Chapter 2

1. Chapter 2 Polar Covalent Bonds: Acid and Bases

2. Chapter 2 - Definitions • Polar Covalent Bonds – electrons are even distributed between two atoms in a molecule. • Electronegativity- the attractiveness of an atom to an electron in a bond. • Dipole Moment – is the total net molecular polarity. • Formal Charge – assigns specific charges to individuals atoms inside molecule, particularly to atoms that have an apparently “abnormal” number of bonds. • Resonance Forms – describes the movement of electrons that accounts for the electron densities of molecules. • Bronsted-Lowery acid – is an proton donator.

3. Chapter 2 - Definitions • Bronsted-Lowery base – is a proton acceptor. • Acidity constant, Ka – determines the strength of an acid. • Lewis acid - is a substance that accepts an electron pair. • Lewis base – are a substance that donates an electron pair.

4. What is Electronegativity? • Electronegativity is the desire of an atom to gain an additional electron to fill its octet or the strength of the atom to pull electrons. • The higher number the stronger the electronegativity. • Largest electronegativity: F = 4.0, Cl = 3.5, O = 3.5, N = 3.0, Br = 2.8, C = 2.5

5. Why is Electronegativity So Important? • The electronegativity of two atoms in a bond determines what type of bond forms. • There are basically 2 types of bonds with a 3rd also considered a type of bond. • Ionic • Covalent • Polar Covalent (between a Ionic and covalent bond.)

6. Similar Electronegativities Same color denotes similar electronegativities and a covalent bond.

7. Some Examples • Na = 0.9, Cl = 3.0 (Ionic) • C = 2.5, C = 2.5 (Covalent) • C = 2.5, O = 3.5 (Polar Covalent)

8. What is the Difference Between Bonding? • Imagine that atoms play tough of war with the electrons that they share in a bond. • There would be three possibilities. • 1) One atom wins and takes the electron (ionic bond) • 2) Both atoms are even matched sharing the electrons evenly. (covalent) • 3) One atom is stronger than the other atoms and has the electron over its side more than the other atom. (polar covalent)

9. Ionic Bonds • Ionic Bonds are where an electron is donated to another atom. This creates two charged species. These charged atoms or molecules are normally free in solution but are held together when solid by electrostatic attractions.

10. Covalent Bond • Describes the sharing of electrons between two atoms. There are two different types of covalent bonds. • Nonpolar covalent bonds is defined as the even distribution of electrons between 2 atoms. • Polar covalent bonds is defined as the uneven distribution of electrons between two atoms.

11. Nonpolar Covalent Bonds • Are bonds that are formed between atoms with similar electronegativities. • Example chains of carbon (C) bonded to hydrogens (H) (Hydrocarbons)

12. Polar Covalent Bonds • Describes the unequal sharing of electrons in a covalent bond. • This makes the oxygen considered partially negative because the electrons around it more. The carbon is considered partially positive because the shared electrons are mainly around the oxygen.

13. Polar Covalent Bonds and Polar Molecules • Polar covalent bonds can be found in both individual bonds and in entire molecules. • To calculate individual bond polarity you need to use the electronegativities of the two atoms. • To calculate the dipole moment of a molecule you need to determine the center of positive and negative charges. If they are not the same then there is an overall polarity of the molecule (called a polar molecule).

14. Dipole Moment • The dipole moment of a molecule describes the region of the molecules where the electron density is highest and lowest. • The dipole moment maintains a vector from low electron density to high electron density. • If the electron density is equal across the molecules then the molecules is nonpolar molecule (evenly distributed)

15. Non-Polar Molecules • Are molecules which either maintain atoms with similar electronegativities or molecules whose dipoles are even in all directions. CH4 (nonpolar) CCl4 (Nonpolar) CH2Cl2 (Polar)

16. Non-Polar Molecules • These molecules are molecules whose dipoles are even spread in all directions or maintain similar electronegativities. Carbon Tetrachloride Methane

17. Questions Carbon Tetrachloride Methane Dichloromethane Pick the polar molecule(s). And its(their) vectors.

18. Polar Covalent Molecules • The electronegativity of the oxygen and nitrogen atoms are different there by causing a overall dipole in the molecule making the molecules polar. Water Ammonia

19. Properties of Covalent Molecules • Remember (LIKE DISSOLVES LIKE) • Polar Covalent Molecules – Water, Methanol, Ethanol. • Nonpolar Covalent Molecules – (oils, hydrocarbons) Propane, decane, etc.. • Oils and water do not mix because one is polar covalent and the other is nonpolar covalent.

20. Formal Charges • Formal charges - assigns specific charges to individuals atoms inside molecule, particularly to atoms that have an apparently “abnormal” number of bonds. • This is used when you see charge separation in a molecules to indicate if the atom is positive or negative.

21. Formal Charge • Formal charge = the charge on each atom in a molecule. Formal # of valence e- # of valence e- charge (free atom) (bound atom) Formal # of valence e- Half of # of Charge bonding e- nonbonding e-

22. Resonance • Resonance is the movement of either free electrons or p electrons to form other possible structures. • Because movement of electron occurs frequently resonance structures try to show how the electrons might move. • Only the movement of double (p bonds) or free electrons are found. Movement of the sigma bonds would breakup the molecule instead of create resonance structures.

23. Examples of Resonance

24. Rules of Resonance • 1) Individual resonance forms are imaginary, not real. • 2) Resonance forms differ only in the placement of their p or nonbonding electrons. • 3) Different resonance forms of a substance don’t have to be equivalent. • 4) Resonance forms obey normal rules of valency. • 5) The resonance hybrid is more stable than any individual resonance form.

25. Resonance Forms 1) Number 1 and 2 are resonance forms 2) Although resonance forms many not be equivalent, 1 and 2 are equal.

26. Question Why does this oxygen have a negative charge? This oxygen has a negative charge because it has taken an electron from a hydrogen giving it a negative electron and a charge of -1. This completes its octet with 8 electrons

27. Bronsted-Lowry Acid Base • Bronsted-Lowry Acid - is a substance that donates a proton (H+). It keeps the electron and becomes negative. • Bronsted-Lowry Base – is a substance that accepts a proton (H+). It gives its proton so it becomes positive.

28. Examples of Bronsted Lowry Acids and Bases 1 2 3

29. Conjugate Acid and Bases • Conjugate Acids and bases are if you look at a reaction as a reversibility, these product would be the acid base found on the opposite side. • If you look at the reaction above the acid an base are clearly defined, however if you switch the reaction then these would be the acid and base, as shown below.

30. Acid Base Strength • Acids differ in their ability to donate an (H+) proton. • Some acids break apart and donate their proton well (~100%), while others acids only gives an proton about 50 percent or less of the time. • The ability to give protons donates the strength of the acid.

31. How do you Calculate the Strength of an Acid? • The way that you normally determine the strength of an acid is to use the equilibrium constant. In dilute solutions this is rewritten as the bottom equation. In Dilute Solutions

32. What is the Ka Range of Acids? • The strongest acids has a range from 1015 and weaker acids are about 10-60. • This is a wide difference so the use of pKa is used. • P = -Log (Number) • So the pKa of an acid whose Ka = 1015 equals • -Log(1015) • -(15) • -15 is the pKa of a strong acid

33. Some Examples of pKa Weakest • CH3CH2OH 16.00 • H2O 15.74 • HCN 9.34 • H2PO4- 7.21 • CH3CO2H 4.76 • HNO3 -1.3 • HCl -7.0 Strongest

34. What is the pKb? • Just the same way that the acid strength can be determined the basic strength can also be determined. This strength is opposite of what the acidic strength would be.

35. Strength of pKb Strongest • CH3CH2O- • HO- • CN- • HPO42- • CH3CO2- • NO3- • Cl- Weakest

36. Predicting Acid Base Reactions • The predict if an acid base reaction will occur you have to determine if the stronger acid and/or stronger base are on the left side. If this is not true then the reaction will not proceed.

37. Organic Acids • Rule of thumb: Organic Acids often are found where oxygen's are. The more oxygen's the stronger the acid. • Some examples are the: alcohols and carboxylic acids Alcohol Carboxylic Acid

38. Which is the stronger acid? Alcohol Carboxylic Acid

39. Organic Bases • Rule of Thumb: Organic Bases can maintain both an oxygen or a nitrogen. Nitrogen almost always functions as a base however oxygen can function as both an acid or base. • When something has a plus charge it is a base.

40. Which Oxygen Functions as a Base?

41. Lewis Acids and Bases • Lewis Acids – is a substance that accepts an electron pair. • Lewis Bases – are a substance that donates an electron pair. • These are much broader and can often be used in both organic and inorganic chemistry.

42. Lewis Acid • Maintains an empty or vacant orbital. If you think of H+ this is a Lewis acid because it has given up its single electron to another atom and has only a proton. • Another example is Mg2+.

43. Lewis Base • Maintains a filled orbital. If you think of Nitrogen containing molecules they maintain 1 filled lone pair of electrons to share making them a Lewis base. Any atom with a filled lone pair of electrons to share. Oxygen, Nitrogen, Sulfur. • Examples are amines and sulfides a sulfur containing molecule.

44. Which of these are considered Lewis Bases? Ammonia Methane Water

45. Non-Covalent Interactions • Just like ionic atoms and molecules pair up due to charge, dipole interaction makes the intermolecular forces occur. • Intermolecular forces use the partial positive and negative charges cause by dipoles to pair. These pair partial positive to partial negative.

46. Types of Non-Covalent Interactions • Hydrogen Bonding – the strongest of these forces shows the attractive forces of a Hydrogen bond atom to an electronegative atom of O and N. • Vander Walls Forces – weaker interactions of non hydrogen bound atoms. For example Cl, Br, etc… • Dispersion Forces – forces other molecules away because the electron distribution is constantly changing in a non-uniform fashion.

47. Things to Know • Electronegativity • Bonding – Ionic, Covalent (nonpolar, polar) • Polar and Non-Polar Molecules • Resonance • Acids and Bases (Bronsted-Lowry) • Acid Strength (Ka, pKa) • Organic Acids and Bases • Know examples of Lewis Acids and Bases • Non-Covalent Interactions (#, how do they work)