UNIT 4 Bonding and Stereochemistry. Stable Electron Configurations All elements on the periodic table (except for Noble Gases) have incomplete outer.

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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 Na+ 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

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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)

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Na+ Cl-

Charge on ion represented by + or – sign Positive ion- cation Negative ion- anion (use suffix –ide) Na+Cl- is sodium chloride (NaCl) 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+Cl- one to one NaCl Mg2+Cl- one to two MgCl2

Crystal Lattices Each ionic compound

makes specific shape based on arrangement

Crystal- solid whose particles are arranged in a lattice structure (NaCl- cubes, ruby- hexagonal)

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)

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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 N2O- 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, C2H2:

Stereochemistry- VSEPR Theory

All molecules have 3D shape Stereochemistry- study of shapes of molecules

VSEPR theory Valence Shell Electron Pair Repulsion Helps to understand and 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

EXAMPLES

Water, H2O

2 bond pairs

2 lone pairs The molecular geometry is

BENT.

H O H••

••

Carbon Dioxide CO2

As a Lewis Dot diagram:

Has two electron pairs - two sets of double bonds – two bonding pairs

Creates a LINEAR shape according to VESPR