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Chemical Bonding

This article provides an introduction to chemical bonding, including the types of bonds (ionic, covalent, and metallic) and how to distinguish between them. It also discusses electronegativity, bond polarity, dipole moments, and the importance of water's dipole moment. The article covers ionic bonding, structures of ionic compounds, covalent bonding, Lewis structures, VSEPR model, metallic bonding, and alloys.

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Chemical Bonding

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  1. Chemical Bonding

  2. I. Introduction • A. Types of Chemical Bonds – forces that hold two atom together • 1. Ionic Bonds – occur b/w a metal & a nonmetal • 2. Covalent Bonds – occur b/w 2 nonmetals & in polyatomic ions • a. Polar Covalent Bonds – • 3. Metallic Bonds – occur b/w 2 metals to form alloys

  3. B. Distinguishing b/w Types of Bonds • 1. Electronegativity – ability of an atom to attract electron’s to itself.

  4. 2. Bond Polarity: You can use the element’s electronegativities to determine the polarity of the bond – Just find the difference b/w the 2 numbers • 0 – 0.5 Covalent bond • 0.51 – 1.6 Polar Covalent bond • > 1.6 Ionic bond

  5. Determine the Polarity! • 1. H – O • 2. C – N • 3. K – F • 4. S – O • 5. Al - P

  6. What Polarity Looks Like

  7. 3. Dipole Moments – When a bond is polar… 1 side of the molecule is more positive and the other side is more negative. • Why? Because the electrons are being pulled toward the more electronegative element.

  8. Dipole Moments

  9. Why is Water’s Dipole Moment So Important? • It has a huge affect on its properties! • It’s so important, it has a specific name, it’s called HYDROGEN BONDING. • It is crucial to life on Earth! • Polar water molecules can surround & attract positive & negative ions which allows materials to dissolve in water!

  10. It’s polarity also means that water molecules are attracted to each other • A LOT of ENERGY is needed to change H2O from a liquid to a gas because the attraction must be overcome to separate 1 H2O molecule from another. • So what?!?! This causes water on Earth, at Earth’s temperatures, to remain a liquid. OTHERWISE, it would all be a gas and the oceans would be empty!!!!

  11. II. Ionic Bonds • A. A strong bond caused by the transfer of electrons from a cation (metal) to an anion (nonmetal). • 1. Why? The driving force behind this bonding is that all elements want to have a completely filled outermost energy level! [OCTET RULE] • a.) These outermost electrons are called the VALENCE ELECTRONS • b.) Metals LOSE valence electrons to be stable. • c.) Nonmetals GAIN valence electrons to be stable.

  12. Valence Electrons

  13. Let’s try it! • 1. Na and O • 2. Al and F • 3. Ca and S • 4. Mg and P

  14. B. Ionic Bonding And Structures of Ionic Compounds • 1. Ionic compounds are • a. very stable, huge amounts of energy necessary to break them apart • b. high melting & boiling points • NaCl has a melting point = ~800°C • NaCl has a boiling point = 1413 °C • c. crystals

  15. 2. Structures of Ionic Compounds • a. When you write the formula for an ionic compound, you are writing its empirical formula. • b. In reality, the actual solid contains tremendous amounts & equal numbers of cations and anions packed together so that the attractions b/w them are maximized. • 1.) Remember that cations are always smaller than anions. WHY?

  16. III. Covalent Bonding • A. Sharing electrons! • 1. All bonding involves valence electrons ONLY!!!!!! • 2. Covalent bonds occur when 2 atoms (usually nonmetals) share electrons. • 3. LEWIS STRUCTURE – a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule. • Thought up by G.N. Lewis while teaching a chemistry class in 1902.

  17. See attached page for writing Lewis Structures!

  18. B. Structures – VSEPR Model • 1. Valence Shell Electron Pair Repulsion Model • a. Useful for predicting the geometric shape of molecules formed from nonmetals! • b. The structure around a given atom is determined by minimizing repulsions between electron pairs.

  19. Metallic Bonding • How atoms are held together in the solid. • Metals hold onto their valence electrons very weakly. • Think of them as positive ions floating in a sea of electrons!

  20. + + + + + + + + + + + + Sea of Electrons! • Electrons are free to move through the solid. • Metals conduct electricity.

  21. + + + + + + + + + + + + Metals are malleable! • Hammered into shape (bend). • Ductile - drawn into wires.

  22. + + + + + + + + + + + Malleable • Electrons allow atoms to slide by.

  23. Alloys • Solutions made by dissolving metal into other elements- usually metals. • Melt them together and cool them. • If the atoms of the metals are about the same size, they substitute for each other • Called a substitutional alloy

  24. Substitutional alloy Metal B Metal A  + Bronze – Copper and Tin Brass- 60 % Copper 39% Zinc and 1%Tin 18 carat gold- 75% gold, 25%Ag or Cu

  25. Alloys • If they are different sizes the small one will fit into the spaces of the larger one • Called and interstitial alloy

  26. Interstitial Alloy Metal A Metal B  + Steel – 99% iron 1 % C Cast iron- 96% Iron, 4%C

  27. Alloys • Making an alloy is still just a mixture • Blend the properties • Still held together with metallic bonding • Most of the metals we use daily are alloys. • Designed for a purpose

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