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Chapter 3 Molecules, Compounds and Chemical Equations

Chapter 3 Molecules, Compounds and Chemical Equations

3.1 3.2 3.3 3.4 3.5 3.6 3.7 3.8 3.9 3.10 3.11 3.12

Hydrogen, Oxygen and Water Chemical Bonds Representing Compounds: Chemical Formulas and Molecular Models An Atomic Level View of Elements and Compounds Ionic Compounds: Formulas and Names Molecular Compounds: Formulas and Names Summary of Inorganic Nomenclature Formula Mass and the Mole Concept for Compounds Composition of Compounds Determining a Chemical Formula from Experimental Data Writing and Balancing Chemical Equations Organic Compounds

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Section 3.1 Hydrogen, Oxygen and Water

Mixtures vs Compounds •  When two or more elements combine to form a compound an entirely new substance results. –  Balloon contains a mixture of hydrogen (H2) and oxygen gas (O2) –  Glass contains a compound composed of hydrogen and oxygen (H2O)

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Section 3.1 Hydrogen, Oxygen and Water

Properties of Hydrogen, Oxygen and Water

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Section 3.1 Hydrogen, Oxygen and Water

Sodium Chloride NaCl When two or more elements combine to form a compound

an entirely new substance results.

Sodium (Na) = Soft reactive metal Chlorine (Cl) = toxic gas

Sodium Chloride (NaCl) = Tasty condiment/preservative 5

Section 3.1 Hydrogen, Oxygen and Water

Ionic vs Covalent Bonds •  Compounds form through Bond formation •  2H2 + O2 → 2H2O

covalent bond

•  2Na + Cl2 → 2NaCl

ionic bond

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Section 3.2 Chemical Bonds

Ionic Bonds • 

Ionic bonds form when one atom loses electrons and another atom gains electrons. –  Metal atoms tend to lose electrons and become positive cations –  Nonmetal atoms tend to gain electrons and become negative anions

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Section 3.2 Chemical Bonds

Ionic Bonds • 

• 

The oppositely charged ions attract one another by electrostatic forces and form an ionic bond The result is an ionic compound, which (in the solid phase) is composed of a lattice – a regular three dimensional array of cations and anions

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Section 3.2 Chemical Bonds

Covalent Bonds •  Covalent bonds form when two atoms share electrons –  Form between non-metals – neither are willing to give up electrons so in order to form a bond – they share.

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Section 3.2 Chemical Bonds

Covalent Bonds •  Water (H2O), ammonia (NH3), and Methane (CH4) all have covalent bonds. A covalent bond is a shared pair of electrons (represented here by a line). Notice covalent compounds are individual units (molecules)

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Section 3.2 Chemical Bonds

Why do bonds form? •  Bonds form to lower the potential energy of the atoms –  In the process each atom ends up with an octet of electrons (8 electrons in outer shell) –  This is the result –  Not the driving force

•  A major driving force in nature is to increase stability (lower potential energy).

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Section 3.3 Representing Compounds: Chemical Formulas and Molecular Models

Types of Chemical Formulas •  Empirical formula –  Simplest – tells what types of atoms are present in a compound –  Empirical formula for hydrogen peroxide is HO

•  Molecular Formula –  Give the actual number of atoms in a compound –  Molecular formula for hydrogen peroxide is H2O2

•  Structural Formula

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Section 3.3 Representing Compounds: Chemical Formulas and Molecular Models

Types of Chemical Formulas •  Structural Formula –  Can be simple (L)

–  Can indicate geometry of the compound (R)

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Section 3.3 Representing Compounds: Chemical Formulas and Molecular Models

Molecular Models •  Ball and Stick Models –  Atoms are balls and bonds are sticks

•  Space filling Models –  Atoms fill the spaces between each other –  Probably the most “accurate” way to imagine a molecule if it was blown up to a size we could see

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Section 3.4 An Atomic-Level View of Elements and Compounds

Elements vs Compounds • 

Pure substances can be elements or compounds (Ch 1)

H2, Ca, Ag

H2O, CaCl2

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Section 3.4 An Atomic-Level View of Elements and Compounds

Elements vs Compounds • 

Elements and Compounds can be further categorized based on the types of units that compose them

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Section 3.4 An Atomic-Level View of Elements and Compounds

Elements vs Compounds • 

Elements can be atomic (exist as individual atoms) or molecular (oxygen is a diatomic molecule). There are even polyatomic elements (next slide).

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Section 3.4 An Atomic-Level View of Elements and Compounds

Molecular Elements •  H2, N2, O2, F2, Cl2, Br2 and I2 •  P4 •  S8, Se8 •  You should learn the 7 diatomic elements 18

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Section 3.4 An Atomic-Level View of Elements and Compounds

Elements vs Compounds • 

• 

• 

Compounds can be molecular (H2O) or ionic (NaCl) The smallest unit of a molecular compound is the molecule The smallest unit of an ionic compound is a formula unit 19

Section 3.4 An Atomic-Level View of Elements and Compounds

Molecules vs Formula Units •  A water molecule is a discrete (individual) stable unit that consists of three atoms tightly bound together. –  Covalent bonds are very strong •  A sodium chloride formula unit is the smallest unit of a sample of NaCl. –  Does not exist on its own as a stable discrete unit. –  Single ionic bond is very weak 20

Section 3.5 and 3.6 Ionic and Molecular Compounds

Types of Compounds • 

Binary Ionic Compounds – 

Metal—nonmetal

NaCl, CaBr2

•  Ionic Compounds with Polyatomic Ions −  Ions with more than one atom

•  Hydrated Ionic Compounds • 

NaNO3, NH3Cl MgSO47 H2O

Binary Covalent Compounds – 

Nonmetal—nonmetal

CO2, H2O, NH3 21

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Section 3.5 Ionic Compounds: Formulas and Names

Writing Formulas for Ionic Compounds •  Binary Ionic Compounds form between a metal and a nonmetal –  Metal forms a cation –  Non metal forms an anion

•  All compounds are charge neutral •  For these two reasons it is fairly simple to reason out the formula of most ionic compounds just by knowing the identity of the component elements. 22

Section 3.5 Ionic Compounds: Formulas and Names

Writing Formulas for Ionic Compounds •  If you have sodium (Na) and chlorine (Cl) –  Na forms a Na+ cation –  Cl forms a Cl– anion –  They combine 1:1 to form NaCl

•  If you have calcium (Ca) and chlorine (Cl) –  Ca forms a Ca2+ cation –  Cl forms a Cl– anion –  They combine 1:2 to form CaCl2

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Section 3.5 Ionic Compounds: Formulas and Names

Learning Check •  Write formulas for ionic compounds formed from the following elements 1. Sodium and Oxygen 2.  Calcium and bromine 3.  Aluminum and Sulfur

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Ionic Compounds •  Ionic compounds have common names and systematic names •  NaCl is sodium chloride or table salt (common name) •  NaHCO3 is sodium bicarbonate or baking soda (common name) •  We will learn the systematic naming system

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Ionic Compounds •  Binary Ionic compounds (metals and non metals) •  Metals are of two types –  Metals that only make one type of cation •  Groups I -3

–  Metals that can from more than one type of cation •  Transition Group Metals and Main Groups 4 and above (basically everything else)

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Binary Ionic Compounds Containing a Metal that Forms Only One Type of Cation 1. The cation is always named first and the anion second. 2. Cation takes its name from the name of the parent element. 3. Anion is named by taking the root of the element name and adding –ide.

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Binary Ionic Compounds Containing a Metal that Forms Only One Type of Cation KCl

Potassium chloride

MgBr2

Magnesium bromide

CaO

Calcium oxide

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Section 3.5 Ionic Compounds: Formulas and Names

Concept Check Write the names of the following compounds. A. MgO B. Al2S3 C. MgF2

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Binary Ionic Compounds Containing a Metal that Form More than One Type of Cation •  •  •  • 

Metals in these compounds form more than one type of positive ion Charge on the metal ion must be specified. Roman numeral indicates the charge of the metal cation. Transition metals and large metals (period IV and higher usually require a Roman numeral • 

Exceptions: Ag+, Cd2+, Zn2+ (form only a single ion) 30

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Binary Ionic Compounds Containing a Metal that Form More than One Type of Cation 1. The cation is always named first and the anion second. 2.  A monatomic cation takes its name from the name of the parent element. 3.  Charge of cation in roman numerals (in parenthesis) 3. A monatomic anion is named by taking the root of the element name and adding –ide. 31

Section 3.5 Ionic Compounds: Formulas and Names

Naming Binary Ionic Compounds Containing a Metal that Form More than One Type of Cation CuBr

Copper (I) bromide

FeS

Iron (II) sulfide

PbO2

Lead (IV) oxide

Figure out the charge on the cation based on the charge of the anion. 32

Section 3.5 Ionic Compounds: Formulas and Names

Learning Check Name the following ionic compounds containing metals that form two kinds of positive ions: A.  Fe2O3 B.  SnCl2 C.  PbI4

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Ionic Compounds Containing Polyatomic Ions •  Naming these compounds is the same only now we use the name of the polyatomic ion wherever it occurs

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Section 3.5 Ionic Compounds: Formulas and Names

Naming Ionic Compounds Containing Polyatomic Ions •  • 

Must be memorized (polyatomic ion handout). Examples of compounds containing polyatomic ions: NaOH Sodium hydroxide Pb(NO3)2 lead (II) nitrate (NH4)2SO4 Ammonium sulfate

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Section 3.5 Ionic Compounds: Formulas and Names

Learning Check Write the names of the following compounds. A. NaC2H3O2 B. Ca3(PO4)2 C.  CuNO2 D. K2Cr2O7 E. Mg(ClO2)2 F.  NH4CN 36

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Section 3.5 Ionic Compounds: Formulas and Names

Hydrated Ionic Compounds •  Some ionic compounds are hydrates •  This means they have a specific number of water molecules associated with each formula unit in the lattice structure •  An example is magnesium sulfate heptahydrate •  MgSO47H2O

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Section 3.5 Ionic Compounds: Formulas and Names

Hydrated Ionic Compounds •  Cobalt chloride hexahydrate before and after heating to drive off the water

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Section 3.5 Ionic Compounds: Formulas and Names

Common hydrate prefixes •  •  •  •  •  •  •  •  • 

Hemi = ½ Mono = 1 Di = 2 Tri = 3 Tetra = 4 Penta = 5 Hexa = 6 Hepta = 7 Octa = 8

These prefixes are also used in

naming molecular compounds

For example: carbon dioxide

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Section 3.6 Molecular Compounds: Formulas and Names

Molecular compounds vs Ionic Compounds •  Ionic compounds form between metal and non metal –  The ions in ionic compounds can only combine one way because compounds are charge neutral –  Na+ and Cl– can only combine one way

•  Molecular compounds form between two nonmetals –  The atoms in molecular compounds don’t form ions – they share electrons. For this reason the same combination of elements can form a number of different molecular compounds. 40

Section 3.6 Molecular Compounds: Formulas and Names

Naming Molecular Compounds 1. The first element in the formula is named first, using the full element name. 2. The second element is named as if it were an anion. 3. Prefixes are used to denote the numbers of atoms present. 4. The prefix mono- is never used for naming the first element.

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Section 3.6 Molecular Compounds: Formulas and Names

Naming Molecular Compounds Prefixes Used to Indicate Number in Chemical Names

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Section 3.6 Molecular Compounds: Formulas and Names

Naming Molecular Compounds CO2

Carbon dioxide

SF6

Sulfur hexafluoride

N2O4

Dinitrogen tetroxide

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Section 3.6 Molecular Compounds: Formulas and Names

Learning Check Write the name of each covalent compound. A. SeO B. NO2 C. PF3 D. CBr4 E. P2O5

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Section 3.6 Molecular Compounds: Formulas and Names

Conceptual Connection •  The compound NCl3 is nitrogen trichloride , but AlCl3 is just aluminum chloride. Why?

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Section 3.6 Molecular Compounds: Formulas and Names

Acids •  Acids can be defined as molecular compounds than release hydrogen ions (H+) when dissolved in water. •  Acids are composed of hydrogen, usually written first in the formula, followed by one or more non metals. •  Two categories of acids for naming –  Binary acids –  Oxy Acids 46

Section 3.6 Molecular Compounds: Formulas and Names

Naming Binary Acids • 

• 

Hydrogen plus a non metal, the acid is named with the prefix hydro– followed by the root of the anion and the suffix –ic. Examples: HCl Hydro chlor ic acid HBr Hydro brom ic acid H 2S Hydro sulfur ic acid

Notice sulfur is kind of a weirdo sulfur instead of sulf 47

Section 3.6 Molecular Compounds: Formulas and Names

Naming OxyAcids • 

Hydrogen plus an oxyanion (anion containing a nonmetal and oxygen)  

 

The suffix –ic is added to the root name if anion name ends in –ate.: HNO3 Nitrate Nitric acid H2SO4 Sulfate Sulfuric acid HC2H3O2 Acetate Acetic acid The suffix –ous is added to the root name if anion name ends in –ite. HNO2 Nitrite Nitrous acid H2SO3 Sulfite Sulfurous acid HClO2 Chlorite Chlorous acid

Notice sulfate is kind of a weirdo sulfur instead of sulf 48

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Section 3.6 Molecular Compounds: Formulas and Names

Flow Chart for Naming Acids

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Section 3.6 Molecular Compounds: Formulas and Names

Learning Check Name the following acids a)  HF (aq) b)  H2CO3 (aq) c)  HI (aq) d)  HClO2 (aq) e)  H2CrO4 (aq)

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Section 3.6 Molecular Compounds: Formulas and Names

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Section 3.6 Molecular Compounds: Formulas and Names

Learning Check Name the following compounds a)  KNO3 b)  TiO2 c)  Sn(OH)4 d)  PBr5 e)  H2SO3 (aq)

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Formula Mass •  In Chapter 2 we defined the average mass of an atom as its atomic mass –  Measured in atomic mass units (amu)

•  The atomic mass of carbon is 12.01 amu •  Now that we have been introduced to ionic and molecular compounds we can build on this definition to something called a Formula Mass

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Formula Mass •  Formula mass is the average mass of a molecule (or formula unit) –  Molecular mass and molecular weight are also common (and interchangeable terms)

•  What does it mean though? •  Formula mass is the sum of all the atomic masses in the chemical formula

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Formula Mass Formula Mass of H2O (2 × 1.008 amu) + 16.00 amu = 18.02 amu Formula Mass of Ba(NO3)2 137.33 g + (2 × 14.01 g) + (6 × 16.00 g) = 261.35

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Molar Mass of a Compound •  Also in chapter 2 we saw that the atomic mass in amu/atom was equal to the molar mass in g/mole. –  This is because the definition for amu and mole are related to each other.

•  The same is true for a compound. •  The formula mass in amu/formula unit = molar mass in g/mole 56

Section 3.8 Formula Mass and the Mole Concept for Compounds

Molar Mass of a Compound •  So if we know a formula mass for water = 18.02 amu/molecule –  18.02 amu in one molecule

•  We also know that the molar mass for water = 18.02 g/mole. –  18.02 grams in one mole

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Using Molar Mass to Count by Weighing •  The fact that amu/molecule is related to g/mole allows us to count molecules by weighing them. •  This is the same thing we did when we counted atoms by weighing them. •  Molar mass – g/mole – is a conversion factor between mass and numbers of molecules.

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Example •  Glucose is one of the end products of photosynthesis, the process that converts CO2 and H2O to complex carbohydrates. The formula for glucose is C6H12O6. Determine the molar mass of glucose. Determine the number of moles in 50.0 g of glucose. Determine the number of molecules in 50.0 g of glucose.

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Example •  Glucose is one of the end products of photosynthesis, the process that converts CO2 and H2O to complex carbohydrates. The formula for glucose is C6H12O6. Determine the molar mass of glucose. (6 x 12.01) + (12 x 1.008) + (6 x 16.00) = 180.16 g/mol

Determine the number of moles in 50.0 g of glucose.

Determine the number of molecules in 50.0 g of glucose.

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Section 3.8 Formula Mass and the Mole Concept for Compounds

Learning Check Ethanol is produced from sugars by yeast in the process called fermentation. The formula for ethanol is C2H5OH. Determine the molar mass of ethanol. Determine the number of moles in 525 g of ethanol. Determine the number of molecules in 525 g of ethanol.

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Section 3.9 Composition of Compounds

Mass Percent •  We use something called Mass Percent to determine the relative amounts of elements in a compound. •  Mass percent takes the different masses of the different elements into consideration in calculating relative amounts.

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Section 3.9 Composition of Compounds

Mass Percent •  Mass Percent of hydrogen in H2O

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Section 3.9 Composition of Compounds

Mass Percent (Mass Percent Composition) •  General Equation

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Section 3.9 Composition of Compounds

How is Mass Percent Composition Useful •  A lot of times we need to know how much of one specific element is present in a compound •  A toxin •  An active ingredient •  Mass Percent (mass percent composition) allows us to calculate this.

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Section 3.9 Composition of Compounds

Learning Check •  Determine the mass percent of each element in calcium carbonate (CaCO3)

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Section 3.9 Composition of Compounds

Mass Percent Composition as a Conversion Factor •  Mass percent can also be used as a conversion factor •  From the last problem •  Mass Percent of Oxygen = 47.96%

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Section 3.9 Composition of Compounds

Mass Percent Composition as a Conversion Factor •  How do we use mass percent as a conversion factor? •  To calculate the mass of an element in a given mass of a compound.

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Section 3.9 Composition of Compounds

Example •  Sodium chloride (table salt) is 39% sodium by mass. How much sodium chloride is allowed per the RDA (recommended daily allowance) if a person is allowed to consume 2.4 g of sodium per day.

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Section 3.9 Composition of Compounds

Example •  Sodium chloride (table salt) is 39% sodium by mass. How much sodium chloride is allowed per the RDA (recommended daily allowance) if a person is allowed to consume 2.4 g of sodium per day.

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Section 3.9 Composition of Compounds

Learning Check •  Calcium carbonate (a common calcium supplement) is 40.04% calcium by mass. How much calcium is present in one tablet of calcium carbonate that has a mass of 500. mg?

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Section 3.9 Composition of Compounds

Conversion Factors from Chemical Formulas •  We just determined the amount of calcium in a given mass of calcium carbonate using mass percent as a conversion factor. •  There is actually another way to do this calculation •  Using the Mole Relationships in the Chemical Formula •  What does this mean and how do we use it?

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Section 3.9 Composition of Compounds

Mole Relationships •  Chemical formulas tell us the number of atoms in a compound •  For a single formula unit of calcium carbonate (CaCO3)

•  But what if we have a mole of CaCO3? •  Mole Relationships for CaCO3 73

Section 3.9 Composition of Compounds

Mole Relationships •  Mole Relationships for CaCO3 •  There is a 1:1 relationship between moles of calcium or carbon atoms and moles of compound •  There is a 3:1 relationship between moles of oxygen atoms and moles of compound

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Section 3.9 Composition of Compounds

Conversion Factors from Chemical Formulas •  How do we use mole relationships to determine composition of a compound? •  The mole relationship is a conversion factor. •  So lets say we want to know how many grams of calcium are in a 500. mg sample of CaCO3

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Section 3.9 Composition of Compounds

Conversion Factors from Chemical Formulas •  How many grams of calcium are in a 500. mg sample of CaCO3

•  We cannot convert directly from grams to grams. •  The relationship between atoms in a compound is moles to moles •  How do we use the mole:mole conversion to do this. 76

Section 3.9 Composition of Compounds

Conversion Factors from Chemical Formulas •  How many grams of calcium are in a 500. mg sample of CaCO3 •  Grams of CaCO3 to moles CaCO3 •  Moles CaCO3 to moles of Ca •  Moles of Ca to grams of Ca •  GMMG

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Section 3.9 Composition of Compounds

Learning Check •  The book uses the compound CCl2F2 (chlorofluorocarbons) for several examples •  Write out the mole relationships for this compound

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Section 3.9 Composition of Compounds

Learning Check •  Use the mole relationships from the last problem to determine how many grams of fluorine (F) are in 1500 g of the compound CCl2F2.

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Section 3.10 Determining a Chemical Formula from Experimental Data

Decomposition Analysis •  How do we determine the formula of a compound that is newly isolated? •  This is commonly done in the lab when new chemicals/drugs are synthesized •  Also common when new compounds are isolated

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Section 3.10 Determining a Chemical Formula from Experimental Data

Decomposition Analysis •  First step is to decompose the compound into its constituent parts •  Weigh them •  But this gives us a ratio of masses (in grams) •  Does not really tell us anything about the relationship of the number of atoms in the compound

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Section 3.10 Determining a Chemical Formula from Experimental Data

Decomposition Analysis •  •  •  • 

If we decompose 7.717 g of water (H2O) we get 6.868 g of oxygen 0.857 g of hydrogen If we divide these it gives an 8:1 ratio of oxygen to water (well we know that’s not right) •  To get the ratio of atoms in the compound we have to correct for the different in their atomic weight •  We have to convert to moles 82

Section 3.10 Determining a Chemical Formula from Experimental Data

Decomposition Analysis •  We have to convert to moles

•  Now this looks more like it •  H0.850 O0.429 •  Divide by the smaller number = H1.98O = H2O 83

Section 3.10 Determining a Chemical Formula from Experimental Data

Empirical vs Molecular Formula •  A decomposition analysis breaks a complex compound down to its elements to determine the relationship between them •  This relationship is called an empirical formula •  The simplest whole number ratio of atoms in the compound •  The form of the compound before it was broken down might actually be a multiple of this empirical formula •  This is called the molecular formula 84

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Section 3.10 Determining a Chemical Formula from Experimental Data

Empirical vs Molecular Formula •  To determine the molecular formula you need the empirical formula and the molecular weight of the original compound. •  For water the empirical formula and molecular formula are the same •  Not true for other compounds •  Molecular formula = empirical formula x n

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Section 3.10 Determining a Chemical Formula from Experimental Data

Example •  The empirical formula for a compound by decomposition analysis is CH2O and the molar mass is determined to be 180.2. What is the molecular formula? •  Molecular formula = (CH2O) x n

•  Molecular formula = C6H12O6 86

Section 3.10 Determining a Chemical Formula from Experimental Data

Learning Check •  A compound with the following percent composition has a molar mass of 60.10 g/mol. Determine its molecular formula. Assume 100 g of the compound. •  C, 39.97% •  H, 13.41% •  N, 46.62%

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Section 3.10 Determining a Chemical Formula from Experimental Data

Combustion Analysis •  One way to determine hydrogen and carbon content of a compound is to react it with oxygen then collect the H2O and CO2 produced

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Section 3.10 Determining a Chemical Formula from Experimental Data

Combustion Analysis •  •  • 

How can this information be used to determine a formula? If a substance has been isolated that contains only C, H and N. If 0.1156 g of this sample is reacted with oxygen and 0.1638 g of CO2 and 0.1676 g of H2O are collected.

•  • 

The nitrogen is what is left over 0.1156 – (0.04470 + 0.01875) = 0.05215 g N

• 

To determine formulas though we need numbers of atoms and grams don’t really help us with that

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Section 3.10 Determining a Chemical Formula from Experimental Data

Combustion Analysis

•  • 

Ratio of moles is the same as ratio of atoms (Section 3.3 The Mole) Divide by the smallest value

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Section 3.10 Determining a Chemical Formula from Experimental Data

Combustion Analysis • 

So we need to convert grams to moles

• 

Empirical Formula = C1H5N1

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Section 3.11 Writing and Balancing Chemical Equations

How to Write Balanced Chemical Equations •  Chemical equations include the states of the reactants and products •  CH4 (g) + O2 (g) → CO2 (g) + H2O (g)

•  All of the reactants and products in this reaction are gases

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Section 3.11 Writing and Balancing Chemical Equations

States of reactants and products •  Notice the difference between (l) liquid and (aq) aqueous •  Aqueous is something dissolved in water. 93

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Section 3.11 Writing and Balancing Chemical Equations

How to Write Balanced Chemical Equations •  Balanced equation •  CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)

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Section 3.11 Writing and Balancing Chemical Equations

Tips for balancing •  Balance polyatomic ions as a group •  Always balance oxygen last •  Can used fractions to balance – then multiply them out at the end •  Practice, practice, and practice.

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Section 3.11 Writing and Balancing Chemical Equations

Example Write a balanced equation for the reaction of aqueous lead (II) nitrate with aqueous sodium phosphate to produce solid lead (II) phosphate and aqueous sodium nitrate. Pb(NO3)2 (aq) + Na3PO4 (aq) → Pb3(PO4)2 (s) + NaNO3 (aq) 3 Pb(NO3)2 (aq) + 2 Na3PO4 (aq) → Pb3(PO4)2 (s) + 6 NaNO3 (aq)

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Section 3.11 Writing and Balancing Chemical Equations

Learning Check Write a balanced equation for the reaction of ammonia (NH3) with the oxygen to produce nitrogen monoxide and water

Section 3.11 Writing and Balancing Chemical Equations

Learning Check Write a balanced equation for the reaction of aqueous cobalt (I) nitrate with sulfuric acid to produce solid cobalt (I) sulfate and nitric acid. Both acids are aqueous solutions

Section 3.11 Writing and Balancing Chemical Equations

Conceptual Connection •  Which quantity or quantities must always be the same on both sides of a chemical equation? •  (a) the number of atoms of each kind •  (b) the number of molecules of each kind •  (c) the number of moles of each kind of molecule •  (d) the sum of the masses of all the substances involved 99

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Section 3.12 Organic Compounds

Organic vs Inorganic Compounds •  Early chemists divided compounds into two types •  Organic - Originated from living things –  sugars

•  Inorganic - Originated from the earth –  salts

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Section 3.12 Organic Compounds

Organic vs Inorganic Compounds •  Organic and inorganic compounds have different chemical properties •  Organic –  Easy to decompose –  Very hard to synthesize – very complex

•  Inorganic – –  Hard to decompose –  Easier to synthesize

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Section 3.12 Organic Compounds

What are organic compounds •  Major components of living organisms –  Proteins, fats, DNA

•  •  •  • 

Smells, tastes, fuels Drugs Food Pretty much everything we consume

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Section 3.12 Organic Compounds

What are organic compounds made of? •  Organic compounds are composed primarily of carbon and hydrogen with a few other elements –  Nitrogen, oxygen and sulfur

•  Key element in an organic compound is carbon

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Section 3.12 Organic Compounds

What are organic compounds made of? •  Carbon makes 4 bonds and often bonds to itself to make chain, branched and ring structures.

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Section 3.12 Organic Compounds

Hydrocarbons •  Organic compounds that contain only carbon and hydrogen. •  These are all the common fuels that we use •  Gasoline, propane, natural gas •  Can have single, double or triple bonds –  Alkane – single bonds –  Alkenes – double bonds –  Alkynes – triple bonds 105

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Section 3.12 Organic Compounds

Naming Hydrocarbons •  Base name determined by a prefix that indicates the number of carbons –  Meth = 1, eth = 2, prop = 3, but = 4, pent = 5 –  hex = 6, hept = 7, oct = 8, non = 9, dec = 10

•  Suffix indicates the presence of a multiple bond –  Single bonds - ane –  Double bond - ene –  Triple bond - yne

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Section 3.12 Organic Compounds

Naming Hydrocarbons •  Butane CH3CH2CH2CH3

•  Butene CH3CH2CH=CH2

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Section 3.12 Organic Compounds

Common Hydrocarbons

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Section 3.12 Organic Compounds

Learning Check •  Write the name for the following hydrocarbons •  CH3CH2CH2CH2CH2CH3

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Section 3.12 Organic Compounds

Functionalized Hydrocarbons •  A functionalized hydrocarbon is a carbon hydrogen chain that has a functional group added to it. •  What is a functional group? •  A functional group is a group of atoms that change the characteristics of the hydrocarbon.

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Section 3.12 Organic Compounds

Functionalized Hydrocarbons •  Example of a functionalized hydrocarbon. •  Ethane

•  Ethanol

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Section 3.12 Organic Compounds

Functional Groups

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Section 3.12 Organic Compounds

Functional Groups

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Section 3.12 Organic Compounds

Learning Check •  For the following compounds, give the family of the functional group (try to name the compound!)

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Section 3.12 Organic Compounds

Suggested Homework •  Chapter 3 •  Review Questions 1 – 5, 8, 14 – 22 •  Odd numbered problems 23 – 118 (skip 25, 63, 69, 73, 79, 87, 89, 97)

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