CHEMISTRY NOTES 6.1 COVALENT BONDS Objectives Explain the role and location of electrons in a covalent bond. Describe the change in energy and stability that takes place as a covalent bond forms. Distinguish between nonpolar and polar covalent bonds based on electronegativity differences. Compare the physical properties of substances that have different bond types, and relate bond types to electronegativity differences. Standards: 1c, 2b, 2g COVALENT BONDS Composed of a nonmetals sharing valence electrons. o Nonmetals have similar attraction for electrons (or similar electronegativity) making it difficult for either nonmetal to remove an electron from the other, causing them to share electrons instead. o The nucleus of one atom is attracted to the electron cloud of a neighboring neutral atom and vice-versa. Electrons are shared and move in the space surrounding the nuclei of the two atoms called the molecular orbital. o Orbitals of the atoms that are sharing electrons overlap The shared electrons are closer to the more electronegative atom o The atom that is more attracted to electrons will have the shared electrons closer to it. o If two nonmetals of the same type of element such as two H s are sharing electrons the electrons will be found somewhere between them but not closer to one or the other. ENERGY AND STABILITY Covalent bonds form when the attraction between two atoms is balanced by the repulsion and the potential energy is at a minimum. Most atoms are not stable (except noble gases) but become more stable when they are part of a compound. o Unstable => far apart, by themselves o Stable => form a compound with another atom. Neutral atoms by themselves far apart from other atoms have high potential energies (maximum potential energy) o potential energy the energy that the atom can release when it becomes part of a compound As neutral atoms get closer they lose their potential energy to it s surroundings. o Atoms come close together to bond. o The closer atoms get to each other the more energy they lose. o Atoms can only get so close before the repulsion forces take over. o Once atoms bond they have minimum potential energy. The amount of potential energy atoms lose determines the bond strength. o more energy released => stronger bond => shorter bond => more energy will be needed to break it! Bond length the distance between two bonded atoms at their minimum potential energy measure in pm (picometers) Atoms/bonds vibrate so the bond distance is the average distance between the atoms. Length of bonds (from longest to shortest): Single, Double, Triple! Bond energy the energy required to break the bonds in 1 mol of a chemical compound. Strength of bonds (from weakest to strongest): Single, Double, Triple
TYPES OF COVALENT BONDING Nonpolar covalent bonding o Occurs when atoms are equally attracted to electrons causing them to share electrons equally. o Atoms have the same or very similar electronegativities.! Weak attractions => Weak bonds o Electrons reside in the middle of the atoms sharing electrons. Polar covalent bonding o Occurs when atoms are not attracted to electrons equally causing them to not share electrons equally. o Atoms have very different electronegativities.! Strong attractions => strong bonds. o Electrons shared will reside closer to the atom that is more electronegative.! One part of the molecule will be relatively more negative or positive than the other side of the molecule.! The part of the molecule that has the electrons closer has a partial negative charge δ-.! The part of the molecule that doesn t have electrons closer to it will have a partial positive charge δ+. When this occurs a dipole is formed and the symbol will be drawn underneath the molecule to show which side is positive and which side is negative. TO DETERMINE WHETHER BOND WILL BE IONIC, POLAR COVALENT AND NONPOLAR COVALENT: Use the electronegativity values and find the difference (subtract the smaller electronegativity value from the larger electronegativity value. o If the difference is 0-0.4 : nonpolar covlant 0.5-2.1: polar covalent >2.1 : ionic PROPERTIES OF SUBSTANCES METALLIC BONDS o Result from the attraction between electrons in the outermost energy level of each metal atom and all of the other atoms in the solid metal o Atoms are held together because all of the valence electrons are attracted to all of the atoms in the solid o Valence electrons can move easily from one atom to another! Since electrons are free to roam around in the solid they are good conductors of electricity. o Have higher melting/boiling points than covalent compound but not higher than ionic compounds. POLARITY IS RELATED TO BOND STRENGTH Greater attractions => stronger bonds Polar molecules are attracted to each other because of the δ+ and δ ends of the molecule, therefore polar bonds are stronger than nonpolar bonds. Polar bonds are stronger so they will have higher melting/boiling points than nonpolar bonds o The strongest covalent bonds and the highest boiling points occur between H and F because the electronegativity difference between them is the highest. Order of strength from strongest to weakest (highest to lowest boiling points): o ionic bonds o metallic bonds o polar covalent bonds o nonpolar covalent bonds
CHEMISTRY NOTES 6.2 DRAWING AND NAMING MOLECULES Objectives: Draw Lewis structures to show the arrangement of valence electrons among atoms in molecule and polyatomic ions Explain the difference between single, double, triple covalent bonds. Draw resonance structures for simple molecules and polyatomic ions, and recognize when they are required. Name binary inorganic covalent compounds using prefixes, roots, and suffixes. Standards: 2a,e LEWIS ELECTRON-DOT STRUCTURES (ELECTRON-DOT DIAGRAMS) A structural formula in which electrons are represented by dots; dot pairs or dashes between two atomic symbols represent pairs in covalent bonds. A two-dimensional shape of a molecule. The dots represent valence electrons Valence electrons electrons found in the outermost shell of an atom and determines the atom s chemical properties. DRAWING LEWIS STRUCTURES FOR SINGLE ATOMS 1. Distribute the valence electrons as dots around the atom one at a time before pairing them up (pair them up as you go around the atom the second time) 2. Lone pairs (unshared electrons) unpaired electrons DRAWING LEWIS STRUCTURES WITH MANY ATOMS 1. Determine the number of valence electrons for each atom 2. When you have two atoms put them next to each other 3. When you have more than two atoms place the least electronegative atom in the center Carbon is usually placed in the center. Halogens and hydrogen are usually placed at the end of the molecule (or around the central atom) 4. Distribute the total number of electrons as dots around each atom and make sure the that the octet rule is satisfied All elements except H, B, and Be need eight electrons o Be needs to have 4 o B needs to have 6 o H needs to have 2 5. Draw the bonds: every two electrons that are in between two atoms forms a bond, so replace every two dots that are between atoms with a slash Single bonds represent two electrons 6. If you run out of dots move the lone pairs in between the atoms to form a double or triple bond. Double bonds represent four electrons Triple bonds represent six electrons 7. Verify that atoms follow the octet rule except for H, Be, and B. DRAWING LEWIS STRUCTURES FOR POLYATOMIC IONS 1. Count up valence electrons 2. If the ion contains a + charge subtract one electron, if ion contain a 2+ charge subtract two and so on; if ion contains a charge add one electron, if it contains a 2- charge add two electrons and so on. 3. Follow steps 2-7 from above 4. Place brackets around the entire molecule and put the charge on the upper right hand corner outside of the bracket that indicates the charge of the polyatomic ion. RESONANCE STRUCTURE When a molecule has more than one possible structure. Draw all possible structures and place a double end arrow ( ) in between.
NAMING COVALENT COMPOUNDS use prefixes to indicate how many atoms of each element are in the molecule. Prefixes: mono 1 atom (used only for the second element) di - 2 atoms tri- 3 atoms tetra - 4 atoms penta 5 atoms (when you have two vowels next to each other drop the first one: P 2O 5 is diphosphorus pentoxide not diphosphorus pentaoxide) hexa - 6 atoms hepta- 7 atoms octa - 8 atoms nona- 9 atoms deca - 10 atoms WRITING FORMULAS FOR COVALENT COMPOUNDS use the prefixes to determine how many atoms of each element are in the molecule mono is not used for the first element, so if an element doesn t have a prefix than there must only be one atom of that element.
CHEMISTRY NOTES 6.3 MOLECULAR SHAPES Objectives: Predict the shape of a molecule using VSEPR theory. Associate the polarity of molecules with the shapes of molecules, and relate the polarity and shape of molecules to the properties of a substance. Standards: 2e,f,h MOLECULAR SHAPES/STRUCTURES The three-dimensional shape of a molecule, or the 3-D arrangement of the atoms in a molecule Uses the VSEPR (Valence Shell Electron Pair Repulsion) theory to predict the molecular shapes based on the idea that pairs of valence electrons surrounding an atom repel each other. DRAWING MOLECULAR STRUCTURES When drawing Molecular Structures make sure you include the following: o 3-dimensional drawing o bond angles (180, 120, 109.5 ) o Name of molecular structure (linear, bent or v-shaped, trigonal planar, trigonal pyramid, or tetrahedral) STRUCTURES/SHAPES o Linear (angle: 180 ) o Bent or v-shaped (angle: 120 ) o Trigonal planar (angle: 109.5 ) o Trigonal pyramid (angle: 109.5 ) o Tetrahedral (angle: 109.5 )
MOLECULAR SHAPE AFFECTS A SUBSTANCE S PROPERTIES Polarity is not only affected by electronegativity values, it is also affected by the molecular shape. Shape affects polarity o Some linear molecules like carbon dioxide have two polar bonds but because of it s molecular shape their polarity cancels out. Linear, trigonal planar and tetrahedral molecules => nonpolar Bent or v-shaped, trigonal pyramid => polar Polarity affects Properties o Polar substances are attracted to other polar substances These will create stronger molecules => high boiling/melting points o Nonpolar substances are attracted to other nonpolar substances These will create weak molecules => low boiling/melting points