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A^ N INTRO TO THE^ D^ ESCRIPTIVE
CHEMISTRY
To this point in our review we’ve worked on methods for establishing a QUANTITATIVE foundation in chemistry.
We can
now manipulate all manner of unit factors to solve problems involving amounts of chemical materials, and while there wassome modest requirement that you understand the nature of the substances involved in the problems, you might just as wellhave applied the techniques you learned to solve problems associated with preparing a good white sauce—make the units canceland you’ve learned the right answer—whatever those numbers meant.
There was nothing particularly relevant to chemistry in the problem solving.^ Well now there is a radical change in material.
For one thing, there aren’t any problems to work. Reading this material is like taking a botany class.
Suddenly our primary focus is on DESCRIBING chemistry. So when you get down to it, in this chapter there is: a whole lot of MEMORIZING going on.
But I will also introduce you to the basic theoretical foundation that^ significantly reduces memorization
by learning the following basic concepts. Concept 1: Periodic Trends Like Metallic Character •^ You learn basic trends in the elements that prompted the creation of a periodic table of those elements.
The first is which elements are metals and which are not? Concept 2:^ Acids Base Chemistry • You will learn what things do and don’t behave as acid and bases. Concept 3: Which Ions dissolve in Water? • You will learn elementary solubility rules for ions in water. Concept 4: Oxidation/Reduction • You will learn to assign oxidation numbers to elements.
The oxidation number is you first indication of
where
ELECTRON DENSITY is located in chemical compounds. Concept 5: Active Metals • You will learn about displacement reactions, which means you will learn to predict what happens when you dump metalsinto water. Concept 6:^ Chemical Nomenclature • You will learn the common way to name binary and ternary ionic compounds.
A hint: don’t memorize compound names, memorize naming rules.
The Periodic Table Long about 1870, Mendeleev was inspired to put together a tabulation of the known elements in such a way as to describe in aperiodic way their physical and chemical properties.
The modern version of Mendeleev’s Periodic Table is an attractive addition to most science lecture halls—it is worth studying during drier moments of lectures if for no other reason than toprepare for chemistry tests.
Families of Elements Note that the table of elements consists of columns (groups or families) of elements in which similar properties are observed.Group IA^ alkali metal
Li, Na, K, Rb, Cs Group IIA^ alkaline earth metals
Be, Mg, Ca, Sr, Ba Group VIIA^ halogens^
F, Cl, Br, I Group 0^ noble or rare gases
He, Ne, Ar, Kr, Xe Periods: Chemical Trends along Rows of Elements You can also look at the table as a collection of horizontal rows called periods:Row 2^
Li, Be, B, C, N, O, F, Ne As we will learn, important trends in chemical and physical properties are observed as we move across a row. For example, .....
Electrolytes: Ions Moving Through Water The chemistry of water is primarily the chemistry of dissolved ions.
Water is a polar solvent with the ability to promote the dissociation of ions. One property of theseions conduct electricity in water. Strong Electrolytes Compounds that dissociate completely in water to form ions are called strong electrolytes.
Examples of these include alkali metal salts and strong acids and bases:+^ NaCl^ Æ^ Na
- Cl + - KNOÆ K+ NO 3 3 + - HCl Æ H+ Cl (^) + - NaOH Æ Na+ OH We assume that essentially 100% of the salts or acids of strong electrolytes dissociate to form ions. Weak electrolytes Some compounds dissociate only slightly, preferring to remain primarily in an undissociated form.
These molecules are referred to as weak electrolytes and include a wide range of sparingly soluble salts as well as weak acids and bases.+AgCl^ Æ^ Ag^
- +^ Cl++^ =CaCO Æ Ca+^ CO 3 3 + - CHCOOH Æ H+^ CH^ COO 3 3
Non-electrolytes Of course, there are molecules that are too tough to have their bonds ripped apart in water.
These are molecules like sugar which possess covalent rather than ionic bonds. Water may be capable of solubilizing these molecules (dissolving them), but noions are formed and electricity is not conducted in solution.SUGAR(s)^
----->^ SUGAR^ (aq)
Note that ions are not formed, but the sugar is dissolved in the water (aq means aqueous.) Acid Base Chemistry What is the big deal with acids and bases? Again, it is because we are working with water, which, as we will see, has its ownchemistry that produces ions like H
+^ -^ and OH, and which as everyone knows, are what Arrhenius called acids and bases. Consequently, we will spend a lot of time looking at what happens with these two ions, the proton or hydronium ion, H
+^ , and
-^ the hydroxide ion, OH.^ In CH 302 we will develop many equations that attempt to determine their concentration in solution,their pH. But again, understand the context, we are working in water, and we are examining electrolytes.
In other +^ -^ environments, Hand OHwould have no significance.
For example, how often do you hear people talk about the pH of gasoline? Bronstead/Lowry Definition Acids and Bases:^ acid: capable of donating a protonbase: capable of accepting a protonThe general equation for the dissociation of an acid is simply:HA^ ---->
+^ - H+^ A^
For example:HCl^ ---->
+^ - H+^ Cl^ orCHCOOH^ ----> H^3
+^ - +^ CH^ COO^3
by an ability to donate a proton or hydroxide, respectively. Whether a compound is a strong or weak acid or base has to do withthe extent to which it dissociates. Strong acids and bases:^ Compounds that are assumed to dissociate completely upon addition to water.
+^ HA Æ H+^
- A
initial^ 100 %^
0%^ 0%
equilibrium^ 0%^
100 %^ 100%
Common Strong Acids^
Anions of These Strong Acids Formula^ Name^
Formula^ Name HCl^ hydrochloric acid^
-^ Cl chloride ion HBr^ hydrobromic acid^ -^ Br bromide ion HI^ hydroiodic acid^ -^ I iodide ion HNOnitric acid^3 -^ NOnitrate ion 3 HClOperchloric acid^4 -^ ClOperchlorate ion 4 HClOchloric acid^3 -^ ClOchlorate ion 3 Hsulfuric acid^ 2SO^4 -^ HSOhydrogen sulfate ion 4 How do we know which compounds are strong acids or bases?
We memorize the table. Weak acids.^ Of course if an acid or base doesn’t dissociate completely, it behaves as a weak electrolyte. Most of the acid staysin the molecular form. Only a small amount dissociates to form a proton and the anion.CHCOOH ----->^3
+^ - H+^ CHCOO^3 initial 100 % 0%^ 0%equilibrium 99% 1 %^ 1% How do we know which acids are weak? Simple. It is any acid that isn’t one of the seven strong acids. Several examples areshown in the table above.
Bases:^ We can describe an analogous collection of strong and weak bases that produce OH
-^ in aqueous solution.^ The strong
bases include the Group IA and IIA metal hydroxides. For example:NaOH^
+^ - -----> Na+^ OHInitial 100 % 0%^ 0%equilibrium 0% 100 %^ 100% Table of Strong Bases LiOH lithium hydroxideCsOH^ NaOH sodium hydroxideKOH potassium hydroxideRbOH rubidium hydroxide cesium hydroxideCa(OH) calcium hydroxide (^2) Sr(OH) strontium hydroxide (^2) Ba(OH) barium hydroxide 2 Weak bases.^ If a base is a weak electrolyte, it does not dissociate significantly in solution.
The most famous weak base is ammonia, NH^3 NH----->^3
+^ - NH+^ OH 4
initial^ 100 %^
0%^ 0%
equilibrium^ 99%^
1 %^ 1%^ Solubility Throw a compound into water.^ Does it dissolve or does it sink to the bottom of the flask?
Whether this happens or not determines the SOLUBILITY of a compound.
A soluble compound dissolves.^
An insoluble compound usually sinks to the bottom of the flask as a PRECIPITATE, although it may form a SUSPENSION in solution. For the most part, we are interestedin what ionic compounds do in water.
How do we know what these ionic materials do in solution?
We learn rules about solubility—see the next page.
There has to be an easier way!! Some cation solubility rules:^ •^ Rule 1:^ All of the alkali metals (K
+^ +^ +^ , Na. etc.) ions plus NH^ are always soluble.^4
-^ Rule 2. Heavy cations (down the periodic table) like Ba
++^ +^ ++^ and Ag^ and Pb^ are insoluble. Some anion solubility rules:^ •^ Rule 3: Anions that are always insoluble except for Rule 1 are mostly the multiply charged species (S
=^ =^ -3^ =^ , CO, PO^ , O,^34
etc.) • Rule 4.^ The anions that are primarily soluble are the conjugate bases of strong acids like Cl
-^ -^ from HCl,^ Br^ from HBr, -^ -^ NOfrom HNO^ from HClO 3 3, ClO^4 etc. 4 These rules aren’t perfect but they are good enough to get an A in CH301 without doing a lot of memorizing.The Best Solubility Rule of All:But the biggest trend of all, the rule of thumb if you only want to remember one thing is this:the singly charged ions are usually solubleandthe multiply charged ions are usually insoluble.The reason for this is something we will learn in CH 302.Solubility Example:We can use this information to decide the types of reactions that occur when we place compounds in solution.Consider the reaction of CaNO
and KCO^3 3. Note from above that nitrates and alkali metals are soluble.
So when these compounds are added to solution, four ions are formed:
-^ ++ +^ CaNOK 3
=CO 3
Note from above that carbonates are generally insoluble, so it is expected that the following precipitation reaction occurs:++^ =^ Ca(aq)^ +^ CO(aq)^3
----->^ CaCO^ (s)^3^ Oxidation Numbers Reactions in which substances undergo changes in oxidation number are called oxidation-reduction reactions or redox reactions.Redox reactions involve the transfer of electrons. Some definitions: •^ oxidation:^ an algebraic increase in oxidation number in which electrons are lost from a compound. •^ reduction:^ an algebraic decrease in oxidation number in which electrons are gained from a compound. •^ oxidizing agent:^ substance that gains electrons and oxidizes other substances by being reduced. •^ reducing agent:^ substance that loses electrons and reduces other substances by being oxidized. Oxidation Rules: Assignment of the oxidation number for a compound follows certain rules:1. The oxidation of a free element is zero.
Na is 0
- The oxidation number of an element in a monatomic ion is the charge on the ion.
+^ Nais +
- The charge of a monatomic ion is usually determined by its group (column on periodic table). See table below.4. In the formula for any compound, the sum of the oxidation numbers of all elements in the compound is zero.
NaCl is 0
- In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge in the ion.-^ NOis -1^3
Wow, what a famous reaction!!There is a special kind of reaction in which one element displaces another element from a compound. For example:Cu(II)SO+^ Zn^4
---->^ ZnSO^ +^ Cu^4 In general, active metals displace less active metals and hydrogen. There are three categories to consider.^ •^ Reactions that displace hydrogen from non-oxidizing acids.^ •^ Reactions involving metals that displace hydrogen from steam.^ •^ Reaction of metals that displace hydrogen from cold water.These results are tabulated below: Type I Elements^
Common^ CommonReduced Form^ Oxidized Forms Li^
+Li Li K^ Displace hydrogen^
+K K
Ca^ from cold water^
2+Ca Ca Na^
+Na Na Note these elements are the ones that explode in water. Type II Elements Mg^
2+Mg Mg Al^
3+Al Al Mn^ Displace hydrogen^
2+Mn Mn Zn^ from steam^
2+Zn Zn Cr^
3+^ 6+Cr Cr , Cr^ Fe^
2+^ 3+Fe Fe, Fe
Type III Elements Cd^
2+Cd Cd Co^ Displace hydrogen^
2+Co Co Ni^ from nonoxidizing acids^
2+Ni Ni Sn^
2+^ 4+Sn Sn , Sn^ Pb^
2+^ 4+Pb Pb , Pb^ Type IV Elements Cu^
+^ 2+Cu Cu , Cu^ Hg^ don’t react^
2+^ Hg Hg , Hg 2 2+ Ag^
+Ag Ag Pt^
2+^ 4+Pt Pt , Pt^ In reading the metal activity table, the metals at the top of the chart displace salts of the less active metals below them. An easyway to think of this, a reaction occurs if the more stable metal (the one lower on the chart) is a reaction product.Again, what are some of the general rules you might use rather than have to memorize everything.The farther to the right, the more reactive the metal.^ •^ Type I: The alkali metals (column 1)explode in cold water^ •^ Type II: The alkali earths (column 2) dissolve in hot water^ •^ Type III: The transition metals dissolve in acid^ •^ Type IV: The coinage metals don’t do anything.Not 100% accurate, but good enough for an A (alright, a low A) in CH301.
sulfur^ S^ sul
sulfide selenium^ Se^ selen
selenide hydrogen^ H^ hydr
hydride fluorine^ F^ fluor
fluoride chlorine^ Cl^ chlor
chloride bromine^ Br^ brom
bromide iodine^ I^ iod
iodide Examples of Binary Compound Names^ Compound^
NameLiBr lithium bromideMgCl magnesium chloride (^2) Li lithium sulfide2SAl aluminum oxide2O (^3) Na sodium phosphide3PMg magensium nitride3N 2 Multiple Oxidation Numbers If only life were this simple.^ Recall that not all elements have a single common oxidation number in ionic form.
The idea of
using a Roman numeral after the name was proposed to name these compounds.
However, there is an older way of naming these ambiguous compounds using -ous and -ic to terminate the metal.
The -ic stood for the higher oxidation state, and -ous was attached to the metal with the lower oxidation number.copper (I) is also
called cuprouscopper (II) is also called cupric
Examplescompound^ older name
modern name FeBr^ ferrous bromide^2
iron (II) bromide FeBr^ ferric bromide^3
iron (III) bromide SnO^ stannic oxide
tin (II) oxide SnOstannous oxide^2
tin (I) oxide TiCl^2
titanium (II) chloride TiCl^3
titanium (III) chloride TiCl^4
titanium (IV) chloride Pseudo Binary Compounds Names (Polyatomic Ions) But what about all the compounds containing polyatomic anions and cations? Ions likehydroxide^ OH
- +ammonium NH 4 -2sulfate SO^4 Well^ these ions behave as simple monatomic anions and cations, so we include them in our naming of binary ionic speciesusing the same rules as above, except that we have to memorize all of the polyatomic ion exceptions. There are only 50 or 100listed below for your memorization pleasure:
Examples of compounds made with pseudobinary ion:^ compound^
nameNasodium sulfate2SO (^4) Cu(NO copper (II) nitrate3) (^2) LiCN lithium cyanideNHCl ammonium chloride 4 So what if it isn’t an ion:
Naming Binary Molecules (Covalent Bonds): Often two non-metals are put together. In these cases neither has a great desire to rip away or give up electrons, so ionic speciesaren’t formed. Moreover, these types of compounds may exist in a variety of oxidation states. Rather than deal with oxidationstate, the number of each atom is listed with a prefix noting the number of atoms of each element.# atoms^2
prefix^ di-^ tri-^ tetra-^
penta-^ hexa-^ hepta-^ octa- Examples of molecules:compound^
name CO^2
carbon dioxide SO^3
sulfur trioxide OF^2
oxygen difluoride P^ tetraphosphorus hexaoxide4O^6 Cl^ 2O^7
dichorine heptaoxide
Ternary compounds (after two comes three) Ternary compounds contain combinations of atoms from three different elements.
Things can get really complicated with the additional element, so we will stick to just a few simple classes.First, in working with ionic compounds there were a variety of polyatomic ions that you were asked to memorize.
These
followed the basic rules of binary ions and are easy to name IF you have them all memorized. Ternary Acids and Their Salts There are a collection of well know ternary acids which consist of hydrogen, oxygen and a nonmetal.The most famous example isH2SO
4, sulfuric acid. Rules for Naming Ternary Acids There are some basic rules for naming such compounds which depends upon the oxidation state of the central atom, XXX :Ternary acid
anion saltdecreasingperXXXic acid^ perXXXate^ decreasingoxidationnumber of X
oxygenson X XXXic^ XXXateXXXous acid^ XXXitehypoXXXous acid^ hypoXXXite
