1 / 60

STRUCTURES AND PROPERTIES OF SUBSTANCES

STRUCTURES AND PROPERTIES OF SUBSTANCES. CHAPTER 4. CHEMICAL BONDING. Only noble gases exist naturally as single, uncombined atoms . All other atoms: combined. CHEMICAL BONDS: Electrostatic forces that hold atoms together in compounds.

gzifa
Download Presentation

STRUCTURES AND PROPERTIES OF SUBSTANCES

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. STRUCTURES AND PROPERTIES OF SUBSTANCES CHAPTER 4

  2. CHEMICAL BONDING • Only noble gases exist naturally as single, uncombinedatoms. • All other atoms: combined. CHEMICAL BONDS: Electrostatic forces that hold atoms together in compounds. - In nature, systems of lower energy (more stable) tend to be favoured over system of higher energy (less stable). - Bonded atoms tend to have lower energies(ie. more favourable!)

  3. LEWIS STRUCTURES (Review) Chemical bonding involves the interaction of valenceelectrons(outermost electrons) • Lets you see exactly how many electrons are involved in the bond  helps you keep track of the number of valence electrons. • Two ways to show bonding pairs of electrons: (Dots represent a lone pair (non-bonding pair) of electrons.)

  4. IONIC BONDING • Force of attraction between oppositely charged ions. • Occurs between atoms with large differences in electronegativity(why is this? Periodic Table!) • Ionic solid: arranged in a specific sequence of repeating units – minimum possible energy. • Atoms of ionic compounds usually come from s and p blocks.

  5. PPs, pg. 165 #1-4

  6. PROPERTIES OF IONIC SOLIDS • Crystalline with smooth, shiny surfaces • Hard but brittle • Non-conductors of electricity and heat • High melting points • Many are soluble in water

  7. LATTICE ENERGY AND IONIC BONDING • Lattice energy: the amount of energy given off when ionic crystal forms from the gaseous ions of its elements. Example: MgF2 has lattice energy= 2957 kJ/mol • Same amount of energy needed to break up the ionic crystal. • Compounds with larger lattice energies have higher melting points.

  8. COVALENT BONDING • Balance between forces of attraction and repulsion that act between the nuclei and electrons of two or more atoms. • Optimum separation for atoms at which nucleus-electron attractions, nucleus-nucleus repulsions, and electron-electron repulsions achieve a balance. • Results in the sharing of pairs of electrons. • Formation of new Orbital: overlapping of atomic orbitals. • Lower energy levels than original atomic orbitals.

  9. Characteristics of Covalent Bonding • In most cases, covalent bonding allows atoms to acquire noble gas configurations. • Hydrogen must fill its s orbital. • Carbon must fill its 2s and three 2p orbitals.

  10. Bond Energy • Energy required to break the force of attraction between two atoms in a bond and to separate them. • Measures the strength of a bond. • Increase in bond energy due to increase of (-) charge in between the two nuclei. Therefore, nuclei are more attracted to the overlap orbital.

  11. PREDICTING IONIC AND COVALENT BONDS Using Physical Properties Covalent (molecular) compounds: • Exist as a soft solid, liquid, or a gas at RT. • Low melting and boiling points • Poor conductors of electricity • May not be soluble in water. Using Formula of the Compound • Two atoms with identical electronegativitiesshare electrons equally. Therefore, the bond is _______________________. • Very different electronegativities– one atom attracts electrons more. For example, sodium chloride – sodium’s valence electron has a very high probability of being found near sodium. This bond is ________________. • Electron is not actually not ‘lost’, ‘gained,’ or transferred. • “bonding contiuum”

  12. PPs 5-8

  13. Metals Properties: • Conduct electricity and heat in both solid and liquid states. • Malleable and ductile. Bonding • Difference in electronegativities not large enough to form ionic bonds. • Do not have a sufficient number of valence electrons to form covalent bonds with one another. • Electrons are shared – but different than covalent.

  14. Metallic Bond • Metals composed of densely packed core of cations. • Electrons are shared and mobile – can move throughout metal. • Force of attraction between the positively charged cations and the pool of valence electrons that move among them  metallic bond.

  15. Properties of Metals and the Free-Electron Model • Conductivity (heat and electricity): electrons can move freely throughout the metallic structure. • Malleability and Ductility: metallic bonds are non-directional. Cations can slide over one another. • Melting and Boiling Points: Group 1 metals have lower melting and boiling points of Group 2  greater number of valence electrons and larger positive charge = stronger metallic bonding forces. • Transition metals: generally high melting and boiling points.

  16. SR, page. 171 #1, 2-7.

  17. 4.2 – Molecular Shape and Polarity

  18. Co-ordinate Covalent Bonds • When one atom contributes both of the electrons to the shared pair. • Occurs when a filled atomic orbital overlaps with an empty atomic orbital. • Behaves in the same way as any other single covalent bond.

  19. Resonance Structures: More than one Possible Lewis Structure • Bonds between S and O are identical: two “one-and-a-half” bonds. • Resonance structures are combinations – hybrids – if its two resonance structures. • Does NOT shift back and forth between bonds. • Resonance structures: models that give the same relative position of atoms as in Lewis structures, but show different places for their bonding and lone pairs.

  20. PPs, page 177 #9-12.

  21. Central Atom with an Expanded Valence Level • Expanded valence energy level: bonding in some molecules is best explained by a model that shows more than eight electrons in the valence energy level of the central atom. • Experimental evidence suggests that larger atoms can accommodate additional valence electrons because of their size. • Sometimes, must violate the octet rule to allow for more than four bonds around a central atom.

  22. PPs, pg. 178 #14-17

  23. Shapes and Polarity of Molecules • Lewis structures do not communicate any information about a molecule’s shape. • VSEPR (Valence-Shell Electron-Pair Repulsion) Theory: • Bonding pairs and lone pairs of electrons repel one another. • Lone pair (LP) will spread out more than a bond pair  repulsion is greatest between lone pairs (LP – LP) • Bonding pairs (BP) are more localized between nuclei, so spread out less.  BP-BP repulsions are smaller than LP-LP repulsions. • Repulsion between BP and LP is intermediate.

  24. Electron-Group Arrangements When all electron groups are bonding pairs, a molecule will have one of these shapes. If there are lone pairs, variations of these result.

  25. Molecular Geometry Shape of molecule depends on electron pairs.

  26. Predicting Molecular Shape

  27. PPs, 185 #18-22

  28. Molecular Shape and Polarity Recall: For diatomic molecules: bond polarity is also the molecule’s polarity. Dipole: term used to describe the charge separation for an entire molecule. Molecular Polarity/Molecular Shape Table: • A: central atom • X: more electronegative than A. • Y: more electronegative than X.

  29. PPs, page 188 # 23-26 SR, page 189. # 1-3, 5, 6

  30. 4.3 – Intermolecular Forces in Liquids and Solids Intramolecular Forces: forces exerted within a molecule or polyatomic ion. Intermolecular Forces: forces of attraction and repulsion that act between molecules or ions. • also called ‘van der Waals forces.’ • Dipole-dipole forces • Ion-dipole forces • Induced dipole forces • Dispersion forces • Hydrogen bonding

  31. Dipole-Dipole Forces • Polar molecules (dipoles) in liquid form orient themselves so that oppositely charged ends are near to one another. • Polar molecules will tend to attract one another more at room temperature than similarly sized non-polar molecules would. • More energy required to separate polar molecules than non-polar of similar molar mass. • Higher melting and boiling points.

  32. Ion-Dipole Forces • Force of attraction between an ion and a dipole. • Reason why most ionic solids are soluble.

  33. Induced Intermolecular Forces • Induced by charge • Ion-induced dipole force: • When an ion in close proximity to a non-polar molecule distorts the electron density of the non-polar molecule. • Molecule becomes temporarily polarized  two species are attracted to each other • Dipole-induced dipole force: • Charge on polar molecule induces the charge on a non-polar molecule.

  34. Dispersion (London) Forces • Shared pairs of electrons in covalent bonds are constantly vibrating. • Bond vibrations cause momentary, uneven distributions of charge. • Act between any particles. • Factors of Magnitude: • Number of electrons: larger molecules, more uneven distribution of charge. Raise boiling point. • Shape of molecule: sphere has less surface area than linear molecule. More linear molecules have higher boiling points than spherical.

  35. Hydrogen Bonding • Strong form of dipole-dipole attraction that exists between a hydrogen atom in a polar-bonded molecule that contains bonds such as H-O, H-N, H-F, and an unshared pair of electrons on another small, electronegative atom such as O, N, or F. • Small, electronegative atom can be on its own, but is usually bonded in a molecule. • H-bond is about 5% as strong as a single covalent bond  strength in numbers. • Ex// DNA

  36. Hydrogen bonding and properties of water • Why is water less dense in a solid state than in a 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. • If water molecules behave as most molecules do, lake water would freeze from the bottom up  floating ice insulates the water beneath it. • Polar covalent compounds are soluble in water • Ex// alcohols (O-H bonds), ammonia (N-H bonds), small mass amines (N-H), etc.

  37. INTRAMOLECULAR AND INTERMOLECULAR FORCES

  38. Bonding in solids • Crystalling solids: organized particle arrangements with distinct shapes: gemstones, etc. • Amorphous solids: indistinct shapes: particle arrangements lack order: glass and rubber.

  39. Crystalline solids: 5 types • Atomic • Molecular • Network • Ionic • Metallic

More Related