The covalent bond Chapter VIII

Why do atoms bond? When atoms form bonds they don’t do it just for the sake of doing it there has to be a reason. Atoms want to be like the noble gases They want stability They want low potential energy When atoms form bonds they are getting these things from the union. They obtain noble gas notation They have full outer shells or they have a lower potential energy Making them more stable

What is a covalent bond? Atoms share a valence electron. Different than an ionic which an electron is given up of accepted When atoms share electrons in a covalent bond a molecule is formed. In these molecules the shared electron is considered to belong to both electrons highest energy level Ionic only the accepter gets credit for this electron. These bonds usually happen with non-metals Ionic is a metal and a non metal usually The atoms in these bonds tend to be close together on the periodic table. Ionic tend to be far apart

How are covalent bonds formed Lets look a Fluorine Fluorine has seven valence electrons If two fluorine atoms approach each other several forces act. Two repulsive forces act on the atoms One form like-charged electrons One from like-charged protons A force of attraction also acts One atoms protons attractions the other atoms electrons As the atoms get closer the attraction of the protons to the electrons increases till it reaches a maximum. When this maximum is achieved the atoms covalently bond. However if the atoms move closer The attraction force is less then the repulsive force The most stable position for these atoms is when the attraction force exceed the net repulsion.

Looks like this

The single covalent bond In the fluorine atom a single covalent bond is formed This can be represented in a Lewis structure with a line or vertical dots Again they do this to be at low energy.

Group 17 and single bonds These atoms have seven valence electrons They will form one single bond to they can achieve a full octet They will form these bonds with other non metals such as carbon Some of these will form bonds with themselves F2, Cl2

Group 16 and single bonds These atoms have six valence electrons They will share two electrons forming to covalent bonds to form an octet For example oxygen will form a covalent bond with two hydrogen to make H2O In this example all of the atoms have noble gas configurations and are stable.

Group 15 and single bonds These atoms have 5 valence electrons and are able to form three single covalent bonds When they form these bonds the complete the octet and are stable Nitrogen for example is able to bond with other non metals to form new molecules NH3 (ammonia), NF3, NCl3, NBr3

Group 14 and single bonds These atoms have 4 valence electrons and therefore need 4 electrons to complete the octet They can make 4 single bonds Carbon for example bond with everything it seems because it needs to make the bonds CH4 (methane)

The Sigma (σ) bond The sigma bond is just a single covalent bond These bonds overlap orbitals between atoms S-S overlap S-P overlap P-P overlap This tells us where the electron is most likely to be found.

Multiple covalent bonds In some molecules, atoms have noble-gas configurations when they share more than one pair of electrons with on or more atoms. Sharing multiple pairs of electron forms multiple covalent bonds Double bonds and triple bonds are examples of multiple bonds.

Multiple covalent bond types Double bonds Two pairs of electrons are shared between atoms Example is O2 each Oxygen has six valence so needs to share two pairs to reach an octet. Triple bonds Three pairs of electrons are shared between atoms N2 is an example

The Pi bond (a fat kids dream) A multiple covalent bond consists of one sigma bond and at least on pi bond The pi bond (π) forms when parallel orbitals overlap and share electrons The shared electron pair a pi bond occupies the space above and below the line that represents where the two atoms are joined together Any multiple covalent bond has both sigma and pi bonds. Double one sigma and one pi Triple one sigma and two pi bonds

Pi bond

Bond strength of covalent bonds Bonds between different atoms have different strengths based upon multiple factors. Distance between the nuclei of the bonding atoms is the biggest factor The shorter the distance the stronger the bond Multiple bonds tend to have shorter bond lengths making them stronger bonds.

Bonds and energy An energy change occurs when a bond between atoms in a molecule forms or breaks. Energy is released (exothermic) when bonds are formed Energy must be added (endothermic) to break a bond. The amount of energy required to break a specific covalent bond is called bond-dissociation energy and is always a positive value. Bond dissociation energy also indicates the strength of the bond Remember bond length is inversely related to bond energy The smaller the bond length the higher the energy. To determine the potential energy of a molecule all one does is add up the bond dissociation energies for all the bonds in said molecule.

Energy in vs. Energy Out Ok so we know that when a bond is formed it releases energy and when a bond is broken it requires or takes in energy, but what happens in a chemical reaction Remember that in a chemical reaction bonds are being broken and made most of the time To figure out total energy one must calculate the difference between the energy taken in vs the energy given out In the energy taken in is more than its an endothermic reaction Think of instant ice packs- take in energy that’s what make them cold If the energy give out is more than its an exothermic reaction Think of hand warmers or MREs or camping food End 8-1

Naming molecules We learned how to name ionic compounds already. Covalent not much harder. Binary compounds 1. the first element in the formula is always named first, using the entire element name. 2. the second element in the formula is named using its root and adding the suffix –ide 3. prefixes are used to indicate the number of atoms of each element that present in the compound Never use mono for the first element If double vowels dropped on one Not monooxide Monoxide

Common prefixes Number of atoms Prefix 1 Mono- 6 Hexa- 2 Di- 7 Hepta- 3 Tri- 8 Octa- 4 Tetra- 9 Nona- 5 Penta- 10 Deca-

Name these CO₂ CO P₂O₅ N₂O SiO₂ CBr₄ SO₃ PBr₅ ICl₃ NI₃ N₂O₃ N₂O₄ SO₂ As₂O₅ PCl₃ CCl₄ H2O SF₆ XeO₃ P₄S₃

Naming acids If a compound produces hydrogen ions (H+) in solution, it is an acid. Two common types Binary Oxyacids Naming binary acids 1. the first word has the prefix hydro- to name the hydrogen part of the compound the rest of the first word consists of a form of the root of the second element plus the suffix –ic 2. the second word is always acid HCl- hydrochloric acid

Naming Oxyacids An acid that contains both a hydrogen atom and an oxyanion is referred to as an oxyacid. Oxyanion- polyatomic ion containing on or more oxygen atoms 1. first, identify the oxyanion present. The first word of an oxyacid's name consists of the root of the oxyanion and the prefix per- or hypo- if it is part of the name, and a suffix. If the oxyanion’s name ends with the suffix –ate, replace if with the suffix –ic. If the name ends with –ite, replace it with –ous. 2. the second word in always acid

Lets name some acids HI HClO3 HClO2 H2SO4

Go the other way Now that we know how to name the acids and other compounds we can write the formulas from the names Just remember the naming rules and its pretty easy

Lets try some Silver monochloride Dihydrogen monoxide Chlorine triflouride Diphosphorus trioxide Disulfur decafluoride End 8-2

Molecular structures Things we need to know Lewis structures Resonance Structural formula Resonance Molecule that can’t be shown by one Lewis structure Coordinate covalent bond End 8.3

Molecular geometry

More shapes BENT

Yet more shapes Square Antiprismatic T-Shaped

Last one

The VSEPR Model VSEPR Valence Shell Electron Pair Repulsion This model is used to determine the shape of a molecule once the Lewis structure of that molecule has been drawn. This model is based on an arrangement that minimizes the repulsion of shared and unshared electron pairs Molecule shape gives molecules many of its properties. Electron densities in overlapping orbitals determine molecular shape.

Bond angles Electron pairs in a molecule arrange so that they minimize interaction between them This creates a natural shape of molecules The unpaired electrons take up slightly more space then the shared electron pairs. So shared bonding orbitals are pushed together by unshared pairs.

Why does shape matter? Molecular shape determines if different molecules are able to react. If the shape isn’t right then the molecules cant get close enough to react. Think about food Food molecules have different shapes The different shape give us different tastes. If they all had the same shape everything would truly taste like chicken.

Hybridization This means the same as it does in biology and engineering the combination of two different things to form something new but has qualities of the old Orbitals can do this also Lets look at carbon Carbon has 4 valence electrons 2 in the s orbital 2 in the p orbital But when carbon bonds with hydrogen the orbitals change into a hybrid called the sp3 orbital Its an sp3 because it forms from one s and 3 p orbitals to form 4 new orbitals Lone pairs also occupy hybrid orbitals. End 8.4

Electronegativity and Polarity The measure of the tendency of an atom to accept an electron No noble gases Don’t bond Xeon will bond with certain elements Increase from bottom left to upper right in general These number were assigned not measure Refer to your flipbook.

Using electronegativity Electronegativity can be used to determine what type of bond will be formed. To do this you calculate the difference in the electronegativity of the elements that are bonding together. Take this difference and see where it falls The problem with bonds is that they are never truly covalent or ionic they are some of both Which ever one is more we call it that.

Lets do one or two Look at Oxygen gas O2 Look up the electronegativity of oxygen and find it to be 3.44 So 3.44-3.44= zero Look at the chart So Oxygen is nonpolar covalent Hint any element bonded only with itself is going to equal zero and be nonpolar covalent.

Cont…. Now lets look at hydrochloric acid What is HCl? HCL Look up the electronegativity Cl=3.16 H=2.20 Do the math 3.16-2.20= .96 Look at the chat What is HCl?

So what's a Polar covalent bond Think about a tug of war between John Cena and Paris Hilton. This is kind of the idea being a polar covalent bond One atom is going to “win” the electron more often than not. Electrons have a negative charge so that atom has a negative charge The losing atom end up with a positive charge The charge is represented with the Greek letter δ with a + or – Depending on the charge

Dipole?????? A dipole is something that has two (di) poles. Like a magnet Just like the magnet a dipole molecule has a positive end and a negative end These are polar covalent bonds With these poles they are attracted by electric fields. Solubility is the ability of a substance to dissolve in another substance Polar and ionic molecules dissolve in polar substances Non polar molecules only dissolve in nonpolar substances Good rule-like dissolves like This tells us why white bears are afraid of water.

Intermolecular forces Nonpolar molecules The force is weak and called a dispersion force Or induced dipole Polar molecules They have charged ends so opposite charges attract Called a dipole-dipole force The more polar a molecule the stronger this force is Finally we have the hydrogen bond Especially strong Forms between hydrogen end of one dipole and A fluorine, oxygen or nitrogen atom on another dipole