Presentation on theme: "Today is Tuesday, March 24 th, 2015 Pre-Class: What kind of bond joins the two oxygen atoms in oxygen gas – O 2 ? What kind of bond joins the carbon and."— Presentation transcript:
Today is Tuesday, March 24 th, 2015 Pre-Class: What kind of bond joins the two oxygen atoms in oxygen gas – O 2 ? What kind of bond joins the carbon and oxygen atoms in carbon dioxide – CO 2 ? Please take a video worksheet from the Turn-In Box and get a paper towel. Also find your bonding overview worksheets (the big grid one). Stuff You Need: See Below In This Lesson: Covalent Bonding, VSEPR Theory, and Intermolecular Forces (Lesson 4 of 4)
Today’s Agenda Covalent/Molecular Bonding Bond/Molecular Polarity VSEPR Theory Intermolecular Forces Where is this in my book? – P. 213 and following…
By the end of this lesson… You should be able to describe the nature of a molecular bond. You should be able to describe the shape of a molecule based on its constituent atoms. You should be able to identify forces between molecules based on the nature of the participating molecules. You should be able to draw resonance structures of molecules.
What you need to know… …is on this checklist.
Reminder HH O This is an atom This is a molecule
Video Introduction Find your Bonding Overview sheets and have your video worksheets ready (on the Covalent Bonding side). It’s another Australian video!
Covalent Bonds: The Important Stuff Covalent bonds are sometimes called molecular bonds. Covalent bonds generally occur between nonmetals and nonmetals (or metalloids). Electron pairs are shared (or fought-over). Bonds can be polar or non-polar. The term “molecule” is used only for covalent compounds. – Ionic compounds sometimes get the term formula unit.
Characteristics of Molecular Compounds Poor conductors of electricity. Low melting points. Polar compounds can dissolve in water. – Non-polar compounds cannot dissolve in water. Many are gases at room temperature.
Ionic or Covalent? But wait, you say. How can you determine whether a compound is ionic or covalent? There are two ways. – The “official” way: By electronegativity. – The “unofficial” way: By location on the periodic table.
The Official Way Remember that electronegativity is the ability of an atom to attract electrons in a covalent compound. – Why are noble gases not on here? – They don’t form compounds!
Determining Bond Type Use the file Periodic Table – Electronegativity. Compare the electronegativity values of the two atoms. Ionic: Difference ≥ 1.67 – Sometimes people use numbers as high as 2.0. Polar Covalent: 1.66 ≥ Difference ≥ 0.4 – This means that one atom (the more electronegative one) “hogs” the shared electrons. Non-Polar Covalent: 0.3 ≥ Difference – The atoms play nice.
The Unofficial Way Because ionic compounds occur between cations and anions, they often occur between elements on opposite sides of the periodic table. Consider both ways when making the distinction between bonds.
Identify the Bond Types NaCl – Ionic (electronegativity difference of 2.1) F 2 – Non-Polar Covalent (electronegativity difference of 0) MgCl 2 – Ionic (electronegativity difference of 1.8) BaO – Ionic (electronegativity difference of 2.6) CH 4 – Polar Covalent (electronegativity difference of 0.4) O 2 – Non-Polar Covalent (electronegativity difference of 0)
Aside: Bond Type Exceptions? You will not need to know them for this class, but keep in mind that there are exceptions to this bond type rule. A notable one is compounds involving Beryllium (Be) – they tend to be covalent despite Beryllium not achieving an octet. – BeCl 2, for example, is a covalent compound as confirmed by experimental evidence involving boiling point and behavior. – However, it still likes to form things that look like crystal lattices when it’s a solid. Weird.
The Octet Rule In ionic compounds, electrons are transferred such that atoms achieve the octet, either by losing or gaining. In covalent compounds, electrons are shared (or fought over) so that each atom has eight. The key? – Whichever type of compound it is, everyone either needs to lose all their valence electrons or share/gain to eight.
Diatomic Molecules Recall that seven elements on the periodic table naturally bond to themselves. Elements that exist diatomically: – Br-I-N-Cl-H-O-F – Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2 These elements are all bonded covalently to themselves.
The HONC Rule* *Can be broken on occasion. The HONC Rule is an easy way to show how many bonds form for an element. Hydrogen and Halogens form one covalent bond. Oxygen (and sulfur) form two covalent bonds. – Either two single bonds or one double bond. Nitrogen (and phosphorus) form three covalent bonds. – Either a triple bond, a double and single bond, or three single bonds. Carbon (and silicon) form four covalent bonds. – Either a triple and a single bond, two double bonds, a double and two single bonds, or four single bonds.
Facebook? Back in the day, Facebook had a “visualizer” for your friends. – Basically, it made a web of friend connections. – Those with more connections tended to be in the inner part of the web, those with fewer were on the outside.
Central Atoms In a molecule, the central atom is the one that forms the most bonds. According to the HONC Rule, Carbon and Silicon (and others in their group) are most likely to be the central atom. Similarly, hydrogens and halogens are rarely the central atoms.
Now, onto the bonding… As you know, Lewis Structures (dot notation? Remember?) show the arrangement of valence electrons in an atom. In covalent bonds, shared/fought over electrons can be shown by two dots ( : ) or a single line ( - ). – IMPORTANT: Each single line (-) connects two dots, so you should count them as two electrons.
Methods of Visualization If you’re using strictly dots, you’re drawing the Lewis Dot Structure. If you’re using lines, you’re drawing the Structural Formula. Throughout this lesson, you’ll learn many ways to accomplish one goal. Use what works for you and/or the problem.
Single Bonds Single bonds involve two shared electrons (or one pair). – Single bonds are sometimes drawn with a single line.
F F 1s 2s 2p Each has seven valence electrons. FF Single Bond Example One pair of electrons is shared. F F
Think of it another way… F F Shared Electrons Lewis Structure
Think of it another way… FF Shared Electrons Structural Formula
Covalent Bonding Practice Try this example: Hydrochloric Acid – HCl HCl H
Covalent Bonding Practice Try this example: Iodine –I2–I2 II I I
Covalent Bonding Practice Try this example: Dihydrogen Monoxide –H2O–H2O H OH O HH
Notice something… In each of the examples we did, I drew in the non-bonding pairs of electrons. As we saw in the video, these electrons have a strong influence on the shape of the molecule even without bonding to another atom. – More to come…
Covalent Bonding Practice Now it’s your turn. Covalent Bonding Worksheet (first page) #2, 4, 5, 7, 9, 10
H Unshared Pairs In our all examples (but let’s return to H 2 O), there are pairs of electrons that are not shared. – There are known as unshared pairs, lone pairs, or non-bonding pairs. IMPORTANT: Unshared pairs only count on the central atom (the one with the most bonds). – You should still draw them, though. OH There are two unshared pairs of electrons.
Unshared Pairs Ammonia (NH 3 ) has one unshared pair of electrons. – The electrons are on the nitrogen atom.
Unshared Pairs Practice How many unshared pairs are on the methane (CH 4 ) molecule? – None.
Unshared Pairs Practice How many unshared pairs are on the carbon tetrachloride molecule? – None (carbon is the central atom). Cl C
Double Bonds Double bonds involve four shared electrons (or two pairs). – Double bonds are sometimes drawn with double lines.
O O 1s 2s 2p Each has six valence electrons. O O Double Bond Example Two pairs of electrons are shared. O O
Covalent Bonding Practice Try this one: Carbon Dioxide – CO 2 CO O O C O
Covalent Bonding Practice Now it’s your turn. Covalent Bonding Worksheet (first page) #3, 12
Triple Bonds Triple bonds involve six shared electrons (or three pairs). – Triple bonds are sometimes drawn with triple lines.
N N 1s 2s 2p Each has five valence electrons. N N Triple Bond Example Three pairs of electrons are shared. N N
Covalent Bonding Practice Now it’s your turn. Covalent Bonding Worksheet #1 Covalent Bonding Worksheet (Reverse) – WARNING: I have done something a little sneaky on Page 2.
Drawing Lewis Structures Bigger molecules can make for complicated Lewis Structure drawings. – Imagine drawing the Lewis Structure for CH 3 Cl. Thankfully, there are two systematic ways to do it.
Covalent Bonding Tips and Tricks Jot it like you mean it! 1.Hydrogen and halogens will only ever make one bond. Connect them with single bonds first. 2.Make sure you don’t use too many electrons! You can only use as many as the atoms bring. 3.Follow the HONC Rule when possible. 4.Single bonds count as 2 electrons, double count as 4, triple count as 6. 5.DO NOT MAKE A QUADRUPLE BOND EVER!
CH 3 Cl: Method #1 H CH H Cl Make carbon the central atom (it wants four bonds). Add up the available valence electrons: – C = 4, H = 1 (3), Cl = 7 – TOTAL = 14. Join peripheral atoms to the central atoms with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons.
Make carbon the central atom (it wants four bonds). Add up the total electrons needed: – C = 8, H = 2 (3), Cl = 8 – TOTAL = 22. Subtract available electrons: – C = 4, H = 1 (3), Cl = 7 – TOTAL = 14. – = 8. Divide by two. (8/2 = 4) Draw in 4 total bonds between atoms, then complete octets. – Remember, double bonds count as two and triple as three. CH 3 Cl: Method #2 H CH H Cl
N Practice Molecule: HCN CH NCH
Practice Covalent Bonding and VSEPR Theory Worksheet – #1-3 Covalent Bonding and VSEPR Theory Worksheet – #4-5
VSEPR Theory So dot structures are just peachy, but they leave out something important. – Something like a third dimension. The thing is, thinking about things in three dimensions can be difficult. So…let’s introduce the next topic with a video: TED: George Zaidan and Charles Morton – What is the Shape of a Molecule?
VSEPR Theory VSEPR Theory (pronounced “vesper”) stands for Valence Shell Electron Pair Repulsion Theory. You can think of it as a theory that helps scientists predict the shape of a molecule. – Because electrons involved in bonding pairs or lone pairs are both negatively charged, they repel one another and get as far apart as possible.
VSEPR Theory Chart Treat this VSEPR Theory Chart like gold. If you need another, it’s on my website. This is invaluable and is basically your lifesaver. – For reals. It summarizes a lot of information that can’t be put into a PowerPoint and I’ll be referring to it frequently. – Unfortunately, it will not be available to you on the test.
VSEPR Theory To use the VSEPR Chart I gave you, do this: 1.Draw the dot structure. 2.Find: – The central atom. – The number of atoms bonded to the central atom. – The number of lone pairs on the central atom. 3.Then use the VSEPR Chart to find its shape.
VSEPR Chart Example: H 2 O We know H 2 O looks like this: 1.According to the HONC rule, O forms two bonds and H only forms one, so O is the central atom. 2.Oxygen has two bonded atoms and two unshared pairs of electrons. 3.The chart tells us that this is molecule takes on the “bent” geometry. H OH
VSEPR Geometry: Linear 1 Bonded Atom or 2 Bonded Atoms with 0 Lone Pairs. – Example: CO 2 CO O
Wait, what? Try drawing the dot structure for AlBr 3. – Notice anything? As you probably could have guessed, there are exceptions to the octet rule occasionally. The most common exceptions are compounds that have only single bonds formed with Group 3A cations. – These compounds do not obey the octet rule and take on trigonal planar geometry. Al Br
Covalent Bonding Worksheet Take a look at #9 on the reverse of your Covalent Bonding Worksheet. If this were Boron (and thus not an ionic compound), this would be a molecular compound and would exhibit the same geometry. – The reason it’s ionic is because of the electronegativity difference between Al and F. – ;) B FF F
VSEPR Geometry: Trigonal Pyramidal 3 Bonded Atoms, 1 Lone Pair. – Example: PH 3 P HH H
VSEPR Geometry: Bent 2 Bonded Atoms, 1 or 2 Lone Pairs. – Example: H 2 O O HH
VSEPR Geometry: Tetrahedral 4 Bonded Atoms, 0 Lone Pairs. – Example: CH 4 C HH H H
VSEPR Summary Bonded AtomsLone PairsShape 10, 1, 2, or 3Linear or 2Bent 30Trigonal Planar 31Trigonal Pyramidal 40Tetrahedral
Practice Find your Covalent Bonding Worksheet and write in the geometry of each covalently-bonded molecule. – Example: #1 (N 2 ) is linear because it has one bonded atom and one lone pair. Now try Covalent Bonding and VSEPR Theory Worksheet. – #6 Now add the geometry to #4 and 5 from Covalent Bonding and VSEPR Theory.
Practice Now try: – Covalent Bonding and VSEPR Theory Worksheet Reverse – Bottom section
Practice Lab – Molecules I
Aside: Empty Space Remember earlier in the year when I told you that we never really touch anything? – As in, repulsive forces between atoms keep us from actually coming in contact with them so we just get really really close? Well, this VSEPR stuff, the repulsion of electrons, is what’s causing that. VSEPR Theory is part of the reason you can’t walk through a wall right now. – You actually have plenty of space to sneak your atoms’ nuclei through the wall’s – if it weren’t for the repulsive forces.
Review Quick – what’s the shape?
Practice Now try: – Covalent Bonding and VSEPR Theory Worksheet Reverse – Top section
Lastly… Finally, let’s do Lewis Structures for covalent ions. – Like polyatomic ions from waaaay back when…
Polyatomic Ion Lewis Structures Or other really tricky ones… Procedure: 1.Determine total valence electrons. 2.Draw single bonds between all atoms. 3.Give bonded atoms octets (except H). 4.Satisfy the octet rule for the central atom(s). Any leftover electrons go to the central atom. If the central atom has six electrons, it double bonds to a bonded atom with a lone pair. If the central atom has four electrons, it triple bonds to a bonded atom with a lone pair, or makes two double bonds. – BIG HUGE IMPORTANT NOTE! The HONC rule can be broken, and will always be for ions but never for hydrogen or halogens.
Electrons Left: 24Electrons Left: 0 Electrons Left: 34 Find the total e - : – Each carbon has 4 and each oxygen has 6, so 32 total. – The overall charge is 2-, so there are 2 more electrons. Grand total 34. Start filling in… Draw a single bond between all atoms. Give bonded atoms octets. If a central atom has only six total electrons, it double bonds to a bonded atom with a lone pair. 2- O C O C O O Example: Oxalate (C 2 O 4 2- ) Then put it all in brackets with a charge indicated.
BIG IMPORTANT NOTE DEUX The method just described works well for challenging dot structures. – In case of emergency, use this technique, but don’t forget to count your electrons.
Practice Covalent Bonding and VSEPR Theory Worksheet (Reverse, Middle Section) – Letters D, E, F, and G are challenging. Leave F for last!
Covalent Bonding Remember that there are two types of covalent bonding: Polar Covalent – Electrons are not shared equally. – Electrons are pulled closer to one atom than another because one atom is more electronegative. Non-Polar Covalent – Electrons are shared equally. – Pulls from both atoms are roughly equal strength.
Water Has Polar Bonds Let’s think of it this way. You tell me which of these two guys is gonna hold onto his electrons more… (draw them too) Meet Hydrogen – Poor guy’s only 2.1 Paulings in electronegativity. OMG OXYGEN! 3.5 PAULINGS ELECTRONEGATIVITY!!!1
But…since oxygen is such a strong atom, it doesn’t share electrons nicely. Having all that negative charge built up around oxygen causes that end of the molecule to take on a negative charge… …and all that’s left on the other end are the hydrogen protons, causing that end of the molecule to take on a positive charge… …which causes neighboring water molecules to form hydrogen bonds with the first one. This is the source of cohesion/surface tension, and why it hurts to do a belly-flop. In a diagram… HH O - Meet water: (say hi) There are 8 total valence electrons in the molecule H H O H H O - ++ Hydrogen Bond Covalent Bond
Polar Covalent Bonds As a result of that big oxygen dude hoggin’ all the negative electrons (and being more electronegative), his end of the water molecule takes on a slightly negative charge. – We use the (lower case) delta symbol to indicate a negative charge: - Because poor ol’ Hydrogen’s doesn’t see much of the electron action (and being less electronegative), his end gets a slightly positive charge (due to the hydrogen protons). – We use the (lower case) delta symbol to indicate a positive charge: +
Polar Covalent Bonds Because water molecules have poles, the positive end of one water molecule (hydrogen) can stick to the negative end of another (oxygen). This is what allows water to undergo adhesion and cohesion.
Non-Polar Covalent Bonds Electrons shared (fought-over?) equally. Example: H 2 (or an H-H bond)
Bonds and Molecules Now we separate the molecules from the bonds. As we have learned, individual bonds can be polar. However, as in the lab, sometimes molecules made only of polar bonds can be non-polar overall. Here’s what I mean…
A figurative look at it? Think of when a child pulls on his/her parent’s arm to lead him/her somewhere: ParentKid This is a polar bond, because “Parent” is likely stronger than “Kid.” Other Kid This is a polar molecule, because the pulls do not cancel each other. Add another kid with equal strength and you have another polar bond. This is a non-polar molecule, because the pulls do cancel each other. But if “Other Kid” is ripped, the pulls may not cancel out. This is a polar molecule, because the pulls do not cancel each other.
Polar Bonds, Non-Polar Molecules Cl C ++ -- -- -- -- Carbon tetrachloride (CCl 4 ) has 4 polar bonds. – Electronegativity difference of 0.5 in each direction. However, because all chlorine atoms are pulling equally on the electrons, the overall molecule remains symmetrical, and thus non-polar.
Identifying Polar Molecules If the “pulls” in a molecule are symmetrical and cancel each other, it’s non-polar. – Even if the bonds are polar. If they are asymmetrical, it is polar. – Example: PH 3 – Electronegativity difference: 2.1 (P) – 2.1 (H) = 0. – So all bonds are non-polar. BUT! There is an unshared pair on P, making the molecule assymetrical (trigonal pyramidal geometry), leading to a polar molecule*. P HH H *Up for debate.
H 2 O – Polar or Non-Polar? Here’s another model of water. We know water’s polar, but it appears to be symmetrical, right? However, when you look at the direction of the pulls, it is clearly not pulling equally in all directions. So, to determine molecular polarity, draw the “pulls” from polar bonds and see if they cancel out. ++ --
Another Look [sketch me] Let’s look at the “pulling” going on in a water molecule again: HH O O pulls downward on electrons more strongly than H pulls upward… …making O slightly negative… …and H slightly positive. KEY: To show this, we add a “plus” to the arrow. -- ++ ++
Polar or Non-Polar? Carbon dioxide (CO 2 ) – Non-polar; linear molecule Ammonia (NH 3 ) – Polar; trigonal pyramidal molecule Chlorine (Cl 2 ) – Non-polar; linear molecule
Practice Lab – Molecules I – Front: #2, 4 – Back: “Molecule Polarity” Column
Aside: Microwave Ovens Microwaves work by heating water molecules in food. – Since the plate usually doesn’t have water, it doesn’t get warm. How do they heat the water? With polarity! – The oven induces an electromagnetic field that causes the polar water molecules to orient in a certain direction. – Then the microwave changes the orientation of the field and the molecules need to change their orientation too. – The shifting back and forth and resulting friction creates heat and deliciousness.
Intermolecular Forces Water in Space!
Intermolecular Forces Time for another Australian video! Find your video worksheets and turn to the reverse side for Intermolecular Forces.
Intermolecular Forces Intermolecular forces are interactive forces between molecules. Here are three, in order of increasing strength: – WEAKEST: London Dispersion Forces Occur: Always Part of Van der Waals Forces – Dipole-Dipole Forces Occur: With two polar molecules. Part of Van der Waals Forces – STRONGEST: Hydrogen Bonding Forces Occur: When H is bonded to O, N, or F on one molecule and there’s an O, N, or F on another molecule.
London Dispersion Forces Demo I need a nucleus or two!
London Dispersion Forces The weakest type of intermolecular force. It results from the motion of electrons around the nucleus. When, momentarily, electrons in one atom happen to be on the same side of the atom, a transient dipole forms. – It’s like a temporary polar molecule. As a result, the molecule has a polar end and non-polar end, and nearby molecules can also gain transient dipoles, causing them to attract each other.
London Dispersion Forces Part of Van der Waals Forces:
Dipole-Dipole Forces Also part of Van der Waals forces, dipole- dipole forces occur when two polar molecules come together. The positive end of one polar molecule is attracted to the negative end of another. Polar molecules are called dipoles, after all!
Van der Waals Forces Where do we see van der Waals forces in nature? – Geckos, for one!
Hydrogen Bonding The strongest force between neutral molecules. – Also known as “when Hydrogen cheats on its molecule.” When is there H-bonding? (NEED BOTH CONDITIONS) – On one molecule, H is bonded to O, N, or F. – Another molecule with O, N, or F is nearby. Mostly because H only has two electrons.
Hydrogen Bonding Explained When is there H-bonding (long answer)? – When H, small little atom that it is, bonds to a strongly electronegative atom like oxygen, nitrogen, or fluorine, none of the electrons hang out near it at all, making it positive. – Similarly, the O, N, or F on another atom usually has all the electrons hanging out near it, making it negative. – As a result, the H on one molecule and the O, N, or F on another are attracted to one another. Like a specific dipole- dipole interaction. Or some kind of extra-molecular couple.
Intermolecular Forces Practice Draw the pair and identify: – The shape of each. – Whether each is polar or non-polar. – The intermolecular forces between the two. H 2 O and BF 3 – H 2 O Bent, BF 3 Trigonal Planar – H 2 O Polar, BF 3 Non-Polar – London Dispersion, H- Bonds HCN and CH 3 F – HCN Linear, CH 3 F Tetrahedral – HCN Polar, CH 3 F Polar – London Dispersion, Dipole-Dipole
Important Note Generally speaking, the more intermolecular forces between two molecules, the higher the boiling and melting points. – This is because the molecules hang onto each other more strongly.
Intermolecular Forces Practice Intermolecular Forces Worksheet
Bond Dissociation Energies For chemical reactions to occur, bonds must be broken. Bond dissociation energy is the amount of energy needed to break a covalent bond. – Of one mole of the substance – more to come. To find the dissociation energy of the whole molecule, add the bond dissociation energies of each bond in the molecule.
Bond Dissociation Examples CH * 4 = 1572 kJ/mol CO * 2 = 1472 kJ/mol H 2 O 464 * 2 = 928 kJ/mol Br kJ/mol I kJ/mol O kJ/mol O—O 145 kJ/mol NH * 3 = 1173 kJ/mol
Polystyrene Dissolution Demo Styrofoam (non-polar) Acetone (non-polar) Permanent marker (to draw the wicked witch or something)
Bond Dissociation Energies Styrofoam (called Polystyrene) Chemical Formula: C 8 H 8 Calculate the total bond dissociation energy for polystyrene. 7 C-H Bonds: 393 * 7 = 2751 kJ/mol 4 C-C Bonds: 347 * 4 = 1388 kJ/mol 4 C=C Bonds: 657 * 4 = 2628 kJ/mol Total 6767 kJ/mol
Resonance Example Let’s draw the nitrate ion: – NO 3 - Add up total electrons available: – 24. Put single bonds everywhere. Complete octets on bonded atoms. Satisfy the octet rule for the central atom. Did you put your double bond somewhere else? N OO O
Resonance in Nitrate – NO 3 - N OO O N OO O N OO O
Resonance Resonance occurs when more than one valid Lewis structure can be drawn for a particular molecule. The actual structure (or bond order) is an average of resonance structures. – Additional bonds are shown as dotted lines, as in this image. N OO O
Resonance in Benzene (C 6 H 6 )
Aside: Coordinate Covalent Bonds On your Covalent Bonding Worksheets, #6 on the front (SO 3 ) is an example of a substance with a coordinate covalent bond. This is when there is a “donation” of an electron pair from one atom to another. – Both electrons are donated from one atom. Here’s how to draw it: S OO O Double Bond Coordinate Bond Coordinate Bond
Aside: Coordinate Covalent Bonds Note that coordinate covalent bonds (also known as dipolar bonds or a dative covalent bond) frequently occur in resonance structures. A good sign of this is when an atom has all its valence electrons shown as lone pairs, yet is still forming a covalent bond. – Usually oxygen, it seems like, but could be others. Let’s redraw nitrate’s resonance structures to show coordinate covalent bonds.
Resonance in Nitrate – NO 3 - N OO O N OO O N OO O
What’s the Big Idea? Coordinate covalent bonds do not function any differently from the kind of covalent bonding you already know how to draw. Drawing a coordinate bond simply allows you to specify the origin of the electrons that are being shared. You do not need to show coordinate bonds for this class, but it’s good to know about them for those of you with AP Chem in your futures.
And the answer to #6? With coordinate bonds: As a single resonance structure: S OO O S OO O
Closure Wow. We’ve done a lot of stuff throughout this lesson. Questions?
Okay, one last thing… Should you decide to take AP Chemistry, you’ll also cover a little bit more of VSEPR Theory called hybridization. Hybridization occurs in covalent bonding, as electron repulsion changes the shape of orbitals (generally just the s and p ones, but sometimes d). When hybridization occurs, the orbitals re-shape to be more like a strange combination of both of them. When those orbitals form bonds with each other, the bond type is renamed as well. Here’s a short summary…
Orbital Hybridization Groups Attached to Central Atom GeometryHybridization 2Linearsp 3Trigonal Planarsp 2 4Tetrahedralsp 3 5 Trigonal Bipyramidal sp 3 d 6Octahedralsp 3 d 2
Hybrid Bonds Bond TypeHybrid Bond Types Single1 Sigma Bond* Double 1 Sigma Bond* 1 Pi Bond** Triple 1 Sigma Bond* 2 Pi Bonds** *Overlapping of two sp hybrid orbitals. **Overlapping of two non-hybridized p orbitals.