ChemistryUnit 1 covered the atomic structure of elements and the periodic table of elements. There are only 92 elements that occur in nature, many of these are extremely rare, and 26 elements have been created in large, expensive experimental sites like the Stanford Linear Accelerator Center and the Large Hadron Collider in Geneva. The millions of other pure substances both in nature and created in labs are compounds. Compounds form when elements combine and their atoms held together by chemical bonds. This unit is about how compounds form and the properties of different types of compounds.

Ionization Energy and ElectronegativityChemical bonds form when electrons are transferred or shared between atoms. The periodic trends, of ionization energy and electronegativity, help explain why metals and nonmetals react by transferring electrons and two nonmetals will share electrons when bonding.

Ionization energy is a measure of how hard it is to remove an electron. Metallic elements on the left of the periodic table have low ionization energies and lose electrons easily. Electronegativity measures the pull of the atomic nucleus for electrons towards itself. The nonmetals in the upper right with a high electronegativity will form bonds by sharing electrons between atoms of similar electronegativity as well as attract electrons from metallic elements with a low electronegativity.

113.6H

2.20

Ionization Energy (eV)&

Electronegativity

24.6He– –

25.4 Li

0.98

9.3Be1.57

8.3B

2.04

11.3C

2.55

14.5N

3.04

13.6O

3.44

17.4F

3.98

21.6Ne– –

35.1Na0.93

7.6Mg1.31

6.0Al

1.61

8.2Si

1.90

10.5P

2.19

10.4S

2.58

13.0Cl

3.16

15.8Ar– –

44.3K

0.82

6.1Ca1.00

6.6Sc1.36

6.8Ti

1.54

6.7V

1.63

6.8Cr1.66

7.4Mn1.55

7.9Fe1.83

7.9Ni

1.88

7.6Co1.91

7.7Cu1.90

9.4Zn1.65

6.0Ga1.81

7.9Ge2.01

9.8As2.16

9.8Se2.55

11.8Br

2.96

14.0Kr– –

54.3Rb0.82

5.7Sr

0.95

6.2Y

1.22

6.6Zr

1.33

6.8Nb1.60

7.1Mo2.16

7.3Tc1.90

7.4Ru2.2

7.5Rh2.28

8.3Pd2.20

7.6Ag1.93

9.0Cd1.69

5.8In

1.78

7.3Sn1.96

8.6Sb2.05

9.0Te

2.10

10.5I

2.66

12.1Xe– –

Ionization Energy (top value) and Electronegativity of Selected Elements

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ChemistryTransferring ElectronsIonic compounds are made by transferring electrons to become ions. A cation, which is a positive ion, forms by losing an electron, and an anion, a negative ion, forms when the other element gains an electron. The two ions will attract each other because of their opposite charge by a bonding force called electrostatic attraction.

The Transfer of an Electron from a Lithium Atom to a Fluorine Atom

A lot of chemistry takes place in this diagram that’s common with most synthesis reactions that combine elements to make an ionic compound. Let’s start with the reaction. The two atoms at the left are called the reactants, which are the starting materials of the reaction, and the two ions at the right are called the products, which are the results of a reaction. The arrow, , means ➔ yields, forms, gives, reacts to make, etc.

The circles in the diagram represent the energy levels or shells of the electron configuration and each dot, •, is an electron. This is called the Shell Model and can be a valuable tool to show the movement of electrons and the formation of chemical bonds. Notice that the electron configuration given below the shell model provides the same information about the number of electrons on each shell.

The low ionization energy of lithium, Li, shows that it takes little energy to remove the electron. The ionization energy of fluorine is higher and fluorine will not lose its electron. The electronegativity of fluorine is high (the highest value actually) and its nucleus has a high attraction for electrons. These two factors combine to cause the transfer of an electron to fluorine.

The products of the reaction are a lithium cation, Li1+, which has lost a negative electron while not changing the three positive protons in the nucleus so it has a positive charge, and a fluorine anion, F1–, which has an extra electron so a negative charge. The superscript, 1+ and 1–, represent the charge of the ion and the number of electrons lost (for +) and electrons gained (for –).

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Li

1s2 2s1

Ionization Energy 5.4 eV

Ionization Energy 17.4 eV

+ ➔• ••

•F

1s2 2s2 2p5

• ••

•••

• • •

•+• ••

°Li1+

1s2

• ••

•••

• • •

••

F1–

1s2 2s2 2p6

Electronegativity 0.98 eV

Electronegativity 3.98 eV

ChemistryFinally, it is important to remember that lithium has a larger radius than fluorine since fluorine has a greater effective nuclear charge (1 shell shielding the valence electrons and the nucleus changes from +3 to +9). This is represented in the shell model. But when fluorine gains an electron there are more electron-electron repulsions and the +9 charge is distributed among ten electrons instead of nine so the radius of the anion increases. Likewise the lithium cation is much smaller than the lithium atom, since the electron in the outer shell is gone and the inner shell experiences an unbalanced positive charge (nucleus is +3 and electrons are -2).

Noble Gas Electron Configuration is StableIn the reaction between lithium and fluorine only one electron was transferred, but as many as four electrons can be transferred in common reactions. The valence electrons, which are the outermost electrons, are most easily removed because of shielding from the inner electrons. In the figure below each atom from lithium to carbon shows the same result when all the valence electrons are removed: two electrons in a 1s2 electron configuration, which is the same electron configuration of helium.

Second Period Metals Losing Electrons to Form a Stable Noble Gas Electron Configuration

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lose 1 e– lose 2 e–    

lose 3 e– lose 4 e–

Li

1s2 2s1

• ••

•

Li1+

°

• ••

1s2

Be

1s2 2s2

• ••

•

•

• ••

Be2+

1s2

°

°

B

1s2 2s2 2p1

• ••

•

•

•

• ••

B3+

1s2

°

°

° • ••

C4+

1s2

°

°

°°

C

1s2 2s2 2p2

• ••

•

•

••

ChemistryMetals with low ionization energies lose electrons until the valence electrons are removed. No inner electrons are removed in the reaction because they are close to the nucleus which has an extra positive charge. Removing these electrons would take too much energy compared to reacting with other atoms that still have valence electrons. Since it is so common for the valence electrons to be lost in a reaction and the inner electrons remain with s and p orbitals filled, noble gas electron configurations are considered stable.

Third Period Metals Losing Electrons to Form a Stable Noble Gas Electron Configuration

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•••

•••

•

• •

•

lose 1 e–

Na

[Ne] 3s1

•

•

••

Na1+

°

• ••

1s2 2s2 2p6

•

••

•

•

•

•••

•••

•

• •

•

lose 2 e–

[Ne] 3s1

•

•Mg

•Mg2+

°

°

• ••

1s2 2s2 2p6

•

••

•••

••

•••

•••

•

• •

•

lose 3 e–

[Ne] 3s1

•

•Al

•

Al3+

•

•••

••

•••

••

•

1s2 2s2 2p6

°

°

°

ChemistryWhen atoms gain electrons, they also stop adding electrons when they get to a noble gas electron configuration, with a filled set of s and p orbitals, because gaining even more electrons would require a new, more distant energy level or shell to hold the electrons. The atoms that gain electrons have high values of electronegativity, which indicates the nucleus has a strong pull toward electrons. Nonmetals on the upper right of the periodic table, other than noble gases, have high electronegativities and form anions easily by gaining electrons.

Second Period Nonmetals Gaining Electrons to form Stable Noble Gas Electron Configurations

Because the electron configurations for families of elements are similar, their chemical reactions are similar. For example, the alkali metals the electron configurations is ns1 , where n is the highest energy level or shell. Thus, all alkali metals easily lose one electron in a reaction where there is a transfer of electrons. Likewise, alkaline earth metals with electron configurations of ns2 will lose two electrons. On the other side of the periodic table the halogens will gain one electron since their electron configuration is ns2 np5, while the oxygen family will gain two electrons because these atoms have electron configurations of ns2 np4.

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F

1s2 2s2 2p5

• ••••

•• • •

•

• ••

•••

• • •

••

F1–

1s2 2s2 2p6

O

1s2 2s2 2p4

• ••

••

• • •

•

•

•• • ••

•

•

••

•O2–

1s2 2s2 2p6

N

1s2 2s2 2p3

• ••

••

• •

•

• ••

••

•

••

•

•

•

N3–

1s2 2s2 2p6

C

1s2 2s2 2p2

• •••

• •

•

• ••

••

•

••

•

•

•

C4–

1s2 2s2 2p6

gain 1 e– gain 2 e– gain 3 e– gain 4 e–

Chemistry

Family Similar Electron Configuration For example

alkali metals ns1 K = [Ne] 3s1 ➔(lose 1e–) [Ne] = 1s➔ 2 2s2 2p6 = K1+

alkaline earth metals

ns2 Ba = [Xe] 6s2 ➔(lose 2e–) [Xe] = Ba➔ 2+

oxygen family ns2 np4 S = [Ne] 3s2 3p4 ➔(gain 2e–) [Ne] 3s➔ 2 3p6 = S2–

halogen ns2 np5 Br = [Ar] 4s2 4p5 ➔(gain 1e–) [Ar] 4s➔ 2 4p6 = Br1–

Most elements gain or lose electrons until a noble gas electron configuration is reached because any further changes are energetically unfavorable. Gaining more electrons requires adding electrons to a more distant energy level and losing more electrons would remove electrons from inner energy levels. Transition metals, however, with their d-orbitals, usually have multiple types of ions (tin has Sn2+ and Sn4+). This is because the d–orbital electrons are in an inner shell, ns2 (n – 1)dx , and have energies very close to the outer s-orbital electrons. So while the outer s-orbital electrons are usually removed, some d-orbital electrons can also be removed. For example, manganese, [Ar] 4s2 3d5, can have cations of Mn2+, Mn3+, Mn4+, Mn6+, and Mn7+.

ChargeUp to now the charge has been added to the cation or anion without explanation. Still the convention for showing charge is straightforward. When aluminum, Al, loses 3 electrons its cation is show as Al3+. The 3 represents the number of electrons lost and the + indicates there is an excess of positive charge. Similarly, when oxygen, O, gains 2 electrons the anion is O2–. The 2 represents the 2 gained electrons and the – is for the negative charge the oxygen gained.

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ChemistryValence Electrons and OctetsThe number of valence electrons, those electrons in the outermost energy level, determines the number of electrons an element will gain or lose. Since the families have the same pattern of outer electrons, the number of valence electrons is the same in each family.

1 2 ⇐ Number of ⇒Valence Electrons 3 4 5 6 7 81 H He

2 Li B *d electrons generally don’t count as B C N O F N3 Na Mg Al Si P S Cl Ar4 K Ca Sc Ti V Cr Mn Fe Ni Co Cu Zn Ga Ge As Se Br Kr5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

1+ 2+ Usual Charge ⇐ ⇒ 3+ 4± 3- 2- 1- NRNR = no reaction

The table above is a result of atoms losing and gaining electrons to form stable noble gas electron configurations. When an element loses electrons it commonly does so until all valence electrons are gone; thus, magnesium with 2 valence electrons will react to become Mg2+, magnesium cation by losing the 2 valence electrons. Likewise, when an element gains electrons, it will gain only enough electrons to fill the outer shell s and p-orbitals; such as, nitrogen with 5 valence electrons will gain 3 electrons to become N3–, nitrogen anion, so that there are 8 electrons in the outer shell. The full outer shell of 8 electrons is called an octet, and ions are said to be stable when they have an octet. Here are two more examples: K = [Ne] 3s1 ➔(lose 1e–)➔ [Ne] = 1s2 2s2 2p6 = K1+ or S = [Ne] 3s2 3p4 ➔(gain 2e–) [Ne] ➔ 3s2 3p6 = S2–.

Noble gases do not react because they already have full shells, ns2 np6. The energy change when electrons are added or removed is too large because the effective nuclear charge is too high for an electron to be removed easily and because adding electrons would require using a new shell which is not favorable.

For nonmetals in the second period that lose electrons, the number of electrons lost will result in the noble gas electron configuration of helium 1s2, so only two valence electrons are in the outer shell not an octet. Hydrogen usually loses one electron to become a single proton in the nucleus without any electrons in the electron cloud, H+; however, it very occasionally gains one electron to become H– with an electron configuration similar to helium.

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ChemistryLewis StructuresTo show the movement of electrons in a reaction, Gilbert Lewis introduced method of writing atoms and their valence electrons using chemical symbols and dots for electrons. Lewis dot diagrams, also called Lewis structures, electron dot structures, electron dot diagrams, etc., simply show the valence electrons about the chemical symbol of an element placed on the top and bottom and on the left and right.

1•H Lewis Dot Structures •He•

2 •Li •Be••  B••

• •C•••• •N••

•• •O••••• 

•F•••• ••

3•Na •Mg•

*d electrons generally don’t count as valence electrons.

•  Al••• •Si••

•• •P•••• •S•••

•• •Cl••••

•• •• Ar••

4•K •Ca•

Sc Ti V Cr Mn

Fe Ni Co Cu Zn•  Ga••

• •Ge•••• •As••

•• •Se••••• 

•Br•••••• •• Kr••

Forming Ionic CompoundsTo form an ionic compound, elements must transfer electrons from one another. So if a metal and nonmetal react, then the metal will lose electrons and the nonmetal will gain electrons to make a cation and an anion.

Model Used to Show the Reaction Reactants Products

Chemical Symbols Be + S ➔ Be2+ + S2–

Shell Model

Electron Configuration 1s2 2s2 + [Ne] 3s2 3p4 ➔ 1s2 + [Ne] 3s2 3p6

Electron Dot Structure •     ••   Be + •S• ➔ • ••

2+      •• 2– Be + •• S•• ••

By convention the cation, usually a metal, is written before the anion. This is particularly important for writing the chemical formula of ionic compounds.

Balancing the Electrons TransferredWhen electrons are transferred, the number of electrons lost and gained must be equal. To do this more atoms are added to provide a balanced number of valence electrons. For the first example below, beryllium, Be, needs to lose both outer

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Chemistryelectrons, but chlorine, Cl, only needs one. To balance the transfer of electrons needed for the atoms to reach a stable noble gas electron configuration another chlorine is added. The other examples demonstrate the balanced transfer of electrons. Here are some examples using electron dot diagrams: •     ••    ••    2+      •• 1– 

Be + •Cl•• + •Cl•• Be + 2 ➔ •• Cl•• • •• •• ••

• • ••      1+     •• 2– K + K + •O• 2 K + ➔ •• O•• •• ••

• • •    ••   ••     2+      •• 3–  Mg + Mg + Mg + •N• + •N• 3 Mg + 2 ➔ •• N•• • • • • • ••

In each of these reactions the number of valence electrons lost from each metallic element must have a place to be gained by the nonmetal element. The total number of electrons lost then gained is the least common multiple of the number of valence electrons (or you can just keep track of the electrons and add elements and valence electrons until the number of electrons lost equals the number of electrons gained).

For example, Identify the number and type of ions formed in a reaction of calcium and carbon. Step 1: Determine the number of valence electrons for each element and whether the element loses or gains electrons. Calcium has 2 valence electrons and is a metal so it will give up two electrons. Carbon is the nonmetal and with 4 valence electrons it will gain 4 electrons (carbon gains and loses electrons depending on what the other element does, in this case calcium wants to give electrons so carbon will gain electrons). Step 2: Determine how many extra elements need to be added so the number of valence electrons gained equals the number lost. Carbon needs more than 2 electrons from calcium so add another calcium. Two calciums will have 4 valence electrons to lose which equals the 4 valence electrons one carbon gains.Step 3: Show the reaction: • • 2+ •• 4– 2Ca + •C• 2 Ca +➔ ••

C•• ••

Electrostatic AttractionsWhen ions are created they can usually dissolve in water (we will discuss solubility in another unit), or will form a solid by sticking together. The opposite charges of cations and anions will force the ions to attract. Electrostatic attraction is the name of the strong bonding force holding the cations and anions together in an ionic compound.

Crystal LatticeThere are, of course, a multitude of cations and anions present in any reaction. So each cation can attract several anions and each anion can attract several cations. The result is a repeating and organized arrangement of cations and

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Chemistryanions called a crystal lattice. A crystal lattice is a very stable bonding for ions and a lattice bonding energy is associated with the repeated arrangement of cations and ions.

Chemical Formulae and Naming.There is a specific method for using element symbols from the periodic table of elements to represent compounds. For ionic compounds the cation is listed first and then the anion. The number of cations and anions needed to balance the charge is listed as subscripts. Two comments before showing some examples: 1) a 1 in chemistry is usually dropped, so that Na1+ equal Na+ or Mg1Cl2 equals MgCl2; and 2) since every solid ionic compound is composed of many, many cations and anions, we say that the chemical formula of an ionic compound represents the lowest whole number ratio of elements making up the compound.

Here are some examples of compounds that form from a pair of elements. The number of elements needed to balance the reaction is not shown.K + S K➔ 2SBa + P Ba➔ 3P2 See if you can determine these compounds using electron dot structures.Sn + F SnF➔ 4

The name of ionic compounds has a similarly straightforward methodology. Name the cation with the same name as the element, then name the anion but add an -ide ending to create a word other than the element so that a reader understands that a compound is being named. For example sodium reacts with chlorine (the two elements) to form sodium chloride (NaCl). Another example barium reacts with phosphorous (see example above) to form barium phosphide (Ba3P2). Note that the number of cations and anions is not listed since every chemist could figure out the number of cations and anions. Transition metalsRecall transition metals form multiple cations. Some important examples are Cu1+ and Cu2+, Fe2+ and Fe3+, Co2+ and Co3+, Pb2+ and Pb4+, Sn2+ and Sn4+, and Hg2

2+ and Hg2+ . To differentiate between the compounds that can form using different ions a roman numeral is used to name each of the cations. For example

Page 10 of 29

The crystal lattice of sodium chloride. Small sodium cations and larger chlorine anions in a

repeating pattern held together by electrostatic attractions.

http://commons.wikimedia.org/wiki/File:Sodium-chloride-unit-cell-3D-ionic.png

Chemistryfor Cu1+ and Cu2+ the two ionic compounds with an oxygen anion, O2–, are Cu2O and CuO. The compounds are copper (I) oxide and copper (II) oxide respectively.

Some metals have only one cation so they do not use roman numerals in their names: Ag forms only Ag1+ , Zn forms only Zn2+, Al forms only Al3+ , Cd forms only Cd2+. Below are examples of how to identifying the chemical formula of a compound from its elements, how to change from a chemical formula to the name of the compound, and how to write the chemical formula if given the name . For the last three examples, the oxidation number of the metal would not be known from just the element, but you can change from chemical formula to name and name to chemical formula

Elements(metal / nonmetal) Ions Chemical Formula of

the Compound Compound Name

Al + Cl Al3+ & Cl1– AlCl3 aluminum chloride

Ca + O Ca2+ & O2– CaO calcium oxide

Na + N Na1+ & N3– Na3N sodium nitride

Cu + F Cu1+ & F1– Cu2+ & F1–

CuFCuF2

copper (I) fluoridecopper (II) fluoride

Pb + S Pb2+ & S2–

Pb4+ & S2–PbSPbS2

lead (II) sulfidelead (IV) sulfide

Mo + O Mo6+ & O2–

Mo4+ & O2–MoO3

MoO2

molybdenum (VI) oxidemolybdenum (IV) oxide

It is necessary to find the ions of the compound before you can determine the chemical formula or name. Furhermore, you may have noticed a pattern for changing from ion to chemical formula called criss-cross. Using the charges of the cation and anion as the subscripts in the chemical formula for the opposite ion, you can write the chemical formula of the compound. This pattern works about 2/3rds of the time, but, as with the last two examples, the criss-cross can give subscripts that are too large. All subscripts for chemical formulae must be written in lowest whole number ratio.

Finally, the compound name can be broken up into the two ions to determine the chemical formula. For example, calcium phosphide would be the ions Ca2+ and P3–

which form Ca3P2. Metallic BondsMetals as either elements or alloys, which are mixtures of metals, have distinct properties. Metals are highly conductive of heat and electricity. They are dense, have high luster (shiny), metallic colors (copper, silver, and gold), can be pulled into wires (ductile), and can be beaten into sheets (malleable). The special bonding of metals is responsible for many of these properties. Metals bond by

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Chemistrysharing their valence electrons and some of the d-orbital electrons between cations of the metal. These shared, freely-moving electrons are called a “sea of electrons” and is how metallic compounds and elements bond as liquids and solids.

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ChemistryCovalent or Molecular CompoundsMost compounds form when elements share valence electrons to form a covalent bond. Covalent compounds include organic compounds, which are compounds formed by living matter (for example, DNA, sugars, vitamins, and proteins), medicines, plastics and polymers, compounds that are gases and liquids at normal temperatures and pressures, acids, and many others. Most covalent compounds are combinations of nonmetallic elements, and may have hundreds even thousands of elements bound together. The smallest individual set of atoms making up each compound that has the properties of that compound is called a molecule.

Sharing of ElectronsIonic bonding takes place by creating ions with filled s and p-orbitals, and sharing of electrons will take place until the atoms reach the same stable electron configuration. For example, each chlorine atom has 7 valence electrons, by sharing 1 valence electron each chlorine will have 8 (sharing electrons does not take away electrons it only adds electrons). Just like ionic compounds, the covalent compounds will have 8 valence electrons around each atom to create a stable electron configuration.

•• •• •• •••• Cl • + •Cl •• ➔ •• Cl •• Cl•• •• •• •• ••

Diatomic ElementsA set of elements, which includes chlorine, form molecules of two atoms at normal temperature and pressures: H2, O2, N2, F2, Cl2, Br2, & I2. This group of elements includes the halogens, hydrogen, oxygen, and nitrogen and can be recognized because they are the only elements that have names ending with -gen or -ine.

•• •• •• ••Fluorine (and other halogens) •• F • + •F •• ➔ •• F •• F•• •• •• •• ••

Hydrogen , H2 H • + •H ➔ H •• H

•• •• •• ••Oxygen •O • + •O• ➔ O•• •• O •• •• •• ••

Page 13 of 29

H He

B

Si

Ge

As

Sb Te

57 72 Po

89 104

Alk

ali M

etal

s

Hal

ogen

s

Semimetals ➷

Alk

alin

e Ea

rth M

etal

s

Nob

le G

ases

Transition Metals

Chemistry •• •• •• ••

Nitrogen •N • + •N• ➔ N•• •• •• N • •

Notice that hydrogen has only 2 valence electrons around each atom after sharing. This is because only the first energy level is filled, which only requires two electrons.Multiple BondsFor oxygen and nitrogen, the atoms must share more than one electron to reach an octet for each atom. When two atoms share two pairs of electron then a double bond forms; when two atoms share three pairs of electrons then a triple bond forms. A single bond forms with only one shared pair of electrons. Notice that for many simple examples of sharing the number of lone electrons an atom has equals the number of shared electrons in the compound.

Unlike ionic compounds, most covalent compounds have more than two atoms. When three, four, five, and more atoms combine they share electrons until each atom reaches an octet (although there are exceptions).

Elements Electron Dot Diagram Molecule Name of Compound

C & H

H • •••C• + H• H ➔ •• C •• H • ••

HCH4 methane

N & H

•• •••N• + H• H ➔ •• N •• H • ••

HNH3 ammonia

O & H •• •••O• + H• H ➔ •• O •• H •• ••

H2O water

C & Cl

•• •• Cl •• • •• •• •• ••

•C• + •• Cl• ➔ •• Cl •• C •• Cl •• • •• •• •• ••

•• Cl•• ••

CCl4 carbon tetrachloride

C & O • •• •• ••• C• + •O• O ➔ •• •• C •• •• O • •• •• ••

CO2 carbon dioxide

P & F

•• •• •• •• •••P• + •• Cl• ➔ •• Cl •• P •• Cl •• • •• •• ••

•• Cl •• ••

PCl3 phosphorous trichloride

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Chemistry

H & C &N

• •• ••H • + •C• + •N• H ➔ •• C •• •• •• N • •

HCN hydrogen cyanide

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ChemistryBonding electrons can be shown as lines connecting atoms and lone pair electrons can be shown as lines above an atom (http://commons.wikimedia.org/wiki/File:H2O.svg#file):

•• H •• O •• H

= ••

Naming of Simple Covalent CompoundsMany covalent compounds have common names like water, ammonia, sugar, acetic acid, vitamin C, hemoglobin, etc. but there is a systematic naming system. Name the first element and name the second with an -ide ending just like ionic compounds, but also tell how many atoms are in the molecule. Some examples are: carbon dioxide is CO2 and carbon monooxide is CO, phosphorous trichloride is PCl3 and phosphorous pentachloride is PCl5; and nitrogen dioxide is NO2 and dinitrogen tetroxide is N2O4. The numbering prefixes are given below.

1 2 3 4 5 6 7 8 9 10

mono- di- tri- tetra- penta- hexa- septa- octa- nona- deca-

Use mono- only with the second atom never with the first.

Complex Molecules and Lewis StructuresOften molecules are complex and it is not easy to figure out the dot structure for the molecule. There are a set of rules that allow you to determine the electron dot diagram for molecules. Steps for Writing Lewis Structures of Covalent Compounds from the Chemical Formula

1. Write the atoms in the order given in the chemical formula, place the atom with the lowest electronegativity in the center (usually the first element) and multiple atoms in a chemical formula surrounding the central atom.

2. Determine and add the valence electrons of all atoms. If there are charges add one electron for each negative charge and subtract one electron for each positive charge.

3. Attach all atoms with a bond.4. Subtract 2e– for each bond from the total number of valence electrons (step

2).5. Add pairs of nonbonding electrons to outer atoms. Continue adding pairs

so that the number of pairs on each outer atom are the same or nearly the same.

6. Continue adding electron pairs until you run out of valence electrons (step 4). If you have octets around every outer electron and still have more electrons to add, then add the rest of the electrons to the central atom(s).

7. All atoms should have octets. If there are too few electrons, then share nonbonding outer electrons with the inner atom to make a double or triple bond.

8. Atoms in 3rd or more row can have expanded octets. Hydrogen will only form a single bond without any outer electrons. Boron has only three pairs of bonding electrons around it.

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Chemistry9. Finally, if there are still problems with octets, then move electron pairs if

needed to deficient outer atoms.

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ChemistryExamples

Molecule NO21– SCl4

Step 1 O N O Cl

Cl S Cl Cl

Step 2

O = 6 e– X 2 = 12N = 5 e– (1–) = 1 e– total e– = 18

S = 6 e–

Cl = 7 e– X 4 = 28total e– = 34 e–

Step 3 O •• N •• O

Cl ••

Cl •• S •• Cl •• Cl

Step 4 18 - 4 = 14 34 - 8 = 26

Step 5 & Step 6 •• •• ••

•• O •• N •• O •• •• ••

•• •• Cl •• •• •• •• + 2e– •• Cl •• S •• Cl •• ——➔ •• •• •• •• Cl •• ••

Step 7

•• •• •• •• •• •• •• O •• N •• •• O •• or O ••

•• N •• O •• •• ••

No change

Step 9 No change There are five pairs of electrons around the S, this is an expanded octet.

Molecular ShapeThe shape of a molecule is important in determining its properties; for example, molecular shape affects a molecule’s solubility in water. A theory called Valence Shell Electron Pair Repulsion theory, VSEPR theory, is used to explain the different shapes that molecules adopt and hybrid molecular orbitals (as apposed to the atomic orbitals from Unit 1) explain how electrons can adopt different geometries.

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ChemistryVSEPR theory states that electron pairs of a central atom shift in the molecular geometry until the electron-electron repulsions between both bonding and nonbonding electron pairs are as small as possible. The table below shows the ideal geometries for most simple covalent compounds.

Molecular Shapes from VSEPR Theoryhttp://commons.wikimedia.org/wiki/File:VSEPR_geometries.PNG#filelinks

In the table above, the steric number refers to the number of electron pairs, both bonding and nonbonding, that surround the central atom. The number of nonbonding pairs called lone pairs, is listed across the top of the table. When you have matched the Lewis structure with the steric number and the number of lone pairs, you can find the geometry of the molecule and the name of the molecular geometry.

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S S = Symmetrical AS = Asymmetrical

S AS

S AS AS

S AS AS S

S AS S AS S

ChemistryLewis Structures and the Corresponding Molecular Geometries

Molecule Complete Lewis Structure Molecular Geometry

H2O ••

H •• O •• H ••

without

withlone pair

lone pair

CCl4

•• •• Cl •• •• •• ••

•• Cl •• C •• Cl •• •• •• •••• Cl••

••

AsCl5

Atomic orbitals, which are spherical for s-orbitals and a figure-8 shape for p-orbitals, cannot explain the tetrahedral shape of carbon tetrachloride, CCl4, or the other molecular shapes. The shared valence electrons in covalent compounds require new wave equations to formulate molecular orbitals that explain the various molecular shapes. Molecular orbitals are hybrid orbitals, which are combinations of atomic orbital. Common molecular orbitals are: 2 sp orbitals from combining 1 s-orbital and 1 p-orbital, three sp2 orbitals result from a combination of 1 s-orbital and 2 p-orbitals, four sp3 orbitals where 1 s-orbital and 3 p-orbitals combine, five sp3d orbitals form as 1 s-orbital, 3 p-orbital, and 1 d-orbital mix, and six sp3d2 orbitals, which is a combination of 1 s-orbital, 3 p-orbital, and 2 d-orbitals.

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Chemistry

The 4 atomic orbitalshybridize to become4 molecular orbitals

The four molecular

orbitals are

sp3

The 3 atomic orbitalshybridize to become3 molecular orbitals

The three molecular

orbitals are

sp2

http://commons.wikimedia.org/wiki/File:S-p-Orbitals.svg & http://commons.wikimedia.org/wiki/File:Sp3-Orbital.svg & http://commons.wikimedia.org/wiki/File:Sp2-Orbital.svg

Polar and Nonpolar Covalent CompoundsMany molecules dissolve in water and some do not. For example sugar, vinegar, and vitamin C all dissolve in water, but wax, oil, and vitamin E do not dissolve well if at all. The reason that some substances dissolve in water and some do not is because some molecules are polar and some are not polar.

Polar molecules have a shift in electron density in the molecular orbitals away from the center of the molecule. This means that the electronegativities of the different atoms are unbalanced and the electrons will be more likely to be found over the more electronegative parts of the molecule (recall that electrons move randomly so they will move everywhere, but the pull of the most electronegative nucleus keeps them near that atom more often). Water, H2O, itself is polar. Each H has an electronegativity of 2.1 and oxygen, O, has an electronegativity of 3.5. The electron density of the shared electrons is shifted towards oxygen since it is most electronegative. The molecule H2O has a dipole moment, which is a shift of negative charge towards one part of the molecule (this is not a transfer of electrons like ionic compounds but a shift of electron density).

Water Molecule, H2OElectron Density of the

molecule shifts to oxygen the most electronegative atom

A Dipole Moment represents a substantial shift in charge with the negative end at the

arrowhead

Polar molecules must have two characteristics: 1) The difference in electronegativity (∆EN is difference in electronegativity) between bonding atoms must be between 0.4 and 1.8. When the electronegativity is greater than 1.8 the

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Chemistrybonding is becoming ionic since the electron is being transferred not shared. When the electronegativity is less than 0.4 then the electron density in a bond is not shifting enough to create a dipole moment. 2) The molecule must be asymmetrical so that the pull of electrons is not cancelled out by equal pulls in opposite directions. For example, CCl4, carbon tetrachloride, is nonpolar because, while the bonds are polar with ∆EN = 0.6 (C = 2.5 and Cl = 3.1), the molecule is tetrahedral and symmetrical so the electron density moves towards each Cl in all directions of the molecule. Polar molecules include: CHCl3, NH3, & COH2 and nonpolar molecules include: CH4, PH3, and SCl5. To figure out why these are polar you must first draw the Lewis structure, determine the VSEPR geometry, find the difference in electronegativity and finally decide whether both the 0.4≤∆EN≤1.8 and the geometry is symmetrical are true.

States of Matter.Ionic compounds are solids. The electrostatic forces hold the ions locked into a solid state. But covalent, or molecular, compounds can be gases, liquids, and solids. For a covalent compound to be a solid the individual molecules must attract each other strongly or the molecular mass of the substance must be high. The attraction between molecules is called intermolecular forces and happens in three different ways with different amounts of attraction. When the electron density shifts in a molecule so that there is a positively charged end and negatively charged end, then the molecules can be attracted to each other through the attractive force of the positive and negative charges. These attractive forces are not bonds because they are 100 of times weaker than the molecular and ionic bonds (sometimes referred to as intramolecular forces)

Hydrogen bonds are the strongest intermolecular bond (not inter- is a prefix meaning between). Hydrogen bonds only occur with a select group of atoms. When HF, HOR, or HNR2, where R is H or more of the molecule, are present then H bonding can occur with a O or N in a molecular compound or with HF. The double helix structure of DNA is due to the hydrogen bonds between the nucleotides. Water also has hydrogen bonds that account for many of its unique properties.

Hydrogen bonds in DNA (the dotted lines)

http://commons.wikimedia.org/wiki/File:DNA_chemical_structure.svg

Dipole-dipole forces are the next strongest intermolecular forces. Polar molecules have shifted electron density that creates partially positive and partially negative charges at opposite ends of the molecule. The molecules will attract each other with the positive end attracting the negative end and visa

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Chemistryversa. Unlike ionic compounds the molecules have strong bonding forces in the molecule but the intermolecular forces are easily broken.

These diatomic (two atom) molecules have dipole-dipole attractions. Because of the

high electronegativity of one of the atoms the electron density has shifted to create a

partial negative charge, δ–, (on top on the first molecule) and partial positive charge, δ+. Intermolecular bonding occurs through attraction of the positive and

negative charges.http://commons.wikimedia.org/wiki/File:Polare_atombindung.png

The weakest intermolecular forces are London Forces, also called Van der Waals forces. This intermolecular force is due to temporary dipoles. Since electrons are in constant motion each nonpolar molecules will form positive and negative end over time (the bonding electrons will randomly show up all on one side). A nearby nonpolar molecule will adjust its electron density because negative charges repel and positive charges attract electrons. This induces a dipole in a neighboring molecule so the two have an attraction. This is the mechanism for intermolecular attraction between nonpolar molecules.

Nonpolar molecules have evenly distributed electron density as shown by the first four molecules. But random motion can create a temporary dipole, the first

molecule to the right of the arrow. The molecules near this temporary dipole react

to the change in electron density and create their own dipoles. They have induced dipoles.

http://de.wikipedia.org/w/index.php?title=Datei:Unpolare_atombindung.png&filetimestamp=20070704231926

London forces are stronger when the valence electrons are dispersed widely in large shells. This is because the electrons will create temporary dipoles that are more difficult to remove when the electrons are so distant from the nucleus and so dispersed in the large electron cloud. The most prominent example of the change in the intermolecular forces due to atomic size is the halogen family. Fluorine, F2, and chlorine, Cl2 are both gases. Like all halogens they are nonpolar because the difference in electronegativity is zero. As gases the molecules of the diatomic molecules do not attract each other and the molecules of gas move randomly and distribute themselves throughout the volume of their container. Bromine, Br2, is a liquid; the only other liquid element at typical temperatures and

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δ –

δ +

δ +

δ –

δ +

δ –

δ –

δ +

Chemistrypressures is mercury, Hg. Bromine has electrons filling the 4th energy level. The molecules of bromine have some attraction for each other because the substance is a liquid, in which the molecules move randomly, but they slide between molecules with small attractive forces holding the particles from separating to become a gas. Finally, iodine, I2, is a solid. The temporary dipole is longer-lived with the outer electrons in the 5th energy level and the molecules have longer, stronger attractions. (Some of the changes in the state of matter—from gas to liquid to solid—are due to the larger molecular weight, but temporary dipole changes have the most prominent role).

Differences in Attractive Forces When molecules have strong intermolecular forces they are most likely to be solids. As solids every molecule is locked in place by the strong attractions. As the intermolecular forces are reduced, the molecules may be liquids where the forces keep the particles near to each other, but the forces are not strong enough to stop random motion. Finally, weak intermolecular forces are present in gases that cannot maintain any hold on each other so they move randomly and in their own path filling the volume of their container. So water, H2O, is a liquid despite its small molecular mass, while CH4, methane, is a gas (the natural gas used in stoves and water heaters).

Remember that molecular weight has something to do with whether molecules are solids. Waxes and your skin are essentially nonpolar, but they are solid, because they have large numbers of atoms in each molecule (wax is C15 H31 CO2 C30

H61) so the molecules are massive and have long, long molecular orbitals so induced and temporary dipoles are long lasting.

Properties Affected by Intermolecular Forces.Intermolecular forces affect the physical changes that substances undergo and their state of matter at standard temperature and pressures. State of matter is determined by the type of intermolecular forces. The stronger the intermolecular force the more likely it is a solid or a liquid (generally, only molecules with high molecular masses will be a solid). Boiling point and melting point (the temperature at which a substance boils or melts) is also affected by intermolecular forces, with higher temperatures related to stronger intermolecular forces. Vapor Pressure, which is the pressure of the gas molecules of a pure substance that evaporate from a pure substance at a temperature below the boiling point, is higher for weaker intermolecular forces.

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ChemistryPolyatomic Ions and Their Ionic CompoundsSome molecular compounds are most stable when they have one or more extra electrons added or removed these are polyatomic ions. Most polyatomic ions are negative and one common one is positive, NH4

+, ammonium. These ions will react with elements or will trade places with another ion to create ionic compounds with other ions. Here is a list of common polyatomic ions.

Some Common Polyatomic Ions

1+ 1– 1– 2– 3–

NH4+

ammoniumOH–

hydroxideC2H3O2

– orCH3COO–

acetateCO3

2–

carbonatePO4

3–

phosphate

NO3–

nitrateClO–

hypochloriteSO4

2–

sulfate

NO2–

nitriteClO2

=

chloriteSO3

2-

sulfite

HCO3–

hydrogen carbonate

ClO3–

chlorateCrO4

2–

chromate

CN–

cyanideClO4

–

perchlorateC2O4

2–

oxalate

Polyatomic ions used just like ions from elements, except parentheses are used to keep the molecule together when you have to have more than one polyatomic ion. For example, potassium carbonate has the ions K+ and CO3

2– which leads to a chemical formula of K2CO3. Or, ammonium sulfide has the ions NH4

+ and S2–and the chemical formula (NH3)2S.

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ChemistryExamples of Chemical Formulae of Compounds with Polyatomic Ions

Name of Ionic Compound Ions Chemical Formula

calcium chlorite Ca2+ ClO2– Ca(ClO2)2

aluminum carbonate Al3+ CO32– Al2(CO3)3

iron (III) nitrite Fe3+ NO2– Fe(NO2)3

Chemical Formula Ions Name of Ionic Compound

Na2SO4 Na+ SO42– sodium sulfate

Mg(ClO)2 Mg2+ ClO– magnesium hypochlorite

(NH4)2CrO4 NH4+ CrO4

2– ammonium chromate

Pt(SO4)2 Pt4+ SO42– platinum (IV) sulfate

This should all be reminiscent of writing chemical formulae for elements and for determining the name. All the polyatomic ions have endings different from the -ide used for ionic compounds of elements (sodium chloride). This makes writing the ion for polyatomic ions as simple as looking on a table to match the name of the ion with its chemical formula and charge. The hardest problem is like the last one, where the cation is a transition metal with unknown charge. But the key is to always balance charge with the number of cations versus anions, or when given an unknown cation charge balance the number of electrons by choosing the appropriate charge.

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ChemistrySummaryCompounds are pure substances that are composed of more than one element. Ionic compounds are held together by electrostatic attractions between cations and anions. Most cations are metal elements, which have electron configurations with three or fewer valence electrons, and most anions are nonmetal elements, which have five to seven valence electrons. The charge of cations is positive and the number of electrons lost and the charge of anions is negative and the number of electrons gained. The number of electrons gained or lost leads to a stable, noble gas electron configuration. After electrons are transferred the ions combine to form a crystal lattice. The number of cations and anions is not always one-to-one in the chemical formula, but must be in a ratio that makes the number of electrons lost and gained be equal. This chemical formula of an ionic compound shows the ratio of elements using subscripts to denote the number of elements.

Another type of compound is covalent compounds or molecular compounds. These compounds are composed of molecules, which are individual groups of atoms bonded with covalent bonds. The molecules form when electrons are shared between elements that are usually nonmetals. Elements in covalent compounds share electrons until they have an octet of electrons around them. This may require double bonds or triple bonds that have 2 pairs of electrons or 3 pairs of electrons shared between two elements. Molecules can be either polar or nonpolar and the shape of the molecule is one factor that determines this property. The shape of a molecule can be determined by using VSEPR theory, which holds that both bonding and nonbonding pairs of electrons will tend to move as far apart as possible because of the electron-electron repulsions. If a shape is symmetrical then the molecule is nonpolar. If the shape of a molecule is asymmetrical and the molecule has bonds that have electronegativity differences of 1.8 to 0.4 then the molecule will be polar.

Polar molecules have intermolecular forces called dipole-dipole attractions. These intermolecular forces hold molecules next to each other and result in the substance being a solid or liquid (compared to a nonpolar compound of the same molecular mass). Nonpolar molecules also exhibit a weak intermolecular attraction called Van der Waal’s forces or London forces. But these weak forces do not hold molecules together strongly so the compounds are gases or easily boiled or melted.

Polyatomic ions are molecules with added or removed electrons which makes them anions or cations respectively. They act like other ions, but the chemical formula of polyatomic ions is listed in a table or needs to be memorized.

Metallic compounds are combinations of metal elements as cations with the valence electrons free to move among the cations. These shared valence electrons form a bond for metals called a “sea of electrons”, which is responsible for many of the properties of metals.

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Chemistry

Type of Bond

Metal or nonmetal elements

Share or Transfer electrons

Melting Point & Boiling

Point (high or low)

State of Matter

Ionic metal + nonmetal Transfer high solid

Covalent nonmetal + nonmetal Share low usually

gas, liquid and solid (oxygen, water, sugar)

Metallic metal + metal Share ; a “sea of electrons”

high but mercury is a

liquid

solid (one liquid)

Type of Bond

Hard or Soft (brittle vs. malleable)

Dissolves in Water

Conducts Electricity Other Notes

Ionic hard Often but not always (like

chalk)

yes if it dissolves, yes if it is a

liquidnot as a solid

The positive ion, cation, combines with a negative ion, anion,

to create a neutral compound.

Minerals are ionic.

Covalent soft usually

Nonpolar covalent

compounds do not, but

Polar covalent compounds do

dissolve.

no

The body’s soft tissue is nearly all covalent

compounds. Organic matter is

covalent.

Metallic

soft (malleable, beaten into

sheets & ductile, pulled into a

wire)

no

Yes, as a solid, liquid, or gas (but it doesn’t

dissolve)

Metals don’t form compounds, they form mixtures, called alloys, or they are elements.

Chemical Bonds2. Biological, chemical, and physical properties of matter result from the ability

of atoms to form bonds from electrostatic forces between electrons and protons and between atoms and molecules. As a basis for understanding this concept:

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Chemistrya. Students know atoms combine to form molecules by sharing electrons to

form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

b. Students know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2, and many large biological molecules are covalent.

c. Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction.

d. Students know the atoms and molecules in liquids move in a random pattern relative to one another because the intermolecular forces are too weak to hold the atoms or molecules in a solid form.

e. Students know how to draw Lewis dot structures.f. * Students know how to predict the shape of simple molecules and their

polarity from Lewis dot structures.g. * Students know how electronegativity and ionization energy relate to bond

formation.h. * Students know how to identify solids and liquids held together by Van der

Waals forces or hydrogen bonding and relate these forces to volatility and boiling/melting point temperatures.

Starred standards are non-tested standards on the California Standards TestContributed by Kenneth PringleEdited by Kathleen DuhlFormatted and Wiki Contribution by Christine Mytko

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