Molecules and Compounds Chapter Three. Wolpa/Advanced Placement Chemistry Molecules Molecules are groups of atoms chemically bonded together. Molecules

Molecules and Compounds

Chapter Three

Wolpa/Advanced Placement Chemistry

Molecules

• Molecules are groups of atoms chemically bonded together.

• Molecules may be elements (the diatomic, sulfur and phosphorus) or compounds.


Allotropes/Molecular Elements

• Hydrogen, fluorine, chlorine, bromine, iodine and oxygen exist as diatomic molecules – HOFBrINCl’s

• Phosphorus exists as a tetratomic molecule• Some elements exist in a variety of forms• Carbon: graphite; diamond; buckminsterfullerine• Phosphorus - red and white

• Sulfur - S6 and S4


Molecules and compounds

• Molecular compounds - molecules containing atoms from two or more different elements

• Covalent bonds - the force holding the atoms together in a molecular compound


Formulas

• A compound is represented by using the symbols for the elements of which it is composed

• Subscripts are used to indicate how many atoms of a particular element exist in the compound

• If there is only one atom of a particular element, the one is assumed


Formulas, con’t

• Changing the subscripts changes the compound– consider H2O and H2O2

• Two different compounds can, however, share the same chemical formula– dimethyl ether and ethyl alcohol both have the

formula C2H6O


Ethyl alcohol and dimethyl ether

• Ethyl alcohol on the left; dimethyl ether on the right• These species are termed geometric isomers• Formulas that show the order and arrangement of

specific atoms are known as structural formulas


Electrical nature of matter

• Electrostatic forces– attraction between opposite charges– repulsion between same charges

• Charged atoms or molecules are known as ions– cations - positively charged– anions - negatively charged


Ionic Compounds

• Compounds consisting of ions are known as ionic compounds.

• The forces holding them together are called ionic bonds


How charged species arise

• Neutral atoms and molecules have the same number of protons and electrons

• Cations have more protons than electrons

• Anions have more electrons than protons


Formulas of Ionic compound

• Formula unit - simplest whole-number ratio of ions in an ionic compound

• For example: Ca2+ + Br-

– you need to have the resulting formula be electrically neutral

– so two Br- are needed for each Ca2+

– the resulting formula is CaBr2


Types of Ionic Compounds

• Ionic compounds will always consist of one of the following combinations: a metal and a nonmetal, a polyatomic ion and a nonmetal, a metal and a polyatomic ion or two polyatomic ions. Ionic compounds can be distinguished from molecular compounds by the kinds of elements they contain.


Polyatomic ions

• cations or anions consisting of groups of atoms that are covalently bonded to each other

• examples are NO3-, SO4

2-

• when more than one appears in a formula unit, the polyatomic ion is put in between parentheses, and a subscript is used to indication the number of the ions that appear in the formula unit

• example: Ba(ClO3)2


Guidelines for determining if a compound is Ionic

• Metals almost always form positive ions and form ionic compounds

• Nonmetals form monatomic ions in ionic compounds only when combined with a metal.

• It is difficult to predict when the metalloids form ions

• The farther apart two elements are in the periodic table, the more likely they are to from an ionic compound on reaction


Ionic Crystal Lattice

• Ionic compounds are generally solids and have their ions arranged in extended three-dimensional networks. This regular array of positive and negative ions is called a crystal lattice.


Properties of Ionic Compounds

• High melting points that correlate with charges on ions

• Most ionic solids do not conduct electricity but molten ionic compounds do.

• Most ionic compounds dissolve in water


Properties cont.

• Solutions of ionic compounds in water conduct electricity (electrolytes)

• In ionic substances, each ion has its own characteristics, and these are different from the characteristics of the atom from which the ion was derived (NaCl)


Binary compound nomenclature

• Five types of binary compounds– Metals exhibiting only one oxidation state

forming a compound with a nonmetal or polyatomic ion

– Metals exhibiting two or more oxidation states forming a compound with a nonmetal or polyatomic ion

– Ammonium ion with nonmetal– Two polyatomic ions– Compounds of nonmetals and nonmetals


Ionic Compound Nomenclature


Metals with only one oxidation state

• Groups of metals with only one common oxidation state– alkali metals - +1– alkaline earths - +2– Zn - +2– Al - +3

• All other metals can exhibit more that one oxidation state


Anions in negative oxidation states

• Nonmetallic anions usually exhibit one negative oxidation state– halogens -1– chalcogens -2– N, P -3


Forming compounds

• Metal and nonmetal combine to neutralize charge

• Consider - Al3+, O2-

– cross multiply charges

– 2 Al3+ + 3 O2- = Al2O3


Naming binary compounds

• Use name of metal with no changes

• Change the name of the anion by taking the “stem” and add the suffix -ide

• Examples– NaCl - sodium chloride

– MgCl2 - magnesium chloride


Metals with multiple oxidation states

• Two systems: Stock and “IUPAC”

• IUPAC system– metal name and the oxidation state in Roman

numbers in parenthesis

– Fe2+ = iron(II)

• Form compound by balance charge of metal with correct number of nonmetals– CoCl3 = cobalt(III) chloride


Stock System Nomenclature

• Metals in multiple oxidation states usually have one or two common oxidation states

• First row transition metals are +2 and +3 (except Cu2+ and Cu+)

• use -ous suffix for lower common oxidation state

• use -ic suffix for higher common oxidation state


Examples

• CoCl3 - cobaltic chloride

• NiCl2 - nickelous chloride

• For metals with Latin names, use them

• CuCl - cuprous chloride

• FeBr3 - ferric bromide


Binary Molecular Nomenclature


Nonmetals + nonmetals

• Name nonmetal further to the left of the periodic table first with no changes

• Name nonmetal further to the right of the periodic table second with the -ide suffix

• Use Greek prefixes to indicate the number of each one


Greek prefixes

Number Prefixes

1 mono

2 di

3 tri

4 tetra

5 penta

6 hexa

7 hepta


Examples

• N2O3 - dinitrogen trioxide

• CO2 - carbon dioxide

• P2O5 - diphosphorus pentoxide


Oxy anions

• anions composed of oxygen and another elements

• other elements can be a metal or a nonmetals

• Examples– SO4

2-, NO2-, MnO4

-


Naming

• Need common oxidation states– most common oxidation state for nonmetals is

the group number (except for the halogens)– next most common oxidation state is the

group number minus one

• use -ate suffix for higher oxidation state and -ite suffix for next higher oxidation state


Examples

• SO42- - sulfate

• SO32- - sulfite

• NO3- - nitrate

• NO2- - nitrite

• Salts with these oxyanions– Na2SO4 - sodium sulfate

– KNO3 - potassium nitrate


Acids

• Binary acids– name begins with hydro– then add stem of nonmetal plus -ic– end with acid

• Examples– HCl - hydrochloric acid

– H2S - hydrosulfuric acid


Oxyacids

• Take oxyanion suffix and convert– change -ate to -ic– change -ite to -ous

• Do not use hydro- in the beginning

• Examples– H2SO4 - sulfuric acid

– H2SO3 - sulfurous acid


Hydrates

• Some ionic compounds can have water molecules attached within the structure

• These compounds are termed hydrates and have properties distinct from the unhydrated form


Naming Hydrates

• Hydrates are named by naming the ionic compound and then using a Greek prefix to indicate the number of water molecules followed by the word hydrate


Chemical Formulas

• Molecular compounds– chemical formula represents a discrete

molecular unit (e. g. CO2)

• Ionic compounds– chemical formula represents a formula unit

(the whole number ratio of cations to anions; e. g. K2SO4)


Empirical Formula

• Simplest whole number ratio of atoms in the compound

• All ionic formulas are empirical

• Molecular formulas are either equal to the empirical or a whole number multiple


The composition of compounds

• Mole composition is the number of moles of each of the elements that make up the compound– CO2 - one mole of C and two moles of O

• Mass composition is the mass of each element in the compound– CO2 - 12.0 g of C and 32.0 g of O


Percent composition

• mass of each element per 100 mass units of compound– in 100 g of NH3, there is 82.0 g of N

– therefore, the mass percentage of N is 82.0% N


CO2

• Calculation of % composition of carbon dioxide requires determining the number of grams of each element (C and O) in one mole


Determination of Empirical formula

– Convert percent composition to an actual mass

– Convert mass to moles of each element– Find the whole number ratio of the moles of

different elements


Laughing gas

• Contains 63.6% N and 36.4% O• Assume 100 g of substance, so you have 63.6 g

of N and 36.4 g of O

• Calculation gives an empirical formula of N2O



Empirical Formula from Experimental Analysis

Combustion Method

• Combust compound in oxygen. Carbon is converted into carbon dioxide, hydrogen is converted into water, remaining element is found by difference.


Example

An acetic acid sample with a mass of 1.000 g combusts to give 1.466 g CO2 and 0.6001 g H2O. The compound is known to contain C, H, and O.

• 0.4001 g C from CO2

• 0.0673 g H from H2O

• 0.533 g O by difference



Empirical Formula From Experiment

• Actual measurement of masses is determined from experiment.

• Mass is converted to moles

• Simplest whole number ratio is determined



Molecular formula

• The actual number of each atom in a formula unit

• Consider acetylene and benzene– both have the empirical formula CH

– acetylene is actually C2H2

– benzene is actually C6H6


Molecular Formula from Empirical

Molecular formula must be integral multiple of empirical formula therefore the mass of the molecular formula must be the same integral multiple of the mass of the empirical formula.


Example

• Acetic acid

• mass of molecular = 60 g/mol

• mass of empirical formula = 30 g/mol

• ratio = 2

• empirical formula = CH2.O

• molecular formula = C2H4O2

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Molecules and Compounds Chapter Three. Wolpa/Advanced Placement Chemistry Molecules Molecules are groups of atoms chemically bonded together. Molecules