UNIT 4 Bonding and Stereochemistry Stable Electron Configurations

UNIT 4 Bonding and Stereochemistry

Stable Electron Configurations All elements on the periodic table (except for Noble Gases) have incomplete outer energy levels Valence electrons- electrons in outer energy level of atom Elements will gain, lose, or share electrons to get full outer levels (octet rule) Eight electrons = STABLE!!! Electron dot diagrams help to visualize valence electrons Symbol represents nucleus and inner electrons Dots represent valence electrons Group # = valence electrons

Drawing Electron Dot Diagrams Determine number of valence electrons from periodic table Draw the symbol for the element Place dots around the symbol, one per side, until all valence electrons are accounted for Example- Aluminum with 3 valence electrons Al

Ions Charged atoms where the number of protons and electrons is not equal Charge indicates how many electrons are added or subtracted Negative charge- ADD electrons Positive charge- SUBTRACT electrons Example Sodium ion Na+ atomic number 11 = 11 electrons subtract one electron = 10 electrons

Chemical Bonds Forces that hold groups of atoms together and make them function as a unit. A bond will form if the energy of the pairing is lower than that of the separate atoms. Some elements have stronger attractions to e- when bonded ELECTRONEGATIVITY (EN) Relative attraction an atom has for shared electrons in a covalent bond Unit- paulings Arbitrary number used for comparison purposes F is 4. 0, Cs/Fr 0. 7 Copyright © Cengage Learning. All rights reserved 8

Increase from left to right in a period- nonmetals higher than metals Decrease from top to bottom in a group Metals on left side and nonmetals on right side most reactive (alkali metals and halogens) Electrons attracted to the higher EN element Using EN to predict bonds Ionic Bonds- Metal + nonmetal Covalent Bonds (nonmetal + nonmetal) Polar- EN is different Nonpolar- EN is same value

Types of Chemical Bonds Ionic Bonds Some elements achieve stable configurations by transferring electrons Example- sodium and chlorine Sodium 1 valence electron Chlorine 7 valence electrons Both want to be stable Sodium will lose the one electron, and chlorine will gain that electron, forming IONS (atoms that have gained or lost electrons) Na+ Copyright © Cengage Learning. All rights reserved Cl 11

Charge on ion represented by + or – sign Positive ion- cation Negative ion- anion (use suffix –ide) Na+Cl- is sodium chloride (Na. Cl) Groups 1, 2, and 3 will lose electrons Groups 5, 6, and 7 will gain electrons Group 4 will go either way- usually share though

Ionic compounds- compounds that contain ionic bonds Can be made with single elements or polyatomic ions Empirical formula- shows ratios of ions contained in the bond Na+Clone to one Na. Cl Mg 2+Cl- one to two Mg. Cl 2

Crystal Lattices Each ionic compound makes specific shape based on arrangement Crystal- solid whose particles are arranged in a lattice structure (Na. Cl- cubes, rubyhexagonal)

Properties of ionic compounds High melting point, boiling point Poor conductor when solid, good when molten/dissolved Crystal structure- shatters when hit

Covalent Bonds Nonmetals have high ionization energy Don’t usually form ions-share electrons to get to stable energy level Covalent bond-chemical bond in which two atoms share a pair of valence electrons May share one (single bond), two (double bond), or three pairs (triple bond) 17

Form molecules Neutral group of atoms that are joined together by one or more covalent bonds May exist as diatomic molecules Made of 2 atoms of same element

May form single or multiple bonds Molecular Formula- expression of the number and type of atoms that are present in a single molecule of a substance. Subscript tells how many of each element are present N 2 O- 2 atoms of N, 1 atom of O

When atoms share electrons, they rarely share equally One element will “attract” electrons more than the others Polar Covalent Bond- a covalent bond in which electrons are not shared equally Atom with greater attraction gets a partial negative charge (δ-), lesser attraction partial positive charge (δ+)

Metallic “bonding” Attraction of metal atoms and the sea of electrons surrounding them Gives metals their properties Malleability Good conductors

Writing Lewis Diagrams for Molecules Steps Count all valence e Draw skeleton structure Put a pair of e- between all atoms to show a covalent bond All should have 8 (exc H which has 2) Distribute lone pairs around atoms (exc H) If an atom needs more e- then move pairs between atoms to get 8 Most multiple bonds are in C, N, and O Double bond (share two), triple bond (share three)

A double covalent bond, or simply a double bond, is a covalent bond in which two pairs of electrons are shared between two atoms. Double bonds are often found in molecules containing carbon, nitrogen, and oxygen. A double bond is shown either by two side-by-side pairs of dots or by two parallel dashes

A triple covalent bond, or simply a triple bond, is a covalent bond in which three pairs of electrons are shared between two atoms. • example 1—diatomic nitrogen: • example 2—ethyne, C 2 H 2:

Stereochemistry- VSEPR Theory All molecules have 3 D shape Stereochemistry- study of shapes of molecules VSEPR theory Valence Shell Electron Pair Repulsion Helps to understand predict molecuar geometry (from Lewis Dot diagrams) Developed by Gillespie and Nyholm in 1956 -57 Rules based on the idea that the arrangement in space of the covalent bonds formed by an atom depends on the arrangement of valence e e- try to push each other far away while still bonding to central atom

Restricted VSEPR rules Valence e- pairs (both shared and lone) arrange themselves around the central atom in a molecule in such a way as to minimize repulsion (as far away from each other as possible) When predicting molecular geometry, double and triple bonds act like single bonds Lone pairs of e- occupy more space than bonding e-

Steps to draw VSEPR molecules Draw Lewis Diagram Determine the central atom (lowest EN) Count the number of bonding and lone pairs surrounding the central atom Multiple bonds count as one pair Shape molecule in order to minimize repulsion Find on VSEPR chart