Intramolecular Bonding Chapters 4, 12 Chemistry Mr. McKenzie
What determines the type of intramolecular bond? An intramolecular bond is any force that holds two atoms together to form a compound; 3 types exist (ionic, covalent, metallic) Ionic Bond - attraction between two oppositely charged ions; occurs when valence electrons are transferred between atoms; involves metals and non-metals
Covalent Bond - formation of a compound when valence electrons are shared between two atoms; involves non-metals Metallic Bond - attraction by the nuclei of the atoms and a sea of de-localized valence electrons; electrons can easily move around (non-directional placement); involves metals
To predict the type of bond that will form, simply determine the difference in electronegativities >2.0 = ionic - electrons stolen 0 (<0.5) = covalent - electrons shared Between 0.5 and 2.0 = polar covalent - electrons shared unevenly
Polar covalent bonds establish dipole moments - a charge distribution of a molecule; electrons spend more time around the more electronegative atom, resulting in... a partial - ( - ) on the more electronegative atom a partial + ( + ) on the less electronegative atom
Why do ionic bonds form? Ionic bonds form when two atoms with large differences in electronegativities are brought together in close proximity The more electronegative atom (non-metal) takes electrons to establish a stable electron configuration (nearest nobel gas configuration) and forms an anion The less electronegative atom (metal) loses electrons to establish a stable electron configuration (nearest nobel gas configuration) and forms a cation
Once the ions are formed, they attract to one another to form an electrically neutral compound Total charge of the cations (+) must be cancelled by the total charge of the anions (-) To determine the ratio of atoms involved, use the criss-cross method Determine charge of the ions Bring the charges down as the subscript to the other ion (if it is a 2:2 or 3:3, reduce it to a 1:1)
What are polyatomic ions? Polyatomic ions are groups of atoms that behave as a single unit; when incorporated into a compound, treat them as a single species and keep them contained in parentheses Ion Name Ion Name NH 4 + ammonium CO 3 2- carbonate PO 4 3- phosphate HCO 3 - hydrogen carbonate (bicarbonate) NO 3 - nitrate SO 4 2- sulfate OH - hydroxide CrO 4 2- Chromate MnO 4 - Permanganate Cr 2 O 7 2- Dichromate
How do you name ionic compounds? Different rules exist for naming compounds Type I (binary compounds) - involves metals that can only form 1 cation Cation is always named first and the anion second Cation is simply the name of the element Anion is the root and the suffix -ide; polyatomics are not changed
Type II (binary compounds) - involves metals that can form multiple cations Cation is always named first and the anion second Cation is simply the name of the element with the charge of the cation written as a Roman numeral (II, III, IV, etc.) Anion is the root and the suffix -ide; polyatomics are not changed
How do covalent bonds form? Just like with ionic compounds, the atoms involved in covalent bonds are also trying to establish a stable electron configuration by sharing valence electrons Electron dot diagrams are used to represent the sharing of the valence electrons and to help determine the ratio of atoms in a molecule The total number of valence electrons are placed around the elemental symbol; if an ion is to be represented, place the symbol(s) inside brackets with the charge outside of the brackets C Na [ Cl ] - O x x xx xx [K] +
Draw the electron dot diagram for the following: Aluminum Fluorine Neon Boron ion Bromine ion
Most stable electron configurations require 8 electrons in the valence shell - octet rule F F One exception is with H as it only needs 1 electron to fill the 1s level - duet rule F F F 2 H H H H H 2 Following these rules helps construct the Lewis structure of the molecule - a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule
When drawing Lewis structures, you must remember that... 1. All valence electrons from all atoms must be shown 2. The octet (or duet) rule must be followed 3. Electrons must come in pairs Lone Pairs If 1 pair of e - shared between atoms - single covalent bond If two pairs of e - shared between atoms - double covalent bond If three pairs of e - shared between atoms - triple covalent bond 4. Electrons not involved in bonds are lone pairs xx xx O xx N xx xxx xx O N xx O x H x O N H O N
To draw Lewis structures... 1. Determine the total number of valence electrons 2. Form the bonds between the atoms 3. Distribute the remaining electrons as lone pairs (may need to incorporate double or triple bonds) 4. For ions, add 1 e - for each negative charge and subtract 1 e - for each positive charge; draw brackets around the structure with the charge outside Draw the Lewis structures for the following: C 2 H 6 C 2 H 4 C 2 H 2 OH - HCN CO 2
How do you name compounds with covalent bonds? 1. The first element in the formula is named first, and the full element name is used. 2. The second element is named as though it were an anion. 3. Prefixes are used to denote the numbers of atoms present. (mono- is never used for naming the first element) Prefix Number mono- 1 di- 2 tri- 3 tetra- 4 penta- 5 hexa- 6 hepta- 7 oxta- 8
Name the following compounds: N 2 O 5 NO BF 3 P 4 O 6 PCl 5 Determine the molecular formulas of the following compounds: dinitrogen trioxide dihydrogen monoxide iodine pentaflouride
What are the structures of molecules? The structural formula is a graphical representation showing how the atoms are arranged; these are important as they help chemists visualize what occurs during chemical reactions; a common type are Lewis structures (others exist)
Draw the Lewis structures for the following. Remember to represent all valance electrons as either bonded or lone pairs. CBr 4 NO 2 NCl 3 Cl 2 O
The geometric (molecular) structure of a molecule refers to the 3-D arrangement of the atoms around an atom; to determine the arrangement, we will use the valence shell electron pair repulsion theory (VSEPR) 1. Negative charge centers (location where electrons are found) around an atom will repel each other 2. Negative charge centers will try to maximize the distance between them 3. A double bond and triple bond behave the same as a single bond (it is an area of electrons...the number of electrons does not matter)
To use VSEPR... 1. Determine Lewis structure of the molecule Negative Charge Centers 2. Count the number of negative charge centers; the # tells the geometry and the bond angle between atoms
When lone pairs are involved, angles are slightly altered (less than if no lone pairs) and the appearance of the molecule is different 0 lone pair = tetrahedral 109.5 1 lone pair = pyramidal 107 2 lone pairs = bent 105
# Negative Charge Centers Geometry # of Negative Charge Centers that are Lone Pairs Appearance Bond Angle 4 Tetrahedral 0 Tetrahedral 109.5 1 Pyramidal ~107 3 Trigonal Planar 2 Bent ~105 0 Trigonal Planar 120 1 Bent <120 2 Linear 0 Linear 180
Draw the Lewis structure, determine the geometry, bond angles, and the overall appearance CH 4 PH 3 H 2 S H 2 O NH 3 NH 4 +