Unit 7 Covalent Bonding Bonding A metal a

Unit 7 Covalent Bonding

Bonding • A metal & a nonmetal transfer electrons – An ionic bond • Two metals mix – An alloy (Metallic “bond”) • What do two nonmetals do?

Covalent Bonding • Nonmetals hold on to their valence electrons • They can’t give away electrons to bond • But still want to be stable! – Need noble gas configuration (octet rule) • Get it by sharing valence electrons with each other. • By sharing, both atoms get to count the electrons toward noble gas configuration.

Covalent Bonding Fluorine has seven valence electrons l A second flourine atom also has seven l By sharing electrons… l Both end with full orbitals l F F

Covalent bonding Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l Both end with full orbitals l F F 8 Valence electrons

Covalent bonding Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals l 8 Valence electrons F F

Ways to Illustrate Covalent Bonds • Molecular formula: shows the number of atoms of each element in a molecule. – Ex. PF 3 • Lewis Structures: uses dots to represent bonding between molecular compounds • Structural Formulas: shows the arrangement of atoms and bonds – Shared electron dots are replaced with a dash • Models: ball and stick (3 -D versions)

Single Covalent Bond • Occurs between 2 nonmetals or a nonmetal & hydrogen • Sharing of two valence electrons (1 pair) • Different from an ionic bond – electrons are SHARED not transferred

An example with dots… • It’s like a jigsaw puzzle • You will be given the formula • You put the pieces together to make everyone stable or happy – Most atoms need an octet • H & He need a duo – Carbon is often the center – Singles are often the center – Save H for last

Water H O Each hydrogen has 1 valence electron and wants 1 more The oxygen has 6 valence electrons and wants 2 more They share to make each other “happy”

Water • Put the pieces together • The first hydrogen is happy • The oxygen still wants one more HO

Water • The second hydrogen attaches • Every atom has full energy levels HO H

Structural formula… • Replace shared dots with a dash H O H

Practice – Dots & Structures • • H 2 • CH 4 • H 2 S • NH 3 • CH 2 Cl 2 • C 2 H 6 • SCl 2 • As. F 3 • Si. H 4 • CHF 3

Multiple Bonds • Sometimes atoms share more than one pair of valence electrons. • A double bond is when atoms share two pair (4) of electrons. • A triple bond is when atoms share three pair (6) of electrons.

Carbon dioxide C O • CO 2 - Carbon is central atom • Carbon has 4 valence electrons • Wants 4 more • Oxygen has 6 valence electrons • Wants 2 more

Carbon dioxide • Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short CO

Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short OC O

Carbon dioxide l The only solution is to share more O CO

Carbon dioxide l The only solution is to share more O C O

Carbon dioxide The only solution is to share more l CO 2 requires two double bonds l Each atom gets to count all the atoms in the bond l O C O

Carbon dioxide The only solution is to share more l CO 2 requires two double bonds l Each atom gets to count all the atoms in the bond 8 valence electrons l O C O

Carbon Dioxide • Replace the shared pairs with dashes O C O

Practice • O 2 • CS 2 • CH 2 O • N 2 F 2 • NO 2 • HCN (triple) • C 2 H 2 (triple)

Exceptions to the Octet Rule/Patterns of Bonding 1. Some elements with odd number of valence electrons – BF 3 – PCl 5 2. Coordinate covalent bonding

Coordinate Covalent Bond • When one atom donates both electrons in a covalent bond • Carbon monoxide (CO) CO

Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide (CO) C O

Summary of Covalent Bonding • Covalent bonds occur by SHARING electrons • Occurs between NONMETALS • End product is called a MOLECULE 1. Molecular compound - formed with different elements 2. Diatomic molecules - 2 of the same atom – There are 7 elements that always form diatomic molecules – H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , and I 2

Diatomic Molecules

Naming Molecular Compounds • Easier than ionic compounds • No balancing charges • 1 mono • 2 di • 3 tri • 4 tetra • 5 penta- • 6 hexa – • 7 hepta – • 8 octa – • 9 nona – • 10 deca –

Naming Molecular Compounds • 1 st element – add the prefix that matches the subscript – Exception – do not add “mono-” if there is 1 atom – No aa, oo, or ao double vowels • 2 nd element – add the prefix that matches the subscript – Still ends in -ide

Naming • CO 2 – 1 st element = Carbon; subscript 1 • Remember exception • “Carbon” – 2 nd element = oxygen; subscript 2 • Prefix for 2 is di • “Dioxide” – Full name = carbon dioxide

Practice Naming • S 2 Cl 2 – Disulfur dichloride • CS 2 – Carbon disulfide • SO 3 – Sulfur trioxide • P 4 O 10 – Tetraphosphorus decoxide

Name Formula • Just look at the prefixes! • Carbon tetrachloride – 1 Carbon, 4 Chlorine atoms – CCl 4 • Iodine heptaflouride • Dinitrogen monoxide • Sulfur dioxide

Common Names • H 2 O – dihydrogen monoxide – Water • NH 3 – Nitrogen trihydride – Ammonia • CH 4 – carbon tetrahydride • Methane • HCl – Hydrogen monochloride – Hydrochloric acid

Molecular Shapes • Lewis diagrams & structural formulas are 2 -dimensional • Real molecules are 3 -D • If there are 2 atoms, the molecule has a LINEAR shape (no other options!) – Carbon monoxide (CO) • If it has more than 2, how do we figure out the shape?

VSEPR Theory • Valence Shell Electron Pair Repulsion Theory • Used to predict shape of a molecule • Negative electrons repel each other and pairs want to be as far apart as possible

Linear • Linear: 2 atoms around central atom, no unshared pairs on central atom • With three atoms the farthest the two outer molecules can get apart is 180º. • Will require 2 double bonds or one triple bond 180º O C O

Trigonal Planar • Trigonal planar: 3 atoms around central atom, no unshared pair on central atom. • Angle = 120°

Tetrahedral H H C H H • 4 molecules around a central atom • All single bonds • Must think in 3 -D!

Tetrahedral • Tetrahedral: 4 atoms around central atom, no unshared pair on central atom • A pyramid with a triangular base. H H 109. 5º C H H

So far… SHAPE # ATOMS FROM THE CENTRAL ATOM # UNSHARED PAIRS ON THE CENTRAL ATOM LINEAR 2 0 TRIGONAL PLANAR 3 0 TETRAHEDRAL 4 0

Molecular Shapes • But what if there are unshared pairs on the central atom? • They still repel each other…

Bent • Bent: 2 atoms around central atom, 1 or 2 unshared pair(s) on central atom. H O H H <109. 5º

Bent • Ball and stick model does not show unshared electron pairs

Pyramidal • Pyramidal: 3 atoms around central atom, 1 unshared pair on central atom H N H H <109. 5º

Pyrimidial • Ball and stick model does not show unshared electron pairs

Trigonal Bipyramidal • Trigonal bipyramidal: 5 atoms around central atom, no unshared pair on central atom • Angles = 90° and 120°

Adding to the chart… SHAPE # ATOMS FROM THE CENTRAL ATOM # UNSHARED PAIRS ON THE CENTRAL ATOM BENT 2 1 or 2 PYRAMIDIAL 3 1 TRIGONAL BIPYRAMIDAL 5 0

There’s more! • Shapes within a larger molecule…. – C 2 H 4

So how do I determine the shape of a given molecule? 1. Draw the Lewis diagram 2. Count the atoms and unshared pairs off the central atom 3. Use the VSEPR Theory to determine the shape

Which type of bond is it? • Also called ionic character • Look at which elements are involved – Metal & nonmetal = ionic bond – 2 nonmetals = covalent bond • Electronegativity – measure of a tendency of an atom to attract a pair of electrons – Influenced by amount of positive charge in the nucleus & electron shielding

Differences in Electronegativity. • Big difference between values (greater than 1. 70) – One atom REALLY wants the electrons and the other…not so much – Ionic bonding ionic compound • Smaller difference between values (less than 1. 70) – Both have “equal” attraction for the e– Covalent bonding molecule

Differences in Electronegativity • Medium difference – Still a bit of a tug of war over e– Unequal sharing of electrons – Results in a POLAR covalent bond

Polar Covalent Bond

Differences in Electronegativity • Very small difference – Share electrons equally – NONPOLAR covalent bond – No positive and negative poles

Nonpolar Covalent Bond

Polar vs. nonpolar molecules • Look at polarity of each bond – All nonpolar bonds = nonpolar molecule (O 2) • Use the structural formula to draw the molecule & look at the overall shape – Symmetrical polar bonds cancel each other out so molecule = nonpolar (CO 2, C 2 H 4, CCl 4) – Nonsymmetrical polar bonds = polar molecule (H 2 O)

Dipole-dipole attraction • Dipole – partial negative on one side, partial positive on the other; has “poles” • Attraction between + part of one dipole and part of another dipole • Hydrogen bond - between an electronegative atom and a hydrogen atom bonded to another electronegative atom – Often involves F, N, or O – Strongest of the intermolecular forces

Hydrogen Bonding

Surface Tension/H Bonding

Van der Waals – London dispersion force • Weak intermolecular force caused by negative electrons on one side of a cloud being attracted to a nearby positive nucleus • Constantly changing

Properties of Molecular Compounds • Poor conductors of heat & electricity • Often found as liquids or gases • Weaker attraction between atoms • Low melting & boiling points

Carbon (Organic) Chemistry • Carbon plays a dominant role in the chemistry of living things • Bonding stability – 4 valence electrons – Very unlikely to form ionic bonds – Can form covalent bonds with LOTS of different elements (especially H & O) • Small molecules link together resulting in the formation of a large variety of structures often with repeating subunits

Examples of carbon-based compounds • Simple hydrocarbons • Small carbon molecules with functional groups • Complex polymers (including biological molecules)

Simple Hydrocarbons • Petrochemicals – Propane, Butane, Octane • Condensed Structural Notation - [CH 2]n

Hydrocarbons • “Saturation” with hydrogens affects shape & reactivity – Saturated vs. unsaturated fats – shape

Alkane vs Alkene vs Alkyne

Alkane vs Alkene vs Alkyne • Easy to identify from full structural fomulas

Functional groups • Specific groups of atoms that are responsible for chemical characteristics of a compound • ALWAYS a close relationship between properties & structure (aspirin, vitamins, insulin)