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Introduction to Chemical Bonding

Introduction to Chemical Bonding.

charlesrivera
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Introduction to Chemical Bonding

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  1. Introduction to Chemical Bonding • WHY?Because as independent particles, atoms are at relatively high potential energy. Nature, however, favors arrangements in which potential energy is minimized. By bonding with each other, atoms decrease in potential energy thereby creating more stable arrangements of matter. Atoms seldom exist as independent particles in nature. Even the air you breathe is made up of molecules held together by chemical bonds.

  2. Introduction to Chemical Bonding • A CHEMICAL BOND is a mutual electrical attraction between the nuclei (protons) of one atom and the valence electrons of another atom that binds the atoms together. IONIC BOND: the electrical attraction between large numbers of cations and anions COVALENT BOND: the sharing of electron pairs between two atoms. METALLIC BOND: the attraction between metal atoms and the surrounding sea of mobile electrons of metals.

  3. IONIC or COVALENT? • Bonds are rarely purely ionic or purely covalent and usually fall somewhere in between. • Recall that electronegativity is the ability of an atom to attract electrons. This property is used to determine if a bond is covalent or ionic. • You can calculate if a bond is more covalent or more ionic by finding the difference between the bonding elements’ electronegativities.

  4. POLAR vs. NON-POLAR • Covalent bonds can be categorized into two major types: polar covalent and non-polar covalent. • Polar covalent bonds involve an UNEQUAL sharing of electrons between the two atoms involved in the bond. • Non-polar bonds involve an EQUAL sharing of electrons between the two atoms involved in the bond.

  5. CALCULATING BOND TYPE • Determining the type of bond formed between two types of atoms is easy. Just subtract their two electronegativity values. Ex: H (electronegativity value: 2.1) 4.0-2.1=1.9 and F (electronegativity value: 4.0)

  6. ELECTRONEGATIVITY DIFFERENCES • If the difference is greater than 1.7, the bond is ionic. • If the difference is 1.7 or less but greater than .3, the bond is polar covalent. • If the difference is .3 or less, the bond is non-polar covalent. • 0-.3 < .31 – 1.7 < 1.71- 3.3 NONPOLAR POLAR IONIC

  7. SECTION REVIEW • What is the main distinction between ionic and covalent bonding? • How is electronegativity used in determining the ionic or covalent character of a bond between two elements? • What type of bonding would be expected between: H and F __________________ • Cu and S __________________ • I and Br __________________

  8. 6.2 Covalent Bonding & Molecular Compounds • Molecule- a neutral group of atoms that are held together by covalent bonds. • Molecular compound - a chemical compound whose simplest units are molecules. • Molecular formula - shows the types and numbers of atoms combined in a single molecule of a molecular compound

  9. 6.2 Covalent Bonding & Molecular Compounds (cont.’d) • Diatomic molecule - a molecule containing only two atoms. • Bond length – the distance between two bonded atoms at their minimum potential energy (the average distance between two bonded atoms) • Bond energy – the energy required to break a chemical bond and form neutral isolated atoms

  10. Bond Lengths vs. Bond Energies UseTable 6-1 Pg. 168 Graph 10 pairs of coordinates Using bond length vs. bond Energy on the grid. What is Relationship between bond Length and bond energy? _________________________ _________________________ _________________________

  11. Binary Molecular Compounds & Their Nomenclature • A binary MOLECULAR compound is composed of two different kinds of nonmetal atoms. Ex. N2O4 • Binary molecular compounds are named by writing the name of the less electronegative element first and the more electronegative element second.

  12. Binary Molecular Compounds & Their Nomenclature PREFIXES are used in front of each element’s name to indicate how many there are in the formula. The second element’s name is replaced with the suffix IDE. N2O4 : dinitrogen tetraoxide Se3Cl6: triselenium hexachloride CO: carbon monoxide **Notice how the prefix “mono” is used on the 2nd element but not the first.

  13. Remember: Molecular Compounds are composed of two non-metals whose electronegativity difference is between .3 and 1.7. That means that they are bonded covalently and share electrons. Also: The prefix system of nomenclature is an older system of nomenclature. A newer system exists and will be covered later. PREFIXES: Mono Hexa Di Hepta Tri Octa Tetra Nona Penta Deca Common nonmetals with “IDE” suffix: Oxide Nitride Sulfide Fluoride Chloride Iodide Bromide Phosphide Selenide Telluride

  14. NO ____________ N2O ____________ NO2 ____________ N2O3 ____________ N2O5 ____________ PCl5 ____________ SF6 ____________ P4O10 ____________ SO2 ____________ sulfur difluoride carbon tetraiodide disulfur dichloride pentaselenium disulfide disilicon octafluoride trihydrogen monophosphide Practice Set Write the name or the formula for the following molecular compounds.

  15. The Octet Rule • Chemical compounds tend to form so that each atom, by gaining, sharing or losing electrons, has an octet of electrons in its highest occupied energy level. • There are few exceptions to the rule: Elements like H and B only have two and three electrons respectively. Hydrogen can only form one bond (2 electrons total) and Boron can form three bonds (6 electrons total).

  16. Lewis Structures Lewis Structures: formulas in which electron dot notation is used to represent molecules. The electrons between the H and the Cl atoms are the shared pair of electrons. The other three pairs around the Cl are called lone pairs or unshared pairs since they are not involved in bonding.

  17. The two dots between hydrogen and chlorine represent a single bond and can be represented with a single dash in place of the dot pair. When only one pair of electrons are shared, the atoms have formed a SINGLE BOND. H Cl The diagram to the left illustrates a STRUCTURAL FORMULA which indicates the kind of atoms, the arrangement and the bonds but NOT THE UNSHARED PAIRS OF ELECTRONS.

  18. Examples of Lewis structures Cl2 O2 N2 H2O NH3 CH4

  19. Independent Practice • SO 3. NBr3 5. H3P 2. CF4 4. CH2Cl2 6. N2

  20. Expanded Valence • Some elements can be surrounded by more than eight electrons when they combine with HIGHLY ELECTRONEGATIVE elements like fluorine, oxygen or chlorine. Ex. SF6 • Lewis structure Couper structure

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