Chapter 4 Molecular Compounds 4.11 Naming Binary Molecular Compounds (No Metals!) When different elements combine, they form a binary compound. The electronegative element is written first. - A nonmetal farther left on the periodic table generally comes before a nonmetal farther right. - The IUPAC ordering scheme is: B, Si, C, P, N, H, S, I, Br, Cl, O, F (you do not need to memorize this.) Molecular formulas require identifying exactly how many atoms of each element are included. Rules: 1. Name the first element in the formula, using a prefix if needed. - Greek prefixes are used to designate the number of atoms of each element present (omit on first element.). - The final o or a is dropped when the prefix is attached to a name starting with the letter 'o', (e.g., CO is carbon monoxide not mono carbon monooxide) 2. Name the second element with a prefix and ending. Ch 4 Page 1
- Examples: N 2 O 3 SF 6 BBr 3 P 2 O 5 CCl 4 Exceptions (common names of binary compounds) - CH 4 methane (and other organic chemicals) - NO nitric oxide (nitrogen monoxide) - N 2 O nitrous oxide (dinitrogen monoxide) (laughing gas) - NH 3 ammonia gas - H 2 O water Ch 4 Page 2
4.1 Covalent Bonds The attractive forces that hold atoms together to form a chemical bond come in 2 general types: Ionic Bond (covered in the last chapter) - One or more valence electrons are from one atom to another. - This allows them to achieve a full. - Atoms are held together by force of (+) to (-). - This is the most common form of bonding between a and a. Covalent Bond - A bond formed by electrons between atoms. - This allows them to achieve a full. - Atoms are held together by connection. - This is the most common form of bonding between a and another nonmetal. - Molecule A group of held together by covalent bonds Ch 4 Page 3
Attractive and Repulsive Forces When two atoms approach we get a mixture of attractive and repulsive forces: - Nuclei each other - Electrons each other - Nuclei electrons If the attractive forces > repulsive forces, a is formed. Covalent Bonding in Hydrogen (H2): - Spherical 1s orbitals to give an egg-shaped region. - H-H, H:H and H 2 all represent a hydrogen. The size of the attractive and repulsive forces depends on the distance between the two nuclei. - Too far - weak - Too close - strong Bond Length: The distance between nuclei in a bond. Ch 4 Page 4
Covalent bonds can also be made by the overlap of orbitals. Chlorine exists as a molecule due to the overlap of 3p orbitals This can also be shown as: As we learned earlier, in addition to H 2 and Cl 2, there are 5 more elements that exist as diatomic molecules. (N 2, O 2, F 2, Br 2, I 2 ) for a total of 4.2 Covalent Bonds and the Periodic Table A molecular compound is one that is made up of not ions In these compounds, each atom shares enough electrons to achieve a noble gas configuration, or. Ch 4 Page 5
Electron Dot Structures (a.k.a. Lewis Dot Structures) for Molecular Compounds These are a way of representing the number of electrons surrounding or shared by each covalently bonded atom. These are cartoons, not realistic structures. - Lewis dot structures are only relevant for representative main group elements. - Remember, of the group number indicates the number of electrons in the valence s and p orbitals. - When atoms combine to form molecules Shared electrons are called Unshared electrons are called lone electrons or Completing the Octet Atoms tend to create enough bonds and share enough electrons in order to form a stable set of 8 electrons around each atom (i.e. OCTET RULE). The Octet Rule is a GUIDELINE that is often broken: Ch 4 Page 6
Exceptions to the octet rule Less than 8: More than 8: Only and below elements (like PCl 5 and SF 6 ) There needs to be orbitals available. Period 1 has only s (2). Period 2 has only s & p (8) Period 3 has s & p & d Problem: Which of the following molecules are likely to exist (y/n)? (a) (b) (c) (d) (a). (b).. (c). (d). Ch 4 Page 7
4.3 Multiple Covalent Bonds The bonding in some molecules cannot be explained by sharing a single pair of electrons. Some molecules will unpaired electrons and share more than 2 e- in order to satisfy the octet rule. Single bond - A covalent bond formed by sharing electron pair. (Represented by a single line: H-H) Double bond - A covalent bond formed by sharing electron pairs. (Represented by a double line: O=O) Example for the formation of a double bond. Triple bond A covalent bond formed by sharing electron pairs. (Represented by a triple line: N N) Example for the formation of a triple bond. Carbon, nitrogen, and oxygen are the elements most often present in multiple bonds. and will form double and triple bonds. only forms double bonds. Multiple covalent bonding is particularly common in organic molecules, which consist predominantly of the element. Note that in compounds containing multiple bonds, C still forms covalent bonds, N still forms covalent bonds, and O still forms covalent bonds. Ch 4 Page 8
4.4 Coordinate Covalent Bonds A coordinate covalent bond is the covalent bond that forms when both electrons are donated by the atom.. Once formed, a coordinate covalent bond is as any other covalent bond. An example of coordinate covalent bond formation is: All 4 N-H bonds in the product are. covalent bonds often result in unusual bonding patterns, such as nitrogen with four covalent bonds, or oxygen with three bonds (H 3 O + ). metals are good at forming coordinate covalent bonds to form coordination compounds. (They can be used to remove toxic metals) 4.5 Characteristics of Molecular Compounds Ionic compounds have - melting and boiling points - oppositely ions - attractive forces between particles. Ch 4 Page 9
Molecular compounds have - melting and boiling points - particles - attractive forces between molecules called forces - don t conduct The strength of the intermolecular forces affects the physical state. Very weak intermolecular forces Þ. Intermediate intermolecular forces Þ. Strongest intermolecular forces Þ. - Molecular solids are rarely soluble in water, and - conduct electricity when melted. Ch 4 Page 10
4.6 Molecular Formulas and Lewis Structures Formulas come in more than one style. Molecular formula A formula that shows the numbers and kinds of atoms in one molecule of a compound. e.g. - A formula gives the number of atoms that are combined in one molecule. - An formula gives only a ratio of ions (for a formula unit). Structural formula - A molecular representation that shows the connections among atoms by using lines to represent covalent bonds. When Dots are added to the structural formula to show any unpaired electrons, we get a complete: Lewis structure - A molecular representation that shows both the connections among atoms and the locations of lone-pair valence electrons. (It does imply the shape of the molecule!) Lone pair - A pair of electrons that is not used for bonding. NH3 has bonding pairs and lone pair. Larger, more complex molecules like isopropyl alcohol (rubbing alcohol) can be represented as: - Molecular formula - Condensed formula - Structural formula Ch 4 Page 11
For very large organic molecules it is often useful to take an even bigger shortcut to writing formulas that still imply structural information. Shorthand (Line-Bond/skeletal) Notation for Aliphatic Hydrocarbons Writing out even the condensed formula for organic molecules can become very burdensome, because of the large size of many molecules. Each bond is represented by a line in a zig-zag. The end of each line shown is assumed to be a carbon atom, unless expressly shown otherwise. Each carbon atom must have bonds. Any bonds not shown are assumed to be C-H bonds. (C atoms will not have lone pairs of electrons.) Halogens, O, and N atoms require a full octet. If they have less than a full octet in the line-bond structure, they fill the octet with lone pairs of electrons. Remember to count the initial carbon! Simple examples: Propane Butane Decane condensed structural line-bond formula formula structure A little bit more complex example. CH 3 CH 2 CHCClCH 3 Ch 4 Page 12
Problem: Write condensed formulas and skeletal line-bond structures for: Problem: Determine the molecular formula for the following: OH Ch 4 Page 13
4.7 Drawing Lewis Structures Rules for Writing Lewis Dot Structures for Molecules 1. Add up valence electrons from all atoms - It does not matter which atom they come from. - 1 e - for each (-) charge - 1 e - for each (+) charge 2. Identify the central and terminal atoms and write the skeletal structure - The central atom will usually be one of the lowest electronegativity (least electron loving) (H excepted) - Electronegativity is the degree to which an atom can draw a shared pair of electrons in a covalent bond towards itself. 3. Draw a bond between each pair of atoms. - Each bond uses up of the available e -. 4. Assign remaining e- to - elements first to fill octet (duet for H) 5. If there are leftover e-, - place them on atoms that have empty d orbitals (period 3 and below). 6. If all e- are used up and the central atoms do not have completed octets, - move one or more lone pairs to form a bond. Ch 4 Page 14
Usually for C, N, O, X (halogen) and H - C forms covalent bonds and lone pairs. - N forms covalent bonds and lone pairs. - O forms covalent bonds and lone pairs. - X forms covalent bonds and lone pairs. - H forms covalent bond Examples for Rows 1 + 2 (Octet rule never exceeded) H2O NH3 HCN Ch 4 Page 15
Problem: Draw the Lewis structure for N2H4 and CO2 Problem: How many double bonds and lone pairs should be shown for the molecule to the right?? 1) Double Bonds a. 1 b. 2 c. 3 d. No double bonds are present 2) Lone Pairs a. 2 b. 3 c. 4 d. 6 Problem: What is the most likely value of x in the molecule CHClx? a. 1 b. 2 c. 3 d. 4 Ch 4 Page 16