Chemical Bonding Ch. 6

What is a chemical bond? mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together Why don’t noble gases do this? Already have filled s and p orbitals stable octet: 8 valence e- (or 2 if, you’re helium) Atoms that don’t have a stable octet are more reactive

Key Point #1: By forming bonds with each other, most atoms reduce their potential energy, becoming more stable. This is a chemical change! All chemical changes involve energy!

What types of bonds can be formed?

Metallic Bonding In a metal, the empty orbitals in the atoms’ outer energy levels overlap Delocalized Electron: outer electron that does not belong to any one atom but can move freely through the metal’s network of empty atomic orbitals. sea of electrons: mobile electrons around the metal atoms, which are packed together in a crystal lattice.

metallic bonding: chemical bonding that results from the attraction between metal atoms and surrounding sea of electrons Key Point: In metallic bonding, valence electrons move freely throughout a network of metal atoms.

Metallic Bonding

Unique Characteristics of Metals Metals have many unique properties because of their sea of electrons Malleability: ability of a substance to be hammered or beaten into thin sheets Ductility: ability of a substance to be pulled into a thin wire Why? atoms can slide past one another along a plane without breaking bonds

Luster: shiny appearance Why? Absorb a wide range of light frequencies, many orbitals separated by small energy differences Conductivity Thermal: ability to conduct heat Electrical: ability to conduct electricity Why? Electrons move easily through network of empty orbitals

Metallic Bonding Strength The strength of metallic bonding is determined by the enthalpy of vaporization: the amount of energy required to vaporize (turn into a gas) 1 mol of a metal In general, the strength of the metallic bond INCREASES moving left to right across the periodic table. Soft metals (less dense) metals  harder (more dense) metals toward right

Properties of Metals: Malleability and Ductility

Properties of Metals: Surface Appearance

Properties of Metals: Electrical and Thermal Conductivity

Types of Bonds What type of bonds can be formed? Ionic bond Covalent bond Nonpolar covalent Polar covalent Ionic bonding: bonds that result from electrical attractions between cations and anions 1 atom losses electrons 1 atom gains electrons Covalent bonding: sharing of electrons between 2 or more atoms

Key Point 2: Rarely is bonding between atoms purely ionic or purely covalent. Instead, it usually falls somewhere between the two extremes. Why? Key Point 3: The extent of ionic or covalent bonding between two atoms can be estimated by calculating the difference in each elements’ electronegativity.

Covalent Bonding Large difference in E.N.: bond has more ionic character Small difference in E.N: bond has more covalent character Types of Covalent Bonds Non-polar covalent bonding: both electrons equally shared between atoms Polar covalent bonding: unequal attraction for the shared electrons

6.1 Practice Worksheet Part 1 The property of electronegativity, which is the measure of an atom’s ability to attract electrons, can be used to predict the degree to which the bonding between atoms of two elements is ionic or covalent. The greater the electronegativity difference, the more ionic the bonding is.

If the calculated electronegative difference is… > 1.7 : ionic bond is formed > 0.3 , < 1.7 : polar-covalent bond 0 – 0.3 : non-polar covalent bond Increasing difference in electronegativity Nonpolar Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e-

Electronegativity Difference Elements Electronegativity Electronegativity Difference Bond Type Element 1 Element 2 Mg to Cl H to O C to Cl N to H C to S K to F Na to Cl H to H 3.0 1.8 Ionic 1.2 2.1 3.5 1.4 Polar covalent 2.5 3.0 .5 Polar covalent 3.0 2.1 .9 Polar covalent 2.5 2.5 Non-polar covalent .8 4.0 3.2 Ionic .9 3.0 2.1 Ionic 2.1 2.1 Non-polar covalent

Polyatomic Ions It is also possible if a compound contains polyatomic ions, for both types of bonding to be present. Monatomic Ions: Fe2+ , Na+, Cl- Polyatomic Ions: PO43-, NH4+ , NO-1 Groups of atoms are bonded covalent together, but because of few or more than expected valence electrons they have an overall charge (so they can also bond ionically with other ions) Ex: Ca2+ and SO42-  CaSO4 (metal & diff. nonmetals)

Classify the following as ionic, covalent, or both 1. CaCl2 = __________ 5. BaSO4 = ___________ (metal & nonmetal) (metal & diff. nonmetals) 2. CO2 = __________ 6. H2O = ____________ (nonmetal & nonmetal) 3. MgO = __________ 7. SO3 = ___________ 4. HCl = ___________ 8. AlPO4 = ___________ Covalent Covalent Ionic Covalent Covalent Both

Section 6.2 Covalent Bonding

What is a molecule? Neutral group of atoms that are held together by covalent bonds. Chemical formula: indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.

Formation of Covalent Bonds The electrons of one atom and protons of the other atom attract each another. The two nuclei and two electrons repel each other. These two forces cancel out to form a covalent bond at a length where the potential energy is at a minimum.

Bond Length vs. Bond Energy Bond length (pm): distance between two bonded atoms at their minimum potential energy Bond energy (kJ/mol): energy required to break a chemical bond and form neutral isolated atoms. Breaking bonds: absorbs (requires) energy Forming bonds: releases energy Key Point: As you increase the number of bonds between 2 atoms the bond energy increases, while the bond length decreases. This is an inverse relationship.

Bond Energies & Bond Lengths A. How many electrons are shared in a single bond: double bond: triple bond: B. Which bond is shorter? C – C or C = C C. Which bond requires more energy to break? In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons

Bond Energies & Bond Lengths A. How many electrons are shared in a single bond: 2 e- double bond: 4 e- triple bond: 6e- B. Which bond is shorter? C – C or C = C C. Which bond requires more energy to break? = In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons

Octet Rule Hydrogen: 2e- Boron: 6e- Octet Rule: chemical compounds tend to form so that each atom has an octet of e-’s in its highest occupied energy level Exceptions to the octet rule: Atoms that cannot fit eight electrons Atoms that can fit more than eight electrons Hydrogen: 2e- Boron: 6e- Phosphorus, Sulfur, & Xenon: expanded valence, more than 8e-

How can we represent molecules? Lewis Structures: formulas in which atomic symbols represent nuclei and inner shells, which are surrounded by dot-pairs/dashes represent valence electrons

Chapter 6.5 Molecular Geometry

VSEPR THEORY Lewis Structures are 2D but we live in a 3D world! molecular geometry: the three-dimensional arrangement of a molecule’s atoms What do those 3D structure/shapes look like?? Follow the Valance Shell Electron Pair Repulsion Theory or VSEPR Repulsion between the sets of valence electrons surrounding an atom causes them to be oriented as far away from each other as possible

Why use VSEPR Theory? Key Point: VSERP Theory is used to predict the shape of molecules based on the fact that electron pairs strongly repel each other. Following VSEPR allows us to predict bond polarity: uneven distribution of electrons

AB2 – Linear Central atom Atoms/group of atoms attached to central atom Atoms bonded to central atom (B) Number of Lone Pairs on central atom (E) Bond Angle 2 180˚

Other Linear Geometries The shape of two atoms bonded together is not given in the chart. Ex: F2 What is the only possible shape a binary compound can have? LINEAR!

Atoms bonded to central atom (B) Number of Lone Pairs (E) AB2E1 – Bent Atoms bonded to central atom (B) Number of Lone Pairs (E) Bond Angle 2 1 <120˚ What happens to the bond angle between atoms as you increase the number of “lone pair electrons” on the central atom? Bond angles decrease!

Atoms bonded to central atom (B) Number of Lone Pairs (E) AB2E2 – Bent Atoms bonded to central atom (B) Number of Lone Pairs (E) Bond Angle 2 104.5˚

Atoms bonded to central atom (B) Number of Lone Pairs (E) AB3 – Trigonal Planar Atoms bonded to central atom (B) Number of Lone Pairs (E) Bond Angle 3 120˚ Shape is often associated with atoms that break octet rule, but doesn’t have to be

AB3E1 – Trigonal Pyramidal Atoms bonded to central atom (B) Number of Lone Pairs (E) Bond Angle 3 1 107˚

Atoms bonded to central atom (B) Number of Lone Pairs (E) AB4 – Tetrahedral Atoms bonded to central atom (B) Number of Lone Pairs (E) Bond Angle 4 109.5˚

Predicting Molecular Geometry Draw Lewis structure for molecule. VSEPR theory: if any lone pairs of electrons are found on the central atom, these electrons decrease the bond angles of atoms attached to it. 2. Draw a revised Lewis structure to show more accurate geometry S F S O AB4E AB2E tetrahedral bent

Predicting Molecular Polarity 3. To indicate the polarity of the bonds, we use this symbol: __________________ , which always points toward the more electronegative element. 4. When multiple bonds are found in a molecule, we must identify polarity of each bond. electron rich region electron poor region H F e- poor e- rich F H d+ d-

5. Observe the overall polarity of the molecule 5. Observe the overall polarity of the molecule. Think of it as “tug-of-war” for valence electrons between the various atoms. Non-polar covalent molecules: If the atom is symmetrical and all atoms have an equal pull on electrons Polar covalent molecules: If the atom is not symmetrical and/or the atoms do not all have an equal chance of winning the tug of war for electrons

Intermolecular Forces

Intermolecular forces: attractive forces between molecules. Intramolecular forces: attractive forces within a molecule (the bonds) Intermolecular Forces Intramolecular Forces (bond) Intramolecular Forces intermolecular forces are much weaker than intramolecular forces

Strength of IMF Hydrogen Bond Dipole – Dipole Induced Dipole London Dispersion Forces strongest weakest

d+ d- Dipoles What is a dipole? A polar molecule Uneven sharing of electrons so there is a separation of charge H F electron poor region electron rich region e- poor F H d+ d-

Dipole-Dipole Forces Attraction between two polar molecules — — + +

Hydrogen Bonding Special type of Dipole – Dipole Attraction between: Hydrogen & Nitrogen/Oxygen/Fluorine

Electrons shift toward positive end of dipole Induced Dipole Attraction between one polar and one nonpolar molecule — + Electrons shift toward positive end of dipole — — + +

London Dispersion Forces Attraction between two nonpolar molecules Electrons become uneven and form a dipole — — + +

What does IMF effect? Viscosity Surface Tension Boiling Point

Stronger IMF  Higher Boiling Point Point at which liquid particles escape the surface of the liquid into the gas phase Stronger IMF  Higher Boiling Point

Stronger IMF  Higher Surface Tension result of an imbalance of forces at the surface of a liquid. Stronger IMF  Higher Surface Tension

Stronger IMF  Higher Viscosity Measures a fluid’s resistance to flow Stronger IMF  Higher Viscosity