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Modern Chemistry Chapter 6 Homework 6-2 Answers

C H A P T E R 6 R E V I E WChemical BondingSECTION 1SHORT ANSWERAnswer the following questions in the space provided.a1.A chemical bond between atoms results from the attraction between the valence electrons and of different atoms.(a)nuclei(c)isotopes(b)inner electrons(d)Lewis structuresb2.A covalent bond consists of(a)a shared electron.(c)two different ions.(b)a shared electron pair.(d)an octet of electrons.a3.If two covalently bonded atoms are identical, the bond is identified asb4.A covalent bond in which there is an unequal attraction for the sharedelectrons isc5.Atoms with a strong attraction for electrons they share with another atomexhibitc6.Bonds that possess between 5% and 50% ionic character are considered to be(a)ionic.(c)polar covalent.(b)pure covalent.(d)nonpolar covalent.a7.The greater the electronegativity difference between two atoms bonded together, thegreater the bond’s percentage of8.The electrons involved in the formation of a chemical bond are called valence electrons.9.A chemical bond that results from the electrostatic attraction between positive andionic bondnegative ions is called a(n) .MODERN CHEMISTRYCHEMICAL BONDING41Copyright © by Holt, Rinehart and Winston. All rights reserved.Name Date Class

Presentation on theme: "CHEMICAL BONDING Chapter 6 – Sections 2 - 5 – Pages 178 - 207."— Presentation transcript:

1 CHEMICAL BONDING Chapter 6 – Sections 2 - 5 – Pages 178 - 207

2 What is a molecule? A neutral group of atoms that are held together by covalent bonds. A single molecule of a chemical compound is an individual unit capable of existing on its own. Bond length and bond energy between atoms in a compound are related. The shorter the bond length between two atoms in a compound, the higher the bond energy will be. If two atoms have a small bond length, then they will be expected to have a high bond energy. This means that the smaller the bond length, then larger energy is required to break this bond - more energy is required when bond length small. When bond length is large, less energy is required to break the bond between the two atoms.

3 Digital Insert “Bond Energy”

4 How to know whether a compound is ionic or covalent Generally, ionic bonds occur between atoms of the elements included in Groups 1 and 2, the alkali metals and alkaline earth metals. As we’ve discussed, these elements form cations meaning that they give up one or two electrons giving them a +1 or a +2 charge. And between atoms of the elements included from groups 16 and 17, the chalcogens and the halogens. These elements form anions meaning that they take on one or two electrons giving them a -1 or -2 charge.

5 Ionic compounds Examples include NaCl, CaO, MgO, LiCl, KCl, MgCl 2, CaF 2 When ionic bonds are formed, there is a balance of charge between the two atoms in the compound. Take NaCl for example: Na forms a +1 ion while Cl forms a -1 ion. Now look at CaF 2 : Calcium forms a 2+ ion while Fluorine only forms a -1 ion meaning that there must be TWO atoms of Fluorine present in the compound to balance the ionic charge between the two atoms of the compound.

6 Covalent Compounds Are usually compounds of two nonmetals – that is, elements that are found on the right hand side of the periodic table.

7 Digital Insert “Comparing Ionic and Covalent Compounds”

8 Resonance Is defined by bonding molecules or ions that cannot be correctly represented by a single Lewis structure.

9 Resonance Is indicated by a double-headed arrow between a molecule’s structures. This is a sulfite ion. From the depiction above, you can see that the structure Exists in three different forms. This is due to resonance.

10 Ionic bonds are different from covalent bonds in that they have transferred valance electrons to satisfy their octet. Their bonds are stronger than covalent forces. Because of this, they exhibit higher melting points and boiling points. They form what is called a crystal lattice. From this crystal lattice, every ionic compound has a lattice energy. When a chemical bond is formed, each individual atom participating in bonding lowers its potential chemical energy by bonding with another atom. All atoms “seek” a noble gas electron configuration because this is most STABLE meaning this is the lowest POTENTIAL ENERGY state.

11 An ionic compound looks like…

12 An ionic crystal lattice

13 How an ionic compound forms…

14 Ionic compounds vs Covalent compounds IONIC SUBSTANCESCOVALENT SUBSTANCES MELTING POINThighlow BOILING POINThighlow SOLUBILITY IN WATERhighlow ODORNon-detectabledetectable ELECTRICAL CONDUCTIVITYconductiveNon-conductive

15 But why do these differences exist? One word…INTERMOLECULAR FORCES. These forces are defined as the forces of attraction ‘between’ molecules. Ionic forces are stronger than covalent forces. “Inter-” = between “Intra-” = within Intramolecular forces include electronegativity and polarity

16 Some INTERmolecular forces Dipole-Dipole forces Hydrogen bonding forces London dispersion forces Refer to your textbook and be sure you are familiar with the definitions of these three intermolecular forces!!!

17 Digital Insert “Dipole-Dipole Forces” “Hydrogen Bonding” “London Dispersion Force”

18 A polar-covalent molecule - water

19 A non-polar covalent bond vs a polar covalent bond

20 MOLECULAR POLARITY What determines whether a molecule exhibits polarity? ________________ When a compound consists of atoms with greater electronegativity than other atoms in the compound, polarity exists. Polyatomic ions are defined as a charged group of covalently bonded atoms.

21 Metallic Bonding The electron-sea model: overlapping of orbitals allows outer electrons of atoms to roam freely throughout metals – which is called electron delocalization – this gives the crystal lattice a very stable structure. This is why metals have high melting points. Good conductors – electron delocalization. Malleable and ductile – all bonds are pulling in the same direction. Lustrous – absorb and reflect light at the same wavelength/frequency due to electron configuration.

22 Digital Insert “Metallic Bonding”

23 MOLECULAR GEOMETRY VSEPR (valence-shell, electron-pair repulsion) theory – repulsion between the sets of valence-level electrons causes the sets to be as far apart as possible. This is because negative charges repel neighboring negative charges. Lone pair electrons have a greater force of repulsion to other electrons than do bonding electrons.

24 Some Common Molecular Geometries

25 Hybridization The second theory explaining molecular geometry. It is the mixing of two or more atomic orbitals of similar energies on the same atom to produce a new hybrid atomic orbital of equal energy. Linear hybridizes sp Trigonal planar hybridizes sp 2 Tetrahedral hybridizes sp 3

26 Digital Insert “Hybrid Orbitals”

27 Chapter 6 Quiz 1. Why do atoms form bonds with other atoms? 2. What is the relationship between bond energy and bond length? 3. What type of bond would you expect to exist between an atom of Calcium, Ca and an atom of Oxygen, O? 4. What type of bond would you expect to exist between two atoms of Hydrogen and an atom of Oxygen? 5. Which type of substance, ionic or covalent, is expected to have a higher melting point? 6. Briefly explain the VSEPR theory. 7. What is the hybridization for a molecule of methane, CH 4 ? 8. Draw the Lewis structure for a molecule of Nitrogen gas, N 2. 9. What are the two theories that help to explain the geometry of molecules? 10. What effect do intermolecular forces have on physical properties such as melting point, boiling point, solubility and conductivity?

28 ANSWERS 1. To lower their individual atomic potential energies. 2. Inverse proportionality – when one goes up, the other goes down. 3. Ionic – Ca is from group 2; O is group 16. 4. Covalent – valence electrons are shared between atoms. 5. Ionic – the intermolecular forces are stronger. 6. Valence electrons will be as far apart as possible about an atom. 7. sp 3 8. see next slide 9. VSEPR and hybridization 10. All of these properties tend to be high in an ionic compound due to strong intermolecular forces.

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