Chapter 5
Chemical Bonding and Nomenclature
Objectives
• Describe covalent bonding
• Explain the energy changes as molecules are formed
Bonds
• Forces that hold atoms together and make them function as a unit
• Two types of bonds– Ionic – Transfer of electrons
• Between metals and nonmetals–Called Salts
– Covalent – Sharing of electrons• Between nonmentals
–Called Molecules
Molecules
• A group of nonmetallic atoms covalently bonded
• Can be different atoms such as H2O• Can be the same atom such as O2
– Diatomic element• H2, N2, O2, F2, Cl2, Br2, I2
• Chemical Formula – Tells the number and types of atoms present– CH4 has 1 Carbon and 4 Hydrogens
Teacher where do molecules come from?
• Well when two atoms are attracted to each other they come together to share and share a special type of bond. Once this bond is established a molecule is born.
• Seriously!
The Making a Molecule of H2
1 A is 1x10-10m
So . . .
10A is 1x10-9m
The Atoms Need a Little Push
They Are Attracted!
What Is Going to Happen?
A Molecule Is Formed!
Bond Length
• Distance between two bonded nuclei
• Bond length changes when different types of atoms are bonded together
• Long bonds tend to be weaker
• Short bonds tend to be stronger
A Molecule Is Formed!
Bond Length
Atomic Interactions
• There are electrostatic interactions that take place within the atom and cause bonds to form
• Proton/Proton
• Electron/Electron
• Proton/Electron
Bonding Forces
Bond Length Diagram
Energy Changes in Reactions
• Breaking Bonds requires an input of energy
– Endothermic
• Forming Bonds releases energy
– Exothermic
• How does this compare to Biology?
How Many Bonds Need To Form
• Enough bonds to give eight electrons
• Octet Rule – Works most of the time – Duet for Hydrogen (2)– Quartet for Beryllium (4)– Sextet for Boron (6)
Specific Example for Hydrogen
Fluorine
Hydrogen Chloride
Homework
• p.201 #37,40,41,43,45,46,48,53
Objectives
• Draw dot diagrams for atoms
• Predict the formulas of compounds
• Draw lewis structures for molecules and ions
Dot Diagrams
• A method for showing the number of valence electrons around an atom
• Place elements symbol in the middle and electrons are represented as dots on the top, bottom, left and right
• Example
Li
Valence Electrons by Group
12 4 5 67
83
• Examples – Draw Dot Diagrams for the following elements Na, Mg, Al, Si, P, S, Cl, Ne
MgNa
Si
Al
PS Cl Ne
Formula Prediction
• Why is water always H2O?
• We can use dot diagrams to predict formulas
• Make formulas that allow the compounds to share electrons so each atom has 8– Or 2 if H, 4 if Be, and 6 is B
• Example – Predict the formula of a compound that is made of only Hydrogen and Oxygen
• Example – Predict the formula of a compound that is made of only Silicon and Hydrogen
• Example – Predict the formula of a compound that is made of only Arsenic and Chlorine
Usual Number of Bonds Formed
14 3 21
Problems
• p201 #56,62,63,64
Lewis Structures
• A method for determining the arrangement of bonds in covalent species
• Similar to dot structures but shows all bonds present
How 2 Draw
• Determine the number of valence electrons
• Determine the central atom. – Usually the single atom, or the one in the middle of
the formula
• Place other atoms around the middle and bond
• Complete octets with remaining electrons
• If each atom does not have 8 electrons – Multiple Bonds my be necessary
Bond Types
• Single Bonds = 2 electrons– Weakest and longest bonds
• Double Bonds = 4 electrons– In the middle
• Triple Bonds = 6 electrons– Strongest and shortest
There are not quadruple bonds!
General Rules
• Hydrogen will only form single bonds
• Halogens usually only form 1 bond. Why?– 7 valence electrons
• Oxygen will have 2 bonds and often forms multiple bonds
• Carbon likes to form chains
• Example – Draw the Lewis Structure for CBr4
• Example – Draw the Lewis Structure for NF3
• Example – Draw the Lewis Structure for O2
• Example – Draw the Lewis Structure for CS2
• Example – Draw the Lewis Structure for BeF2
• Example – Draw the Lewis Structure for C6H14
Polyatomic Ions
• Atoms that are covalently bonded together and have a charge
• Lewis structure rules– Negatively charged add electrons– Positively charged subtract electrons– Place Lewis structure in brackets when you
are finished
• Example – Draw the Lewis Structure for NO+
Resonance
• Species where equivalent Lewis structures exist
• Electron density is spread out evenly between resonant bonds– Delocalized – Spread out
• Often present in polyatomic ions
• Example – Draw the Lewis Structure for CO3
-2
• Example – Draw the Lewis Structure for NO2
-
• Example – Draw the Lewis Structure for AsF5
Homework
• p202 #71,74,75,76,77,82
• #71 C2H4
• #74 SO2
• #75 O3
• #76 Any Similarities b/t 74 and 75
• #77 Justify that that bonds are 1.5 of a bond
• #82 NO3-
Objectives
• Explain ionic bonding
• Name ionic compounds, acids, and covalent compounds
• Write formulas for ionic compounds, acids and covalent compounds
Ionic Bonds
• Involves the transfer of electrons between atoms to form ions
• Usually takes place between metals and nonmetals
• The ions are oppositely charged and strongly attract each other
• Large numbers of ions come together to form a crystal of the ionic compound
Electron Transfer
Na Cl
Properties of Ionic Compounds
• High melting points
– Solids at room temperature
• Hard
• Brittle
• Do not conduct current as solids
• Conduct current as liquids, gas, solutions
• Why?
Properties of Ionic Compounds
• High melting points
– Solids at room temperature
• Hard
• Brittle
• Do not conduct current as solids
• Conduct current as liquids, gas, solutions
• Why?
Ion Formation
• Atoms gain or lose electrons to achieve the same number of electrons as noble gases
• Metals lose electrons – Become positive ions (cations)
• Nonmetals gain electrons– Become negative ions (anions)
• Many groups form specific ions of specific charge
Ion Charge by Group+1
+2 -3 -2 -1+3
Ionic Compounds
• Total charge on the compound is zero
• Formulas are reduced to lowest ratios
• Formulas indicate the ratio of ions present
Naming Ionic Compounds
• Cation is named first, anion second• Cation keeps its name, anion ends in –ide
– Ex. Sulfide, Chloride, Oxide
• CaCl2 – Calcium Chloride
• NaBr– Sodium Bromide
• Ba3P2
– Barium Phosphide
Naming Ionic Compounds II
• Some cations can have more than one charge– Transition metals, Tin, Lead
• A roman numeral is placed after the cation name to indicate charge
• FeI2
– Iron (II) Iodide
• FeI– Iron (I) Iodide
Cont.
• Exceptions to the transition metal rule– Silver, Zinc, Chromium only form one ion
• Silver only +1• Zinc only +2• Chromium only +2
– Therefore roman numeral is not required• AgBr
– Silver Bromide not Silver (I) Bromide
Polyatomic Ions
• A covalently bonded group of atoms that possess a charge.
• Table 2.5 MUST be memorized!
• One more oxygen that normal Per- -ate
• Normal number of oxygens -ate
• One less oxygen than normal -ite
• Two less oxygens than normal Hypo- -ite
Polyatomic Ions p. 193
Naming Rules
• -ate to –ite one less oxygen
• Per- one more oxygen that –ate
• Hypo- one less oxygen that –ite
• Hydrogen added to ion increases charge by +1
• Phosphate PO4-3
• Hydrogen Phosphate HPO4-2
• Dihydrogen Phosphate H2PO4-1
Naming Ionic Compounds III
• Name cation as normal – Include roman number if necessary
• Name polyatomic ion
• NaNO3
– Sodium Nitrate
• NiSO4
– Nickel (II) Sulfate
Naming Acids
• Two types of acids1. Hydrogen and halogen (Hyrdohalic Acids)
– Named Hydroelementic Acid– HCl Hydrochloric Acid
2. Hydrogen and polyatomic ions– Name depends on name of polyatomic
ion• -ate becomes –ic• -ite becomes – ous
Cont.
• Acid formulas with start with H
• H2SO4
– Sulfuric Acid
• HNO3
– Nitric Acid
• H2SO3
– Sulfurous Acid
• HNO2
– Nitrous Acid
Covalent Compounds
• Only have nonmetals
• Use Number prefixes
• 1 = Mono 6 = Hexa
• 2 = Di 7 = Hepta
• 3= Tri 8 = Octa
• 4 = Tetra 9 = Nona
• 5 = Penta 10= Deca
Cont.
• The first element gets its name– If something other than one use a prefix
• The second element ends in –ide– Use prefix no matter what
• CO– Carbon Monoxide
• CO2
– Carbon Dioxide
Writing Formulas from Names
• Work backwards from the name
– Determine charges of the ions
– Make ratio so charges balance
• Reduce formulas to lowest ratio
• Except formulas that are fixed
– Peroxide O2-2
Examples
• Calcium Fluoride– CaF2
• Calcium Nitrate– Ca(NO3)2
• Iron (III) Sulfate– Fe2(SO4)3
• Titanium (IV) Oxide– TiO2
Cont.
• Chloric Acid– HClO3
• Hydroiodic acid– HI
• Carbon Tetrabromide– CBr4
• Diphosphorus Tetrafluoride– P2F4
Odd Compounds
• Cobalt (IV) Oxide
– CoO2
– Just being reduced
– Peroxides tend to form w/ Group 1
• Mercury (I) Bromide
– Hg2Br2
– Mercury (I) must be Hg2+2
Homework
• p203 #’s 92,94,96,97, 121, 122, 124 new only,125,126,127,135
Top Related