UNIT 7: BONDING Why do elements form bonds? How can we explain and draw ionic bonds? How can we explain and draw covalent bonds? What are metallic bonds and why are they good conductors? What is the difference between bond polarity and molecule polarity? How do we predict shapes of covalent molecules? What are the different forces that hold molecules together?

Aim: What previous knowledge will help us understand the bonding unit? 1. Why do atoms become ions? 2.How do atoms become ions?  3. How do metals form ions? 4. How do nonmetals form ions? 5. Conductivity: 6. Electricity: 7. Melting Point: 8. Boiling Point:

Aim # 1- Why do elements form bonds? A chemical bond is - the force of attraction between the atoms of a compound - the proton of one atom is attracted to the electron of another atom

the combining atoms either lose, gain, or share electrons in order to complete their outer shells

Energy and Chemical Bonds Endothermic- energy is absorbed, bonds break Energy is consumed for the bond to break Ex) AB + energy  A + B Exothermic- energy is released, bonds form Creating bonds creates stability Ex) A + B  AB + energy “BARF”

Types of bonds They differ in the types of elements involved. Also, how the valence electrons are handled.

Types of Bonds Ionic Bond: Covalent Bond: electrons are shared electrons are transferred from one atom to another Covalent Bond: electrons are shared Metallic electrons are mobile Polar Covalent: Unequal sharing Nonpolar Covalent: Equal sharing

Aim # 2: How can we explain and draw ionic bonds?  An IONIC bond is The Transfer of electrons Attraction between oppositely charged ions Bond between a metal and a nonmetal For example when Na and Cl atoms come together

When metal atoms react: They lose electrons They become + charged ions (cations) They acquire a complete octet Their radii decrease (become smaller) “MELPS”

Ionic bond: When nonmetal atoms react: They gain electrons They become – charged ions (anions) They acquire a complete octet Their radii increase

Properties of Ionic solids High MP and BP Hard substances Conduct electricity in the liquid phase and in solutions ONLY Good solubility Atoms have an electronegativity difference of 1.7 or greater

 DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS There are several steps to follow in order to draw the Lewis dot structure for ion compounds. Write the metal symbol with no dots in brackets (brackets are optional) Place the charge at the top right of the bracket Write the nonmetal symbol with 8 dots around it (except H!) Draw brackets around the symbol and place the charge of the ion at the top right of the bracket

Drawing Lewis Dot Structures for Ionic Bonds NaF

**The exception to the rule ** Polyatomic ions are composed of multiple atoms (table E) They have both covalent and ionic bonds

Compound Bond Type Dot Structure KF MgI2 BeS AlBr3 BaCl2 SrI2

AIM# 3: How can we explain and draw covalent bonds? These compounds are formed when two or more nonmetals share electrons nonmetals have the ability to have single, double and triple bonds Each bond contain 2 electrons

Types of Covalent bonding Nonpolar Covalent Bond Electrons are shared equally Occurs between atoms of the same element Little or no difference in electronegativity Polar Covalent Bond Electrons are shared unequally One atom is pulling on electrons more strongly Electronegativity difference is less than 1.7 but greater than 0

COVALENT BONDING (Nonpolar) All Diatomic Molecules are covalent. They are Br2, I2, N2, Cl2 ,H2, O2, F2, N2 - Is the only diatomic molecule that has a triple bond at room temperature.

Coordinate Covalent Bonds -A type of covalent bond in which BOTH electrons come from the same atom -Exists in polyatomic ions -Examples: NH4+ H3O+

Properties exist in gas, liquid, or solid state Good insulators Poor conductors of electricity in any phase Low melting points Soft substances Poor solubility

Network Solids These are covalent compounds that are extremely hard and have very high melting and boiling points. Strong bonds and strong intermolecular forces Cannot conduct electricity Ex. Diamonds, silicon dioxide (SiO2)

   CONSTRUCTING LEWIS DOT STRUCTURE FOR (SINGLE BONDS) COVALENT MOLECULAR COMPOUNDS Determine valence electrons in total (add them up for each element in the compound) Divide by 2 to determine the number of pairs of electrons in total for the compound Place first pair between the two elements (use a dash – to represent the shared pair) Place remaining pairs around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

H2 Cl2 Br2 HCl

 CONSTRUCTING LEWIS DOT STRUCTURES FOR (SINGLE BONDS) COVALENT MOLECULAR COMPOUNDS (more than two elements involved)   Determine valence electrons in total (add them up for each element in the compound Divide by 2 to determine the number of pairs of electrons in total for the compound Determine the most electronegative element and place it in the middle Place the other elements around it Start placing pairs (as dash lines) between the central atom and the terminal atoms Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

NH3 CH4 CCl4

 CONSTRUCTING LEWIS DOT STRUCTURES FOR MULTIPLE COVALENT MOLECULAR COMPOUNDS (more than two elements involved)   Determine valence electrons in total (add them up for each element in the compound Divide by 2 to determine the number of pairs of electrons in total for the compound Determine the most electronegative element and place it in the middle Place the other elements around it Start placing pairs (as dash lines) between the central atom and the terminal atoms Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons) If octet rule is not yet reached you can make additional pairs of electrons into double or triple bonds until octet rule is obeyed by all elements *can only be done with CNOPS

CO2 O2 N2

AIM #4 : What are metallic bonds and why are they good conductors? Metallic atoms lose their valence electrons easily (remember they have low ionization energy) Arranged in fixed positions called a crystal lattice. Therefore their bonds are a force of attraction between the negatively charged electrons and positively charged nucleus. The electrons are not attached to a particular nucleus mobile see of electrons (MSOME)

Properties of Metals Malleability High MP Good conductors in any phase Insoluble Luster

Aim # 5 what is the difference between bond polarity and molecule polarity BOND POLARITY (1.7 Rule) To determine the bond type, this is based on the difference in electronegativity The higher the difference the more polar the bond is Ionic bonds will have an electronegativity difference of greater then 1.7 Remember “BONDED”

Polar Covalent Bonds HCl H2O NH3 Have electronegativity differences of 0.4 and 1.7 HCl H2O NH3

Nonpolar covalent bond Have electronegativity difference of 0 and 0.3 Cl2 Br2

Nonpolar covalent bond Ionic bond Covalent bonds

Molecular Polarity (remember “SNAP”) - asymmetrical - unequal sharing of electrons - asymmetrical distribution of charge - lone electrons on an atom HCl H2O NH3

Nonpolar Molecules - symmetrical - equal sharing of electrons - 2 atoms are the same, no lone electrons Cl2 CH4

Compounds with polar bonds and nonpolar shapes CCl4

Aim # 6 How do we predict shapes of covalent molecules?  MOLECULAR SHAPES To determine the shapes of molecules we use Lewis Structures and the VSEPR theory VSEPR: Valence shell electron pair repulsion will predict the shapes of molecules from electron pairs on the central atom Linear – line, “HX” HCl CO2 N2

Bent (angular) – Hydrogen and Group 16 H2O H2Se Pyramidal- going to have one side unoccupied (NX3, PX3) NF3 PI3 Tetrahedral- CX4 CCl4 CH3Br

Aim # 7: What are the different forces that hold molecules together?  INTERMOLECULAR FORCES (IMF) The attractions are ONLY found in covalent compounds IMF’s between molecules are not nearly as strong as the intramolecular attraction that hold compounds together These are considered weak forces in comparison to bonds They are, however, strong enough to control physical properties such as boiling and melting points and vapor pressures There are four types of IMF’s

1. London Dispersion Forces The weakest of all of the IMF’s Only important for the NONPOLAR molecules, such as the diatomics and noble gases. This repulsion creates brief dipoles in atoms

2. Dipole-Dipole Stronger version of LDF Dipole: are interactions between molecules that have a positive end and a negative end - this is a result from the unequal sharing of electrons (think dipole, polar) These forces are important when the molecules are close to each other The more polar the molecule, the higher is its boiling point

Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point. © 2009, Prentice-Hall, Inc.

3. Hydrogen Bonds Not a BOND, but a force of attraction The strongest type of IMF Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen and fluorine “FON” These attractions are responsible for the high boiling point of water

4. Molecule- Ion Attraction How ionic compounds dissolve in water Example: NaCl in H2O

Bond Type Conductivity BP/MP Examples Ionic Only in liquids and aqueous states (l) (aq) HIGH! Strong bonds and strong IMF CuCl2 NaCl Covalent (Molecular) None! Low! Weak Bonds and low IMF’s CO2 H2O Metallic Solid and Liquid Phase Only! (s) (l) Strong Bonds and Strong IMF’s Al K Ca *Network Solids NONE! Strong bonds and IMF’s SiO2 and Diamond