740 likes | 868 Views
Bonding. How do compounds benefit society? Which type of compound is the most useful to you and why?. Octet Rule. In compounds, atoms tend to achieve noble gas electron configurations Metals lose to achieve octet in lower energy level Non-metals gain to get 8 if combine with metal
E N D
Bonding How do compounds benefit society? Which type of compound is the most useful to you and why?
Octet Rule • In compounds, atoms tend to achieve noble gas electron configurations • Metals lose to achieve octet in lower energy level • Non-metals gain to get 8 if combine with metal • or can share with other non-metals • Metal cation + non-metal anion = ionic bond (chapter 7) • 2 non-metals = covalent bond (chapter 8) • 2 metals = metallic bond (also chapter 7)
8.1 Molecular Compounds • Covalent bonds • Sharing of electrons to hold atoms together • Between 2 non-metals • Neutral • Makes molecules • Bond dissociation energy is the amount of energy required to break a covalent bond between atoms • A large bond dissociation energy corresponds to a strong covalent bond • Double and triple bonds are stronger than single cov bonds
Molecular compounds • Chemical formula = molecular formula • Not lowest whole number, just actual number of atoms of each element • Properties • Comprised of two or more non-metals • Not good conductors of heat and electric current • Lower melting and boiling points than ionic • Diatomic molecules • Molecule of 2 identical atoms • H2 O2 N2 Cl2 Br2 I2 F2
Nomenclature • The same elements can combine in multiple ways via different numbers of atoms to form different molecules • Ex CO vs CO2 • –ide suffix on second element just like ionic • Use prefixes to indicate number of each • Mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca- • No mono- for first element • Ex: carbon monoxide and carbon dioxide
Naming molecular compounds • Practice • SiO2 • Silicon dioxide • SO3 • sulfur trioxide • XeF4 • Xenon tetrafluoride • Phosphorus pentachloride • PCl5 • Dinitrogen monoxide • N2O
8.2 Nature of covalent bonding • Still obeys octet rule • Share as many electrons as needed for 8 • Single covalent bond • Share one pair of electrons (1 from each atom) • Ex H-H • Double covalent bond • Share two pairs of electrons (2 from each) • Ex O=O • Triple covalent bond • Share 3 pairs of electrons (3 from each) • Ex N≡N • Coordinate covalent bond • Share pair of electrons (usually only 1) but both from 1 atom • Ex :C: + :Ö: :C≡O: (often drawn :C=O: )
Recall electron configurations • Valence electrons • Highest occupied energy level • Largely determine chemical properties • Corresponds to American group number • Exception = He • Used for chemical bonds • Shown in electron-dot structures/diagrams • Ex: H·
Electron Dot Diagrams • Practice: Na F He Cl Mg S B Li Ne C Kr N Si O Al P Ca Find a trend?
Structures • Dot structures can be made for covalent bonds • Place shared electrons between the 2 symbols either as pair of dots or dashes • Ex H-H or :Ö=Ö: or :N≡N: • Each dash represents a single covalent bond • Unshared pairs / lone pairs / non-bonding pairs = valence e- pairs not in cov bond • Structural formulas use dashes, but also show spatial arrangement of atoms
How Covalent Bonds Form • Draw the Lewis dot diagrams for all the atoms in CH4 • Draw 1 carbon diagram • Draw 4 hydrogen diagrams
Bonding in Methane C H H H H
Bonding in Methane C H H H H
H Bonding in Methane C H H H
H Bonding in Methane C H H H
H Bonding in Methane C H H H
H H Bonding in Methane C H H
H H Bonding in Methane C H H
H H Bonding in Methane C H H
H H H Bonding in Methane C H
H H H Bonding in Methane C H
H H H H Bonding in Methane C
Bonding in Methane H C H H H
Draw the following compounds: PCl3 H2O SF2 AlCl3
Do Now: Get out your homework to be stamped and be prepared to answer: • What is a covalent bond? • What types of elements do covalent bonds form between? • What are valence electrons? • How many electrons are in a bond?
Single Bond • Two electrons are shared between two atoms • 1 electron pair • Represented by 1 line between the atoms
Let’s Make F2 F F
Let’s Make F2 F F
Let’s Make F2 F F
Double Bond • Four electrons are shared between two atoms • 2 electron pairs • Represented by 2 lines between the atoms
Let’s Make O2 O O
Let’s Make O2 O O
Let’s Make O2 O O
Let’s Make O2 O O
Double bonds Try It! • Draw the Lewis dot diagram for • O3 • CO2 • CF2S
Triple Bond • Six electrons are shared between two atoms • 3 electron pairs • Represented by 3 lines between the atoms
Let’s Make N2 N N
Let’s Make N2 N N
Let’s Make N2 N N
Let’s Make N2 Let’s Make N2 N N
Let’s Make N2 N N
Triple bonds Try It! • Draw the Lewis Dot Diagram for • HCN • N2O • C2H2
Lewis Structures summary • Place shared electrons between the 2 symbols either as pair of dots or dashes • Ex H-H or :Ö=Ö: or :N≡N: • Each dash represents a single covalent bond • Unshared pairs / lone pairs / non-bonding pairs = valence e- pairs not in covalent bond • Structural formulas use dashes, but also show spatial arrangement of atoms • Oriented to maximize distance between e- pairs (bonded or unshared)
Polyatomic ions • Even though overall charged, polyatomic ions are covalently bonded together • Overall charge comes from extra electrons either gained or lost in order to achieve an octet for all constituent atoms • Ex: SO32- O S O O • Use brackets to denote total charge
Polyatomic Ions • Covalently bonded • With additional or removed electrons • Cations lose and anions gain in e- count NO31- SO32- PO43- CN1-
Resonance structures • Occur when possible to draw two or more valid electron dot structures for a molecule • Originally thought that electrons would flip back and forth (resonate) between the options • Actual bonding is a hybrid (mixture) of the options • All valid structures are separated by double headed arrow • Ex: O3 (ozone) :Ö = Ö – Ö: ↔ :Ö - Ö = Ö: ˙˙ ˙˙
Exceptions to Octet Rule • If total # of valence e- for compound is odd, the octet rule cannot be satisfied • Nitrogen dioxide • Phosphorus pentachloride • Sulfur hexafluoride • *Boron trifluoride • readily reacts with NH3
VSEPR Theory • Valence-shell electron-pair repulsion theory • States: the repulsion between electron pairs cause molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible • Non-bonding pairs are important because they will repel the other e- pairs more • No other atom stakes claim, therefore held closer to the central nucleus and will squeeze out the others