Understanding Bonding and Geometry in Chemistry: Metals, Nonmetals, and Bond Types

1. BONDING AND GEOMETRY Unit 10 Chemistry Langley **Corresponds to Chapter 7 and 8 (pages 186-247) in the Prentice Hall Chemistry textbook

2. PERIODIC TABLE REVIEW • Location of Metals and Nonmetals on the periodic table: • Metals are to the left of the “staircase” • Nonmetals are to the right of the “staircase” • For bonding, the 7 metalloids will treated as metals • All though hydrogen is to the left of the “staircase”, it is not, nor has it ever been a metal. IT IS A NONMETAL!

3. ATOMS AND IONS REVIEW • Atoms are neutral • They have the same number of protons and electrons • Number of positives = number of negatives • Example: Na 11 protons, 11 electrons  11 – 11 = 0 • Ions have a charge • They have a different number of protons and electrons • Example: Na+111 protons, 10 electrons 11 – 10 = +1 • If an atom GAINS an electron  becomes negatively charged  ANION • If an atom LOSES an electron  becomes positively charged  CATION

4. TYPES OF BONDS • Bonding occurs because every element is either trying to get to 0 electrons in the valence or 8 electrons in the valence (zero and 8 are both stable) • Valence is the outer electron shell—place where bonding occurs • Ionic – Bonding between a metal and a nonmetal • Metallic – Bonding between two metals • Covalent – Bonding between two nonmetals

5. IONIC BONDING • Very stable and strong • Strongest possible bond • Requires a large amount of energy to break an ionic bond • Forms compounds known as “ionic compounds” • All ionic compounds will dissolve in water and carry a current (electrolyte) • Generally have high melting and boiling points • Compounds are generally hard and brittle

6. IONIC BONDING • Draw the dot diagram for Na AND Cl Na has 1 valence electron, wants to give that 1 away and get to zero and be stable Cl has 7 valence electrons, wants to get 1 electron so it can get to eight and be stable Na give an electron to Cl and Cl takes that electron from Na

7. METALLIC BONDING • Metal atoms are pieces of metal that consist of closely packed cations (positively ions) • Cations are surrounded by mobile valence electrons that are free to drift from one part of the metal to another • Metal atoms are arranged in very compact and orderly (crystalline) patterns • Metallic bonding is the electrostatic attraction between conduction electrons, and the metallic ions within the metals, because it involves the sharing of free electrons among a lattice of positively-charged metal ions • Occurs between 2 or more metals • Result of the attraction of free floating valence electrons for the positive ion • These bonds hold metals together

8. METALLIC BONDING • Properties of metallic bonds • Good conductors of electricity • Electrons are free flowing • Malleablehammered into sheets • Ductiledrawn into wires • Alloy-two metals are bonded together to get the benefits of each • 14 karat gold

9. COVALENT BONDING • Covalent: • Prefix “co” means share, together • “valent” means valence • Covalent bonds are when atoms SHARE VALENCE electrons • A covalent compound is called a molecule • Covalent bond ALWAYS occurs between 2 nonmetals

10. TYPES OF COVALENT BONDS • Single Bond • Covalent bond where one pair of electrons (2 electrons total) are shared between 2 atoms • Atoms share electrons so that each has a full octet (8 valence) • Electrons that are shared count as valence electrons for both atoms • Examples • HCl • Cl2

11. COVALENT BONDING • Double Bonds • Bond in which two pairs of electrons (4 electrons total) are shared between 2 atoms • Examples • O2 • C2F2 • Triple Bonds • Bond in which 3 pairs of electrons (6 total electrons) are shared between atoms • Examples • N2 • AsP

12. COVALENT BONDING • Covalent Bonds with more than 2 atoms • Examples • CH4 • OF2 • Electron Pairs • Electron pairs involved in the actual bond are called BONDING PAIR or SHARED PAIR electrons • Electrons not involved in the actual bond, those surrounding the rest of each element are called LONE PAIR electrons

13. POLAR BONDS AND MOLECULES • Covalent bonds are formed by sharing electrons between two atoms • The bonding pair of electrons is shared between both elements, but each atom is tugging on the bonding pair • When atoms in a molecule are the same (diatomic) the bonding pair is shared equallythis bond is called non polar covalent • When atoms in a molecule are different, the bonding pair of electrons are not shared equallythis is called a polar covalent bond

14. POLAR BONDS AND MOLECULES • Why is the bonding pair not shared equally? • The answer lies within electronegativity • One of the elements is more electronegative than the other and therefore has a greater desire for the shared pair • The MORE electronegative element tends to pull the electrons closer and thus has a slightly negative charge • The LESS electronegative element has a slightly positive charge since the shared pair is being pulled away

15. POLAR BONDS AND MOLECULES • Drawing/Indicating Polarity

16. POLAR BONDS AND MOLECULES • Polar Molecules • Molecule in which one end of the molecule is slightly negative and the other end is slightly positive • Just because a molecule contains a polar bond DOES NOT mean the entire molecule is polar • The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds

17. POLAR BONDS AND MOLECULES • Example: CO2O = C = O • Carbon and Oxygen lie along the same axis. • Bond polarities are going to cancel out because they are in opposite directions • Carbon dioxide is a nonpolar molecule even though there are two polar bonds present • Would cancel out if the polarities moved towards each other as well • When polarities cancel out, the molecule is non-polar

18. POLAR BONDS AND MOLECULES • Example: H2O • Example: CH3Br

19. FORCES IN A MOLECULE • Dipole-Dipole Forces • Dipoles are created when equal but opposite charges are separated by a short distance • Have to have a positive and a negative end so that one of the elements is pulling on the electron • Only happens in polar molecules • Dipole forces are extremely strong and lead to high melting and boiling points

20. FORCES IN A MOLECULE • Hydrogen Bonding • Very strong type of dipole force • Only occurs when hydrogen is covalenty bonded to a highly electronegative atom • Always involves hydrogen • Example: HF, HCl

21. FORCES IN A MOLECULE • London Dispersion Forces • Electrons are in constant motion around a nucleus • At any given time there might be more electrons on one side of an atom than on the other • For a split second, the side with more electrons is negative, and the side with less electrons is positive

22. FORCES IN A MOLECULE • London Dispersion Forces • Recall that Noble Gases have a full outer shell and you have been told they are unreactive BUT due to London Dispersion Forces, they COULD bond for an instant • Example: Ar2 • London Forces are very weak • The smaller the mass of the atom, the smaller the London Force

23. BOND DETAILS • Terminology • Bond strength-energy required to break a bond • Bond axis-imaginary line joining two bonded atoms (example: C-C) • Bond length-the distance between two bonded atoms at their minimum potential enery; the average distance between two bonded atoms • Bond energy-energy required to break a chemical bond and form neutral isolated atoms • Chemical compound tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

24. BOND DETAILS • Comparison of Bond Length/Strength for Covalent Bond Types: • Longer bond = less bond strength • Rating 1-3 (with 3 as the largest and 1 as the smallest)

25. BOND DETAILS • Coordinate Covalent Bonds • Very rare • Tend to form harmful molecules • Occurs when both of the bonding pair of electrons in a covalent bond come from only ONE of the atoms • Example: CO

26. BOND DETAILS • Resonance • Occurs when there are more than one possible structures for a molecule • Refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure • Example: CO2 • To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures • Even though all of the structures are different, the number of bonding pair of electrons and lone pair of electrons stay the same in each structure

27. VSEPR THEORY • Valence Shell Electron Pair Repulsion Theory • Allows us to picture molecules in 3 dimensions • Centers around the fact that electrons have negative charges and repel one another • So electron pairs within a structure try to arrange themselves to be as far away from other pairs as possible

28. VSEPR THEORY • Tetrahedral • Central atom bonds to 4 atoms and has zero lone pairs • CH4

29. VSEPR THEORY • Pyramidal • The central atom bonds to 3 atoms and has 1 lone pair of electons • NH3

30. VSEPR THEORY • Trigonal Planar • The central atom bonds to 3 atoms and has zero lone pairs • CO3-2

31. VSEPR THEORY • Bent Triatomic • The central atom bonds to 2 atoms and has 2 lone pair of electrons • H2O

32. VSEPR THEORY • Linear Triatomic • The central atom bonds to 2 atoms and has zero lone pair of electrons • CO2

33. VSEPR THEORY • Linear • One bond between 2 atoms • HCl • N2