Unit 3: Chemical Bonding and Nomenclature Part 2

Electronegativity The ability of an atom in a molecule to attract electrons to itself

Electronegativity Since atoms have varying electronegativities they are pulling electrons towards them with different amounts of force Polar bond occurs when one atom is pulling electrons towards it with more force than the other atom involved in the bond

Electronegativity This value can be read from the periodic table We use the electronegativity (ΔEN) difference to determine if a bond is polar or not Find the EN values in the table then subtract the smallest from the largest number Compare this value to the ranges for different types of bonds

Polar and Non-polar Bonds

Polar and Non-polar Bonds Electronegativity Difference (ΔEN) Type of Bond >2.0 Ionic 0.4-2.0 Polar Covalent <0.4 Purely (non-polar) Covalent

Polar and Non-polar Bonds Determine whether the following bonds are polar, non-polar, or ionic K-Cl H-P B-F

Partial Charges In polar covalent bonds, one atom is pulling electrons closer to them Higher electronegativity This causes the electrons to spend more time closer to that atom The atom with higher electronegativity has a partially negative charge The atom with the lower electronegativity has a partially positive charge

Partial Charges Partial charges are represented by: δ+ and δ- Polar bonds are said to have a dipole Shown by an arrow with a positive (+) end and the arrow head pointing to the more electronegative atom

Partial Charges

Partial Charges

Polarity of Molecules Determined by: Presence of lone pairs of electrons on the central atom in a molecule Presence of different elements surrounding the central atom Binary molecules involving two different elements

VSEPR Theory Valence Shell Electron Pair Repulsion Based on idea of like charges repelling each other Electron domains (bonds or lone pairs) are negatively charged and therefore repel one another

VSEPR Theory We can predict the 3D shape a molecule will take on by looking at how many electron domains the central atom has Remember the atom that needs the most electrons is placed in the middle of our Lewis dot structures

VSEPR Theory There are 6 shapes you are responsible for Found on page 2 of your Data Sheets booklet First we must figure out how many electron domains surround the central atom How many bonding pairs? How many lone pairs? From here we can figure out the shape the molecule will take on

VSPER Shapes Linear 2 electron domains 2 bonding pairs 0 lone pairs

VSEPR Shapes Trigonal planar 3 electron domains 3 bonding pairs 0 lone pairs

VSPER Shapes Bent or V-shaped 3 electron domains 2 bonding pairs 1 lone pair

VSPER Shapes Tetrahedral 4 electron domains 4 bonding pairs 0 lone pairs

VSEPR Shapes Trigonal pyramidal 4 electron domains 3 bonding pairs 1 lone pair

VSEPR Shapes Bent or V-shaped 4 electron domains 2 bonding pairs 2 lone pairs

Molecular Polarity We know how to find out if a bond is polar For molecules we must look at all bonds and then the molecule as a whole To find out if something is polar we must figure out if all sides are the same or different

Molecular Polarity If all sides of the shape are the same partial charge (all negative or all positive) then the molecule is non-polar If two or more sides have different partial charges (some positive some negative) then the molecule is polar

Molecular Polarity Molecules that contain polar bonds can be non-polar as a whole Examples: CO2 CH4 CH3Cl NH3

Inter and Intra Molecular Forces Inter = between Intermolecular forces are forces between molecules Influence boiling and melting points Intra = within Intramolecular forces are forces within a molecule Ionic, covalent, metallic

Intermolecular Forces London dispersion forces Weakest Dipole-dipole forces Hydrogen bonding Strongest

Intermolecular Forces Found in molecular (NOT ionic) compounds Allows us to organize compounds based on melting and boiling points Tells how tightly molecules of a substance are held together How much energy needs to be added to change states

Intermolecular Forces If a molecule has a strong intermolecular force (hydrogen bonding) it will also have the weaker forces (dipole-dipole forces and London dispersion forces) Substances with stronger forces will boil and melt at higher temperatures than those with weaker forces

London Dispersion Forces There is always at least a slight electrostatic (+ and -) attraction between molecules Explains why non-polar gases can be liquified Remember electrons are floating around a positive nucleus in a cloud at all times Instantaneous distribution of electrons can vary from average distribution

London Dispersion Forces Present between all molecules Significant only when molecules are very close together Increases with increasing atomic and molecular size

London Dispersion Forces

Dipole-dipole Forces Permanent dipole is present Must have polar molecules δ+ and δ- ends of the molecules align so the opposite charges are close

Dipole-dipole Forces

Hydrogen Bonding Attraction between H atom attached to highly electronegative atom and nearby small electronegative atom in another molecule Strongest of the intermolecular forces Molecules must have H bonded to: O F N

Hydrogen Bonding

Intermolecular Forces What is the strongest intermolecular force present in the following: NH3 HCl CCl4

Intermolecular Forces As the strength of intermolecular forces increase so do boiling and melting points If the same intermolecular forces are present in two molecules the force is increased as the size of the molecule increases C2H6 has a higher melting/boiling point than CH4

Hydrogen Bonding - Water Unusual because solid form is less dense than liquid form In liquid form, H-bonds between molecules When frozen there are more H-bonds made and a lattice is formed Mass does not change When frozen volume increases (takes up for space) which decreases density

Hydrogen Bonding - Water

Hydrogen Bonding-Water Since H-bonding is such a strong intermolecular force water has a high melting and boiling point Molecules are held close together Requires a lot of energy to make molecules more mobile

Acids and Bases Arrhenius concept was useful but limited Bonsted-Lowry Acids were developed Acid-base reactions involved the transfer of H+ ions from one substance to another

Bronsted-Lowry Acids and Bases Acids donate the proton (H+) Bases accept the proton (H+) When identifying compounds as acids for naming in this course we will look for H at the beginning of the formula

Acids and Bases Acid Base Low pH High pH Sour taste Bitter taste Turns litmus paper red Turns litmus paper blue Increases concentration of H+ ions in solution (Arrhenius) Increases concentration of OH- ions in solution (Arrhenius) Proton (H+) donors (Bronsted-Lowry) Proton (H+) acceptors (Bronsted-Lowry)

Acids Formula to Name Look for the H at the beginning of the formula It’s an acid! Is there two elements or more than two elements in the formula? Is it binary or polyatomic?

Binary Acids Formula to Name Binary acids have H and one other element in their formula Always “hydro______ic acid” HF Hydrofluoric acid HBr Hydrobromic acid

Polyatomic Acids Formula to Name Polyatomic acids contain polyatomic ions attached to H Remember: Ate  ic Ite ous Then add acid to the end We also don’t need to show H is there, drop the hydro!

Polyatomic Acids Formula to Name H3PO4 Phosphoric acid H2SO3 Sulfurous acid CH3COOH (also HCH3COO, CH4CO2) Acetic acid

Acids Name to Formula “Acid” must start with an H Does the name begin with hydro? Is it binary? Work backwards

Acids Name to Formula Polyatomics: Binary: Balance your charges!! Ic  ate Ous  ite Binary: Look for the root element Balance your charges!!

Polyatomic Acids Formula to Name Hydroiodic acid HI Chloric acid HClO3 Hydrochloric acid HCl Dichromic acid H2Cr2O7