UNIT 7: BONDING 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? What are the different forces that hold molecules together?

Review of Prior Knowledge ENDOTHERMIC: EXOTHERMIC: POTENTIAL ENERGY: Why do atoms become ions? How do atoms become ions?  How do metals form ions? How do nonmetals form ions?

Aim # 1- Why do elements form bonds? A chemical bond is - force of attraction between the atoms of a compound

the electrons of one atom are attracted to the protons of another 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 Exothermic- energy is released, bonds form Remember: “BARF”

Bonding and Stability When bonds are formed their products are more stable The bonded elements of a compound become stable because their valence shell is full Atoms become ions to become stable Remember metals lose electrons to become positively charged Nonmetals gain electrons and become negatively charged

Types of bonds There are three types of bonds Ionic Covalent Metallic They differ in the elements involved and how their handle their valence electrons

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

Aim # 2: How can we explain and draw ionic bonds? An ionic bond is Transfer of electrons Attraction between oppositely charged ions Bond between a metal and a nonmetal

Ionic Bond: 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 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

Drawing of aluminum oxide is MgCl2

PRACTICE

AIM# 3: How can we explain and draw covalent bonds? Occurs between nonmetals Nonmetals hold on to their valence electrons (don’t want to give them up) They still want a noble gas configuration They get it this sharing 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.

Properties of covalent compounds (molecular solids) Soft Good insulators Low MP Poor conductors of electricity in any phase May conduct electricity in solution if they are polar

Network Solids Covalent compounds that are extremely hard and have very high melting and boiling points. Cannot conduct electricity Ex. Diamonds, silicon dioxide (SiO2- quartz), and silicon carbide (SiC)

   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

Polyatomic Ions Table E Polyatomic ions are held together by covalent bonds but form ionic bonds with other ions. ***Any compound containing a polyatomic ion has both covalent and ionic bonds!!!!!!

AIM #4 : What are metallic bonds and why are they good conductors? Bonds between metals that form as a result of ions immersed in a “sea” of mobile electrons

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

Conductivity Bond Type MP and BP Hardness Solid Liquid Aqueous Metallic High Hard Yes Covalent Low Soft No Ionic

< 1.7 but not “0”, then polar covalent Aim # 5 what is the difference between bond polarity and molecule polarity BOND POLARITY (1.7 Rule) 1.7 rule is applied primarily to binary compounds. Determine the electronegativities of all atoms in the bond. Take the difference between the bonded atoms The higher the difference the more polar the bond is If the difference is: >1.7 then ionic < 1.7 but not “0”, then polar covalent =0 non polar covalent

Nonpolar covalent bond Ionic bond Covalent bonds

Molecular Polarity Molecular polarity can be determined by the shape of the molecule and the distribution of charge Molecular shapes Linear (X2 HX CO2) Bent (H2O) Pyramidal (NH3) Tetrahedral (CH4 CCl4) A polar molecule is called a dipole It has a positive side and a negative side

Nonpolar Polar H2   HCl Symmetry:

MOLECULE POLARITY: Nonpolar/polar shapes SNAP : Symmetrical Nonpolar Asymmetrical Polar

Compound Polar or Nonpolar NH3   H20

How can a molecule be both polar and non polar? CX4 tetrahedral shape. They are non polar. The individual ligands or bonds are polar Conclusion – nonpolar/polar.

Aim # 6 What are the different forces that hold molecules together? Hydrogen Bonding VERY STRONG forces of attraction between specific polar molecules Bond between an atom of H from one molecule with FON(fluorine, oxygen, nitrogen) in another molecule Sometimes called dipole-dipole forces

Dipole-Dipole Interactions Molecules that have permanent dipoles are attracted to each other. The positive end of one is attracted to the negative end of the other and vice-versa. These forces are only important when the molecules are close to each other.

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

Intermolecular Forces The stronger the intermolecular forces, the higher the boiling points and melting points Ionic Solids Molecules with Hydrogen bonds H must be bonded to N, O, or F Polar molecules Non polar molecules For non polar molecules, the greater the mass, the greater the force of attraction

Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities. © 2009, Prentice-Hall, Inc.

Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces. © 2009, Prentice-Hall, Inc.

Ion-Dipole Interactions Ion-dipole interactions are important in solutions of ions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

Dipole- Dipole Interactions Molecules that have permanent dipoles are attracted to each other. The positive end of one is attracted to the negative end of the other and vice-versa. These forces are only important when the molecules are close to each other. The more polar the molecule, the higher is its boiling point.

London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole. © 2009, Prentice-Hall, Inc.

London Dispersion Forces These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability. © 2009, Prentice-Hall, Inc.

Hydrogen Bonds   Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

Summarizing Intermolecular Forces