Name: Class: Date: Multiple Choice Identify the letter of the choice that best completes the statement or answers the question.

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1 Name: Class: Date: SCH4U Chapter 4 Formative Test Multiple Choice Identify the letter of the choice that best completes the statement or answers the question. 1. Which of the following statements about ionic bonding is INCORRECT. A) The force of attraction between oppositely charged ions (cations and anions) makes up the ionic bond. B) Ionic bonding occurs between atoms of elements that have large differences in electronegativity. C) The units of ionic compounds can be separated easily by direct heating of the crystal salt. D) The ions that make up the ionic solid are arranged in a specific array of repeating units. E) The ions that make up the ionic solid are arranged in a rigid lattice structure so that the cations and anions are arranged so that the system has the minimum possible energy. 2. Which one of the following statements about covalent bonding is CORRECT. A) Covalent bonding involves an imbalance between the forces of attraction and repulsion that act between the nuclei and electrons of two or more atoms. B) There is an optimum separation between the individual atoms involved in a covalent bond at which the nucleus-electron attractions, nucleus-nucleus repulsions, and electron-electron repulsions achieve a balance. C) Lithium Sulfide is an example of a molecule with covalent bonds. D) A covalent bond involves the formation a new orbital, caused by the overlapping of atomic orbitals. The new orbital has energy levels that are higher than those of the original atomic orbitals. E) In many cases, electron sharing enables each atom in a covalent bond to acquire full d orbital. 3. Covalent (molecular) compounds typically have the following properties: A) exist as a soft solid, a liquid, or a gas at room temperature B) have low melting points and boiling points C) are poor conductors of electricity, even in solution D) may not be soluble in water E) All of the above 4. In the liquid state, polar molecules (dipoles) orient themselves so that oppositely charged end of the molecules are near to the one another. The attraction between these opposite charges are called A) dipole-dipole forces D) dipole induced forces B) ion dipole forces E) intramolecular forces C) ion induced dipole forces 5. Non-polar gases such as oxygen and nitrogen dissolve sparingly, in water because of these forces. A) dipole - dipole forces B) ion dipole forces C) ion induced dipole forces D) dipole induced dipole forces E) london dispersion forces 6. Which of the following is NOT an intermolecular bond. A) ionic C) hydrogen D) dispersion E) dipole-dipole B) ion-dipole 7. Which of the following statements is CORRECT for two bonded atoms that have very different electronegativities. A) The atom with the lower electronegativity attracts the shared electrons more strongly than the atom with the higher electronegativity. B) The atom with the higher electronegativity attracts the shared electrons more strongly than the atom with the lower electronegativity. C) If the EN is greater than 0.4 and less than 1.7 the bond is considered ionic. D) both A and C are correct E) both B and C are correct 1

2 Name: 8. The factor(s) that effect the magnitude of dispersion forces. A) number of electrons D) Both A and B B) the shape of the molecule E) Both A and C C) the electron group arrangement 9. The factor(s) that effect the magnitude of dispersion forces. A) number of electrons D) Both A and B B) the shape of the molecule E) Both A and C C) the electron group arrangement 10. A molecule with the VSEPR notation of AX 2 E. A) will never be polar B) will always be polar C) will only be polar if the bond dipoles are in the direction of the central atom D) will only be polar is the bond dipoles are in the direction of the peripheral atoms E) will only be polar if the bond dipoles are in opposite directions 11. The reason the atoms form chemical bonds is that A) bonded atoms tend to have lower energy than single, uncombined atom B) bonded atoms tend to have higher energy than single, uncombined atom C) bonded atoms need to obtain an octect valence shell of electrons D) bonded atoms need to lose or gain electrons E) bonded atoms tend to have stable ions 12. The free-electron model explains the following properties of metals. A) Conductivity B) Malleability and Ductility C) Melting and Boiling Points D) Only A and B E) All of the above 13. The energy required to break the force of attraction between two atoms in a bond and to separate them: A) Bond energy B) AFE (Attractive Force Energy) C) BFE (Break Force Energy) D) Separation Energy E) None of the above 14. The number of lone pairs on this molecule would be: A) 8 B) 4 C) 0 D) 2 E) Can tell from this diagram 15. Which of the following represents ionic bonding? A) correct D) B) E) C) 2

3 Name: Short Answer Answer in the space provided. If you need further space please use full scape provided. Grading will be based upon the quality, not quantity, of your answer and how clearly you present the information within your answers. You are expected to incorporate key terms, definitions and concepts you have learned to fully develop your answer. For math questions proper format and a concluding statement are required. Each answer will be out of 5 marks. 16. Use VSEPR theory to draw the Lewis structure for HCN and SiF 6 2- Show the math, don t forget to consider resonance structures and fill in the table below. Molecule Num of electron groups Geometric Arrangement of electron groups Types of electron pairs VSEPR notation Name of Molecular Shape HCN linear 2BP AX 2 linear 2 SiF octahedral 6BP AX 6 octaheddral 3

4 Name: 17. Draw orbital diagrams and Lewis structures to show how the following pairs of elements can combine. In each case, write the chemical formula for the product. Na and N 18. The graph below shows the effect of one of the intramolecular forces that we studied this chapter. a) Add the appropriate labels to the graph. b) What is the intermolecular force? c) Explain the trend in group 14. d) Why is there no hydrogen bonding in H2S and H2Se, which have unshared electron pairs on their central atoms? a) x = period number no units y = boiliing point (degrees celcius) b) Hydrogen Bonding c) The increase in boiling point from methane, CH4, to tin(iv) hydride, SnH4, is due to an increase in the number of electrons and size of the molecules. With an increase in size and number of electrons there are larger dispersion forces, since more electrons are able to temporarily shift from one part of the molecule to another. d) Neither S nor Se is small enough or electronegative enough to support hydrogen bonding between the lone pair of electrons and an H on the other S or Se. 4

5 Name: 19. Discuss the validity of the statement using specific example molecules: All polar molecules must have polar bonds and all non-polar molecules must have non-polar bonds. All polar molecules must have polar bonds, however not all non-polar molecules have non-polar bonds. (i.e. some non polar molecules can and do have polar bonds). If the bonds in the molecule are polar but the molecule is symetrical (for example in a trigonal planner or tetrahedral molecule) the dipoles of individual bonds cancel each other to give a net zero dipole the molecule will be non-polar. Ex CCl4 is non polar molecule with four polar bonds. 20. Explain how hydrogen bonding explains waters density change from a liquid to a solid. Hydrogen bonding also explains water s unusual property of being less dense in the solid state than in the liquid state. As you know from kinetic molecular theory, lower temperatures cause liquid molecules to move slower than in their gaseous states. When liquids solidify, the kinetic energy of the particles is no longer enough to prevent the intermolecular forces of attraction from causing particles to stick together causing their volumes to decrease for a given mass and therefore cause an increase in density. For most substances, their solid states are denser than their liquid states. When water freezes, however, the water molecules pack in such a way that the solid state is less dense than the liquid state. Water molecules align in a specific pattern so that hydrogen atoms of one molecule are oriented toward the oxygen atom of another molecule causing the ice to expand (i.e. volume increases for a given mass and therefore density decreases) 21. Explain the intermolecular and intramolecular forces in HBr (l). The electronegativity difference of 0.76 between H (2.0) and Br (2.96) suggests that the H Br bond is polar covalent. The intramolecular force is the simultaneous attraction on a pair of electrons by the hydrogen and bromine nuclei. Since bromine has a greater electronegativity, the bromine atom pulls the pair of electrons closer to itself, giving the molecule a permanent dipole. In HBr(l), the positive end of the molecule lines up with the negative end of another molecule, thus setting up an electrostatic attraction (dipole-dipole force) between the two molecules. Therefore, the intermolecular forces in HBr(l) are dipoledipole and dispersion forces. 5

6 Name: 22. Explain the following two diagrams. (i.e. what do they represent? What does it tell the person that is looking at it?) For full marks all parts of the diagram will need to be labeled correctly. This is a hydrogen bond. When a hydrogen is bonded to oxygen, fluorine, or nitrogen, (A or B) the strong force of attraction exerted by any of these three very electronegative atoms (shown as the (1std-)) draws the electron density from the hydrogen atom in the polar-bonded molecule, leaving the hydrogen atom with a partial positive charge. (as shown as d+) This dipole is easily attracted to the partially negative lone pair (shown as the (2ndd-) (on a nearby electronegative atom: N, O, or F. This picture shows metallic bonding via the free-electron model. Shown is a typical metal as being composed of a densely packed core of metallic cations, within a delocalized region of shared, mobile valence electrons. The force of attraction between the positively charged cations and the pool of valence electrons that moves among them constitutes a metallic bond. 6