Reactions:

A substitution reaction occurs when a hydrogen atom is removed from the hydrocarbon and replaced by a halide substituent.

Structural isomerism is the existence of two or more structural formulas for one chemical formula. Single step or multi-step substitution reactions can result in the production of structural isomers. Consider the reaction between propane and chlorine. If a hydrogen on carbon #1 is replaced, then the product is 1-chloropropane; however, if a carbon #2 hydrogen is replaced, then the product is2-chloropropane.

Production of Organic Halides: Addition Reactions: Alkenes and alkynes are unsaturated hydrocarbons containing at least one double or triple bond respectively. They do not undergo substitution reactions; instead, they undergo addition - a reaction in which substituents are added to both carbons involved in the multiple bond. Alkenes and alkynes are chemically more reactive than alkanes because of the presence of the multiple carbon to carbon bonds.

Example :Predict the products of addition reactions involving ethene and chlorine.

Elimination Reactions: Halogen substituents can be removed from an alkyl halide using a base.

Hydrocarbon derivatives are defined on the basis of their functional groups - atoms or groups of atoms that give compounds their unique chemical and physical properties. Alcohols are defined by the functional group called hydroxyl (-OH). The general formula for an alcohol is R-OH where R represents an alkyl groupsEthers are defined by the functional group known as ether (-O-). The general formula for an ether is R-O-R'  where R and R' represent alkyl groups.

Naming and Drawing Structural Formulas for Alcohols: The -ol suffix in a chemical name identifies a compound as an alcohol; in other words, it signals the presence of a hydroxyl group. Count to find the number of carbon atoms in the longest continuous chain. Change the -e ending of the alkane name to –ol. Indicate the location of the hydroxyl group using the lowest possible number. If the alkyl group is branched, priority in the numbering of the parent goes to the location of the hydroxyl group.

2-methyl-2-butanol

Ethers are hydrocarbon derivatives that contain an oxygen atom bonded to two alkyl groups. They have the general formula of R-O-R' where R and R' are alkyl groups. Write alkyl’s in alphabetical order

ethylpropylether

Reactions of Alcohols. Addition Reactions: Consider this example:

Elimination Reactions of Alcohols: Alkenes can be produced by elimination of a water molecule from an alcohol. This reaction involves the use of an acid catalyst, with the symbol H+, or H2SO4.

Properties of Alcohols and Ethers: The properties of alcohols are a function of the hydrogen bonding associated with the highly polar "OH" bond. Short chain alcohols like methanol, ethanol, and propanols are soluble in nonpolar and polar solvents. This makes them very useful for cleaning oily, greasy, or waxy materials. The alkyl component of these alcohols dissolves in nonpolar oils, grease, or wax while the hydroxyl end dissolves easily in water.  The higher melting and boiling points of alcohols compared to corresponding aliphatic hydrocarbons is due to the strong hydrogen bonds between alcohol molecules and the slightly greater London dispersion forces due to the higher number of electrons. For example, ethane boils at -88°C but ethanol boils at 78°C.

The properties of ethers are a function of the stable ether link between the alkyl groups. Aside from being flammable, ethers are generally unreactive. Ethers are volatile - they evaporate more easily than alcohols because they lack hydrogen bonding. This property makes them useful as anaesthetics.

Aldehydes and Ketones : Aldehydes and ketones both contain the functional group carbonyl (-C=O),the location of the carbonyl group in a carbon chain determines whether the hydrocarbon derivative is an aldehyde or a ketone. A functional group gives a compound its unique chemical and physical properties. Based on this definition, you would think that aldehydes and ketones have similar properties.

Aldehydes have a terminal carbonyl group. A good way to remember this fact is that the name aldehyde begins with "al" and the word terminal ends in "al".  The general formula of an aldehyde is R–CHO where R represents a single hydrogen atom or a chain of carbon atoms.

propanal

Ketones: Ketones also contain the functional group called carbonyl; however, unlike the aldehydes, the carbonyl group in ketones is located on a non-terminal carbon. This means the simplest possible ketone is propanone: CH3COCH3. 

2-Propanone

Amines: Amines are derived from ammonia (NH3) when more of the hydrogen atoms is replaced by a hydrocarbon group. A nitrogen atom is the functional group. To name an amine, identify the alkyl group and add the suffix -amine to its name.

Example: The alkyl group consists of four carbon atoms, so its name is butyl. It is bonded to an amino group, so its name is butylamine.

Carboxlyic Acids and Esters:End in oic acid and esters end in oate.

REVIEW PROBLEMS

1.Make the following conversions:

a) 1.75 g to mg b) 2.40 mmol to mol c) 0.057 L to mL d) 364 mL to L

2. Do the following operations and give the answer to the correct number of significant digits:

a. 1.23 m + 3.674 m + 8.2 m b) 74.372 g – 23.4 g

b. 6.43 m X 0.27 m d) 0.474 g ÷ 0.020 mol

3. Convert the following measurements to scientific notation, each with 2 significant digits.

a. 2730 m b) 347 g c) 0.00456 L

4. Find the average atomic mass for oxygen from the data provided

Isotope Atomic mass (u) % Abundance

Oxygen-16 15.995 99.759

Oxygen-17 16.995 0.037

Oxygen-18 17.999 0.204

5. Complete the following table:

Isotope Isotope Symbol Mass Number Atomic Number

Magnesium-24

14

47

6. How many moles are 3.42 X 1024 atoms of copper?

7. How many molecules are in 2.65 mol of CO2?

8. How many formula units are in 0.67 mol of NaCl?

9. How many calcium ions are in 2.17 mol of calcium phosphate?

10. Calculate the moles present in each of the following:

a) 32.0 g of methane b) 168.0g of mercury(II)sulfide

11. Calculate the volume, at STP, of 8.356 mol of oxygen.

12. Find the number of moles of helium gas present in a 3.75 L balloon at STP.

13. Find the number of atoms in 2.50 L of carbon dioxide gas, at STP.

14. Determine the percent composition of the following:

a) ethanol b) lithium nitrate

15.Find empirical formula of a copper sulfide ore if 7.68 g of the compound contains 6.13 g of Cu

16. Determine the molecular formula of a compound that has a molar mass of 94.92 g/mol and has a percent composition of 25.3% carbon and 74.7% chlorine.

17. A 2.78 g sample of hydrated iron(III)sulfate was heated to remove all the water. The mass of the anhydrous compound was found to be 1.52 g. Calculate the number of water molecules associated with each formula unit.

18.When 20.4 g of sodium metal are combined with chlorine gas, what mass of sodium chloride is produced?

19. What is the concentration of a potassium hydroxide solution if 12.8 mL of this solution is required to react with 25.0 mL of 0.110 mol/L sulfuric acid?

20. If 1.41 g of zinc are added to 1.85 g of hydrochloric acid, identify the limiting and excess reagents (give reasons for your choice) and determine how much (moles) of the excess reagent remains?

21.What volume of 0.125 mol/L sodium hydroxide solution is required to completely react with 15.0 mL of 0.100 mol/L silver sulfate? Give your answer in mL.22.Coal can undergo an incomplete combustion in the absence of a plentiful supply of oxygen to produce deadly carbon monoxide gas (the only product). What volume of carbon monoxide is produced at STP by the incomplete combustion of 120 kg of coal (assume pure carbon)?

23.If 12.5 g of copper are reacted with an excess of chlorine gas, then 25.4 g of copper(II)chloride is obtained. Calculate the theoretical yield and the percent yield.

24. Ethane, C2H6(g), burns completely in the presence of oxygen. How many molecules of carbon dioxide gas can be produced when 0.14 L of ethane is burned?

25. Write dissociation equations for each of the following compounds. Assume all are soluble in water.

26. Solid barium nitrate b) hydrogen iodide gas

c) liquid hydrogen nitrite d) solid strontium hydroxide

27. Write dissociation equations for the 6 strong acids.

28. Indicate whether each of the following compounds acts as an electrolyte when dissolved in water.

a)sulphuric acid b) ethanol c) glucose d) sodium nitrate

29.Know the meaning of each of the following terms:

solution, solute, solvent, aqueous solution, dilute solution, concentrated solution, saturated solution, unsaturated solution, supersaturated solution, dynamic equilibrium, solubility,

acidic solution, basic solution and neutral solution.

30.What is the molar concentration of a solution in which 120.0 g of sodium hydrogen sulfite, NaHSO3, is dissolved in water to form 8.00 L of solution?

31. Find the number of moles of sodium phosphate in 6.00 L of a 0.0125 mol/L cleaning solution.

32. What volume of 2.50 mol/L sulfuric acid solution would contain 2.00 mol of sulfuric acid?

33.What mass of sodium hydroxide is required to prepare 400.0 mL of a 0.200 mol/L cleaning solution?

34.What volume of a 1.0 mol/L calcium chloride solution could be prepared from 1.00 kg of CaCl 2

35.What is the concentration of a lithium nitrate solution that contains 2.50 mol of lithium nitrate dissolved in water to a final volume of 750.0 mL?

36.An ammonia solution was made by diluting 150.0 mL of a 4.50 mol/L concentrated solution until the final volume reached 1000.0 mL. What was the final molar concentration?

37.What volume of 0.250 mol/L nitric acid solution is required to prepare 320.0 mL of a 0.030 mol/L solution of this acid?

38.What is the molar concentration of a sodium hydrogen carbonate solution which has 3.57 mmol of sodium hydrogen carbonate dissolved in water to make 240.0 mL of solution?

39.In a chemical analysis, a 25.0 mL sample was diluted to 500.0 mL and analyzed. If the diluted solution had a molar concentration of 0.108 mol/L, what was the molar concentration of the original sample?

40.Find the concentration of each ion in the following aqueous solutions(assume they dissociate):

a) 0.25 mol/L aluminum sulftae b) 2.35 mol/L calcium nitrate

41.What is the concentration of ions in a CaBr2 solution prepared by dissolving 18.2g of the solid in water to a final volume of 300.0 mL?

42. What mass of cobalt(II)chloride is needed to prepare a 1.50L of solution that contains 0.109 mol/L Cl -(aq) ions?

43. Find the mass of solid silver chromate that forms when 50.0 mL of 0.100mol/L silver nitrate reacts with 25.0 mL of 0.150 mol/L sodium chromate.

44. Find the maximum mass of precipitate that can form when 8.76g of sodium sulphide is added to 350.0 mL of 0.250 mol/L lead(II)nitrate solution. If 16.2g are recovered, what is the % yield?