NOTE PACKET: PERIODICITY AND THE MOLE by x11D89e

VIEWS: 24 PAGES: 15

									Unit 4 notes: BONDING BASICS AND NOMENCLATURE
OXIDATION STATES (common ion charges)
        The elements on the periodic table are all ___________ as atoms thus having an
oxidation state of ___________. When combined in chemical compounds, however,
the elements are no longer considered neutral. The oxidation state determines the
___________ that an element carries when combined with another element. Certain
elements only have one option for oxidation. Their position on the periodic table (due to
their electron configuration) can be used to determine these oxidation numbers.
Group 1 metals:
Group 2 metals:
Group 3-12 metals (transition metals):
Group 13:
Group 14: various choices
Group 15: various choices
Group 16: various choices
Group 17 (halogens):
Group 18 (noble gases):
For all compounds, whether covalent, polar covalent, or ionic, we treat as __________
for counting electrons and for determining oxidation numbers.
        Rule 1: ________ of the oxidation numbers of all the atoms in the chemical
        species equals the __________ on the species.
        Neutral compounds: Sum of oxidation numbers = 0
        Ionic species: Sum of oxidation numbers = charge of the ion
        Rule 2: In Binary Compounds, the more ___________________ element is
        assigned to have a negative oxidation number. (See EN trends.)
        Rule 3: Atoms may have only certain oxidation numbers (see above). The
        maximum is ______________.
Examples:
CO: (Sum will equal 0 since it is a neutral molecule)
O will have a -2 ox. number.
               1C+1O=0            (C?) + (-2) = 0                C? = +2
               Oxidation number of C in CO is +2
               Oxidation number of O in CO is -2 (known)
Cr2O72-: (Sum of all oxidation numbers will equal -2 since it is an ion.)
               2 Cr + 7 O = -2             2(Cr?) + 7(-2) = -2
               2(Cr?) + (-14) = -2         2(Cr?) = +12
                                     Cr? = +6
               Oxidation number of each Cr in Cr2O72- is +6
               Oxidation number of each O in Cr2O72- is -2 (known)




Unit 4 notes                              -1-
OK, Now you try:

CS2                                                          NH4+
NaH                                                          ClO3-
ClO2-                                                        H2SO4
BOND TYPES: We will look at each in more detail
        Ionic bonds- formed from a _________________________ from one atom to
another thus creating charged particles known as ions. The ionic bond is a strong
attraction between a ___________ and an _____________due to the opposite charges.
Ionic bonds are found between ___________ and nonmetals. Compounds containing
ionic bonds are called ____________ compounds. The smallest unit of an ionic
compound is called a ___________________.

      Covalent bonds- formed from a __________________ between two atoms.
Covalent bonds are found between nonmetals. Compounds containing all covalent
bonds are called ______________________ compounds. The smallest unit of a
molecular compound is called a ________________.

Determine if the following contain an ionic or covalent bonds:
1. Be3N2                        4. H2O
2. NaCl                         5. CaF2
3. KOH                          6. NH3

COMPOUNDS:
A compound is a chemical combination of two or more _______________. The overall
compound is ___________, however, depending on the elements involved they may
have regions that are charged. We will learn different ways to name compounds
depending on which type they are. The following represents a basic list of the types of
compounds you will be able to name. Let’s first be able to identify the types.

BINARY IONIC COMPOUNDS: IONIC COMPOUNDS CONTAINING ONLY 2
TYPES OF ATOMS.                                 Ex. NaCl
BINARY MOLECULAR COMPOUNDS: MOLECULAR COMPOUNDS
CONTAINING ONLY 2 TYPES OF ATOMS.              Ex. CO
POLYATOMIC IONIC COMPOUNDS: IONIC COMPOUNDS CONTAINING
POLYATOMIC IONS.                               Ex. Mg(OH)2
ACIDS: COMPOUNDS PRODUCING HYDROGEN IONS.      Ex. HNO3
HYDRATES: COMPOUNDS THAT HOLD WATER WITHIN THE CRYSTAL
STRUCTURE.                                     Ex. CuSO4•5H2O
ORGANIC COMPOUNDS: COVALENTLY BONDED MOLECULES CONTAINING
CARBON BACKBONES.                              Ex. C3H8
Decide what TYPE of compound the following represent:
1.   C6H12O6                            5.   H3PO4
2.   H2O                                6.   C4H7O
3.   KCl                                7.   NO
4.   Na2SO4                             8.   O2
Unit 4 notes                           -2-
IONIC (ELECTROVALENT) BONDING (chapter 15)
This page explains what ionic (electrovalent) bonding is. It starts with a simple picture of
the formation of ions, and then modifies it slightly for basic chemistry purposes.

A simple view of ionic bonding
The importance of noble gas structures
At a simple level, a lot of importance is attached to the electronic structures of noble gases
like neon or argon which have eight electrons in their outer energy levels (or two in the
case of helium). These noble gas structures are thought of as being in some way a
"desirable" thing for an atom to have.
You may well have been left with the strong impression that when other atoms react, they
try to organize things such that their outer levels are either completely full or completely
empty.

  Note: The central role given to noble gas structures is very much an over-simplification. Higher level
  chemistry will have to spend some time later on demolishing the concept!

Ionic bonding in sodium chloride Sodium (2,8,1) has 1 electron more than a stable
noble gas structure (2,8). If it gave away that electron it would become more stable.
Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain
an electron from somewhere it too would become more stable.
The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both
become more stable.




The sodium has lost an electron, so it no longer has equal numbers of electrons and
protons. Because it has one more proton than electron, it has a charge of 1+. If electrons
are lost from an atom, positive ions are formed. Positive ions are sometimes called
cations.
The chlorine has gained an electron, so it now has one more electron than proton. It
therefore has a charge of 1-. If electrons are gained by an atom, negative ions are formed.
A negative ion is sometimes called an anion.
The nature of the bond
The sodium ions and chloride ions are held together by the strong electrostatic attractions
between the positive and negative charges. LDS or Electron Dot Symbols:
The formula of sodium chloride
You need one sodium atom to provide the extra electron for one chlorine atom, so they
combine together 1:1. The formula is therefore NaCl.

 Unit 4 notes                                   -3-
Some other examples of ionic bonding
magnesium oxide




Again, noble gas structures are formed, and the magnesium oxide is held together by very
strong attractions between the ions. The ionic bonding is stronger than in sodium chloride
because this time you have 2+ ions attracting 2- ions. The greater the charge, the greater
the attraction.
The formula of magnesium oxide is MgO.
calcium chloride




This time you need two chlorines to use up the two outer electrons in the calcium. The
formula of calcium chloride is therefore CaCl2.
potassium oxide




Again, noble gas structures are formed. It takes two potassiums to supply the electrons the
oxygen needs. The formula of potassium oxide is K2O.

THE BASIC VIEW OF IONIC BONDING
      Electrons are transferred from one atom to another resulting in the formation of
       positive and negative ions.
      The electrostatic attractions between the positive and negative ions hold the
       compound together.
So what's new? What needs modifying is the view that there is something magic about
noble gas structures. There are far more ions which don't have noble gas structures than
there are which do.



 Unit 4 notes                            -4-
Some common ions which don't have noble gas structures These are all perfectly
stable , but not one of them has a noble gas structure.
                               Iron (III)    Fe3+        [Ar]3d5
                               Copper (II)   Cu2+        [Ar]3d9
                               Zinc          Zn2+        [Ar]3d10
                               Silver        Ag+         [Kr]4d10
                               Lead (II)     Pb2+        [Xe]4f145d106s2
Remember, Noble gases (apart from helium) have an outer electronic structure ns2np6.

Apart from some elements at the beginning of a transition series (scandium forming Sc 3+
with an argon structure, for example), all transition elements and any metals following a
transition series (like tin and lead in Group 4, for example) will have structures like those
above.
That means that the only elements to form positive ions with noble gas structures (apart
from odd ones like scandium) are those in groups 1 and 2 of the Periodic Table and
aluminum in group 3 (boron in group 3 doesn't form ions).
Negative ions are easier to predict charges for. Those elements in Groups 5, 6 and 7
which form simple negative ions all have noble gas structures.
Adapted from site: http://www.chemguide.co.uk/atoms/bonding/ionic.html

 Explore Learning Activity!!!!!Exploration Guide: Ionic Bonds
 Sodium is a soft, silvery metal that reacts violently with water and air. Chlorine is a poisonous
 gas used to kill bacteria in swimming pools and drinking supplies. If ingested, each could cause a
 quick and painful death. But together, sodium and chlorine form salt, a necessary nutrient for all
 living things! How can these two dangerous substances combine to form a harmless crystalline
 solid?
 The answer lies in ionic bonds. Sodium is highly reactive because it has a single outer electron,
 called a valence electron, which is easily lost. Chlorine, on the other hand, has a strong tendency
 to gain an electron. When atoms gain or lose electrons, they become charged atoms called ions.
 Ions with opposite charges are attracted to one another and form ionic bonds. When sodium and
 chlorine form an ionic bond, the result is the stable, non-reactive substance we know as salt.
 Forming Ionic Bonds: In this activity, you will discover how the electron
 arrangements of atoms determine their ability to form bonds.
     1. In the Gizmo™, observe the red sodium atom and blue chlorine atom in the
        SIMULATION pane. Click Play (       ), and observe the electrons orbiting the nucleus of
        each atom.
                1. In the Gizmo, only the outermost electrons, or valence electrons, are shown. How
                   many electrons are visible on the sodium atom?__ On the chlorine atom? __
                2. Sodium has a total of 11 electrons, and chlorine has a total of 17. How many
                   electrons are NOT shown on each atom?___ Where are these electrons?___

 Unit 4 notes                                  -5-
               3. Some atoms give up electrons quite easily, and some do not. Try clicking and
                  dragging an electron away from each atom. Which atom easily gives up its
                  valence electron(s)?____. What does this mean about the ionization energy of
                  metals? __________
    2. Drag the electron from the sodium atom to the chlorine atom. (If the electrons are moving
       too fast, adjust the Speed slider or click Pause ( ).)
               1. The tendency of an atom to attract electrons is called electron affinity which is
                  similar to electronegativity. Which atom has a greater electron affinity?___
               Atoms that have a strong electron affinity (and therefore tend to gain electrons) are
               called nonmetals. Metals have a low electron affinity, which also explains their
               tendency to lose valence electrons.
               2. How many valence electrons does the sodium atom appear to have now?___
               3. How many valence electrons does the chlorine atom have now? ___
               4. Atoms are most stable when their outermost energy level is full of electrons. The
                  third energy level of the sodium atom is now empty. Therefore, the second level is
                  now the outermost, with eight electrons. Are both atoms currently in a stable
                  configuration? ___
    3. A normal sodium atom has eleven protons, each with a positive charge. It also has eleven
       electrons, each negatively charged so that the atom is neutral.
               1. After losing an electron, what do you think is the charge of the sodium atom?___
               2. Before gaining an electron, the chlorine atom had 17 protons and 17 electrons.
                  What is the charge of the chlorine atom now? ___
                Write these down, then check your hypotheses by clicking the Show charge
               checkbox. Were you correct? Charged atoms are called ions.
               3. Ions with opposite charges are attracted to one another, forming an ionic bond. To
                  represent this bond, click on the nucleus of the sodium ion and drag it closer to
                  the chlorine ion. Click Check. What is the name of the substance you
                  created?___________________________
               4. Click Show formula. What is the chemical formula of sodium chloride? ________
               5. Click Show completed compounds. A list of formulas that you have completed
                  correctly will appear here.
    4. Choose Lithium (Li) from the Select a metal menu and Oxygen (O) from the Select a
       nonmetal menu. Count the number of valence electrons in each atom.
               1. Transfer the valence electron from the lithium atom to the oxygen atom. How
                  many valence electrons does the oxygen atom have now? ____ Is the oxygen
                  atom stable? ____ When you have finished, click Check.



Unit 4 notes                                   -6-
               2. The oxygen atom needs 8 valence electrons to reach stability. To provide an
                  additional electron, click Add metal. When the new lithium atom appears, transfer
                  the remaining electron. Is the oxygen atom stable now?___ Why or why not?____
               3. What is the formula of lithium oxide? __________
               4. Click the Show charge checkbox. Notice that each lithium ion has a 1+ charge,
                  while the oxygen ion has a 2− charge. What is the total charge of the three
                  ions?____
               Because the positive charges of each lithium atom repel one another, the lithium
               atoms will be found on opposite sides of the oxygen atom.
    5. Try several combinations of metals and nonmetals. Transfer electrons from the metal
       atoms to the nonmetal atoms. If the metal atom has one or more electrons left over, click
       Add nonmetal to add an additional nonmetal atom. If the nonmetal has additional spaces
       to fill, click Add metal. Extra atoms can be dragged to the Trash icon and eliminated, or
       you can start over by clicking Reset ( ). Once all electrons have been transferred,
       rearrange the atoms to show the attraction between positive ions and negative ions.
               1. What is the formula when the metal has 2 valence electrons and the nonmetal has
                  7? ___________ (For example, what is the formula of magnesium fluoride?)
               2. What is the formula when the metal has 2 valence electrons and the nonmetal has
                  5? ___________ (For example, beryllium and nitrogen.)
               3. Try the remaining combinations of metals and nonmetals. For each that you do,
                  write down a predicted formula based on the numbers of valence electrons. Then,
                  use the Gizmo to check your answer.
Extension Activity: Chemical Families
In the periodic table, elements are grouped by their chemical properties. These groups, also
called families, are determined by the valence electrons.
    1. Obtain a periodic table. (If you don't have one handy, you can print one from the Electron
       Configuration Gizmo.) Each column in the periodic table is a family of elements with
       similar properties.
               1. In the Gizmo, select Lithium (Li) from the list of metals. How many valence
                  electrons does lithium have? _____
               2. According to the periodic table, what other element (from the Select a metal
                  menu) is in the same chemical family? ______
               3. How many valence electrons does this other element have? ____ Use the Gizmo
                  to check your answer. How many valence electrons would you expect for
                  potassium (K), rubidium (Rb), and cesium (Cs)? _____
               4. Try finding other pairs of elements from the same family in the Gizmo. What
                  patterns do you see (explain)?
                  __________________________________________________________________
                  __________________________________________________________________

Unit 4 notes                                  -7-
               5. Notice that the elements in group 18 of the periodic table are not found in this
                  Gizmo. These elements are called noble gases and do not readily react to form
                  compounds. Based on the periodic table and what you have seen in the Gizmo,
                  how many valence electrons do most noble gases have? ____ Why don't these
                  gases form ionic bonds? ____
    2. When Dmitri Mendeleev first published his periodic table in 1869, nothing was known
       about valence electrons. Instead, Mendeleev used chemical properties to group elements
       together. In the Gizmo, create a compound of lithium and oxygen.
               1. What is the formula of lithium oxide? _____
               2. Create a compound of sodium and oxygen. What is the formula of sodium
                  oxide?____ How is this formula similar to lithium oxide? ______________
               3. Create a compound of beryllium and oxygen. Based on this formula, what do you
                  think the formula of magnesium oxide will be? ________ Use the Gizmo to check
                  your answer.
               4. Create a compound of aluminum and oxygen. What is the formula of aluminum
                  oxide? ______What would you predict the formula of gallium oxide to be? _____
               5. Based on the patterns you observed, what is the formula of strontium bromide (Sr
                  and Br)?______ Gallium phosphide (Ga and P)?_____ Cesium telluride (Cs and
                  Te)? ______
               6. In general, what do elements from the same family have in common? ______
NOMENCLATURE:
PREREQUISITE SKILLS (what you absolutely have to know before you can do this
unit!)
            Element names…see list
                                                            +
            Common ion charges …ex. A sodium ion = Na
            Polyatomic ions (and charges)…see list
            Transition metal charges…use Roman Numerals
            Ionic versus molecular compounds
            Recognize acids and hydrates
NAMING BINARY IONIC COMPOUNDS
       First name -cation name
       Last name -anion name with “ide” ending
              Don’t use prefixes (mono, di, tri etc…)
              Ex. NaCl                          sodium chloride
You try:
       1. KF
       2. MgCl2
       3. Na2S
       4. CaO
WRITING BINARY IONIC FORMULAS
       Criss-Cross method
              Write cation with charge
              Write anion with charge
              Criss-cross and reduce to find subscripts
Unit 4 notes                                   -8-
       Ex. aluminum oxide                         zinc hydride
             Al 3+ O 2-                           Zn 2+ H -
             Al2O3                                ZnH2
You try:
   1. calcium fluoride                      4. silver chloride
   2. aluminum nitride                      5. strontium sulfide
   3. cadmium nitride                       6. zinc bromide

BUT, WHAT IF...
       What if there is more than one possible cation?
Remember, copper, manganese and many of the transition metals have more than one
possible charge. You can’t say copper chloride, for example, you need to tell which
copper. Use Roman numerals to indicate the charge. (hint…use the anion and/or
reverse the criss-cross method to help you determine the charge)
       Ex.    FeO          iron (II) oxide
              Fe2O3        iron (III) oxide
You try:
   1. chromium (II) fluoride                                3. CuO
   2. iron (III) sulfide                                    4. Mn2O3

NAMING COMPOUNDS CONTAING POLYATOMIC IONS
       First name -cation name
       Last name -anion name
              Ex.  NH4Cl       ammonium chloride
                   Na2SO3      sodium sulfite
You try:
   1. K3PO4                          4. NaNO3
   2. Hg2O                           5. Sr(ClO4)2
   3. Ca(lO)2                        6. LiOH

WRITING FORMULAS WITH POLYATOMIC IONS
Criss-Cross method
              Write cation with charge
              Write anion with charge
              Criss-cross and reduce to find subscripts
       Ex. aluminum sulfate
              Al 3+ SO4 2-
              Al2(SO4)3
You try:
   1. calcium hydroxide                  4. ammonium chloride
   2. copper (II) nitrate                5. strontium sulfite
   3. iron (II) hydroxide                6. chromium (III) nitrate

IONIC COMPOUND TRAITS TO KNOW:
               o Crystalline solids
               o Conduct electricity when dissolved or melted
               o Solid at room temp with very high melting points (around 800◦C)

Ionic practice game link (also on the moodle): http://chemistrygeek.com/rainbow/launch.htm
Unit 4 notes                              -9-
      COVALENT BONDING - SINGLE BONDS
      A simple view of covalent bonding
      The importance of noble gas structures
      As well as achieving noble gas structures by transferring electrons from one
      atom to another as in ionic bonding, it is also possible for atoms to reach
      these stable structures by sharing electrons to give covalent bonds.
      Some very simple covalent molecules
      Chlorine (diatomic Cl2)
      For example, two chlorine atoms could both achieve stable structures by
      sharing their single unpaired electron as in the diagram.




      The fact that one chlorine has been drawn with electrons marked as crosses
      and the other as dots is simply to show where all the electrons come from. In
      reality there is no difference between them.
      The two chlorine atoms are said to be joined by a covalent bond. The reason
      that the two chlorine atoms stick together is that the shared pair of electrons
      is attracted to the nucleus of both chlorine atoms.
      Hydrogen




      Hydrogen atoms only need two electrons in their outer level to reach the
      noble gas structure of helium. Once again, the covalent bond holds the two
      atoms together because the pair of electrons is attracted to both nuclei.
      Hydrogen chloride




      The hydrogen has a helium structure, and the chlorine an argon structure.




Unit 4 notes                            - 10 -
      Covalent bonding
      Most of the simple molecules you draw do in fact have all their atoms with
      noble gas structures. (Octet rule)
      For example:




      Even with a more complicated molecule like PCl3, there's no problem. In this
      case, only the outer electrons are shown for simplicity. Each atom in this
      structure has inner layers of electrons of 2,8. Again, everything present has a
      noble gas structure.




      Cases where the simple view throws up problems
      Boron trifluoride, BF3




      A boron atom only has 3 electrons in its outer level, and there is no
      possibility of it reaching a noble gas structure by simple sharing of electrons.
      Is this a problem? No. The boron has formed the maximum number of bonds
      that it can in the circumstances, and this is a perfectly valid structure.
      Energy is released whenever a covalent bond is formed. Because energy is
      being lost from the system, it becomes more stable after every covalent bond
      is made. It follows, therefore, that an atom will tend to make as many
      covalent bonds as possible. In the case of boron in BF3, three bonds is the
      maximum possible because boron only has 3 electrons to share.

Unit 4 notes                            - 11 -
Explore Learning Activity! Exploration Guide:Covalent Bonds

"I'll have my bond, speak not against my bond, I have sworn an oath that I will have my bond,"
said Shylock in Shakespeare's The Merchant of Venice1. Though Shylock was not an atom of
fluorine, he could be speaking for one. His words ring true for fluorine and its strong inclination
to form chemical bonds.
Single fluorine atoms are even more unstable than Shylock's character. In fact, single fluorine
atoms are never found alone in nature. Fluorine atoms have a strong tendency to form bonds and
will react strongly with other substances to form ionic or covalent compounds.
An atom is most stable when it has a full outer energy level, or valence. For most atoms, a full
outer energy level contains 8 valence electrons. In a covalent bond, two atoms share a pair of
electrons. The two electrons in that pair act as if they are in the outer energy level of both atoms
at the same time, helping each atom to achieve a full outer energy level.
Creating Covalent Molecules
In this Gizmotm, you will build covalent molecules. You will choose a molecule to build, and
then share two electrons between a pair of atoms to create a covalent chemical bond.
    1. In the Gizmotm, check that Fluorine is selected from the Select a substance dropdown
       menu at top right. From the left fluorine atom on the SIMULATION pane, choose an
       electron from the outer energy level, also known as a valence electron. Click on the
       electron and drag it to the other fluorine atom. Click Play ( ).
               1. Observe what happens, and write down any changes you notice in the orbits of the
                  electrons. How many electrons are in the shared orbit (the oblong one)? ______
               2. Click Pause ( ), and move an electron from the outer energy level of the right
                  atom to the left atom. Click Play. How many electrons are in the shared orbit
                  now? _____
               3. An atom is most stable when its outer energy level is filled with electrons. For
                  most atoms, the outer energy level can hold 8 electrons. Look at the left-hand
                  fluorine atom. How many electrons are orbiting only the left-hand atom's nucleus
                  (on the circular path)? Including the shared electrons, how many total electrons
                  are orbiting the left-hand nucleus?
               4. Count the total electrons orbiting the right-hand fluorine nucleus. Do both atoms
                  have a full complement of 8 orbiting electrons? ____ Click Check to verify that
                  you have correctly created a molecule of fluorine (F2). In reality, the orbits of
                  electrons are more complex than what is shown, but the effects on valence are the
                  same.
               5. A useful way to represent covalent bonds is a Lewis diagram. In the
                  DESCRIPTION pane, select the Show Lewis diagram checkbox. The diagram
                  represents a stable fluorine molecule. How many dots are drawn around the left-
                  hand F? ____ What do you think these dots represent? ______________________
                  What is represented by the line between the two F's? _______________________



Unit 4 notes                                 - 12 -
    2. Click Pause, and choose Hydrogen from the Select a substance menu. Hydrogen is a
       unique element because it forms stable covalent bonds with only two electrons orbiting
       the nucleus. Be sure the Show Lewis diagram box is checked.
               1. Examine the Lewis diagram in the DESCRIPTION pane. Compared to the
                  diagram of the fluorine molecule, what seems to be missing? _____________
                  How many electrons are represented by the line between the two H's? ____
               2. Drag electrons between the atoms to create a hydrogen molecule, H2, and click
                  Play. Click Check to test your configuration. If necessary, repeat the process until
                  you are correct. (You can click Reset ( ) to start over.)
               3. How many electrons are orbiting each hydrogen atom? ____ How is that different
                  from what you observed when you built a fluorine molecule?
                  ______________________________________________________________
    3. Uncheck the Show Lewis diagram box. Click Pause, and choose Oxygen from the Select
       a substance menu. Like fluorine and hydrogen, oxygen is a diatomic element, meaning
       that it most commonly occurs as two atoms combined into a covalent molecule (O2).
               1. Observe the atoms displayed in the SIMULATION pane. How many electrons
                  are in orbit around each atom? ______ How many additional electrons does each
                  oxygen atom need to share to make the atom stable? _______
               2. Transfer electrons until both atoms are orbited by 8 electrons, and check your
                  answer. How many electrons are shared? ______
               3. Each pair of shared electrons represents a single covalent bond. How many
                  covalent bonds are there in a molecule of O2? ______
               4. Based on the molecule, draw what you think the Lewis diagram would look like.
                  Click Show Lewis diagram to check your answer.
    4. Build covalent bonds and create stable molecules for the remaining substances. As you
       build, remember that each covalent bond must be symmetrical; it must have one electron
       from each atom that shares the bond. Make sure each atom has enough electrons orbiting
       it to fill its outermost energy level. (Challenge: Try to draw the Lewis diagram before you
       start transferring electrons, and then click Show Lewis Diagram to check your answer.)
    Draw the LDS for these below:




1
 William Shakespeare (1564–1616), British dramatist, poet. Shylock, in The Merchant of Venice,
act 3, sc. 3, l. 4–5.

Unit 4 notes                                  - 13 -
NAMING BINARY MOLECULAR COMPOUNDS
     First name -element name with prefix for subscript
     Last name -element name with prefix for subscript and end with “ide”

                 Ex.   CO          carbon monoxide
                       CO2         carbon dioxide
You try:
1. P2O5                                               3. N2O
2. SCl6                                               4. CCl4

WRITING BINARY MOLECULAR FORMULAS
     Use prefixes to indicate how many of each.
           Ex. silicon dioxide        SO2

        1.     dinitrogen monoxide
        2.     sulfur monoxide
        3.     pentaphosphorus decaoxide
        4.     trisulfur hexachloride
        5.     silicon trifluoride

Prefixes:
      1-mono               2-di                 3-tri               4-tetra
      5-penta              6-hexa               7-hepta             8-octa
      9-nano               10- deca
*mono only used for last name

NAMING ACIDS
BINARY ACIDS
       First name -hydro – root name – ic
       Last name -acid
              Ex. HCl           hydrochloric acid
“OXYACIDS” (with oxygen)
       First name -anion name with “special” ending (ic and ate, ous and ite)
       Last name -acid
              Ex. H2SO4 sulfuric acid
                  H2SO3 sulfurous acid
You try:
   1. HF                              4. H3PO4
   2. HNO3                            5. H2S
   3. HBrO2                           6. HClO

WRITING FORMULAS FOR ACIDS
       Look at name carefully. If it has hydro, then it does not contain oxygen. If it
       doesn’t have hydro, then it does contain oxygen. Look at the anion and
       determine its name then write the acid accordingly. Use criss-cross to help with
       charges and subscripts!
             Ex.     nitric acid                  HNO3
                     perchloric acid              HClO4
                     Hydrobromic acid HBr
You try:
Unit 4 notes                           - 14 -
    1. hydrofluoric acid                 4. hypochlorous acid
    2. phosphoric acid                   5. hydrosulfuric acid
    3. acetic acid                       6. perbromic acid

NAMING HYDRATES
       First name the compound, then add the prefix for the number of waters attached.
              Ex. CuSO4•5H2O
                      copper (II) sulfate pentahydrate
You try:
   1. Na2SO4 • 10 H2O
   2. MgSO4 • 7 H2O
   3. Na2CO3 • 10 H2O
   4. CaSO4 • 2 H2O
WRITING FORMULAS FOR HYDRATES
       First use the name to write the formula of the compound, then write a dot and the
number of waters attached.
You try:
   1. manganese (II) chloride tetrahydrate
   2. cobalt (II) sulfate hexahydrate




Unit 4 notes                          - 15 -

								
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