2.  no simple, and yet complete, way to define this  forces that hold groups of atoms together and make them function as a unit  a bond will form if the energy of the bonded atoms is lower than that of the separated atoms 2

3. CCl 4 - Covalent C Cl HCl - Covalent H Cl MgF 2 - Ionic [ F ] 2 – [Mg] 2+ H 2 O - Covalent H O H NH 3 - Covalent H N H H NaCl - Ionic [ Cl ] – [Na] + OH – - Covalent O H H 2 - Covalent HH

4. HCl - Covalent H Cl H CO 2 - Covalent COO Na 2 O - Ionic [ O ] 2– [Na] 2 + H N H H H N H H OO OO O 2 - Covalent OO C II II I 2 - Covalent [ O ] 3 2– [Al] 2 3+ Al 2 O 3 - Ionic NH 3 - Covalent OO O OOO O 3 - Covalent H C H H H H C H H H

5.  bond strength from smallest to largest ◦ single ⇒ double ⇒ triple  bond length from smallest to largest ◦ triple ⇒ d o u b l e ⇒ s i n g l e

6.  VSEPR Theory YouTube (4:52) VSEPR Theory YouTube (4:52)  the structure around a given atom is determined mostly by minimizing electron pair repulsions  electron pairs (IB calls negative charge centres) are…  lone pairs – pairs of electrons around a central atom not in a bond with another atom  bonding pairs – pairs of electrons being shared found in the space between the atoms (can be single, double, or triple bonds)

8.  the repulsion between electron pairs (negative charge centres) causes molecular shapes to adjust  electron pairs arrange themselves around the central atom so that they are as far apart from each other as possible ◦ explains the three dimensional shape of a molecule

9. SL level HL level

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13.  2 bonding pairs  0 lone pairs linear Shapes for species with 2, 3, and 4 negative charge centres on the central atom

14.  3 bonding pairs  0 lone pairs  2 bonding pairs  1 lone pairs trigonal planar bent

15.  4 bonding pairs  0 lone pairs  3 bonding pairs  1 lone pair  2 bonding pairs  2 lone pairs tetrhedral trigonal pyramid bent

19.  shared bonding electrons pairs are pulled (as in a “tug-of-war”) between atoms ◦ results in an equal OR unequal sharing

21.  atoms in the bond pull shared pair of electrons equally  always the case in diatomic molecules ◦ H 2 O 2 N 2 Cl 2 ….

22.  atoms in the bond pull the shared pair of electrons unequally  results in a dipole because it has two poles  Ex: HCl, H 2 O

23. BF 3 non-polar molecule

24. more electronegative atoms have a greater attraction for electrons a number is assigned to each element to quantify how much is “wants/likes” electrons atoms with the higher electronegativity give that “side” of the molecule a slightly negative charge (δ -) less electronegative atoms therefore have a slightly positive charge (δ +)

29. covalent, non-polar covalent, polar ionic

30. 0.1 – 1.0 1.1 – 1.8 >1.8 0.0covalent, nonpolar covalent, slightly polar covalent, very polar ionic electronegativty difference probable type of bond

33. Arrange the following bonds from most to least polar: a) N–FO–FC–F a) C–F, N–F, O–F b)C–FN–OSi–F b) Si–F, C–F, N–O c)Cl–Cl, B–Cl, S–Cl c) B–Cl, S–Cl, Cl–Cl

34. H2OH2O  CO 2  CH 3 Cl  CCl 4

35. non-polar polar

36. Which of the following molecules have a dipole moment? H 2 O, CO 2, SO 2, and CH 4 O H H dipole moment polar molecule S O O CO O no dipole moment nonpolar molecule dipole moment polar molecule C H H HH no dipole moment nonpolar molecule