Chemical Bonding

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

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

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 – 0.4 Covalent bond 0.41 – 1.3 Polar Covalent bond ≥ 1.4 Ionic bond

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

What Polarity Looks Like

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.

How do you know how many valence electrons an element has? Look at their e- configuration: which electrons are in the higest orbital? Ex. Li 1s2 2s1 Be 1s2 2s2 B 1s2 2s22p1 C 1s2 2s22p2 N 1s2 2s22p3

Is there an easier way to determine the # of valence e-? YES!!!!

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

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 D. conduct electricity when dissolved in water or molten.

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?

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.

Properties of Covalent Compounds Low Melting and Boiling Points Methane (CH4) melts at -183°C and boils at -163°C Sugar melts at 186°C Soft or brittle solids, or gases or liquids Poor electrical and thermal conductors

See attached page for writing Lewis Structures!

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.

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!

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

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

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

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

 + 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

Alloys If they are different sizes the small one will fit into the spaces of the larger one Called an interstitial alloy

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

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