Chemical bonding

1. Chemical bonding Unit 9, Chapter 6

2. Electron-dot notation • An electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbols

4. Unshared or lone pair • A pair of electrons that is not involved in bonding and that belongs exclusively to one atom  Unshared Pair Shared Electrons

5. Lewis structures • Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs, or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons End Day 1

6. review • What is needed to draw a Lewis structure of either a molecule or atom? • How many double bonds are in the Lewis structure for carbon dioxide, CO₂? • Draw the Lewis structures for: • CO₂ • HCl • CF₄ • H₂O

7. bonding • A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together • When an atom forms a chemical bond, there is a decrease in its potential energy

8. bonding • Bond length – the distance between two bonded atoms at their minimum potential energy that is, the average distance between two bonded atoms • We know that main-group metals tend to lose electrons to form positive ions, or cations, and nonmetals tend to gain electrons to form negative ions, or anions. Therefore, bond time depends on electronegativity.

9. Octet rule • Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level • s and p orbitals must be completely filled by the total eight electrons

10. Ionic bonding • Ionic Bonding is chemical bonding that results from the electrical attraction between large numbers of cations and anions • A well-known ionic bond is NaCl. In sodium chloride the ions combine in a one-to-one ratio. • The type of bond can be estimated by calculating the difference in the element’s electronegativities. Ionic nature of a bond increases as electronegativity difference between 2 atoms increases.

11. Formation of ionic bonds

12. Characteristics of ionic bonds • Lattice energy – the energy released when one mole of an ionic crystalline compound is formed from gaseous ions • Lattice energy is an indication of the strength of the ionic bond

13. Polyatomic ions • A charged group of covalently bonded atoms • Polyatomic ions combine with ions of opposite charge to form ionic compounds • The charge of a polyatomic ion results from an excess of electrons (negative charge) or a shortage of electrons (positive charge)

14. Ionic compounds • Ionic compound - is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal • Formula unit – is the simplest collection of atoms from which an ionic compound’s formula can be established

15. molecules • Molecule – a neutral group of atoms that are held together by covalent bonds • Molecular compound – a chemical compound whose simplest units are molecules. Low boiling point is a property of molecular compounds.

16. formulas • Chemical formula – indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts • Molecular formula – shows the types and numbers of atoms combined in a single molecule of a molecular compound • Diatomic molecule – a molecule containing only two atoms Ex: O₂

17. Covalent bonds • Covalent bonds – results from the sharing of electron pairs between two atoms • Nonpolar-covalent bonds – a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. It is unlikely when two different atoms join because the atoms are likely to differ in electronegativity.

18. Polar bonds • Polar bonds – bonds with an uneven distribution of charge • Polar-covalent bonds – a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons

19. Resonance structures • Refers to the bonding in molecules or ions that cannot be correctly represented by a single Lewis structure

20. Structural formula • Indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule

21. Molecular geometry • Molecular polarity – the uneven distribution of molecular charge • Stereochemistry – the study of the spatial arrangement of atoms in a molecule • VSEPR Theory – states that the repulsion between sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible, i.e. move away from each other

22. VSEPR & Unshared electrons • Although unshared electrons occupy space around the central atoms, the shapes of the molecules depend only on the position of the molecules’ atoms

23. Metallic bonding • The bond formed from the attraction between positive ions and surrounding mobile electrons

24. Dipole-dipole forces • Forces of attraction between polar molecules

25. Hydrogen bonding • The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule

26. London dispersion forces • The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles