Presentation is loading. Please wait.

Presentation is loading. Please wait.


Similar presentations

Presentation on theme: "COVALENT COMPOUNDS Chapter 6. 1) HOW DO COVALENT COMPOUNDS DIFFER FROM IONIC COMPOUNDS? Unit Essential Question:"— Presentation transcript:




4 Section 1: Covalent Bonds Covalent- electrons are shared. Simplest covalent bonds are diatomic molecules. Always 2 atoms; atoms can be different or the same (diatomic elements). Ways to remember diatomic elements: BrINClHOF, HON & the halogens, #7 Ex: H 2 Each H atom shares its valence electron so each can have 2. Called a shared pair of electrons..

5 Molecular Orbitals Where shared electrons are located. Region of high probability of finding the shared electrons- recall quantum model. Form when covalent bonds occur and atomic orbitals overlap.

6 Energy and Stability Non-bonded atoms have high PE and low stability. Become more stable when bonded. E is released when bonds form, PE decreases. This determines the bond length. Ideal length is when the two bonded atoms are at their lowest PE.

7 Energy and Stability Cont. Covalent bonds are flexible. Bond length is an average distance. Bond length oscillates like a spring between the two nuclei. Bond energy – energy required to break a bond. Different from lattice energy! More E is needed to break a shorter bond. Shorter bond length = stronger bond.

8 Re-visiting Electronegativity How do we know if atoms will transfer or share electrons in a bond? Check electronegativity values! If the difference is greater than 2.1 it is considered ionic. Ex: KCl Electronegativities: K= 0.8 and Cl= 3.2 Subtract: = 2.4, so the bond is ionic.

9 Re-visiting Electronegativity If the difference is less than 2.1 it is considered covalent. Ex: CO 2 Electronegativities: C=2.6 and O=3.4 Subtract: 3.4 – 2.6 = 0.8, so the bond is covalent. You do not need to memorize electronegativity values! They will be given to you on tests and quizzes. pg.194 Figure 6 shows electronegativity values

10 Re-visiting Electronegativity It is important to note that this cut-off value of 2.1 is arbitrary! Therefore, it is not 100% accurate. Properties of the compound must be investigated for better classification. Ex: magnesium chloride Electronegativities: Mg=1.3 and Cl=3.2 Subtract: 3.2 – 1.3 = 1.9 It looks like it should be a covalent bond, but properties indicate it is actually ionic.

11 Polar vs. Nonpolar Covalent If the electronegativity difference is less than 0.5 we will consider the bond to be nonpolar covalent. Means the atoms are essentially sharing the electrons equally. If the difference is between 0.5 and 2.1, we will consider it to be polar covalent. Means one atom attracts the electrons more towards itself; there is unequal sharing of electrons.

12 Determining Bond Type

13 Practice Classifying Bonds A) Classify the following bonds as ionic, polar covalent, or nonpolar covalent: 1) Li – Cl 2) B – C 3) N – O 4) Mg – Br 5) C – F B) Since F has the highest electronegativity value, can F ever form a nonpolar covalent bond? Why or why not? 2.2 ionic 0.6 polar covalent 0.4 nonpolar covalent 1.7 ionic 1.4 polar covalent

14 Simple Polar Molecules Ions have full charges (electrons are transferred). Polar molecules have partial charges (unequally shared electrons). Use Greek symbol delta δ (and +/-) to indicate partial charges. Dipole – molecule that contains both positive and negative partial charges. Ex: HCl

15 Simple Polar Molecules Continued Use electronegativity values to determine partial charges. Atom with larger e - neg value = δ - Atom with lower e - neg value = δ + Ex: HCl Elecronegativities: H=2.2 andCl=3.2 So: H is δ+ and Cl is δ- In other words, the electrons are more likely to be found at the Cl atom than the H because it has a larger electronegativity; this makes Cl more negative.

16 Different Properties for Different Bonds Besides electronegativity, you can predict bond type by the type of elements involved in the bond. Metallic = metal atoms (K, Cu, etc.) Ionic = 1 metal (typically) + 1 nonmetal Covalent = 2 or more nonmetals Properties of these compounds are determined by bond type. Recall ionic compound properties and how they stem from the crystal lattice.

17 Covalent Properties Some are soluble in water and some are not. Depends if the bonds are polar or not. Poor conductivity. Covalent compounds can exist as a solid, liquid, or gas. Depends upon the polarity of bonds. Tend to have low melting and boiling points. *Remember: covalent compounds are made of molecules, not ions!

18 Bond typeMetallicIonicCovalent Model Example PotassiumPotassium chloride Chlorine Melting point 63 o C770 o C-101 o C Boiling point 760 o C1500 o C (sublimes) o C Properties Soft, grey lustrous Conductor as solid Crystalline white solid Conductor when molten or in solution Green-yellow gas Insulator


20 Section 2: Drawing and Naming Molecules Lewis structures (electron dot diagrams) use dots and lines to show valence electrons. Atoms only use dots: Each dot represents a valence e -. Molecules use both dots and dashes: Each line represents a bond. Help to determine and show bonds that will form in covalent compounds.

21 Lewis Structures- Atoms The 4 sides of the element symbol represent the s and three p orbitals. s = 2 e- and p = 6 e- so total = 8 No more than 8 dots per symbol Put one dot on each side of the atom before pairing: Draw the Lewis structure for the following: Al, Br, N, Ne

22 Lewis Structures- Molecules When drawing Lewis structures for molecules, give each atom an octet. Except for H! Looking at the Lewis structures for each atom can help to determine how many bonds it will form. Ex: CH 4 Needs 4 e -, so it will form 4 bonds. Each H needs 1 e -, so they will form 1 bond.

23 Electron Pairs Shared pairs: e- are shared between two atoms, forming a bond. 1 shared pair = single bond or 2 shared pairs = double bond 3 shared pairs = triple bond Unshared/lone pairs: e- are not shared between two atoms, and are not involved in a bond.

24 Steps for Drawing Lewis Structures 1) Find the total # of valence electrons for all atoms in the compound. Ex: NH 3 N: 5e- H: 1e- total: 5 + 1(3) = 8e- 2) Arrange atoms/determine ‘backbone’ of the compound. H and halogens tend to be on the end. They don’t like to share more than 1 pair of e -. C is almost always in the center, if present. Likes to form 4 bonds. Except for C, the atom with the lowest electronegativity will be the center. H will NEVER be in the center! Ex: NH 3 N will be in the center:

25 3) Keep track of how many e- are used and how many remain. Each bond used to build the backbone uses 2 electrons. NH 3 has 3 bonds, so 6 e - used. 8 e e - = 2 e - left. 4) Distribute remaining e- by giving each atom an octet. Give the most electronegative atoms e- first. Ex: 2 e - left; N is the only one that needs an octet, so it gets the remaining 2 e -. 5) Verify the structure! Make sure all atoms have an octet and check that all valence electrons have been used. Steps for Drawing Lewis Structures

26 Polyatomic Ions 1)The bonds that hold together polyatomic ions are covalent bonds. 2)When drawing the Lewis structure for a polyatomic ion, you must take into account its charge when determining the total number of valence electrons you’re working with: negative ion: add electrons to the total positive ion: subtract electrons from the total 3) Finally, put brackets around the ion and include its charge. Ex: Draw the Lewis structure for the sulfate ion.

27 Practice Problem #2: Ions Draw the Lewis structure for the ammonium ion.

28 Practice Problem #4: Multiple Bonds 1) It is possible to run out of remaining electrons to give to atoms as lone pairs. But they still need an octet! 2) If this happens, you take lone pairs from a neighboring atom and form another bond to the atom that needs an octet. The neighboring atom donates both electrons to be shared in this bond. This is how double and triple bonds are formed. Ex: Draw the Lewis structure for O 2.

29 Practice Problems #5-6: Multiple Bonds Draw the Lewis structure for CO 2. Draw the Lewis structure for N 2.

30 Resonance Structures 1) Some compounds can be represented with more than one Lewis structure. All Lewis structures are drawn. A double-headed arrow is put between them to indicate all possibilities. 2) Compound is an average of the possible Lewis structures (resonance hybrid).

31 Naming Binary Covalent Compounds 1) Name the first element listed in the formula. The name does not change. 2) Second element always has ‘-ide’ ending. 3) Prefixes are used to indicate numbers of atoms of both elements. Exception: do not use ‘mono-’ for the first element listed if there is only one.

32 Naming Binary Covalent Compounds Take off the ‘a’ at the end of the prefix if the element begins with a vowel. Ex: pentoxide, not pentaoxide. Ex #1: CCl 4 Ex #2: P 2 O 5 Ex #3: H 2 O Ex #4: CO carbon tetrachloride diphosphorus pentoxide dihydrogen monoxide carbon monoxide


34 Section 3: Molecule Shapes 1) VSEPR theory: Valence Shell Electron Pair Repulsion Predicts shape based on electron repulsion. e- want to be as far apart as possible! Look at the central atom to determine the shape. 2) Unshared pairs repel more than shared pairs. All electrons repel each other, whether they are bonded or lone pairs. 3) In addition to bond polarity, shape also plays a role in determining the properties of a substance.

35 Predicting Shapes Count the number of shared and unshared pairs on the central atom in the Lewis structure. Count double/triple bonds as 1 shared pair. Total pairsShared pairs Unshared pairs Shape name 440Tetrahedral 431 Trigonal pyramidal 422Bent Linear 330Trigonal planar

36 Angles Bond angles also show repulsion of electrons in molecules. Tetrahedral angles = o Base all other angles off of this. From here, we can predict >, <, or = o Depends upon shared vs. lone pairs. (Additional factors that affect bond angles: atom size, multiple bonds take up more space. We won’t worry about these.)

37 Angles Cont. Trigonal Pyramidal: <109.5 o Remember- tetrahedral was 109.5, now we’ve replaced a shared pair with a lone pair that repels more! So angles decrease! Linear: > o (180 o is much bigger!) Trigonal Planar: >109.5 o (120 o is bigger!) Bent – depends on # of lone pairs Bent with one lone pair is like taking trigonal planar and replacing a shared pair with a lone pair, so angle is Bent with two lone pairs is like taking tetrahedral and replacing two shared pairs with two lone pairs, so angles are much < This is most common!

38 Polarity Tetrahedral, trigonal planar, and linear molecules can be nonpolar. Polar bonds can ‘cancel’ out other polar bonds because of the ‘symmetrical’ shape. This happens if the bonds are all the same. Ex: CH 4 & CO 2 They can also be polar if the bonds are not all the same. Bonds do not completely cancel out. Ex: CH 3 F

39 Polarity Cont. Trigonal pyramidal and bent molecules tend to be polar. Polar bonds are not cancelled out because of the ‘asymmetrical’ shape. Example: H 2 O Even if the bonds are not polar, the lone pair makes it polar. Let’s relate this back to covalent compound properties from section 1. Methanol is soluble in water. Oil is not soluble in water.


Similar presentations

Ads by Google